Optical Activity
Optical activity describes the phenomenon by which chiral molecules are observed to rotate polarized light in either a clockwise or counterclockwise direction. This rotation is a result of the properties inherent in the interaction between light and the individual molecules through which it passes. Material that is either achiral or equal mixtures of each chiral configuration (called a racemic mixture) do not rotate polarized light, but when a majority of a substance has a certain chiral configuration the plane can be rotated in either direction.
What Is Plane Polarized Light?
Polarized light consists of waves of electromagnetic energy in the visible light spectrum where all of the waves are oscillating in the same direction simultaneously. Put simply, imagine a ray of light as a water wave, with crests and peaks. All the peaks of a water wave point in the same direction (up, against gravity) pretty much at the same time. Light is not usually this way - it's peaks and troughs are often in random array, so one ray of lights peaks might point in a direction 90o opposite of another ray. When all of the rays have their peaks pointing in the same direction - like all the waves in the ocean have peaks pointing up - then those rays of light are said to be polarized to one another.
Why Polarized Light Is Affected
So why do chiral molecules affect only polarized light, and not unpolarized? Well, they do affect unpolarized light, but since the rays have no particular orientation to one another, the effect can not be observed or measured. We observe the polarized light rays being rotated because we knew their orientation before passing through the chiral substance, and so we can measure the degree of change afterwards.
What happens is this; when light passes through matter, e.g. a solution containing either chiral or achiral molecules, the light is actually interacting with each molecule's electron cloud, and these very interactions can result in the rotation of the plane of oscillation for a ray of light. The direction and magnitude of rotation depends on the nature of the electron cloud, so it stands to reason that two identical molecules possessing identical electron clouds will rotate light in the exact same manner. This is why achiral molecules do not exhibit optical activity.
In a chiral solution that is not a racemic mixture, however, the chiral molecules present in greater numbers are configurationally equivalent to each other, and therefore each possesses identical electron clouds to its molecular twins. As such, each interaction between light and one of these 'majority' molecule's electron clouds will result in rotations of identical magnitude and direction. When these billions of billions of interactions are summed together into one cohesive number, they do not cancel one another as racemic and achiral solutions tend to do - rather, the chiral solution as a whole is observed to rotate polarized light in one particular direction due to its molecular properties.
Enantiomers
It is just such specificity that accounts for the optical isomerism of enantiomeric compounds. Enantiomers possess identical chemical structures (i.e. their atoms are the same and connected in the same order), but are mirror images of one another. Therefore, their electron clouds are also identical but actually mirror images of one another and not superimposable. For this reason, enantiomeric pairs rotate light by the same magnitude (number of degrees), but they each rotate plane polarized light in opposite directions. If one chiral version has the property of rotating polarized light to the right (clockwise), it only makes sense that the molecule's chiral mirror image would rotate light to the left (counterclockwise).
Equal amounts of each enantiomer results in no rotation. Mixtures of this type are called racemic mixtures, and they behave much as achiral molecules do.
History
Via a magneto-optic effect, the (-)-form of an optical isomer rotates the plane of polarization of a beam of polarized light that passes through a quantity of the material in solution counterclockwise , the (+)-form clockwise. It is due to this property that it was discovered and from which it derives the name optical activity. The property was first observed by J.-B. Biot in 1815, and gained considerable importance in the sugar industry, analytical chemistry, and pharmaceuticals.
Louis Pasteur deduced in 1848 that the handedness of molecular structure is responsible for optical activity. He sorted the chiral crystals of tartaric acid salts into left-handed and right-handed forms, and discovered that the solutions showed equal and opposite optical activity.
Artificial composite materials displaying the analog of optical activity but in the microwave regime were introduced by J.C. Bose in 1898, and gained considerable attention from the mid-1980s.
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Web sites
http://www.goiit.com/chapters/tutorial/chemistry/isomerism.htm
Sunday, December 23, 2007
Saturday, December 22, 2007
Tips to Tame the IIT-JEE
The primary emphasis is to be on depth of knowledge, analytical and comprehension skills and attitude.
If you are a class 11th/12th student, try to synergise your school study with the JEE study. While doing a new topic, always do it first from your school textbook followed by higher level books.
Self Study Plan: School Students: A student going to the school should follow the 4/10 plan, that is self-study for at least 4 hrs. on school days and at least 10 hrs. on holidays. There must be a 7 day or 10 day revision plan as well. 12th Pass Students: A student not going to the school should follow the 10 hrs. plan, that is self-study for at least 10 hrs. everyday. There must be a 7 day or 10 day revision plan as well.
It is more important to do a question completely rather than trying to do more half-done questions.
Note that too much of test taking does not help. Only a deep understanding and other personal attributes can get you through JEE and not blind test taking. The frequency of test taking may be less for JEE 2007 aspirants in class 11th but needs to be higher when they reach class 12th.
Do approximations in calculations keeping an eye on the error.
In objective type of questions, method of elimination of options may work to your advantage in many questions. Using dimensional analysis, putting boundary conditions, putting values of variables, working backwards and many more techniques may work in such scenario.
http://iitjee-aieee-cbse.blogspot.com/2007/11/12-points-to-tame-new-pattern-iit-jee.html
If you are a class 11th/12th student, try to synergise your school study with the JEE study. While doing a new topic, always do it first from your school textbook followed by higher level books.
Self Study Plan: School Students: A student going to the school should follow the 4/10 plan, that is self-study for at least 4 hrs. on school days and at least 10 hrs. on holidays. There must be a 7 day or 10 day revision plan as well. 12th Pass Students: A student not going to the school should follow the 10 hrs. plan, that is self-study for at least 10 hrs. everyday. There must be a 7 day or 10 day revision plan as well.
It is more important to do a question completely rather than trying to do more half-done questions.
Note that too much of test taking does not help. Only a deep understanding and other personal attributes can get you through JEE and not blind test taking. The frequency of test taking may be less for JEE 2007 aspirants in class 11th but needs to be higher when they reach class 12th.
Do approximations in calculations keeping an eye on the error.
In objective type of questions, method of elimination of options may work to your advantage in many questions. Using dimensional analysis, putting boundary conditions, putting values of variables, working backwards and many more techniques may work in such scenario.
http://iitjee-aieee-cbse.blogspot.com/2007/11/12-points-to-tame-new-pattern-iit-jee.html
Monday, December 17, 2007
Another good website
Chemistry lessons Videos
Chemistry E books
Chemistry Lecture notes
http://www.learnerstv.com
http://www.learnerstv.com/lecturenotes/lecturenotes.php?note=15&cat=Chemistry
Chemistry E books
Chemistry Lecture notes
http://www.learnerstv.com
http://www.learnerstv.com/lecturenotes/lecturenotes.php?note=15&cat=Chemistry
Sunday, December 16, 2007
IIT JEE Preparation Online Material
I am putting in effort to put in useful material as number of persons are visiting the blog and visiting number of pages.
I hope they are getting an additional version on the topic apart from their regular institute faculty, coaching instiute faculty, the books they are referring. In my personal study, I realise that reading a variety of books is necessary to clarify certain concepts.
I hope persons who are visiting the blog are getting similar benefit. Some concept is now more clear, as it is explained in a slightly different manner.
I started a blog to develop chemistry glossary
www.chemgloss.blogspot.com
I hope they are getting an additional version on the topic apart from their regular institute faculty, coaching instiute faculty, the books they are referring. In my personal study, I realise that reading a variety of books is necessary to clarify certain concepts.
I hope persons who are visiting the blog are getting similar benefit. Some concept is now more clear, as it is explained in a slightly different manner.
I started a blog to develop chemistry glossary
www.chemgloss.blogspot.com
IIT JEE Chemistry Ch.16A. Coordination Compounds
See for a set of questions on this topic the post
http://iit-jee-chemistry-ps.blogspot.com/2007/10/iit-jee-chemistry-questions.html
----------------------
Syllabus
Nomenclature of mononuclear coordination compounds, XII 10.1,2,3
Cis-trans and ionisation isomerisms, 10.4
Hybridization and geometries of mononuclear coordination compounds (linear, tetrahedral, square planar and octahedral).10.7
The topics are covered in detail Jauhar's XII book. The section numbers are given beside the topic.
-----------------------
Coordination compounds are a special class of compounds in which the centgral metal atom is surrounded by ions or molecules beyond their valency.
There are also referred to as coordination complexes or complexes.
Haemoglobin, Chlrophyll, and vitamin B-12 are coordinatio compounds of iron, magnesium and cobalt respectively.
The interesting thing of coordination compound is that these are formed from apparently saturated molecules capable of independent existence.
for example, when acqueous ammonia is addedt o green solution of nickel chloride, NiCl2, the colour changes to purple. The ni^2+ ions almost diappear from the solution. The solution on evaporation yields purple crystals corresponding to the formula [Ni(NH-3)-6]Cl-2. such a compound is called coordinatin compound. When this compound is now dissolved in water, there is hardly any evidence of Ni^2+ ions or NH-3 molecules. It ionizes to give a new species [Ni(NH-3)-6]^2+. the species in the square brackets does not ionise further. It remains as a single entity as an ion.
This is the unique feature of coordination compounds.
---------------
Nomenclature rules
From
http://www.iupac.org/publications/books/principles/principles_of_nomenclature.pdf
IUPAC booklet available for download at the above page id.
A coordination entity is composed of a central atom or atoms to which are attached
other atoms or groups of atoms, which are termed ligands. A central atom occupies a
central position within the coordination entity. The ligands attached to a central
atom define a coordination polyhedron. Each ligand is assumed to be at the vertex of
an appropriate polyhedron. The usual polyhedra are shown in Table 3.3 and they are
also listed in Table 4.4. Note that these are adequate to describe most simple
coordination compounds, but that real molecules do not always fall into these simple
categories. In the presentation of a coordination polyhedron graphically, the lines
defining the polyhedron edges are not indicative of bonds.
However, many ligands do not behave as donors of a single electron pair. Some
ligands donate two or more electron pairs to the same central atom from different
donor atoms. Such ligands are said to be chelating ligands, and they form chelate
rings, closed by the central atom. The phenomenon is termed chelation.
The number of electron pairs donated by a single ligand to a specific central atom
is termed the denticity. Ligands that donate one pair are monodentate, those that
donate two are didentate, those that donate three are tridentate, and so on.
Sometimes ligands with two or more potential donor sites bond to two (or more)
different central atoms rather than to one, forming a bridge between central atoms. It
may not be necessary for the ligand in such a system to be like ethane-l ,2-diamine,
with two distinct potential donor atoms. A donor atom with two or more pairs of
non-bonding electrons in its valence shell can also donate them to different centralatoms. Such ligands, of whatever type, are called bridging ligands. They bond to two
or more central atoms simultaneously. The number of central atoms in a single
coordination entity is denoted by the nuclearity: mononuclear, dinuclear, trinuclear,
etc. Atoms that can bridge include 5, 0 and Cl.
The original concepts of metal—ligand bonding were essentially related to the
dative covalent bond; the development of organometallic chemistry has revealed a
further way in which ligands can supply more than one electron pair to a central
atom. This is exemplified by the classical cases of bis(benzene)chromium and
bis(cyclopentadienyl)iron, trivial name ferrocene. These molecules are characterised
by the bonding of a formally unsaturated system (in the organic chemistry sense, but
expanded to include aromatic systems) to a central atom, usually a metal atom.
4.4.3 Mononuclear coordination compounds
4.4.3.1 Formulae. The central atom is listed first. The formally anionic ligands appear next,listed in alphabetical order of the first symbols of their individual formulae. The neutral ligands follow, also in alphabetical order. Polydentate ligands are included in alphabetical order, the formula to be presented as discussed in Chapter 3. The formula for the entire coordination entity, whether charged or not, is enclosed in square brackets. For coordination formulae, the nesting order of enclosing marks is as given on p. 13. The charge on an ion is indicated in the usual way by use of a right superscript. Oxidation states of particular atoms are indicated by an appropriate roman numeral as a right superscript to the symbol of the atom in question, and not in parentheses on the line. In the formula of a salt containing coordination entities, cation always precedes anion, no charges are indicated and there is no space between the formulae for cation and anion.
Examples
1. [Co(NH3)6]Cl
2. [PtC14]2
3. [CoC1(NH3)5]C1
6. [Cr"(NCS)4(NH3)2]
4. Na[PtBrC1(N02)(NH3)J
7. [Fe"(CO)4]2
5. [CaC12{OC(NH2)2}2]
The precise form of a formula should be dictated by the needs of the user.
The precise form of a formula should be dictated by the needs of the user. For
example, it is generally recommended that a ligand formula within a coordination
formula be written so that the donor atom comes first, e.g. [TiCl3(NCMe)3], but this
is not mandatory and should not affect the recommended order ofligand citation. It
may also be impossible to put all the donor atoms first, e.g. where two donors are
present in a chelate complex: [Co(NH2CH2CH2NH2)3]3. Whether the ethane-l,2-
diamine is displayed as shown, or simply aggregated as [Co(C2H8N2)3]3t is a matter
of choice. Certainly there is a conflict between this last form and the suggestion that
the donor atoms be written first. The aim should always be clarity, at the expense of
rigid adherence to recommendations.
It is often inconvenient to represent all the ligand formulae in detail. Abbreviations
are often used and are indeed encouraged, with certain provisos. These are: the
abbreviations should all be written in lower case (with minor exceptions, such as Me,
Et and Ph) and preferably not more than four letters; with certain exceptions of wide
currency, abbreviations should be defined in a text when they first appear; in a
formula, the abbreviation should be enclosed in parentheses, and its place in the
citation sequence should be determined by its formula, as discussed above; and
particular attention should be paid to the loss of hydrons from a ligand precursor.
This last proviso is exemplified as follows. Ethylenediaminetetraacetic acid
should be rendered H4edta. The ions derived from it, which are often ligands in
coordination entities, are then (H3edta), (H2edta)2, (Hedta)3 and (edta)4. This
avoids monstrosities such as edta-H2 and edtaH_2 which arise if the parent acid is
represented as edta. A list of recommended abbreviations is presented in Table 4.5.
4.4.3.2 Names. The addition of ligands to a central atom is paralleled in name construction.
The names of the ligands are added to that of the central atom. The ligands are listed
in alphabetical order regardless of ligand type. Numerical prefixes are ignored in this
ordering procedure, unless they are part of the ligand name. Charge number and
oxidation number are used as necessary in the usual way.
Of the two kinds of numerical prefix (see Table 4.2), the simple di-, tn-, tetra-, etc.
are generally recommended. The prefixes bis-, tris-, tetrakis-, etc. are to be used only
with more complex expressions and to avoid ambiguity. They normally require
parentheses around the name they qualify. The nesting order of enclosing marks is as
cited on p. 13. There is normally no elision in instances such as tetraammine and the
two adjacent letters 'a' are pronounced separately.
The names of ligands recommended for general purposes are given in Table 4.6.
The names for anionic ligands end in -o. If the anion name ends in -ite, -ate or -ide, the ligand name is changed to -ito, -ato or -ido. The halogenido names are, by
custom, abbreviated to halo. Note that hydrogen as a ligand is always regarded as
anionic, with the name hydride. The names of neutral and cationic ligands are never
modified. Water and ammonia molecules as ligands take the names aqua and
ammine, respectively. Parentheses are always placed around ligand names, which
themselves contain multiplicative prefixes, and are also used to ensure clarity, but
aqua, ammine, carbonyl (CO) and nitrosyl (NO) do not require them.
The names of all cationic and neutral entities end in the name of the element,
together with the charge (if appropriate) or the oxidation state (if desired). The
names of complex anions require modification, and this is achieved by adding the
termination -ate. All these recommendations are illustrated in the following examples.
----------------
Examples
1. Dichloro(diphenylphosphine)(thiourea)platinum(ii)
2. K4[Fe(CN)6]
3. [Co(NH3)6]C13
4. [CoCl(NH3)5}Cl2
5. [CoC1(N02)(NH3)4]Cl
6. [PtCl(NH2CH3)(NH3)2]Cl
7. [CuC12{OC(NH2)2}2]
8. K2[PdC14]
9. K[OsCl5N]
10. Na[PtBrC1(N02)(NH3)]
11. [Fe(CNCH3)6]Br2
12. [Ru(HSO3)2(NH3)4]
13. [Co(H20)2(NH3)4]C13
14. [PtC12(C5H5N)(NH3)]
15. Ba[BrF4]2
16. K[CrF4O]
17. [Ni(H20)2(NH3)4]S04
potassium hexacyanoferrate(ii)
potassium hexacyanoferrate(4—)
tetrapotassium hexacyanoferrate
hexaamminecobalt(iii) chloride
pentaamminechlorocobalt(2+) chloride
tetraamminechloronitrito-N-cobalt(iii) chloride
diamminechloro(methylamine)platinum(ii)
chloride
dichlorobis(urea)copper(ii)
potassium tetrachloropalladate(ii)
potassium pentachloronitridoosmate(2—)
sodium amminebromochloronitrito-
N-platinate( 1—)
hexakis(methyl isocyanide)iron(ii) bromide
tetraamminebis(hydrogensulfito)ruthenium(ii)
tetraamminediaquacobalt(iii) chloride
amminedichloro(pyridine)platinum(ii)
barium tetrafluorobromate(iii)
potassium tetrafiuorooxochromate(v)
tetraamminediaquanickel(ii) sulfate
***Table to be reformatted
-----------------------------
Designation of donor atom. In some cases, it may not be evident which atom in a
ligand is the donor. This is exemplified by the nitrito ligand in Examples 5 and 10,
p. 59. This can conceivably bind through an 0 or N atom. In simple cases, the donor
atom can be indicated by italicised element symbols placed after the specific ligand
name and separated from it by a hyphen, as demonstrated in those particular
examples. More complex examples will be dealt with below. With polydentate
ligands, this device may still be serviceable. Thus, dithiooxalate ion may be attached through S or 0, and formulations such as dithiooxalato-S,S' and dithiooxalato-0,0'should suffice. It could be necessary to use superscripts to the donor atom symbols if these need to be distinguished because there is more than one atom of the same kind to choose from.
Complicated examples are more easily dealt with using the kappa convention,
and this is particularly useful where a donor atom is part of a group that does not
carry a locant according to organic rules. The two oxygen atoms in a carboxylato
group demonstrate this. The designator i is a locant placed after that portion of the
ligand name that denotes the particular function in which the ligating atom is found.
The ligating atoms are represented by superscript numerals, letters or primes affixed
to the donor element symbols, which follow i without a space. A right superscript to
i denotes the number of identically bound ligating atoms.
Inclusion of structural information. The names described so far detail ligands and
central atoms, but give no information on stereochemistry. The coordination number
and shape of the coordination polyhedron may be denoted, if desired, by a
polyhedral symbol. These are listed in Table 4.4. Such a symbol is used as an affix in
parentheses, and immediately precedes the name, separated from it by a hyphen.
This device is not often used.
Geometrical descriptors, such as cis, trans, mer (from meridional) and fac (from
facial), have found wide usage in coordination nomenclature. The meaning is
unequivocal only in simple cases, particularly square planar for the first two and
octahedral for the others
More complex devices have been developed that are capable of dealing with all
cases. The reader is referred to the Nomenclature of Inorganic Chemistry, Chapter
10.
------------------------
Ionisation Isomers
Ionisation isomers:
Molecular structural formula is same. But different isomers give different ions in solution.
one isomer [PtBr(NH3)3]NO2 -> gives NO2- anions in solution
another isomer [Pt(NH3)3(NO2)]Br -> gives Br- anions in solution
Notice that both anions are necessary to balance the charge of the complex, and that they differ in that one ion is directly attached to the central metal but the other is not.
Geometric Isomers or Cis-Trans isomers
Geometric isomers are two or more coordination compounds which contain the same number and types of atoms, and bonds (i.e., the connectivity between atoms is the same), but which have different spatial arrangements of the atoms.
Not all coordination compounds have geometric isomers.
For example, in the square planar molecule, Pt(NH3)2Cl2, the two ammonia ligands (or the two chloride ligands) can be adjacent to one another or opposite one another.
Note that these two structures contain the same number and kinds of atoms and bonds but are non-superimposable. The isomer in which like ligands are adjacent to one another is called the cis isomer. The isomer in which like ligands are opposite one another is called the trans isomer.
For the common structures which contain two or more different ligands, geometric isomers are possible only with square planar and octahedral structures (i.e., geometric isomers cannot exist for linear and tetrahedral structures).
cis-[Co(NH3)4Cl2]+
Note that the two chloride ligands are adjacent to one another in this octahedral complex ion. In aqueous solution, this complex ion has a violet color.
trans-[Co(NH3)4Cl2]+
Note that the two chloride ligands are opposite one another in this complex ion. In aqueous solution, this complex ion has a green color.
------------------
Material about all isomers - In syllabus only ionic and cis-trans are specially mentioned
http://www.chem.purdue.edu/gchelp/cchem/whatis2.html
Coordination Isomers
Coordination isomers are two or more coordination compounds in which the composition within the coordination sphere (i.e., the metal atom plus the ligands that are bonded to it) is different (i.e., the connectivity between atoms is different).
Not all coordination compounds have coordination isomers.
Coordination isomers have different physical and chemical properties.
Example
[Cr(NH3)5(OSO3)]Br
Note that the sulfate group is bonded to the Cr atom (via an O atom) and is within the coordination sphere. Note also the octahedral structure. The bromide counterion is needed to maintain charge neutrality with the complex ion (i.e., [Cr(NH3)5(OSO3)]+) and is not shown in the structure.
[Cr(NH3)5Br]SO4
Note that the bromine atom is bonded to the Cr atom and is within the coordination sphere. Note also the octahedral structure. The sulfate counterion is not shown in the structure.
Linkage Isomers
Linkage isomers are two or more coordination compounds in which the donor atom of at least one of the ligands is different (i.e., the connectivity between atoms is different).
This type of isomerism can only exist when the compound contains a ligand that can bond to the metal atom in two (or more) different ways. Some ligands that can form linkage isomers are shown below.
Not all coordination compounds have linkage isomers.
Linkage isomers have different physical and chemical properties.
[Co(NH3)4(NO2)Cl]+
Note that the N atom of the nitrite group is bonded to the Co atom. The nitrite group is written as "NO2" in the molecular formula (rather than "ONO") with the N atom nearest to the Co symbol to indicate that the N atom (rather than an O atom) is the donor atom. Note also the octahedral structure.
[Co(NH3)4(ONO)Cl]+
Note that one of the O atoms of the nitrite group is bonded to the Co atom. The nitrite group is written as "ONO" in the molecular formula (rather than "NO2") with the O atom nearest to the Co symbol to indicate that the O atom is the donor atom. Note also the octahedral structure.
Geometric Isomers
Geometric isomers are two or more coordination compounds which contain the same number and types of atoms, and bonds (i.e., the connectivity between atoms is the same), but which have different spatial arrangements of the atoms.
Not all coordination compounds have geometric isomers.
For example, in the square planar molecule, Pt(NH3)2Cl2, the two ammonia ligands (or the two chloride ligands) can be adjacent to one another or opposite one another.
Note that these two structures contain the same number and kinds of atoms and bonds but are non-superimposable. The isomer in which like ligands are adjacent to one another is called the cis isomer. The isomer in which like ligands are opposite one another is called the trans isomer.
For the common structures which contain two or more different ligands, geometric isomers are possible only with square planar and octahedral structures (i.e., geometric isomers cannot exist for linear and tetrahedral structures).
cis-[Co(NH3)4Cl2]+
Note that the two chloride ligands are adjacent to one another in this octahedral complex ion. In aqueous solution, this complex ion has a violet color.
trans-[Co(NH3)4Cl2]+
Note that the two chloride ligands are opposite one another in this complex ion. In aqueous solution, this complex ion has a green color.
Optical Isomers
Optical isomers are two compounds which contain the same number and kinds of atoms, and bonds (i.e., the connectivity between atoms is the same), and different spatial arrangements of the atoms, but which have non-superimposable mirror images. Each non-superimposable mirror image structure is called an enantiomer. Molecules or ions that exist as optical isomers are called chiral.
Not all coordination compounds have optical isomers.
The Two Enantiomers of CHBrClF
Note that the molecule on the right is the reflection of the molecule on the left (through the mirror plane indicated by the black vertical line). These two structures are non-superimposable and are, therefore, different compounds.
Pure samples of enantiomers have identical physical properties (e.g., boiling point, density, freezing point). Chiral molecules and ions have different chemical properties only when they are in chiral environments.
Optical isomers get their name because the plane of plane-polarized light that is passed through a sample of a pure enantiomer is rotated. The plane is rotated in the opposite direction but with the same magnitude when plane-polarized light is passed through a pure sample containing the other enantiomer of a pair.
---------------------
web sites
http://www.chem.purdue.edu/gchelp/cchem/whatis2.html
http://iit-jee-chemistry-ps.blogspot.com/2007/10/iit-jee-chemistry-questions.html
----------------------
Syllabus
Nomenclature of mononuclear coordination compounds, XII 10.1,2,3
Cis-trans and ionisation isomerisms, 10.4
Hybridization and geometries of mononuclear coordination compounds (linear, tetrahedral, square planar and octahedral).10.7
The topics are covered in detail Jauhar's XII book. The section numbers are given beside the topic.
-----------------------
Coordination compounds are a special class of compounds in which the centgral metal atom is surrounded by ions or molecules beyond their valency.
There are also referred to as coordination complexes or complexes.
Haemoglobin, Chlrophyll, and vitamin B-12 are coordinatio compounds of iron, magnesium and cobalt respectively.
The interesting thing of coordination compound is that these are formed from apparently saturated molecules capable of independent existence.
for example, when acqueous ammonia is addedt o green solution of nickel chloride, NiCl2, the colour changes to purple. The ni^2+ ions almost diappear from the solution. The solution on evaporation yields purple crystals corresponding to the formula [Ni(NH-3)-6]Cl-2. such a compound is called coordinatin compound. When this compound is now dissolved in water, there is hardly any evidence of Ni^2+ ions or NH-3 molecules. It ionizes to give a new species [Ni(NH-3)-6]^2+. the species in the square brackets does not ionise further. It remains as a single entity as an ion.
This is the unique feature of coordination compounds.
---------------
Nomenclature rules
From
http://www.iupac.org/publications/books/principles/principles_of_nomenclature.pdf
IUPAC booklet available for download at the above page id.
A coordination entity is composed of a central atom or atoms to which are attached
other atoms or groups of atoms, which are termed ligands. A central atom occupies a
central position within the coordination entity. The ligands attached to a central
atom define a coordination polyhedron. Each ligand is assumed to be at the vertex of
an appropriate polyhedron. The usual polyhedra are shown in Table 3.3 and they are
also listed in Table 4.4. Note that these are adequate to describe most simple
coordination compounds, but that real molecules do not always fall into these simple
categories. In the presentation of a coordination polyhedron graphically, the lines
defining the polyhedron edges are not indicative of bonds.
However, many ligands do not behave as donors of a single electron pair. Some
ligands donate two or more electron pairs to the same central atom from different
donor atoms. Such ligands are said to be chelating ligands, and they form chelate
rings, closed by the central atom. The phenomenon is termed chelation.
The number of electron pairs donated by a single ligand to a specific central atom
is termed the denticity. Ligands that donate one pair are monodentate, those that
donate two are didentate, those that donate three are tridentate, and so on.
Sometimes ligands with two or more potential donor sites bond to two (or more)
different central atoms rather than to one, forming a bridge between central atoms. It
may not be necessary for the ligand in such a system to be like ethane-l ,2-diamine,
with two distinct potential donor atoms. A donor atom with two or more pairs of
non-bonding electrons in its valence shell can also donate them to different centralatoms. Such ligands, of whatever type, are called bridging ligands. They bond to two
or more central atoms simultaneously. The number of central atoms in a single
coordination entity is denoted by the nuclearity: mononuclear, dinuclear, trinuclear,
etc. Atoms that can bridge include 5, 0 and Cl.
The original concepts of metal—ligand bonding were essentially related to the
dative covalent bond; the development of organometallic chemistry has revealed a
further way in which ligands can supply more than one electron pair to a central
atom. This is exemplified by the classical cases of bis(benzene)chromium and
bis(cyclopentadienyl)iron, trivial name ferrocene. These molecules are characterised
by the bonding of a formally unsaturated system (in the organic chemistry sense, but
expanded to include aromatic systems) to a central atom, usually a metal atom.
4.4.3 Mononuclear coordination compounds
4.4.3.1 Formulae. The central atom is listed first. The formally anionic ligands appear next,listed in alphabetical order of the first symbols of their individual formulae. The neutral ligands follow, also in alphabetical order. Polydentate ligands are included in alphabetical order, the formula to be presented as discussed in Chapter 3. The formula for the entire coordination entity, whether charged or not, is enclosed in square brackets. For coordination formulae, the nesting order of enclosing marks is as given on p. 13. The charge on an ion is indicated in the usual way by use of a right superscript. Oxidation states of particular atoms are indicated by an appropriate roman numeral as a right superscript to the symbol of the atom in question, and not in parentheses on the line. In the formula of a salt containing coordination entities, cation always precedes anion, no charges are indicated and there is no space between the formulae for cation and anion.
Examples
1. [Co(NH3)6]Cl
2. [PtC14]2
3. [CoC1(NH3)5]C1
6. [Cr"(NCS)4(NH3)2]
4. Na[PtBrC1(N02)(NH3)J
7. [Fe"(CO)4]2
5. [CaC12{OC(NH2)2}2]
The precise form of a formula should be dictated by the needs of the user.
The precise form of a formula should be dictated by the needs of the user. For
example, it is generally recommended that a ligand formula within a coordination
formula be written so that the donor atom comes first, e.g. [TiCl3(NCMe)3], but this
is not mandatory and should not affect the recommended order ofligand citation. It
may also be impossible to put all the donor atoms first, e.g. where two donors are
present in a chelate complex: [Co(NH2CH2CH2NH2)3]3. Whether the ethane-l,2-
diamine is displayed as shown, or simply aggregated as [Co(C2H8N2)3]3t is a matter
of choice. Certainly there is a conflict between this last form and the suggestion that
the donor atoms be written first. The aim should always be clarity, at the expense of
rigid adherence to recommendations.
It is often inconvenient to represent all the ligand formulae in detail. Abbreviations
are often used and are indeed encouraged, with certain provisos. These are: the
abbreviations should all be written in lower case (with minor exceptions, such as Me,
Et and Ph) and preferably not more than four letters; with certain exceptions of wide
currency, abbreviations should be defined in a text when they first appear; in a
formula, the abbreviation should be enclosed in parentheses, and its place in the
citation sequence should be determined by its formula, as discussed above; and
particular attention should be paid to the loss of hydrons from a ligand precursor.
This last proviso is exemplified as follows. Ethylenediaminetetraacetic acid
should be rendered H4edta. The ions derived from it, which are often ligands in
coordination entities, are then (H3edta), (H2edta)2, (Hedta)3 and (edta)4. This
avoids monstrosities such as edta-H2 and edtaH_2 which arise if the parent acid is
represented as edta. A list of recommended abbreviations is presented in Table 4.5.
4.4.3.2 Names. The addition of ligands to a central atom is paralleled in name construction.
The names of the ligands are added to that of the central atom. The ligands are listed
in alphabetical order regardless of ligand type. Numerical prefixes are ignored in this
ordering procedure, unless they are part of the ligand name. Charge number and
oxidation number are used as necessary in the usual way.
Of the two kinds of numerical prefix (see Table 4.2), the simple di-, tn-, tetra-, etc.
are generally recommended. The prefixes bis-, tris-, tetrakis-, etc. are to be used only
with more complex expressions and to avoid ambiguity. They normally require
parentheses around the name they qualify. The nesting order of enclosing marks is as
cited on p. 13. There is normally no elision in instances such as tetraammine and the
two adjacent letters 'a' are pronounced separately.
The names of ligands recommended for general purposes are given in Table 4.6.
The names for anionic ligands end in -o. If the anion name ends in -ite, -ate or -ide, the ligand name is changed to -ito, -ato or -ido. The halogenido names are, by
custom, abbreviated to halo. Note that hydrogen as a ligand is always regarded as
anionic, with the name hydride. The names of neutral and cationic ligands are never
modified. Water and ammonia molecules as ligands take the names aqua and
ammine, respectively. Parentheses are always placed around ligand names, which
themselves contain multiplicative prefixes, and are also used to ensure clarity, but
aqua, ammine, carbonyl (CO) and nitrosyl (NO) do not require them.
The names of all cationic and neutral entities end in the name of the element,
together with the charge (if appropriate) or the oxidation state (if desired). The
names of complex anions require modification, and this is achieved by adding the
termination -ate. All these recommendations are illustrated in the following examples.
----------------
Examples
1. Dichloro(diphenylphosphine)(thiourea)platinum(ii)
2. K4[Fe(CN)6]
3. [Co(NH3)6]C13
4. [CoCl(NH3)5}Cl2
5. [CoC1(N02)(NH3)4]Cl
6. [PtCl(NH2CH3)(NH3)2]Cl
7. [CuC12{OC(NH2)2}2]
8. K2[PdC14]
9. K[OsCl5N]
10. Na[PtBrC1(N02)(NH3)]
11. [Fe(CNCH3)6]Br2
12. [Ru(HSO3)2(NH3)4]
13. [Co(H20)2(NH3)4]C13
14. [PtC12(C5H5N)(NH3)]
15. Ba[BrF4]2
16. K[CrF4O]
17. [Ni(H20)2(NH3)4]S04
potassium hexacyanoferrate(ii)
potassium hexacyanoferrate(4—)
tetrapotassium hexacyanoferrate
hexaamminecobalt(iii) chloride
pentaamminechlorocobalt(2+) chloride
tetraamminechloronitrito-N-cobalt(iii) chloride
diamminechloro(methylamine)platinum(ii)
chloride
dichlorobis(urea)copper(ii)
potassium tetrachloropalladate(ii)
potassium pentachloronitridoosmate(2—)
sodium amminebromochloronitrito-
N-platinate( 1—)
hexakis(methyl isocyanide)iron(ii) bromide
tetraamminebis(hydrogensulfito)ruthenium(ii)
tetraamminediaquacobalt(iii) chloride
amminedichloro(pyridine)platinum(ii)
barium tetrafluorobromate(iii)
potassium tetrafiuorooxochromate(v)
tetraamminediaquanickel(ii) sulfate
***Table to be reformatted
-----------------------------
Designation of donor atom. In some cases, it may not be evident which atom in a
ligand is the donor. This is exemplified by the nitrito ligand in Examples 5 and 10,
p. 59. This can conceivably bind through an 0 or N atom. In simple cases, the donor
atom can be indicated by italicised element symbols placed after the specific ligand
name and separated from it by a hyphen, as demonstrated in those particular
examples. More complex examples will be dealt with below. With polydentate
ligands, this device may still be serviceable. Thus, dithiooxalate ion may be attached through S or 0, and formulations such as dithiooxalato-S,S' and dithiooxalato-0,0'should suffice. It could be necessary to use superscripts to the donor atom symbols if these need to be distinguished because there is more than one atom of the same kind to choose from.
Complicated examples are more easily dealt with using the kappa convention,
and this is particularly useful where a donor atom is part of a group that does not
carry a locant according to organic rules. The two oxygen atoms in a carboxylato
group demonstrate this. The designator i is a locant placed after that portion of the
ligand name that denotes the particular function in which the ligating atom is found.
The ligating atoms are represented by superscript numerals, letters or primes affixed
to the donor element symbols, which follow i without a space. A right superscript to
i denotes the number of identically bound ligating atoms.
Inclusion of structural information. The names described so far detail ligands and
central atoms, but give no information on stereochemistry. The coordination number
and shape of the coordination polyhedron may be denoted, if desired, by a
polyhedral symbol. These are listed in Table 4.4. Such a symbol is used as an affix in
parentheses, and immediately precedes the name, separated from it by a hyphen.
This device is not often used.
Geometrical descriptors, such as cis, trans, mer (from meridional) and fac (from
facial), have found wide usage in coordination nomenclature. The meaning is
unequivocal only in simple cases, particularly square planar for the first two and
octahedral for the others
More complex devices have been developed that are capable of dealing with all
cases. The reader is referred to the Nomenclature of Inorganic Chemistry, Chapter
10.
------------------------
Ionisation Isomers
Ionisation isomers:
Molecular structural formula is same. But different isomers give different ions in solution.
one isomer [PtBr(NH3)3]NO2 -> gives NO2- anions in solution
another isomer [Pt(NH3)3(NO2)]Br -> gives Br- anions in solution
Notice that both anions are necessary to balance the charge of the complex, and that they differ in that one ion is directly attached to the central metal but the other is not.
Geometric Isomers or Cis-Trans isomers
Geometric isomers are two or more coordination compounds which contain the same number and types of atoms, and bonds (i.e., the connectivity between atoms is the same), but which have different spatial arrangements of the atoms.
Not all coordination compounds have geometric isomers.
For example, in the square planar molecule, Pt(NH3)2Cl2, the two ammonia ligands (or the two chloride ligands) can be adjacent to one another or opposite one another.
Note that these two structures contain the same number and kinds of atoms and bonds but are non-superimposable. The isomer in which like ligands are adjacent to one another is called the cis isomer. The isomer in which like ligands are opposite one another is called the trans isomer.
For the common structures which contain two or more different ligands, geometric isomers are possible only with square planar and octahedral structures (i.e., geometric isomers cannot exist for linear and tetrahedral structures).
cis-[Co(NH3)4Cl2]+
Note that the two chloride ligands are adjacent to one another in this octahedral complex ion. In aqueous solution, this complex ion has a violet color.
trans-[Co(NH3)4Cl2]+
Note that the two chloride ligands are opposite one another in this complex ion. In aqueous solution, this complex ion has a green color.
------------------
Material about all isomers - In syllabus only ionic and cis-trans are specially mentioned
http://www.chem.purdue.edu/gchelp/cchem/whatis2.html
Coordination Isomers
Coordination isomers are two or more coordination compounds in which the composition within the coordination sphere (i.e., the metal atom plus the ligands that are bonded to it) is different (i.e., the connectivity between atoms is different).
Not all coordination compounds have coordination isomers.
Coordination isomers have different physical and chemical properties.
Example
[Cr(NH3)5(OSO3)]Br
Note that the sulfate group is bonded to the Cr atom (via an O atom) and is within the coordination sphere. Note also the octahedral structure. The bromide counterion is needed to maintain charge neutrality with the complex ion (i.e., [Cr(NH3)5(OSO3)]+) and is not shown in the structure.
[Cr(NH3)5Br]SO4
Note that the bromine atom is bonded to the Cr atom and is within the coordination sphere. Note also the octahedral structure. The sulfate counterion is not shown in the structure.
Linkage Isomers
Linkage isomers are two or more coordination compounds in which the donor atom of at least one of the ligands is different (i.e., the connectivity between atoms is different).
This type of isomerism can only exist when the compound contains a ligand that can bond to the metal atom in two (or more) different ways. Some ligands that can form linkage isomers are shown below.
Not all coordination compounds have linkage isomers.
Linkage isomers have different physical and chemical properties.
[Co(NH3)4(NO2)Cl]+
Note that the N atom of the nitrite group is bonded to the Co atom. The nitrite group is written as "NO2" in the molecular formula (rather than "ONO") with the N atom nearest to the Co symbol to indicate that the N atom (rather than an O atom) is the donor atom. Note also the octahedral structure.
[Co(NH3)4(ONO)Cl]+
Note that one of the O atoms of the nitrite group is bonded to the Co atom. The nitrite group is written as "ONO" in the molecular formula (rather than "NO2") with the O atom nearest to the Co symbol to indicate that the O atom is the donor atom. Note also the octahedral structure.
Geometric Isomers
Geometric isomers are two or more coordination compounds which contain the same number and types of atoms, and bonds (i.e., the connectivity between atoms is the same), but which have different spatial arrangements of the atoms.
Not all coordination compounds have geometric isomers.
For example, in the square planar molecule, Pt(NH3)2Cl2, the two ammonia ligands (or the two chloride ligands) can be adjacent to one another or opposite one another.
Note that these two structures contain the same number and kinds of atoms and bonds but are non-superimposable. The isomer in which like ligands are adjacent to one another is called the cis isomer. The isomer in which like ligands are opposite one another is called the trans isomer.
For the common structures which contain two or more different ligands, geometric isomers are possible only with square planar and octahedral structures (i.e., geometric isomers cannot exist for linear and tetrahedral structures).
cis-[Co(NH3)4Cl2]+
Note that the two chloride ligands are adjacent to one another in this octahedral complex ion. In aqueous solution, this complex ion has a violet color.
trans-[Co(NH3)4Cl2]+
Note that the two chloride ligands are opposite one another in this complex ion. In aqueous solution, this complex ion has a green color.
Optical Isomers
Optical isomers are two compounds which contain the same number and kinds of atoms, and bonds (i.e., the connectivity between atoms is the same), and different spatial arrangements of the atoms, but which have non-superimposable mirror images. Each non-superimposable mirror image structure is called an enantiomer. Molecules or ions that exist as optical isomers are called chiral.
Not all coordination compounds have optical isomers.
The Two Enantiomers of CHBrClF
Note that the molecule on the right is the reflection of the molecule on the left (through the mirror plane indicated by the black vertical line). These two structures are non-superimposable and are, therefore, different compounds.
Pure samples of enantiomers have identical physical properties (e.g., boiling point, density, freezing point). Chiral molecules and ions have different chemical properties only when they are in chiral environments.
Optical isomers get their name because the plane of plane-polarized light that is passed through a sample of a pure enantiomer is rotated. The plane is rotated in the opposite direction but with the same magnitude when plane-polarized light is passed through a pure sample containing the other enantiomer of a pair.
---------------------
web sites
http://www.chem.purdue.edu/gchelp/cchem/whatis2.html
Thursday, December 13, 2007
Tutorials Organic Chemistry
On Prentice Hall site, there are tutorials for organic Chemistry for Wades book
http://wps.prenhall.com/esm_organic_wade_5/5/1361/348631.cw/index.html
http://wps.prenhall.com/esm_organic_wade_5/5/1361/348631.cw/index.html
Sunday, December 9, 2007
TMH Ch. 34 Exercises in Organic Chemistry
Practical organic chemistry:
JEE Syllabus
Detection of elements (N, S, halogens);
Detection and identification of the following functional groups:
hydroxyl (alcoholic and phenolic),
carbonyl (aldehyde and ketone),
carboxyl, amino and nitro;
Chemical methods of separation of mono-functional organic compounds from binary mixtures.
Detection of Elements (N,S, halogens)
Lassaigne test is used for this.
First sodium extract of the given organic compound is prepared.
A pea sized dry piece of sodium metal (freshly cut) is taken in an ignition tube and the tube is heated to melt the sodium metal piece to shining globule. Then a pinch of the organic compound is introduced in the tube. Then the tube is further heated first gently and strongly till the lower end of the tube becomes red hot. The tube is then plunged and broken in about 20 ml. of distilled water taken in porcelain dish. The solution is boiled for about 5 minutes and filtered. The filtrate obtained is known as sodium extract.
Test for Nitrogen
To 2ml of sodium extract, a little ferrous sulphate is added. The solution is boiled and a little dilute sulphuric acid is added. Appearance of prussian blue or green colouration indicates the presence of nitrogen in the compound.
JEE Syllabus
Detection of elements (N, S, halogens);
Detection and identification of the following functional groups:
hydroxyl (alcoholic and phenolic),
carbonyl (aldehyde and ketone),
carboxyl, amino and nitro;
Chemical methods of separation of mono-functional organic compounds from binary mixtures.
Detection of Elements (N,S, halogens)
Lassaigne test is used for this.
First sodium extract of the given organic compound is prepared.
A pea sized dry piece of sodium metal (freshly cut) is taken in an ignition tube and the tube is heated to melt the sodium metal piece to shining globule. Then a pinch of the organic compound is introduced in the tube. Then the tube is further heated first gently and strongly till the lower end of the tube becomes red hot. The tube is then plunged and broken in about 20 ml. of distilled water taken in porcelain dish. The solution is boiled for about 5 minutes and filtered. The filtrate obtained is known as sodium extract.
Test for Nitrogen
To 2ml of sodium extract, a little ferrous sulphate is added. The solution is boiled and a little dilute sulphuric acid is added. Appearance of prussian blue or green colouration indicates the presence of nitrogen in the compound.
Chapter 13A. Halogens
The Halogens are non-metals and form the 7th Group in the Periodic Table.
'Halogens' means 'salt formers' and the most common compound is sodium chloride which is found from natural evaporation as huge deposits of 'rock salt' or the even more abundant 'sea salt' in the seas and oceans.
Physical properties
• typical non-metals with relatively low melting points and boiling points.
• They are all poor conductors of heat and electricity - typical of non-metals.
• When solid they are brittle and crumbly e.g. iodine.
•
Colour
F Fluorine--- pale yellow gas
Cl Chlorine--- green gas
Br Bromine--- dark red liquid, brown vapour
I Iodine---- dark crumbly solid, purple vapour
At Astatine--- black solid, dark vapour
•
•
important trends down the Group with increasing atomic number
• The melting points and boiling increase steadily down the group (so the change in state at room temperature from gas ==> liquid ==> solid), this is because the weak electrical intermolecular attractive forces increase with increasing size of atom or molecule.
• They are all coloured non-metallic elements and the colour gets darker down the group.
• The size of the atom gets bigger as more inner electron shells are filled going down from one period to another.
Chemical Properties
• The atoms all have 7 outer electrons,
o they form singly charged negative ions e.g. chloride Cl- because they are one electron short of a noble gas electron structure. They gain one negative electron (reduction) to be stable and this gives a surplus electric charge of -1. These ions are called the halide ions, the bromide Br- and iodide I- ions.
o they form ionic compounds with metals e.g. sodium chloride Na+Cl-. (ionic bonding revision page)
o they form covalent compounds with non-metals and with themselves.
o The bonding in the molecule involves single covalent bonds e.g. hydrogen chloride HCl or H-Cl.
• The elements all exist as X2 or X-X, diatomic molecules where X represents the halogen atom.
• A more reactive halogen can displace a less reactive halogen from its salts .
• The reactivity decreases down the group .
• they are all TOXIC elements .
• Astatine is very radioactive, so difficult to study
Reaction with hydrogen H-2
• Halogens readily combine with hydrogen to form the hydrogen halides which are colourless gaseous covalent molecules. e.g. hydrogen + chlorine ==> hydrogen chloride
• H-2(g) + Cl-2(g) ==> 2HCl(g)
• The hydrogen halides dissolve in water to form very strong acids with solutions of pH1 e.g. hydrogen chloride forms hydrochloric acid in water HCl(aq) or H+Cl-(aq) because they are fully ionised in aqueous solution even though the original hydrogen halides were covalent. An acid is a substance that forms H+ ions in water.
• Bromine forms hydrogen bromide gas HBr(g), which dissolved in water forms hydrobromic acid HBr(aq). Iodine forms hydrogen iodide gas HI(g), which dissolved in water forms hydriodic acid HI(aq).
Reaction with Group 1 Alkali Metals Li, Na, K etc.
• Alkali metals burn very exothermically and vigorously when heated in chlorine to form colourless crystalline ionic salts e.g. NaCl or Na+Cl-.
• e.g. sodium + chlorine ==> sodium chloride
• 2Na(s) + Cl2(g) ==> 2NaCl(s)
• The sodium chloride is soluble in water to give a neutral solution pH 7, universal indicator is green. The salt is a typical ionic compound i.e. a brittle solid with a high melting point. Similarly potassium and bromine form potassium bromide KBr, or lithium and iodine form lithium iodide LiI.
•
Reaction with other metals
• If aluminium or iron is heated strongly in a stream of chlorine (or plunge the hot metal into a gas jar of chlorine carefully in a fume cupboard) the solid chloride is formed.
• aluminium + chlorine ==> aluminium chloride(white):
o 2Al(s) + 3Cl2(g) ==> 2AlCl3(s)
• iron + chlorine ==> iron(III) chloride(brown):
o 2Fe(s) + 3Cl2(g) ==> 2FeCl3(s)
• If the iron is repeated with bromine the reaction is less vigorous, with iodine there is little reaction, these also illustrate the halogen reactivity series.
'Halogens' means 'salt formers' and the most common compound is sodium chloride which is found from natural evaporation as huge deposits of 'rock salt' or the even more abundant 'sea salt' in the seas and oceans.
Physical properties
• typical non-metals with relatively low melting points and boiling points.
• They are all poor conductors of heat and electricity - typical of non-metals.
• When solid they are brittle and crumbly e.g. iodine.
•
Colour
F Fluorine--- pale yellow gas
Cl Chlorine--- green gas
Br Bromine--- dark red liquid, brown vapour
I Iodine---- dark crumbly solid, purple vapour
At Astatine--- black solid, dark vapour
•
•
important trends down the Group with increasing atomic number
• The melting points and boiling increase steadily down the group (so the change in state at room temperature from gas ==> liquid ==> solid), this is because the weak electrical intermolecular attractive forces increase with increasing size of atom or molecule.
• They are all coloured non-metallic elements and the colour gets darker down the group.
• The size of the atom gets bigger as more inner electron shells are filled going down from one period to another.
Chemical Properties
• The atoms all have 7 outer electrons,
o they form singly charged negative ions e.g. chloride Cl- because they are one electron short of a noble gas electron structure. They gain one negative electron (reduction) to be stable and this gives a surplus electric charge of -1. These ions are called the halide ions, the bromide Br- and iodide I- ions.
o they form ionic compounds with metals e.g. sodium chloride Na+Cl-. (ionic bonding revision page)
o they form covalent compounds with non-metals and with themselves.
o The bonding in the molecule involves single covalent bonds e.g. hydrogen chloride HCl or H-Cl.
• The elements all exist as X2 or X-X, diatomic molecules where X represents the halogen atom.
• A more reactive halogen can displace a less reactive halogen from its salts .
• The reactivity decreases down the group .
• they are all TOXIC elements .
• Astatine is very radioactive, so difficult to study
Reaction with hydrogen H-2
• Halogens readily combine with hydrogen to form the hydrogen halides which are colourless gaseous covalent molecules. e.g. hydrogen + chlorine ==> hydrogen chloride
• H-2(g) + Cl-2(g) ==> 2HCl(g)
• The hydrogen halides dissolve in water to form very strong acids with solutions of pH1 e.g. hydrogen chloride forms hydrochloric acid in water HCl(aq) or H+Cl-(aq) because they are fully ionised in aqueous solution even though the original hydrogen halides were covalent. An acid is a substance that forms H+ ions in water.
• Bromine forms hydrogen bromide gas HBr(g), which dissolved in water forms hydrobromic acid HBr(aq). Iodine forms hydrogen iodide gas HI(g), which dissolved in water forms hydriodic acid HI(aq).
Reaction with Group 1 Alkali Metals Li, Na, K etc.
• Alkali metals burn very exothermically and vigorously when heated in chlorine to form colourless crystalline ionic salts e.g. NaCl or Na+Cl-.
• e.g. sodium + chlorine ==> sodium chloride
• 2Na(s) + Cl2(g) ==> 2NaCl(s)
• The sodium chloride is soluble in water to give a neutral solution pH 7, universal indicator is green. The salt is a typical ionic compound i.e. a brittle solid with a high melting point. Similarly potassium and bromine form potassium bromide KBr, or lithium and iodine form lithium iodide LiI.
•
Reaction with other metals
• If aluminium or iron is heated strongly in a stream of chlorine (or plunge the hot metal into a gas jar of chlorine carefully in a fume cupboard) the solid chloride is formed.
• aluminium + chlorine ==> aluminium chloride(white):
o 2Al(s) + 3Cl2(g) ==> 2AlCl3(s)
• iron + chlorine ==> iron(III) chloride(brown):
o 2Fe(s) + 3Cl2(g) ==> 2FeCl3(s)
• If the iron is repeated with bromine the reaction is less vigorous, with iodine there is little reaction, these also illustrate the halogen reactivity series.
One more orkut JEE community for chemistry
Came across the community today and joined the community.
CHEMISTRY- FOR ALLHome > Communities > CHEMISTRY- FOR ALL
description:
This community is created to help the students who are preparing for different competitive exams viz ENGG/MEDICAL.Chemistry is always supposed to be more scoring in jee but there is great misconception that-IT IS A SUBJECT TO RATTOFY (CRAMMING).
Its more of understanding LETS ACCEPT IT AND MOVE TOWARDS REACHING GREATER HEIGHTS IN TERMS OF CONCEPTS AND TO EASE THE PREPARATION OF JEE.
This forum also welcome those people who have keen interest in chemistry and wish to get guidance.
Your queries and doubts of any level is most WELCOME and will be taken care of on priority basis.
HAPPY LEARNING.
http://www.orkut.com/Community.aspx?cmm=35456617
SUGGESTED READINGS FOR IIT-JEE:-
CHEMISTRY:NCERT-Xl AND Xll
GENERAL CHEMISTRY- EBBING
ORGANIC CHEMISTRY:-SOLOMON/CAREY/BRUICE ANY ONE
QUALITATIVE ANALYSIS:- VOGEL
DON'T CONFUSE WITH MANY MORE........
owner:
SANJAY K. THAKUR
NEW DELHI, DELHI, 110075, India
created:
July 8, 2007 4:28 PM
Join the community and benefit from prof Sanjay's initiative
CHEMISTRY- FOR ALLHome > Communities > CHEMISTRY- FOR ALL
description:
This community is created to help the students who are preparing for different competitive exams viz ENGG/MEDICAL.Chemistry is always supposed to be more scoring in jee but there is great misconception that-IT IS A SUBJECT TO RATTOFY (CRAMMING).
Its more of understanding LETS ACCEPT IT AND MOVE TOWARDS REACHING GREATER HEIGHTS IN TERMS OF CONCEPTS AND TO EASE THE PREPARATION OF JEE.
This forum also welcome those people who have keen interest in chemistry and wish to get guidance.
Your queries and doubts of any level is most WELCOME and will be taken care of on priority basis.
HAPPY LEARNING.
http://www.orkut.com/Community.aspx?cmm=35456617
SUGGESTED READINGS FOR IIT-JEE:-
CHEMISTRY:NCERT-Xl AND Xll
GENERAL CHEMISTRY- EBBING
ORGANIC CHEMISTRY:-SOLOMON/CAREY/BRUICE ANY ONE
QUALITATIVE ANALYSIS:- VOGEL
DON'T CONFUSE WITH MANY MORE........
owner:
SANJAY K. THAKUR
NEW DELHI, DELHI, 110075, India
created:
July 8, 2007 4:28 PM
Join the community and benefit from prof Sanjay's initiative
JEE 2008 Physical Chemistry Syllabus
Physical chemistry
General topics: Concept of atoms and molecules; Dalton’s atomic theory; Mole concept; Chemical formulae; Balanced chemical equations; Calculations (based on mole concept) involving common oxidation-reduction, neutralisation, and displacement reactions; Concentration in terms of mole fraction, molarity, molality and normality.
Gaseous and liquid states: Absolute scale of temperature, ideal gas equation; Deviation from ideality, van der Waals equation; Kinetic theory of gases, average, root mean square and most probable velocities and their relation with temperature; Law of partial pressures; Vapour pressure; Diffusion of gases.
Atomic structure and chemical bonding: Bohr model, spectrum of hydrogen atom, quantum numbers; Wave-particle duality, de Broglie hypothesis; Uncertainty principle; Qualitative quantum mechanical picture of hydrogen atom, shapes of s, p and d orbitals; Electronic configurations of elements (up to atomic number 36); Aufbau principle; Pauli’s exclusion principle and Hund’s rule; Orbital overlap and covalent bond; Hybridisation involving s, p and d orbitals only; Orbital energy diagrams for homonuclear diatomic species; Hydrogen bond; Polarity in molecules, dipole moment (qualitative aspects only); VSEPR model and shapes of molecules (linear, angular, triangular, square planar, pyramidal, square pyramidal, trigonal bipyramidal, tetrahedral and octahedral).
Energetics: First law of thermodynamics; Internal energy, work and heat, pressure-volume work; Enthalpy, Hess’s law; Heat of reaction, fusion and vapourization; Second law of thermodynamics; Entropy; Free energy; Criterion of spontaneity.
Chemical equilibrium: Law of mass action; Equilibrium constant, Le Chatelier's principle (effect of concentration, temperature and pressure); Significance of DG and DGo in chemical equilibrium; Solubility product, common ion effect, pH and buffer solutions; Acids and bases (Bronsted and Lewis concepts); Hydrolysis of salts.
Electrochemistry: Electrochemical cells and cell reactions; Standard electrode potentials; Nernst equation and its relation to DG; Electrochemical series, emf of galvanic cells; Faraday's laws of electrolysis; Electrolytic conductance, specific, equivalent and molar conductivity, Kohlrausch's law; Concentration cells.
Chemical kinetics: Rates of chemical reactions; Order of reactions; Rate constant; First order reactions; Temperature dependence of rate constant (Arrhenius equation).
Solid state: Classification of solids, crystalline state, seven crystal systems (cell parameters a, b, c, alpha, beta, gamma), close packed structure of solids (cubic), packing in fcc, bcc and hcp lattices; Nearest neighbours, ionic radii, simple ionic compounds, point defects.
Solutions: Raoult's law; Molecular weight determination from lowering of vapour pressure, elevation of boiling point and depression of freezing point.
Surface chemistry: Elementary concepts of adsorption (excluding adsorption isotherms); Colloids: types, methods of preparation and general properties; Elementary ideas of emulsions, surfactants and micelles (only definitions and examples).
Nuclear chemistry: Radioactivity: isotopes and isobars; Properties of alpha, beta and gamma rays; Kinetics of radioactive decay (decay series excluded), carbon dating; Stability of nuclei with respect to proton-neutron ratio; Brief discussion on fission and fusion reactions.
General topics: Concept of atoms and molecules; Dalton’s atomic theory; Mole concept; Chemical formulae; Balanced chemical equations; Calculations (based on mole concept) involving common oxidation-reduction, neutralisation, and displacement reactions; Concentration in terms of mole fraction, molarity, molality and normality.
Gaseous and liquid states: Absolute scale of temperature, ideal gas equation; Deviation from ideality, van der Waals equation; Kinetic theory of gases, average, root mean square and most probable velocities and their relation with temperature; Law of partial pressures; Vapour pressure; Diffusion of gases.
Atomic structure and chemical bonding: Bohr model, spectrum of hydrogen atom, quantum numbers; Wave-particle duality, de Broglie hypothesis; Uncertainty principle; Qualitative quantum mechanical picture of hydrogen atom, shapes of s, p and d orbitals; Electronic configurations of elements (up to atomic number 36); Aufbau principle; Pauli’s exclusion principle and Hund’s rule; Orbital overlap and covalent bond; Hybridisation involving s, p and d orbitals only; Orbital energy diagrams for homonuclear diatomic species; Hydrogen bond; Polarity in molecules, dipole moment (qualitative aspects only); VSEPR model and shapes of molecules (linear, angular, triangular, square planar, pyramidal, square pyramidal, trigonal bipyramidal, tetrahedral and octahedral).
Energetics: First law of thermodynamics; Internal energy, work and heat, pressure-volume work; Enthalpy, Hess’s law; Heat of reaction, fusion and vapourization; Second law of thermodynamics; Entropy; Free energy; Criterion of spontaneity.
Chemical equilibrium: Law of mass action; Equilibrium constant, Le Chatelier's principle (effect of concentration, temperature and pressure); Significance of DG and DGo in chemical equilibrium; Solubility product, common ion effect, pH and buffer solutions; Acids and bases (Bronsted and Lewis concepts); Hydrolysis of salts.
Electrochemistry: Electrochemical cells and cell reactions; Standard electrode potentials; Nernst equation and its relation to DG; Electrochemical series, emf of galvanic cells; Faraday's laws of electrolysis; Electrolytic conductance, specific, equivalent and molar conductivity, Kohlrausch's law; Concentration cells.
Chemical kinetics: Rates of chemical reactions; Order of reactions; Rate constant; First order reactions; Temperature dependence of rate constant (Arrhenius equation).
Solid state: Classification of solids, crystalline state, seven crystal systems (cell parameters a, b, c, alpha, beta, gamma), close packed structure of solids (cubic), packing in fcc, bcc and hcp lattices; Nearest neighbours, ionic radii, simple ionic compounds, point defects.
Solutions: Raoult's law; Molecular weight determination from lowering of vapour pressure, elevation of boiling point and depression of freezing point.
Surface chemistry: Elementary concepts of adsorption (excluding adsorption isotherms); Colloids: types, methods of preparation and general properties; Elementary ideas of emulsions, surfactants and micelles (only definitions and examples).
Nuclear chemistry: Radioactivity: isotopes and isobars; Properties of alpha, beta and gamma rays; Kinetics of radioactive decay (decay series excluded), carbon dating; Stability of nuclei with respect to proton-neutron ratio; Brief discussion on fission and fusion reactions.
Saturday, December 8, 2007
JEE 2008 Syllabus Inorganic Chemistry
I am studying class XI and XII CBSE books of Jauhar along with TMH JEE Book of 2007. I am recording the sections and pages numbers where the specified topics are dealt with in these books for ready reference.
Such a reference seems to be essential for inorganic chemistry because the syllabus is not fitting into a simple chapter schemes of texts. For physical chemistry and organic chemistry each topic is distinct chapter in the books.
Inorganic Chemistry
Isolation/preparation and properties of the following non-metals:
Boron,XI Bk, Section 13.1, p.800-802
silicon, TMH JEE 2007 (TJ) page 13.4
nitrogen, XI Bk, Section 13.5,p.827-30
phosphorus, XII Bk, Sec 8.15
oxygen, XI, sec 13.7
sulphur and XII, 8.23,
halogens; XII, 8.26
Properties of allotropes of:
carbon (only diamond and graphite), XI, 812-814, XII, s 8.6.d.2, p 352
phosphorus and XII, 8.16, 369-70
sulphur.XII, 8.24, 389
Preparation and properties of the following compounds:
of sodium,
Oxides, XI, sec 12.6, 755-57, 763,764
peroxides, XI, 764
hydroxides, XI, 765
carbonates, XI, 767
bicarbonates,
chlorides and
sulphates
potassium,
Oxides,
peroxides,
hydroxides,
carbonates,
bicarbonates,
chlorides and
sulphates of
Compound of Sodium and Potassium are covered in unit 14 in TJ
magnesium
Oxides, XI, p774-76
peroxides,
hydroxides, XI 776
carbonates,
bicarbonates,
chlorides and XI 783
sulphates XI, 785
calcium;
Oxides, XI, 774, 786
peroxides,
hydroxides, XI, 787
carbonates,
bicarbonates,
chlorides and
sulphates XI, 779, 788
Compound of Magnesium and Calcium are covered in unit 14 of TJ
Boron:
diborane, XI, 13.2 806-10
boric acid and 804-5
borax; 802-4
Aluminium:
alumina,
aluminium chloride and
alums; - all in unit 14 TJ
Carbon:
oxides and XI , 13.4 818-21
oxyacid (carbonic acid);
Silicon:
silicones, XII, 8.10
silicates and 8.9
silicon carbide;
Nitrogen: XI 13.6
oxides, 833-36
oxyacids and 836-41
ammonia; 830-32
Phosphorus:
oxides, XII, p 367
oxyacids (phosphorus acid, phosphoric acid) and p 368, TJ p.15.17
phosphine; XII Sec 8.17, p370
Oxygen:
ozone and Xi sec 13.9, 848-853
hydrogen peroxide; XI sec 11.9, 729-738
XII, p 385
Sulphur:
hydrogen sulphide, XII p 364
oxides, XII sec 8.22.1 p 375
sulphurous acid, sec 8.22.2.1
sulphuric acid and 8.22.2.2, 8.25, p389
sodium thiosulphate; TJ p.15.21
Halogens:
hydrohalic acids, XII, 8.28.1 p 394
oxides and 394
oxyacids of chlorine, 95
bleaching powder; 8.29, 396
Xenon fluorides.XII, sec 8.33.A,
Transition elements (3d series):
Definition, XII, 9.1, 422
general characteristics, sec 9.2, 9.3
oxidation states and their stabilities, sec 9.3.6
colour (excluding the details of electronic transitions) and sec 9.3.7, Table 9.8 p430
calculation of spin-only magnetic moment; sec 9.3.9, p430
Preparation and properties of the following compounds:
Oxides and chlorides of tin and lead; XII, sec 8.11
Oxides, chlorides and sulphates of Fe2+, Cu2+ and Zn2+; XII sec 9.4, 9.11
Potassium permanganate, XII sec 9.11.6
potassium dichromate, 9.11.5
silver oxide,
silver nitrate, TJ p.14.9
silver thiosulphate.
Coordination compounds: XII, section 10
nomenclature of mononuclear coordination compounds, XII 10.1,2,3
cis-trans and ionisation isomerisms, 10.4
hybridization and geometries of mononuclear coordination compounds (linear, tetrahedral, square planar and octahedral).10.7
Ores and minerals: For this topic and topic of extraction metallurgy - refer Chapter X of Class XI book of Jauhar.
Commonly occurring ores and minerals of
iron,
copper,
tin,
lead,
magnesium,
aluminium,
zinc and
silver.
XI, sec 10.5
Extractive metallurgy: Chemical principles and reactions only (industrial details excluded); XI, 10.10
Carbon reduction method (iron and tin); 10.10.3.B.I.(i)
Self reduction method (copper and lead); 10.10.3.B.I.(vi)
Electrolytic reduction method (magnesium and aluminium); 10.10.3.B.III
Cyanide process (silver and gold). TJ p.17.6 (*gold is not there)
Principles of qualitative analysis: Ch. 18 of TMH JEE Book
Groups I to V (only Ag+, Hg2+, Cu2+, Pb2+, Bi3+, Fe3+, Cr3+, Al3+, Ca2+, Ba2+, Zn2+, Mn2+ and Mg2+);
Group I - TJp.18.3
Group II - 18.4
Group III- 18.5
Group IV - 18.5
Group V -- 18.6
Nitrate, 18.2
halides (excluding fluoride), 18.2
sulphate and 18.2
sulphide. 18.1
Except two or three topics all the topics are covered in the books. The specific items not covered can be accessed from one of the sites on the internet. I shall do it and post in the concerned chapter.
Such a reference seems to be essential for inorganic chemistry because the syllabus is not fitting into a simple chapter schemes of texts. For physical chemistry and organic chemistry each topic is distinct chapter in the books.
Inorganic Chemistry
Isolation/preparation and properties of the following non-metals:
Boron,XI Bk, Section 13.1, p.800-802
silicon, TMH JEE 2007 (TJ) page 13.4
nitrogen, XI Bk, Section 13.5,p.827-30
phosphorus, XII Bk, Sec 8.15
oxygen, XI, sec 13.7
sulphur and XII, 8.23,
halogens; XII, 8.26
Properties of allotropes of:
carbon (only diamond and graphite), XI, 812-814, XII, s 8.6.d.2, p 352
phosphorus and XII, 8.16, 369-70
sulphur.XII, 8.24, 389
Preparation and properties of the following compounds:
of sodium,
Oxides, XI, sec 12.6, 755-57, 763,764
peroxides, XI, 764
hydroxides, XI, 765
carbonates, XI, 767
bicarbonates,
chlorides and
sulphates
potassium,
Oxides,
peroxides,
hydroxides,
carbonates,
bicarbonates,
chlorides and
sulphates of
Compound of Sodium and Potassium are covered in unit 14 in TJ
magnesium
Oxides, XI, p774-76
peroxides,
hydroxides, XI 776
carbonates,
bicarbonates,
chlorides and XI 783
sulphates XI, 785
calcium;
Oxides, XI, 774, 786
peroxides,
hydroxides, XI, 787
carbonates,
bicarbonates,
chlorides and
sulphates XI, 779, 788
Compound of Magnesium and Calcium are covered in unit 14 of TJ
Boron:
diborane, XI, 13.2 806-10
boric acid and 804-5
borax; 802-4
Aluminium:
alumina,
aluminium chloride and
alums; - all in unit 14 TJ
Carbon:
oxides and XI , 13.4 818-21
oxyacid (carbonic acid);
Silicon:
silicones, XII, 8.10
silicates and 8.9
silicon carbide;
Nitrogen: XI 13.6
oxides, 833-36
oxyacids and 836-41
ammonia; 830-32
Phosphorus:
oxides, XII, p 367
oxyacids (phosphorus acid, phosphoric acid) and p 368, TJ p.15.17
phosphine; XII Sec 8.17, p370
Oxygen:
ozone and Xi sec 13.9, 848-853
hydrogen peroxide; XI sec 11.9, 729-738
XII, p 385
Sulphur:
hydrogen sulphide, XII p 364
oxides, XII sec 8.22.1 p 375
sulphurous acid, sec 8.22.2.1
sulphuric acid and 8.22.2.2, 8.25, p389
sodium thiosulphate; TJ p.15.21
Halogens:
hydrohalic acids, XII, 8.28.1 p 394
oxides and 394
oxyacids of chlorine, 95
bleaching powder; 8.29, 396
Xenon fluorides.XII, sec 8.33.A,
Transition elements (3d series):
Definition, XII, 9.1, 422
general characteristics, sec 9.2, 9.3
oxidation states and their stabilities, sec 9.3.6
colour (excluding the details of electronic transitions) and sec 9.3.7, Table 9.8 p430
calculation of spin-only magnetic moment; sec 9.3.9, p430
Preparation and properties of the following compounds:
Oxides and chlorides of tin and lead; XII, sec 8.11
Oxides, chlorides and sulphates of Fe2+, Cu2+ and Zn2+; XII sec 9.4, 9.11
Potassium permanganate, XII sec 9.11.6
potassium dichromate, 9.11.5
silver oxide,
silver nitrate, TJ p.14.9
silver thiosulphate.
Coordination compounds: XII, section 10
nomenclature of mononuclear coordination compounds, XII 10.1,2,3
cis-trans and ionisation isomerisms, 10.4
hybridization and geometries of mononuclear coordination compounds (linear, tetrahedral, square planar and octahedral).10.7
Ores and minerals: For this topic and topic of extraction metallurgy - refer Chapter X of Class XI book of Jauhar.
Commonly occurring ores and minerals of
iron,
copper,
tin,
lead,
magnesium,
aluminium,
zinc and
silver.
XI, sec 10.5
Extractive metallurgy: Chemical principles and reactions only (industrial details excluded); XI, 10.10
Carbon reduction method (iron and tin); 10.10.3.B.I.(i)
Self reduction method (copper and lead); 10.10.3.B.I.(vi)
Electrolytic reduction method (magnesium and aluminium); 10.10.3.B.III
Cyanide process (silver and gold). TJ p.17.6 (*gold is not there)
Principles of qualitative analysis: Ch. 18 of TMH JEE Book
Groups I to V (only Ag+, Hg2+, Cu2+, Pb2+, Bi3+, Fe3+, Cr3+, Al3+, Ca2+, Ba2+, Zn2+, Mn2+ and Mg2+);
Group I - TJp.18.3
Group II - 18.4
Group III- 18.5
Group IV - 18.5
Group V -- 18.6
Nitrate, 18.2
halides (excluding fluoride), 18.2
sulphate and 18.2
sulphide. 18.1
Except two or three topics all the topics are covered in the books. The specific items not covered can be accessed from one of the sites on the internet. I shall do it and post in the concerned chapter.
JEE 2008 Organic Chemistry Syllabus
Organic Chemistry
Concepts:
Hybridisation of carbon;
Sigma and pi-bonds;
Shapes of simple organic molecules;
Structural and geometrical isomerism;
Optical isomerism of compounds containing up to two asymmetric centres, (R,S and E,Z nomenclature excluded);
IUPAC nomenclature of simple organic compounds (only hydrocarbons, mono-functional and bi-functional compounds);
Conformations of ethane and butane (Newman projections);
Resonance and hyperconjugation;
Keto-enol tautomerism;
Determination of empirical and molecular formulae of simple compounds (only combustion method);
Hydrogen bonds: definition and their effects on physical properties of alcohols and carboxylic acids;
Inductive and resonance effects on acidity and basicity of organic acids and bases; Polarity and inductive effects in alkyl halides;
Reactive intermediates produced during homolytic and heterolytic bond cleavage; Formation, structure and stability of carbocations, carbanions and free radicals.
Preparation, properties and reactions of alkanes: Homologous series, physical properties of alkanes (melting points, boiling points and density); Combustion and halogenation of alkanes; Preparation of alkanes by Wurtz reaction and decarboxylation reactions.
Preparation, properties and reactions of alkenes and alkynes: Physical properties of alkenes and alkynes (boiling points, density and dipole moments); Acidity of alkynes; Acid catalysed hydration of alkenes and alkynes (excluding the stereochemistry of addition and elimination); Reactions of alkenes with KMnO4 and ozone; Reduction of alkenes and alkynes; Preparation of alkenes and alkynes by elimination reactions; Electrophilic addition reactions of alkenes with X2, HX, HOX and H2O (X=halogen); Addition reactions of alkynes; Metal acetylides.
Reactions of benzene: Structure and aromaticity; Electrophilic substitution reactions: halogenation, nitration, sulphonation, Friedel-Crafts alkylation and acylation; Effect of o-, m- and p-directing groups in monosubstituted benzenes.
Phenols: Acidity, electrophilic substitution reactions (halogenation, nitration and sulphonation); Reimer-Tieman reaction, Kolbe reaction.
Characteristic reactions of the following (including those mentioned above):
Alkyl halides: rearrangement reactions of alkyl carbocation, Grignard reactions, nucleophilic substitution reactions;
Alcohols: esterification, dehydration and oxidation, reaction with sodium, phosphorus halides, ZnCl2/concentrated HCl, conversion of alcohols into aldehydes and ketones;
Ethers:Preparation by Williamson's Synthesis;
Aldehydes and Ketones: oxidation, reduction, oxime and hydrazone formation; aldol condensation, Perkin reaction; Cannizzaro reaction; haloform reaction and nucleophilic addition reactions (Grignard addition);
Carboxylic acids: formation of esters, acid chlorides and amides, ester hydrolysis;
Amines: basicity of substituted anilines and aliphatic amines, preparation from nitro compounds, reaction with nitrous acid, azo coupling reaction of diazonium salts of aromatic amines, Sandmeyer and related reactions of diazonium salts; carbylamine reaction; Haloarenes: nucleophilic aromatic substitution in haloarenes and substituted haloarenes (excluding Benzyne mechanism and Cine substitution).
Carbohydrates: Classification; mono- and di-saccharides (glucose and sucrose); Oxidation, reduction, glycoside formation and hydrolysis of sucrose.
Amino acids and peptides: General structure (only primary structure for peptides) and physical properties.
Properties and uses of some important polymers: Natural rubber, cellulose, nylon, teflon and PVC.
Practical organic chemistry: Detection of elements (N, S, halogens); Detection and identification of the following functional groups: hydroxyl (alcoholic and phenolic), carbonyl (aldehyde and ketone), carboxyl, amino and nitro; Chemical methods of separation of mono-functional organic compounds from binary mixtures.
Concepts:
Hybridisation of carbon;
Sigma and pi-bonds;
Shapes of simple organic molecules;
Structural and geometrical isomerism;
Optical isomerism of compounds containing up to two asymmetric centres, (R,S and E,Z nomenclature excluded);
IUPAC nomenclature of simple organic compounds (only hydrocarbons, mono-functional and bi-functional compounds);
Conformations of ethane and butane (Newman projections);
Resonance and hyperconjugation;
Keto-enol tautomerism;
Determination of empirical and molecular formulae of simple compounds (only combustion method);
Hydrogen bonds: definition and their effects on physical properties of alcohols and carboxylic acids;
Inductive and resonance effects on acidity and basicity of organic acids and bases; Polarity and inductive effects in alkyl halides;
Reactive intermediates produced during homolytic and heterolytic bond cleavage; Formation, structure and stability of carbocations, carbanions and free radicals.
Preparation, properties and reactions of alkanes: Homologous series, physical properties of alkanes (melting points, boiling points and density); Combustion and halogenation of alkanes; Preparation of alkanes by Wurtz reaction and decarboxylation reactions.
Preparation, properties and reactions of alkenes and alkynes: Physical properties of alkenes and alkynes (boiling points, density and dipole moments); Acidity of alkynes; Acid catalysed hydration of alkenes and alkynes (excluding the stereochemistry of addition and elimination); Reactions of alkenes with KMnO4 and ozone; Reduction of alkenes and alkynes; Preparation of alkenes and alkynes by elimination reactions; Electrophilic addition reactions of alkenes with X2, HX, HOX and H2O (X=halogen); Addition reactions of alkynes; Metal acetylides.
Reactions of benzene: Structure and aromaticity; Electrophilic substitution reactions: halogenation, nitration, sulphonation, Friedel-Crafts alkylation and acylation; Effect of o-, m- and p-directing groups in monosubstituted benzenes.
Phenols: Acidity, electrophilic substitution reactions (halogenation, nitration and sulphonation); Reimer-Tieman reaction, Kolbe reaction.
Characteristic reactions of the following (including those mentioned above):
Alkyl halides: rearrangement reactions of alkyl carbocation, Grignard reactions, nucleophilic substitution reactions;
Alcohols: esterification, dehydration and oxidation, reaction with sodium, phosphorus halides, ZnCl2/concentrated HCl, conversion of alcohols into aldehydes and ketones;
Ethers:Preparation by Williamson's Synthesis;
Aldehydes and Ketones: oxidation, reduction, oxime and hydrazone formation; aldol condensation, Perkin reaction; Cannizzaro reaction; haloform reaction and nucleophilic addition reactions (Grignard addition);
Carboxylic acids: formation of esters, acid chlorides and amides, ester hydrolysis;
Amines: basicity of substituted anilines and aliphatic amines, preparation from nitro compounds, reaction with nitrous acid, azo coupling reaction of diazonium salts of aromatic amines, Sandmeyer and related reactions of diazonium salts; carbylamine reaction; Haloarenes: nucleophilic aromatic substitution in haloarenes and substituted haloarenes (excluding Benzyne mechanism and Cine substitution).
Carbohydrates: Classification; mono- and di-saccharides (glucose and sucrose); Oxidation, reduction, glycoside formation and hydrolysis of sucrose.
Amino acids and peptides: General structure (only primary structure for peptides) and physical properties.
Properties and uses of some important polymers: Natural rubber, cellulose, nylon, teflon and PVC.
Practical organic chemistry: Detection of elements (N, S, halogens); Detection and identification of the following functional groups: hydroxyl (alcoholic and phenolic), carbonyl (aldehyde and ketone), carboxyl, amino and nitro; Chemical methods of separation of mono-functional organic compounds from binary mixtures.
Sunday, December 2, 2007
Blog-Status
Today, material is added to the chapter on Chemical Kinetics.
http://iit-jee-chemistry.blogspot.com/2007/10/study-guide-ch10-chemical-kinetics.html
Material will be added to Chemical Equilibrium chapter and questions will be added in the practice sets blog.
http://iit-jee-chemistry.blogspot.com/2007/10/study-guide-ch10-chemical-kinetics.html
Material will be added to Chemical Equilibrium chapter and questions will be added in the practice sets blog.
Friday, October 26, 2007
Study Notes for IIT JEE
I am trying to develop study notes in the form of Study guides to a standard book.
Study Guide TMH Chemistry 34. Exercises in Organic Chemistry
-----------------
JEE question 2007 Paper II
Riemer-Tiemann reaction introduces an aldehyde group, on to the aromatic ring of phenol, ortho to the hydroxyl group. This reaction involves electrophilic aromatic substitution. This is a general method for the synthesis of substituted salicylaldehydes.
1. Which one of the following reagents is used in the above reaction?
(A) aq.NaOH+ CH-3Cl
(B) aq.NaOH+ CH-2Cl-2
(C) aq.NaOH+ CHCl-3
(D) aq.NaOH+ CCl-4
Answer C
Reimen-Tiemann reaction takes place is presence of aq. NaOH+ CHCl3
2. The electrophile in this
(A) :CHCl
(B) +CHCl-2
(C) :CCl-2
(D) :CCl-3
Answer: C
-------------------------------
JEE question 2007 Paper II
Riemer-Tiemann reaction introduces an aldehyde group, on to the aromatic ring of phenol, ortho to the hydroxyl group. This reaction involves electrophilic aromatic substitution. This is a general method for the synthesis of substituted salicylaldehydes.
1. Which one of the following reagents is used in the above reaction?
(A) aq.NaOH+ CH-3Cl
(B) aq.NaOH+ CH-2Cl-2
(C) aq.NaOH+ CHCl-3
(D) aq.NaOH+ CCl-4
Answer C
Reimen-Tiemann reaction takes place is presence of aq. NaOH+ CHCl3
2. The electrophile in this
(A) :CHCl
(B) +CHCl-2
(C) :CCl-2
(D) :CCl-3
Answer: C
-------------------------------
Thursday, October 25, 2007
IIT JEE 2008 Chemistry Syllabus
JEE 2008
Chemistry Syllabus
Physical chemistry
General topics: Concept of atoms and molecules; Dalton’s atomic theory; Mole concept; Chemical formulae; Balanced chemical equations; Calculations (based on mole concept) involving common oxidation-reduction, neutralisation, and displacement reactions; Concentration in terms of mole fraction, molarity, molality and normality.
Gaseous and liquid states: Absolute scale of temperature, ideal gas equation; Deviation from ideality, van der Waals equation; Kinetic theory of gases, average, root mean square and most probable velocities and their relation with temperature; Law of partial pressures; Vapour pressure; Diffusion of gases.
Atomic structure and chemical bonding: Bohr model, spectrum of hydrogen atom, quantum numbers; Wave-particle duality, de Broglie hypothesis; Uncertainty principle; Qualitative quantum mechanical picture of hydrogen atom, shapes of s, p and d orbitals; Electronic configurations of elements (up to atomic number 36); Aufbau principle; Pauli’s exclusion principle and Hund’s rule; Orbital overlap and covalent bond; Hybridisation involving s, p and d orbitals only; Orbital energy diagrams for homonuclear diatomic species; Hydrogen bond; Polarity in molecules, dipole moment (qualitative aspects only); VSEPR model and shapes of molecules (linear, angular, triangular, square planar, pyramidal, square pyramidal, trigonal bipyramidal, tetrahedral and octahedral).
Energetics: First law of thermodynamics; Internal energy, work and heat, pressure-volume work; Enthalpy, Hess’s law; Heat of reaction, fusion and vapourization; Second law of thermodynamics; Entropy; Free energy; Criterion of spontaneity.
Chemical equilibrium: Law of mass action; Equilibrium constant, Le Chatelier's principle (effect of concentration, temperature and pressure); Significance of DG and DGo in chemical equilibrium; Solubility product, common ion effect, pH and buffer solutions; Acids and bases (Bronsted and Lewis concepts); Hydrolysis of salts.
Electrochemistry: Electrochemical cells and cell reactions; Standard electrode potentials; Nernst equation and its relation to DG; Electrochemical series, emf of galvanic cells; Faraday's laws of electrolysis; Electrolytic conductance, specific, equivalent and molar conductivity, Kohlrausch's law; Concentration cells.
Chemical kinetics: Rates of chemical reactions; Order of reactions; Rate constant; First order reactions; Temperature dependence of rate constant (Arrhenius equation).
Solid state: Classification of solids, crystalline state, seven crystal systems (cell parameters a, b, c, alpha, beta, gamma), close packed structure of solids (cubic), packing in fcc, bcc and hcp lattices; Nearest neighbours, ionic radii, simple ionic compounds, point defects.
Solutions: Raoult's law; Molecular weight determination from lowering of vapour pressure, elevation of boiling point and depression of freezing point.
Surface chemistry: Elementary concepts of adsorption (excluding adsorption isotherms); Colloids: types, methods of preparation and general properties; Elementary ideas of emulsions, surfactants and micelles (only definitions and examples).
Nuclear chemistry: Radioactivity: isotopes and isobars; Properties of alpha, beta and gamma rays; Kinetics of radioactive decay (decay series excluded), carbon dating; Stability of nuclei with respect to proton-neutron ratio; Brief discussion on fission and fusion reactions.
Inorganic Chemistry
Isolation/preparation and properties of the following non-metals: Boron, silicon, nitrogen, phosphorus, oxygen, sulphur and halogens; Properties of allotropes of carbon (only diamond and graphite), phosphorus and sulphur.
Preparation and properties of the following compounds: Oxides, peroxides, hydroxides, carbonates, bicarbonates, chlorides and sulphates of sodium, potassium, magnesium and calcium; Boron: diborane, boric acid and borax; Aluminium: alumina, aluminium chloride and alums; Carbon: oxides and oxyacid (carbonic acid); Silicon: silicones, silicates and silicon carbide; Nitrogen: oxides, oxyacids and ammonia; Phosphorus: oxides, oxyacids (phosphorus acid, phosphoric acid) and phosphine; Oxygen: ozone and hydrogen peroxide; Sulphur: hydrogen sulphide, oxides, sulphurous acid, sulphuric acid and sodium thiosulphate; Halogens: hydrohalic acids, oxides and oxyacids of chlorine, bleaching powder; Xenon fluorides.
Transition elements (3d series): Definition, general characteristics, oxidation states and their stabilities, colour (excluding the details of electronic transitions) and calculation of spin-only magnetic moment; Coordination compounds: nomenclature of mononuclear coordination compounds, cis-trans and ionisation isomerisms, hybridization and geometries of mononuclear coordination compounds (linear, tetrahedral, square planar and octahedral).
Preparation and properties of the following compounds: Oxides and chlorides of tin and lead; Oxides, chlorides and sulphates of Fe2+, Cu2+ and Zn2+; Potassium permanganate, potassium dichromate, silver oxide, silver nitrate, silver thiosulphate.
Ores and minerals:Commonly occurring ores and minerals of iron, copper, tin, lead, magnesium, aluminium, zinc and silver.
Extractive metallurgy: Chemical principles and reactions only (industrial details excluded); Carbon reduction method (iron and tin); Self reduction method (copper and lead); Electrolytic reduction method (magnesium and aluminium); Cyanide process (silver and gold).
Principles of qualitative analysis: Groups I to V (only Ag+, Hg2+, Cu2+, Pb2+, Bi3+, Fe3+, Cr3+, Al3+, Ca2+, Ba2+, Zn2+, Mn2+ and Mg2+); Nitrate, halides (excluding fluoride), sulphate and sulphide.
Organic Chemistry
Concepts: Hybridisation of carbon; Sigma and pi-bonds; Shapes of simple organic molecules; Structural and geometrical isomerism; Optical isomerism of compounds containing up to two asymmetric centres, (R,S and E,Z nomenclature excluded); IUPAC nomenclature of simple organic compounds (only hydrocarbons, mono-functional and bi-functional compounds); Conformations of ethane and butane (Newman projections); Resonance and hyperconjugation; Keto-enol tautomerism; Determination of empirical and molecular formulae of simple compounds (only combustion method); Hydrogen bonds: definition and their effects on physical properties of alcohols and carboxylic acids; Inductive and resonance effects on acidity and basicity of organic acids and bases; Polarity and inductive effects in alkyl halides; Reactive intermediates produced during homolytic and heterolytic bond cleavage; Formation, structure and stability of carbocations, carbanions and free radicals.
Preparation, properties and reactions of alkanes: Homologous series, physical properties of alkanes (melting points, boiling points and density); Combustion and halogenation of alkanes; Preparation of alkanes by Wurtz reaction and decarboxylation reactions.
Preparation, properties and reactions of alkenes and alkynes: Physical properties of alkenes and alkynes (boiling points, density and dipole moments); Acidity of alkynes; Acid catalysed hydration of alkenes and alkynes (excluding the stereochemistry of addition and elimination); Reactions of alkenes with KMnO4 and ozone; Reduction of alkenes and alkynes; Preparation of alkenes and alkynes by elimination reactions; Electrophilic addition reactions of alkenes with X2, HX, HOX and H2O (X=halogen); Addition reactions of alkynes; Metal acetylides.
Reactions of benzene: Structure and aromaticity; Electrophilic substitution reactions: halogenation, nitration, sulphonation, Friedel-Crafts alkylation and acylation; Effect of o-, m- and p-directing groups in monosubstituted benzenes.
Phenols: Acidity, electrophilic substitution reactions (halogenation, nitration and sulphonation); Reimer-Tieman reaction, Kolbe reaction.
Characteristic reactions of the following (including those mentioned above): Alkyl halides: rearrangement reactions of alkyl carbocation, Grignard reactions, nucleophilic substitution reactions; Alcohols: esterification, dehydration and oxidation, reaction with sodium, phosphorus halides, ZnCl2/concentrated HCl, conversion of alcohols into aldehydes and ketones; Ethers:Preparation by Williamson's Synthesis; Aldehydes and Ketones: oxidation, reduction, oxime and hydrazone formation; aldol condensation, Perkin reaction; Cannizzaro reaction; haloform reaction and nucleophilic addition reactions (Grignard addition); Carboxylic acids: formation of esters, acid chlorides and amides, ester hydrolysis; Amines: basicity of substituted anilines and aliphatic amines, preparation from nitro compounds, reaction with nitrous acid, azo coupling reaction of diazonium salts of aromatic amines, Sandmeyer and related reactions of diazonium salts; carbylamine reaction; Haloarenes: nucleophilic aromatic substitution in haloarenes and substituted haloarenes (excluding Benzyne mechanism and Cine substitution).
Carbohydrates: Classification; mono- and di-saccharides (glucose and sucrose); Oxidation, reduction, glycoside formation and hydrolysis of sucrose.
Amino acids and peptides: General structure (only primary structure for peptides) and physical properties.
Properties and uses of some important polymers: Natural rubber, cellulose, nylon, teflon and PVC.
Practical organic chemistry: Detection of elements (N, S, halogens); Detection and identification of the following functional groups: hydroxyl (alcoholic and phenolic), carbonyl (aldehyde and ketone), carboxyl, amino and nitro; Chemical methods of separation of mono-functional organic compounds from binary mixtures.
Source: http://www.iitkgp.ernet.in/jee/chemistry.htm
Chemistry Syllabus
Physical chemistry
General topics: Concept of atoms and molecules; Dalton’s atomic theory; Mole concept; Chemical formulae; Balanced chemical equations; Calculations (based on mole concept) involving common oxidation-reduction, neutralisation, and displacement reactions; Concentration in terms of mole fraction, molarity, molality and normality.
Gaseous and liquid states: Absolute scale of temperature, ideal gas equation; Deviation from ideality, van der Waals equation; Kinetic theory of gases, average, root mean square and most probable velocities and their relation with temperature; Law of partial pressures; Vapour pressure; Diffusion of gases.
Atomic structure and chemical bonding: Bohr model, spectrum of hydrogen atom, quantum numbers; Wave-particle duality, de Broglie hypothesis; Uncertainty principle; Qualitative quantum mechanical picture of hydrogen atom, shapes of s, p and d orbitals; Electronic configurations of elements (up to atomic number 36); Aufbau principle; Pauli’s exclusion principle and Hund’s rule; Orbital overlap and covalent bond; Hybridisation involving s, p and d orbitals only; Orbital energy diagrams for homonuclear diatomic species; Hydrogen bond; Polarity in molecules, dipole moment (qualitative aspects only); VSEPR model and shapes of molecules (linear, angular, triangular, square planar, pyramidal, square pyramidal, trigonal bipyramidal, tetrahedral and octahedral).
Energetics: First law of thermodynamics; Internal energy, work and heat, pressure-volume work; Enthalpy, Hess’s law; Heat of reaction, fusion and vapourization; Second law of thermodynamics; Entropy; Free energy; Criterion of spontaneity.
Chemical equilibrium: Law of mass action; Equilibrium constant, Le Chatelier's principle (effect of concentration, temperature and pressure); Significance of DG and DGo in chemical equilibrium; Solubility product, common ion effect, pH and buffer solutions; Acids and bases (Bronsted and Lewis concepts); Hydrolysis of salts.
Electrochemistry: Electrochemical cells and cell reactions; Standard electrode potentials; Nernst equation and its relation to DG; Electrochemical series, emf of galvanic cells; Faraday's laws of electrolysis; Electrolytic conductance, specific, equivalent and molar conductivity, Kohlrausch's law; Concentration cells.
Chemical kinetics: Rates of chemical reactions; Order of reactions; Rate constant; First order reactions; Temperature dependence of rate constant (Arrhenius equation).
Solid state: Classification of solids, crystalline state, seven crystal systems (cell parameters a, b, c, alpha, beta, gamma), close packed structure of solids (cubic), packing in fcc, bcc and hcp lattices; Nearest neighbours, ionic radii, simple ionic compounds, point defects.
Solutions: Raoult's law; Molecular weight determination from lowering of vapour pressure, elevation of boiling point and depression of freezing point.
Surface chemistry: Elementary concepts of adsorption (excluding adsorption isotherms); Colloids: types, methods of preparation and general properties; Elementary ideas of emulsions, surfactants and micelles (only definitions and examples).
Nuclear chemistry: Radioactivity: isotopes and isobars; Properties of alpha, beta and gamma rays; Kinetics of radioactive decay (decay series excluded), carbon dating; Stability of nuclei with respect to proton-neutron ratio; Brief discussion on fission and fusion reactions.
Inorganic Chemistry
Isolation/preparation and properties of the following non-metals: Boron, silicon, nitrogen, phosphorus, oxygen, sulphur and halogens; Properties of allotropes of carbon (only diamond and graphite), phosphorus and sulphur.
Preparation and properties of the following compounds: Oxides, peroxides, hydroxides, carbonates, bicarbonates, chlorides and sulphates of sodium, potassium, magnesium and calcium; Boron: diborane, boric acid and borax; Aluminium: alumina, aluminium chloride and alums; Carbon: oxides and oxyacid (carbonic acid); Silicon: silicones, silicates and silicon carbide; Nitrogen: oxides, oxyacids and ammonia; Phosphorus: oxides, oxyacids (phosphorus acid, phosphoric acid) and phosphine; Oxygen: ozone and hydrogen peroxide; Sulphur: hydrogen sulphide, oxides, sulphurous acid, sulphuric acid and sodium thiosulphate; Halogens: hydrohalic acids, oxides and oxyacids of chlorine, bleaching powder; Xenon fluorides.
Transition elements (3d series): Definition, general characteristics, oxidation states and their stabilities, colour (excluding the details of electronic transitions) and calculation of spin-only magnetic moment; Coordination compounds: nomenclature of mononuclear coordination compounds, cis-trans and ionisation isomerisms, hybridization and geometries of mononuclear coordination compounds (linear, tetrahedral, square planar and octahedral).
Preparation and properties of the following compounds: Oxides and chlorides of tin and lead; Oxides, chlorides and sulphates of Fe2+, Cu2+ and Zn2+; Potassium permanganate, potassium dichromate, silver oxide, silver nitrate, silver thiosulphate.
Ores and minerals:Commonly occurring ores and minerals of iron, copper, tin, lead, magnesium, aluminium, zinc and silver.
Extractive metallurgy: Chemical principles and reactions only (industrial details excluded); Carbon reduction method (iron and tin); Self reduction method (copper and lead); Electrolytic reduction method (magnesium and aluminium); Cyanide process (silver and gold).
Principles of qualitative analysis: Groups I to V (only Ag+, Hg2+, Cu2+, Pb2+, Bi3+, Fe3+, Cr3+, Al3+, Ca2+, Ba2+, Zn2+, Mn2+ and Mg2+); Nitrate, halides (excluding fluoride), sulphate and sulphide.
Organic Chemistry
Concepts: Hybridisation of carbon; Sigma and pi-bonds; Shapes of simple organic molecules; Structural and geometrical isomerism; Optical isomerism of compounds containing up to two asymmetric centres, (R,S and E,Z nomenclature excluded); IUPAC nomenclature of simple organic compounds (only hydrocarbons, mono-functional and bi-functional compounds); Conformations of ethane and butane (Newman projections); Resonance and hyperconjugation; Keto-enol tautomerism; Determination of empirical and molecular formulae of simple compounds (only combustion method); Hydrogen bonds: definition and their effects on physical properties of alcohols and carboxylic acids; Inductive and resonance effects on acidity and basicity of organic acids and bases; Polarity and inductive effects in alkyl halides; Reactive intermediates produced during homolytic and heterolytic bond cleavage; Formation, structure and stability of carbocations, carbanions and free radicals.
Preparation, properties and reactions of alkanes: Homologous series, physical properties of alkanes (melting points, boiling points and density); Combustion and halogenation of alkanes; Preparation of alkanes by Wurtz reaction and decarboxylation reactions.
Preparation, properties and reactions of alkenes and alkynes: Physical properties of alkenes and alkynes (boiling points, density and dipole moments); Acidity of alkynes; Acid catalysed hydration of alkenes and alkynes (excluding the stereochemistry of addition and elimination); Reactions of alkenes with KMnO4 and ozone; Reduction of alkenes and alkynes; Preparation of alkenes and alkynes by elimination reactions; Electrophilic addition reactions of alkenes with X2, HX, HOX and H2O (X=halogen); Addition reactions of alkynes; Metal acetylides.
Reactions of benzene: Structure and aromaticity; Electrophilic substitution reactions: halogenation, nitration, sulphonation, Friedel-Crafts alkylation and acylation; Effect of o-, m- and p-directing groups in monosubstituted benzenes.
Phenols: Acidity, electrophilic substitution reactions (halogenation, nitration and sulphonation); Reimer-Tieman reaction, Kolbe reaction.
Characteristic reactions of the following (including those mentioned above): Alkyl halides: rearrangement reactions of alkyl carbocation, Grignard reactions, nucleophilic substitution reactions; Alcohols: esterification, dehydration and oxidation, reaction with sodium, phosphorus halides, ZnCl2/concentrated HCl, conversion of alcohols into aldehydes and ketones; Ethers:Preparation by Williamson's Synthesis; Aldehydes and Ketones: oxidation, reduction, oxime and hydrazone formation; aldol condensation, Perkin reaction; Cannizzaro reaction; haloform reaction and nucleophilic addition reactions (Grignard addition); Carboxylic acids: formation of esters, acid chlorides and amides, ester hydrolysis; Amines: basicity of substituted anilines and aliphatic amines, preparation from nitro compounds, reaction with nitrous acid, azo coupling reaction of diazonium salts of aromatic amines, Sandmeyer and related reactions of diazonium salts; carbylamine reaction; Haloarenes: nucleophilic aromatic substitution in haloarenes and substituted haloarenes (excluding Benzyne mechanism and Cine substitution).
Carbohydrates: Classification; mono- and di-saccharides (glucose and sucrose); Oxidation, reduction, glycoside formation and hydrolysis of sucrose.
Amino acids and peptides: General structure (only primary structure for peptides) and physical properties.
Properties and uses of some important polymers: Natural rubber, cellulose, nylon, teflon and PVC.
Practical organic chemistry: Detection of elements (N, S, halogens); Detection and identification of the following functional groups: hydroxyl (alcoholic and phenolic), carbonyl (aldehyde and ketone), carboxyl, amino and nitro; Chemical methods of separation of mono-functional organic compounds from binary mixtures.
Source: http://www.iitkgp.ernet.in/jee/chemistry.htm
Chemistry JEE Matrix Match Questions
JEE 2007 paper I
Match the chemical substances in column I with type of polymers/type of bonds in column II.
Indicate your answer by darkening the appropriate bubbles of the 4 × 4 matrix given in the ORS.
------Column I----------------------- Column II
(A) Cellulose------------ (p) natural polymer
(B) nylon-6,6------------ (q) synthetic polymer
(C) Protein-------------- (r) amide linkage
(D) Sucrose--------------- (s) glycoside linkage
Solution:
A — p, A — s
B — q, B — r
C — r, C — p
D — s
---------------------------------
JEE 2007 Paper II
Match gases under specified conditions listed in column I with their properties/laws in column II.
Indicate your answer by darkening the appropriate bubbles of the 4 × 4 matrix given in the ORS.
Column I------------------ Column II
(A) hydrogen gas
(P = 200 atm,
T = 273 K) ---------------(p) compressibility factor ¹ 1
(B) hydrogen gas
(P = 0, T = 273 K)------- (q) attractive forces are dominant
(C) 2 CO (P = 1 atm, T = 273 K)--- (r) PV = nRT
(D) real gas with very
large molar volume------------ (s) P(V – nb) = nRT
Solution:
A — p, s,
B — r,
C — q, p,
D — p, q
------------------
JEE question 2007 paper II
Match the complexes in column I with nature of the reactions/type of the products in column II. Indicate your answer by darkening the appropriate bubbles of the 4 × 4 matrix given in the ORS.
(A) O-2^- → O-2+O-2^2- -------(p) redox reaction
====================================================
(B) [CrO-4]^2- +H^+ → --------(q) one of the products
----------------------------------has trigonal-planar
----------------------------------structure
=====================================================
(C)[MnO-4]^- +[NO-2]^- +H^+ → (r) dimeric bridged
---------------------------------tetrahedral metal ion
=====================================================
(D)[NO-3]^- + H-2SO-4+Fe^2+ → (s) disproportionation
====================================================
Solution
A – p, s
B – r
C – p, q
D – p
----------------------------------------
JEE Question 2007 Paper II
Match the compounds/ions in Column I with their properties/reactions in Column II. Indicate your answer by darkening the appropriate bubbles of the 4 × 4 matrix given in the ORS.
Column I------------------- Column II
===============================================
(A) C-6H-5CH0------- (p) gives precipitate
-------------------------with 2, 4-
-------------------------dinitrophenylhydrazine
=================================================
(B) CH-3C≡CH-------- (q) gives precipitate with AgNO3
===================================================
(C) CN^--------------(r) is a nucleophile
===================================================
(D) I^- -------------(s) is involved in
--------------------------cyanohydrin formation
====================================================
Solution
A – p, q, s
B – q
C – r, s
D – q, r
---------------------------------------------------
Match the chemical substances in column I with type of polymers/type of bonds in column II.
Indicate your answer by darkening the appropriate bubbles of the 4 × 4 matrix given in the ORS.
------Column I----------------------- Column II
(A) Cellulose------------ (p) natural polymer
(B) nylon-6,6------------ (q) synthetic polymer
(C) Protein-------------- (r) amide linkage
(D) Sucrose--------------- (s) glycoside linkage
Solution:
A — p, A — s
B — q, B — r
C — r, C — p
D — s
---------------------------------
JEE 2007 Paper II
Match gases under specified conditions listed in column I with their properties/laws in column II.
Indicate your answer by darkening the appropriate bubbles of the 4 × 4 matrix given in the ORS.
Column I------------------ Column II
(A) hydrogen gas
(P = 200 atm,
T = 273 K) ---------------(p) compressibility factor ¹ 1
(B) hydrogen gas
(P = 0, T = 273 K)------- (q) attractive forces are dominant
(C) 2 CO (P = 1 atm, T = 273 K)--- (r) PV = nRT
(D) real gas with very
large molar volume------------ (s) P(V – nb) = nRT
Solution:
A — p, s,
B — r,
C — q, p,
D — p, q
------------------
JEE question 2007 paper II
Match the complexes in column I with nature of the reactions/type of the products in column II. Indicate your answer by darkening the appropriate bubbles of the 4 × 4 matrix given in the ORS.
(A) O-2^- → O-2+O-2^2- -------(p) redox reaction
====================================================
(B) [CrO-4]^2- +H^+ → --------(q) one of the products
----------------------------------has trigonal-planar
----------------------------------structure
=====================================================
(C)[MnO-4]^- +[NO-2]^- +H^+ → (r) dimeric bridged
---------------------------------tetrahedral metal ion
=====================================================
(D)[NO-3]^- + H-2SO-4+Fe^2+ → (s) disproportionation
====================================================
Solution
A – p, s
B – r
C – p, q
D – p
----------------------------------------
JEE Question 2007 Paper II
Match the compounds/ions in Column I with their properties/reactions in Column II. Indicate your answer by darkening the appropriate bubbles of the 4 × 4 matrix given in the ORS.
Column I------------------- Column II
===============================================
(A) C-6H-5CH0------- (p) gives precipitate
-------------------------with 2, 4-
-------------------------dinitrophenylhydrazine
=================================================
(B) CH-3C≡CH-------- (q) gives precipitate with AgNO3
===================================================
(C) CN^--------------(r) is a nucleophile
===================================================
(D) I^- -------------(s) is involved in
--------------------------cyanohydrin formation
====================================================
Solution
A – p, q, s
B – q
C – r, s
D – q, r
---------------------------------------------------
Sunday, October 21, 2007
Blog Structure
Tbe blog now has important points related to each topic in JEE syllabus in the form of revision points.
They can be accessed through label heads of the chapters.
The points are updated whenever I study these chapters.
Separately for each chapter, I collected the syllabus, some past JEE questions and some materials that explains some concepts in the chapter, web sites that are giving material related to the chapter. These can be accessed through the TMH Study guide label.
They can be accessed through label heads of the chapters.
The points are updated whenever I study these chapters.
Separately for each chapter, I collected the syllabus, some past JEE questions and some materials that explains some concepts in the chapter, web sites that are giving material related to the chapter. These can be accessed through the TMH Study guide label.
Study Guide ch 1. THE CONCEPT OF ATOMS AND MOLECULES
JEE Syllabus
General topics:
The concept of atoms and molecules;
Dalton's atomic theory;
Mole concept;
Chemical formulae;
Balanced chemical equations;
Calculations (based on mole concept) involving common oxidation-reduction, neutralisation, and displacement reactions;
Concentration in terms of mole fraction, molarity, molality and normality.
-------------------------
Main topics in TMH Book Chapter
ATOM AND MOLECULAR MASSES
LAWS OF CHEMICAL COMBINATION
QUANTITATIVE INFORMATION FROM A CHEMICAL EQUATION
EXPRESSING CONCENTRATION OF A SUBSTANCE
CONVERSION OF CONCENTRATION UNITS
CONCEPT OF EQUIVALENT
BALANCING CHEMICAL EQUATION
RULES TO COMPUTE OXIDATION NUMBER
BALANCING REDOX REACTIONS VIA OXIDATION NUMBERS
BALANCING REDOX REACTIONS VIA ION ELECTRON (OR HALF EQUATION) METHOD
-----------------
Sub-topics and concepts
_______________________
ATOM AND MOLECULAR MASSES
* Relative Atomic Mass of an Element
Relative Molecular Mass of a Compound
Atomic Mass Unit
Atomic Mass
Molecular Mass
Mole of a Substance
Amount of a Substance: is expressed in moles.
Molar Mass: The average mass per unit amount of a substance.
----------
LAWS OF CHEMICAL COMBINATION
____________________________
Law of Conservation fo Masses
Law of Constant Composition
Law of Multiple Proportion
QUANTITATIVE INFORMATION FROM A CHEMICAL EQUATION
_________________________________________________
Stochiometric coefficients or numbers: The numbers which appear before the chemical symbols in a chemical equation.
Chemical equation gives information about moles of various reactants and products. Hence molar masses involved in the reaction and molar masses of products.
EXPRESSING CONCENTRATION OF A SUBSTANCE
_______________________________________
Mass percentage of substance in a system
Mole fraction of a substance in a system
Molarity = Amount of a substance (in mol)/Volume of solution expressed in dm^3
It is applicable to solutions only.
The unit of molarity is mol dm^-3. It is commonly abbreviated by the symbol M and is spelled as molar.
Molality = Amount of a a substance (in mol)/Mass of solvent expressed in kg
It is also applicable to solutions only
CONVERSION OF CONCENTRATION UNITS
_________________________________
Mole fraction into molarity
Mole fraction inot molality
Molality not molarity
Molarity into mole fraction
Molality into mole fraction
Molarity into molality
--------------------
CONCEPT OF EQUIVALENT
__________________________
In TMH Book it is given: "One equivalent of a substance in a reaction is defined as the amount of substance which reacts or liberates 1 mol of electrons (or H^+ or OH^- ions). In a reaction, a substance always reacts with another substance in equivalent amounts."
More detailed treatment is given in Madan and Bisht, ISC Chemistry.
Equivalent weight of a substance (element or compound) is defined as 'the number of parts by weight of it, that will combine with or displace directly or indirectly, 1.008 parts by weight of hydrogen, 8 parts by weight of oxygen, 35.5 parts by weight of chlorine, 108 parts by weight of silver or the equivalent parts by weight of any other elements."
The meaning will be more clear as we study experiments that determine equivalent weights. Since substances react in the ratio of their equivalent weights, their knowledge and experimental determination is extremely important. A number of methods, with different methods being applicable to different materials are available. Two methods are mentioned below.
a. Hydrogen displacement method: The equivalent weight of metals like calcium, zinc, tin, magnesium etc., which react with dilute acids to produce hydrogen can be determined by this method. In this method, a known quantity of the metal is treated with excess of the dilue sulphuric acid or hydrochloric acid till whole of the metal disappears by way of reaction with the acid. the volume of hydrogen gas so produced is determined and reduced to STP. The weight of the hydrogen at STP is determined from the volume of hydrogen at STP by making use of molar volume concept. From the weight of hydrogen produced and the weight of metal taken for the experiment, equivalent weight of the metal is determined.
Example: o.205 g of a metal on treatment with a dilute acid gave 106.6 ml of hydrogen(dry) at 740.6 mm pressure and 17 degrees centrigrade temperature. Calculate the equivalent weight of the metal.
volume of hydrogen at 740.6 mm pressure and 17 degree centigrade (290 K) = 106.6 ml
Volume of hydrogen at STP 760 mm and 273 K = (740.6/290)*(273/760)*106.6
= 97.72 ml
22.4 litres of H-2 = 22400 ml of H-2 at STP weighs 2g
So 97.72 ml of hydrogen weights (97.72/22400)*2 g
0.205 g of metal gives (97.72/22400)*2 g = .008725 g
To get one g of hydrogen, metal required = (1/.008725)*.205 = 23.5 g
Hence equivalent weight of metal = 23.5
23.5 parts of metal by weight react with 1 part by weight of hydrogen.
b. Oxide Formation Method
The equivalent weight of metals like copper, magnesium, mercury, zinc etc., which form their oxides relatively easily can be found by this method.
In this method a known weight of the metal is taken an converted into its oxide say by heating the metal in the atmosphere of oxygen.
Calculation: Let the weight of the metal = x g.
Let the weight of oxide formed = y g
Therefore y - x g of oxygen combines with x g of the metal.
1 g of oxygen combines with [1/(y-x)]*x g of metal
8 g of oxygen combines with [1/(y-x)]*8x g of metal. this quantity is the equivalent weight of metal.
----------------------
BALANCING CHEMICAL EQUATION
RULES TO COMPUTE OXIDATION NUMBER
BALANCING REDOX REACTIONS VIA OXIDATION NUMBERS
BALANCING REDOX REACTIONS VIA ION ELECTRON (OR HALF EQUATION) METHOD
--------------
web sites
Calculations - Mole concept and some more
http://www.docbrown.info/page04/4_73calcs.htm
------------------------------
JEE Question 2007 paper I Linked Comprehension
Chemical reactions involve interaction of atoms and molecules. A large number of atoms/molecules (approximately 6.023 ×10^23 ) are present in a few grams of any chemical compound varying with their atomic/molecular masses. To handle such large numbers conveniently, the mole concept was introduced. This concept has implications in diverse areas such as analytical chemistry, biochemistry electrochemistry and radiochemistry. The following example illustrates a typical case, involving
chemical/electrochemical reaction, which requires a clear understanding of the mole concept.
A 4.0 molar aqueous solution NaCl is prepared and 500 mL of this solution is electrolysed. This leads to the evolution of chlorine gas at one of the electrodes (atomic mass : Na = 23, Hg = 200;1 Faraday = 96500 coulombs).
1. The total number of moles of chlorine gas evolved is
(A) 0.5
(B) 1.0
(C) 2.0
(D) 3.0
Answer: B
2. If the cathode is a Hg electrode, the maximum weight (g) of amalgam formed from this solution is
(A) 200
(B) 225
(C) 400
(D) 446
Answer: D
In presence of Hg cathode sodium ion will discharge in place of hydrogen gas due to over voltage in the form of amalgams.
Weight of amalgam = 2 × (23 + 200) = 446 g
3. The total charge (coulombs) required for compete electrolysis is
(A) 24125
(B) 48250
(C) 96500
(D) 193000
Answer: D
Total charge = 2 × 96500 = 193000 C
This question is better answered after studying the electro chemistry chapter.
--------------------------
JEE Question 2007 paper II
Consider a titration of potassium dichromate solution with acidified Mohr’s salt solution using diphenylamine as indicator. The number of moles of Mohr’s salt required per mole of dichromate is
(A) 3
(B) 4
(C) 5
(D) 6
Answer D
From Compounds of metals chapter
Mohr's salt is a double salt. FeSO-4.(NH-4)-2SO-4.6H-2O
It is obtained by mixing freshly prepared ferrous sulphate in solution with equal molar amounts of ammonium sulphate and then allow the solution to crystallize.
----------------
General topics:
The concept of atoms and molecules;
Dalton's atomic theory;
Mole concept;
Chemical formulae;
Balanced chemical equations;
Calculations (based on mole concept) involving common oxidation-reduction, neutralisation, and displacement reactions;
Concentration in terms of mole fraction, molarity, molality and normality.
-------------------------
Main topics in TMH Book Chapter
ATOM AND MOLECULAR MASSES
LAWS OF CHEMICAL COMBINATION
QUANTITATIVE INFORMATION FROM A CHEMICAL EQUATION
EXPRESSING CONCENTRATION OF A SUBSTANCE
CONVERSION OF CONCENTRATION UNITS
CONCEPT OF EQUIVALENT
BALANCING CHEMICAL EQUATION
RULES TO COMPUTE OXIDATION NUMBER
BALANCING REDOX REACTIONS VIA OXIDATION NUMBERS
BALANCING REDOX REACTIONS VIA ION ELECTRON (OR HALF EQUATION) METHOD
-----------------
Sub-topics and concepts
_______________________
ATOM AND MOLECULAR MASSES
* Relative Atomic Mass of an Element
Relative Molecular Mass of a Compound
Atomic Mass Unit
Atomic Mass
Molecular Mass
Mole of a Substance
Amount of a Substance: is expressed in moles.
Molar Mass: The average mass per unit amount of a substance.
----------
LAWS OF CHEMICAL COMBINATION
____________________________
Law of Conservation fo Masses
Law of Constant Composition
Law of Multiple Proportion
QUANTITATIVE INFORMATION FROM A CHEMICAL EQUATION
_________________________________________________
Stochiometric coefficients or numbers: The numbers which appear before the chemical symbols in a chemical equation.
Chemical equation gives information about moles of various reactants and products. Hence molar masses involved in the reaction and molar masses of products.
EXPRESSING CONCENTRATION OF A SUBSTANCE
_______________________________________
Mass percentage of substance in a system
Mole fraction of a substance in a system
Molarity = Amount of a substance (in mol)/Volume of solution expressed in dm^3
It is applicable to solutions only.
The unit of molarity is mol dm^-3. It is commonly abbreviated by the symbol M and is spelled as molar.
Molality = Amount of a a substance (in mol)/Mass of solvent expressed in kg
It is also applicable to solutions only
CONVERSION OF CONCENTRATION UNITS
_________________________________
Mole fraction into molarity
Mole fraction inot molality
Molality not molarity
Molarity into mole fraction
Molality into mole fraction
Molarity into molality
--------------------
CONCEPT OF EQUIVALENT
__________________________
In TMH Book it is given: "One equivalent of a substance in a reaction is defined as the amount of substance which reacts or liberates 1 mol of electrons (or H^+ or OH^- ions). In a reaction, a substance always reacts with another substance in equivalent amounts."
More detailed treatment is given in Madan and Bisht, ISC Chemistry.
Equivalent weight of a substance (element or compound) is defined as 'the number of parts by weight of it, that will combine with or displace directly or indirectly, 1.008 parts by weight of hydrogen, 8 parts by weight of oxygen, 35.5 parts by weight of chlorine, 108 parts by weight of silver or the equivalent parts by weight of any other elements."
The meaning will be more clear as we study experiments that determine equivalent weights. Since substances react in the ratio of their equivalent weights, their knowledge and experimental determination is extremely important. A number of methods, with different methods being applicable to different materials are available. Two methods are mentioned below.
a. Hydrogen displacement method: The equivalent weight of metals like calcium, zinc, tin, magnesium etc., which react with dilute acids to produce hydrogen can be determined by this method. In this method, a known quantity of the metal is treated with excess of the dilue sulphuric acid or hydrochloric acid till whole of the metal disappears by way of reaction with the acid. the volume of hydrogen gas so produced is determined and reduced to STP. The weight of the hydrogen at STP is determined from the volume of hydrogen at STP by making use of molar volume concept. From the weight of hydrogen produced and the weight of metal taken for the experiment, equivalent weight of the metal is determined.
Example: o.205 g of a metal on treatment with a dilute acid gave 106.6 ml of hydrogen(dry) at 740.6 mm pressure and 17 degrees centrigrade temperature. Calculate the equivalent weight of the metal.
volume of hydrogen at 740.6 mm pressure and 17 degree centigrade (290 K) = 106.6 ml
Volume of hydrogen at STP 760 mm and 273 K = (740.6/290)*(273/760)*106.6
= 97.72 ml
22.4 litres of H-2 = 22400 ml of H-2 at STP weighs 2g
So 97.72 ml of hydrogen weights (97.72/22400)*2 g
0.205 g of metal gives (97.72/22400)*2 g = .008725 g
To get one g of hydrogen, metal required = (1/.008725)*.205 = 23.5 g
Hence equivalent weight of metal = 23.5
23.5 parts of metal by weight react with 1 part by weight of hydrogen.
b. Oxide Formation Method
The equivalent weight of metals like copper, magnesium, mercury, zinc etc., which form their oxides relatively easily can be found by this method.
In this method a known weight of the metal is taken an converted into its oxide say by heating the metal in the atmosphere of oxygen.
Calculation: Let the weight of the metal = x g.
Let the weight of oxide formed = y g
Therefore y - x g of oxygen combines with x g of the metal.
1 g of oxygen combines with [1/(y-x)]*x g of metal
8 g of oxygen combines with [1/(y-x)]*8x g of metal. this quantity is the equivalent weight of metal.
----------------------
BALANCING CHEMICAL EQUATION
RULES TO COMPUTE OXIDATION NUMBER
BALANCING REDOX REACTIONS VIA OXIDATION NUMBERS
BALANCING REDOX REACTIONS VIA ION ELECTRON (OR HALF EQUATION) METHOD
--------------
web sites
Calculations - Mole concept and some more
http://www.docbrown.info/page04/4_73calcs.htm
------------------------------
JEE Question 2007 paper I Linked Comprehension
Chemical reactions involve interaction of atoms and molecules. A large number of atoms/molecules (approximately 6.023 ×10^23 ) are present in a few grams of any chemical compound varying with their atomic/molecular masses. To handle such large numbers conveniently, the mole concept was introduced. This concept has implications in diverse areas such as analytical chemistry, biochemistry electrochemistry and radiochemistry. The following example illustrates a typical case, involving
chemical/electrochemical reaction, which requires a clear understanding of the mole concept.
A 4.0 molar aqueous solution NaCl is prepared and 500 mL of this solution is electrolysed. This leads to the evolution of chlorine gas at one of the electrodes (atomic mass : Na = 23, Hg = 200;1 Faraday = 96500 coulombs).
1. The total number of moles of chlorine gas evolved is
(A) 0.5
(B) 1.0
(C) 2.0
(D) 3.0
Answer: B
2. If the cathode is a Hg electrode, the maximum weight (g) of amalgam formed from this solution is
(A) 200
(B) 225
(C) 400
(D) 446
Answer: D
In presence of Hg cathode sodium ion will discharge in place of hydrogen gas due to over voltage in the form of amalgams.
Weight of amalgam = 2 × (23 + 200) = 446 g
3. The total charge (coulombs) required for compete electrolysis is
(A) 24125
(B) 48250
(C) 96500
(D) 193000
Answer: D
Total charge = 2 × 96500 = 193000 C
This question is better answered after studying the electro chemistry chapter.
--------------------------
JEE Question 2007 paper II
Consider a titration of potassium dichromate solution with acidified Mohr’s salt solution using diphenylamine as indicator. The number of moles of Mohr’s salt required per mole of dichromate is
(A) 3
(B) 4
(C) 5
(D) 6
Answer D
From Compounds of metals chapter
Mohr's salt is a double salt. FeSO-4.(NH-4)-2SO-4.6H-2O
It is obtained by mixing freshly prepared ferrous sulphate in solution with equal molar amounts of ammonium sulphate and then allow the solution to crystallize.
----------------
Study Guide Ch.2. GASEOUS, LIQUID AND SOLID STATES
JEE syllabus
Gaseous and liquid states:
Absolute scale of temperature,
ideal gas equation;
Deviation from ideality,
van der Waals equation;
Kinetic theory of gases, average,
root mean square and most probable velocities and their relation with temperature; Law of partial pressures;
Vapour pressure;
Diffusion of gases.
----------------------
Main Topics in TMH Book Chapter
SECTION I GASEOUS STATE
KINETIC THEORY OF GASES
EXPRESSIONS OF SOME USEFUL PHYSICAL QUANTITIES
VAN DER WAALS EQUATION OF STATE
REDUCTION OF VAN DER WAALS TO VIRIAL EQUATION
SECTION II LIQUID STATE
VAPOUR PRESSURE
SURFACE TENSION
VISCOCITY
SECTION III SOLID STATE
CLASSIFICATION OF CRYSTALS BNASED ON BOND TYPE
SEVEN CRYSTAL TYPES
PERCENTAGE OF VOID VOLUME
PONT DEFECTS
-------------------
web sites
Part I States of Matter gas-liquid-solid
http://www.docbrown.info/page03/3_52states.htm
------------------------------
JEE Question 2007 paper II
Matrix matching
Match the chemical system/units cells mentioned in column I with their characteristic features mentioned in column II. Indicate your answer by darkening the appropriate bubbles of the 4 × 4 matrix given in the ORS.
Column I------------------------- Column II
(A)simple cubic and ---------(p)have these cell
face-centred cubic------------parameters a = b = c and
------------------------------ α =β = γ
(B)cubic and rhombohedral-- (q) are two crystal systems
(C) cubic and tetragonal (r) Have only two
-----------------------------crystallographic angles of
------------------------------90°
(D) hexagonal and monoaclinic-- (s) Belong to same
------------------------------------crystal system
Solution
A – p, s
B – p, q
C – q
D – r, q
----------------------------------
Gaseous and liquid states:
Absolute scale of temperature,
ideal gas equation;
Deviation from ideality,
van der Waals equation;
Kinetic theory of gases, average,
root mean square and most probable velocities and their relation with temperature; Law of partial pressures;
Vapour pressure;
Diffusion of gases.
----------------------
Main Topics in TMH Book Chapter
SECTION I GASEOUS STATE
KINETIC THEORY OF GASES
EXPRESSIONS OF SOME USEFUL PHYSICAL QUANTITIES
VAN DER WAALS EQUATION OF STATE
REDUCTION OF VAN DER WAALS TO VIRIAL EQUATION
SECTION II LIQUID STATE
VAPOUR PRESSURE
SURFACE TENSION
VISCOCITY
SECTION III SOLID STATE
CLASSIFICATION OF CRYSTALS BNASED ON BOND TYPE
SEVEN CRYSTAL TYPES
PERCENTAGE OF VOID VOLUME
PONT DEFECTS
-------------------
web sites
Part I States of Matter gas-liquid-solid
http://www.docbrown.info/page03/3_52states.htm
------------------------------
JEE Question 2007 paper II
Matrix matching
Match the chemical system/units cells mentioned in column I with their characteristic features mentioned in column II. Indicate your answer by darkening the appropriate bubbles of the 4 × 4 matrix given in the ORS.
Column I------------------------- Column II
(A)simple cubic and ---------(p)have these cell
face-centred cubic------------parameters a = b = c and
------------------------------ α =β = γ
(B)cubic and rhombohedral-- (q) are two crystal systems
(C) cubic and tetragonal (r) Have only two
-----------------------------crystallographic angles of
------------------------------90°
(D) hexagonal and monoaclinic-- (s) Belong to same
------------------------------------crystal system
Solution
A – p, s
B – p, q
C – q
D – r, q
----------------------------------
Study Guide Ch.3. ATOMIC STRUCTURE
JEE Syllabus
Atomic structure and chemical bonding:
Bohr model, spectrum of hydrogen atom, quantum numbers;
Wave-particle duality, de Broglie hypothesis;
Uncertainty principle;
Quantum mechanical picture of hydrogen atom (qualitative treatment), shapes of s, p and d orbitals;
Electronic configurations of elements (up to atomic number 36);
Aufbau principle;
Pauli's exclusion principle and Hund's rule;
------------
Main topics in TMH Book Chapter
CHARACTERISTICS OF ATOMS
RUTHERFORD'S SCATTERING EXPERIMENT
SPRECTRUM OF HYDROGEN ATOM
BOHR'S ATOMIC MODEL
QUANTUM NUMBERS
ELECTRONIC CONFIGURATION OF ELEMENTS
------------------------------
In 1913, Niels Bohr proposed a model of the atom. He proposed that the electrons in an atom could only be in certain orbits, or energy levels, around the nucleus. Refinement of Bohr theory led to the modern theory of atomic structure based on quantum mechanics. The energy levels proposed by Bohr were calculated and it was further modeled that each energy level or orbit has sublevels and each sublevels has one or more orbitals.
Orbit - sublevel – Orbital
Electrons are moving around the nucleus and orbitals represent a space where a pair of electrons is most likely to be found. This means that electrons may sometimes go out of the orbital but most of the time they are found in the orbital.
The quantum mechanics calculations gave out the result that there is a limit to the number of electrons that can occupy a given orbit or energy level.
Orbits
The orbits are called as shells. The energy level of orbits or shells increases as they increase in distance from the nucleus of the atom. The orbits or shells are represented by numbers as 1,2,3,4,5,6 or 7. They are represented by letters as K,L,M,N,O,P,Q.
It is found that the maximum number of electrons in each energy level is equal to 2n2 where n is the number of energy level.
Therefore energy level 1 will have 2 electrons.
Energy level 2 will have 2*4 = 8 electrons
Energy level 3 will have 2*9 = 18 electrons.
Sublevel of an Orbit
The energy levels, or orbits or shells are further divided into sublevels, or subshells. These subshells are designated by letters: s for the first possible sublevel, p for the second possible sublevel, d for the third, f for the fourth, g for the fifth, and from here on they simply go in alphabets.
The number of sublevels of each energy level is equal to the number of the energy levels. This means energy level 1, the K shell will have only one sub levels – s sublevel. The energy level 2, the L shell will have 2 sub levels – s and p.
Orbitals
Sublevels have further divisions called orbitals. Electrons are found in these orbitals. Each orbital contains two electrons.
“s” sublevel has only one orbital. “p” sublevel has 3 orbitals. “d” sublevel has 5 orbitals. “f” sublevel has 7 orbitals.
As each orbital can hold two electrons, orbitals of s can hold two electrons. The orbitals are of p sublevel are named as px and py and pz. The orbitals of p contain 6 electrons. The orbitals of d are 5. The orbitals of d are named as dxy, dxz,dyz,dx2-y2 and dz2. The d sublevel orbitals contain 10 electrons.
The two electons in each orbital spin in different directions.
Shape of Orbitals
Each type of orbital( s, px and py and pz, dxy, dxz,dyz,dx2-y2 and dz2 ) has a unique shape.
1. Spherical shape for s.
2. Dumbbell shape for orbitals of p.
3. Four-lobed shape for orbitals of d.
4. Complex shape for all orbitals of higher sublevels.
Electron Structure or Configuration
In the orbital-sub level-orbit structure, energy level of orbit 1 is less than that of 2 and so on. In each orbit, the sublevel s if of lower energy than the p sublevel, and p is of lower energy than the d sublevel and so on. Orbital of a sublevel are all of equal energy.
Electrons occupy the lowest energy sublevels that are available. This is known as ‘aufbau’ order or principles. In the case of an atom having atomic number of 1, the lone electron occupied the s orbital of sublevel s of orbit 1(represented as 1s1). In case of an atom having atomic number 3 the electrons first occupy the sublevel of orbit 1(this can hold only two electrons) and then occupy p sublevels of orbit 2 (represented as 1s2,2s1).
Hund’s rule says that, for any set of orbitals of equal energy say p orbitals of orbit 2, there is one electron is each orbital before the second electron enters or occupies an orbital.
The energy level of some sublevels at higher orbits is less than the some sublevels at lower orbitals. This fact is to be kept in mind when electron configuration is determined for any atom. The increasing order of energy levels of sublevels is:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f*, 5d, 6p, 7s, 5f*, 6d, 7p, 8s
* In this case one electron one electron goes into 5d and then 4f fills completely and then rest of 5d. Similar thing happens in 5f and 4d.
-------------
Atomic particles are said to have spin. What does that imply?
Solution:
A good explanation for electron spin is provided by
Britannica:
Spin is an intrinsic property of an electron, like its mass or charge. In elementary treatments, spin is often visualized as an actual spinning motion. However, it is a quantum mechanical property without a classical counterpart, and to picture spin in this way can be misleading. Nevertheless, for the present discussion, such a picture is useful. An electron has a fixed amount of spin, in the sense that every electron in the universe is continually spinning at exactly the same rate. Although the spin of an electron is constant, the orientation of the axis of spin is variable, but quantum mechanics restricts that orientation to only two possibilities. The two possible spin states of an electron are represented by the arrows and and are distinguished by the spin magnetic quantum number, ms, which takes the values +1/2 (for the spin) or -1/2 (for the spin). Because of its spin, an electron must obey a fundamental requirement known as the Pauli exclusion principle. This principle (which is a consequence of the more fundamental Pauli principle; see the article atom: Electrons) states that no more than two electrons may occupy a given orbital and, if two electrons do occupy one orbital, their spins must be paired (denoted ; that is, one electron must be and the other must be ). The Pauli exclusion principle is responsible for the importance of the electron pair in the formation of covalent bonds. It is also, on a more cosmic scale, the reason why matter has bulk; that is to say, all electrons cannot occupy the orbitals of lowest energy but are instead located in the many shells that are centred on the nucleus. Also owing to the existence of spin, two objects do not simply blend into one another when they are in contact; the electrons of adjacent atoms cannot occupy the same space, thereby prohibiting the combining of two atoms into one. Here again is an example of a seemingly trivial property, in this case spin, having consequences of profound and macroscopic importance. In this instance, the spin of the electron is responsible for the existence of identifiable forms of matter.
Source: http://www.emsb.qc.ca/laurenhill/science/99scsoln.html#dec
-----------------------------
web sites
Atomic Structure, Isotopes, Periodic Table and Electronic Structure of Atoms
http://www.docbrown.info/page04/4_71atom.htm
Atomic structure and chemical bonding:
Bohr model, spectrum of hydrogen atom, quantum numbers;
Wave-particle duality, de Broglie hypothesis;
Uncertainty principle;
Quantum mechanical picture of hydrogen atom (qualitative treatment), shapes of s, p and d orbitals;
Electronic configurations of elements (up to atomic number 36);
Aufbau principle;
Pauli's exclusion principle and Hund's rule;
------------
Main topics in TMH Book Chapter
CHARACTERISTICS OF ATOMS
RUTHERFORD'S SCATTERING EXPERIMENT
SPRECTRUM OF HYDROGEN ATOM
BOHR'S ATOMIC MODEL
QUANTUM NUMBERS
ELECTRONIC CONFIGURATION OF ELEMENTS
------------------------------
In 1913, Niels Bohr proposed a model of the atom. He proposed that the electrons in an atom could only be in certain orbits, or energy levels, around the nucleus. Refinement of Bohr theory led to the modern theory of atomic structure based on quantum mechanics. The energy levels proposed by Bohr were calculated and it was further modeled that each energy level or orbit has sublevels and each sublevels has one or more orbitals.
Orbit - sublevel – Orbital
Electrons are moving around the nucleus and orbitals represent a space where a pair of electrons is most likely to be found. This means that electrons may sometimes go out of the orbital but most of the time they are found in the orbital.
The quantum mechanics calculations gave out the result that there is a limit to the number of electrons that can occupy a given orbit or energy level.
Orbits
The orbits are called as shells. The energy level of orbits or shells increases as they increase in distance from the nucleus of the atom. The orbits or shells are represented by numbers as 1,2,3,4,5,6 or 7. They are represented by letters as K,L,M,N,O,P,Q.
It is found that the maximum number of electrons in each energy level is equal to 2n2 where n is the number of energy level.
Therefore energy level 1 will have 2 electrons.
Energy level 2 will have 2*4 = 8 electrons
Energy level 3 will have 2*9 = 18 electrons.
Sublevel of an Orbit
The energy levels, or orbits or shells are further divided into sublevels, or subshells. These subshells are designated by letters: s for the first possible sublevel, p for the second possible sublevel, d for the third, f for the fourth, g for the fifth, and from here on they simply go in alphabets.
The number of sublevels of each energy level is equal to the number of the energy levels. This means energy level 1, the K shell will have only one sub levels – s sublevel. The energy level 2, the L shell will have 2 sub levels – s and p.
Orbitals
Sublevels have further divisions called orbitals. Electrons are found in these orbitals. Each orbital contains two electrons.
“s” sublevel has only one orbital. “p” sublevel has 3 orbitals. “d” sublevel has 5 orbitals. “f” sublevel has 7 orbitals.
As each orbital can hold two electrons, orbitals of s can hold two electrons. The orbitals are of p sublevel are named as px and py and pz. The orbitals of p contain 6 electrons. The orbitals of d are 5. The orbitals of d are named as dxy, dxz,dyz,dx2-y2 and dz2. The d sublevel orbitals contain 10 electrons.
The two electons in each orbital spin in different directions.
Shape of Orbitals
Each type of orbital( s, px and py and pz, dxy, dxz,dyz,dx2-y2 and dz2 ) has a unique shape.
1. Spherical shape for s.
2. Dumbbell shape for orbitals of p.
3. Four-lobed shape for orbitals of d.
4. Complex shape for all orbitals of higher sublevels.
Electron Structure or Configuration
In the orbital-sub level-orbit structure, energy level of orbit 1 is less than that of 2 and so on. In each orbit, the sublevel s if of lower energy than the p sublevel, and p is of lower energy than the d sublevel and so on. Orbital of a sublevel are all of equal energy.
Electrons occupy the lowest energy sublevels that are available. This is known as ‘aufbau’ order or principles. In the case of an atom having atomic number of 1, the lone electron occupied the s orbital of sublevel s of orbit 1(represented as 1s1). In case of an atom having atomic number 3 the electrons first occupy the sublevel of orbit 1(this can hold only two electrons) and then occupy p sublevels of orbit 2 (represented as 1s2,2s1).
Hund’s rule says that, for any set of orbitals of equal energy say p orbitals of orbit 2, there is one electron is each orbital before the second electron enters or occupies an orbital.
The energy level of some sublevels at higher orbits is less than the some sublevels at lower orbitals. This fact is to be kept in mind when electron configuration is determined for any atom. The increasing order of energy levels of sublevels is:
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f*, 5d, 6p, 7s, 5f*, 6d, 7p, 8s
* In this case one electron one electron goes into 5d and then 4f fills completely and then rest of 5d. Similar thing happens in 5f and 4d.
-------------
Atomic particles are said to have spin. What does that imply?
Solution:
A good explanation for electron spin is provided by
Britannica:
Spin is an intrinsic property of an electron, like its mass or charge. In elementary treatments, spin is often visualized as an actual spinning motion. However, it is a quantum mechanical property without a classical counterpart, and to picture spin in this way can be misleading. Nevertheless, for the present discussion, such a picture is useful. An electron has a fixed amount of spin, in the sense that every electron in the universe is continually spinning at exactly the same rate. Although the spin of an electron is constant, the orientation of the axis of spin is variable, but quantum mechanics restricts that orientation to only two possibilities. The two possible spin states of an electron are represented by the arrows and and are distinguished by the spin magnetic quantum number, ms, which takes the values +1/2 (for the spin) or -1/2 (for the spin). Because of its spin, an electron must obey a fundamental requirement known as the Pauli exclusion principle. This principle (which is a consequence of the more fundamental Pauli principle; see the article atom: Electrons) states that no more than two electrons may occupy a given orbital and, if two electrons do occupy one orbital, their spins must be paired (denoted ; that is, one electron must be and the other must be ). The Pauli exclusion principle is responsible for the importance of the electron pair in the formation of covalent bonds. It is also, on a more cosmic scale, the reason why matter has bulk; that is to say, all electrons cannot occupy the orbitals of lowest energy but are instead located in the many shells that are centred on the nucleus. Also owing to the existence of spin, two objects do not simply blend into one another when they are in contact; the electrons of adjacent atoms cannot occupy the same space, thereby prohibiting the combining of two atoms into one. Here again is an example of a seemingly trivial property, in this case spin, having consequences of profound and macroscopic importance. In this instance, the spin of the electron is responsible for the existence of identifiable forms of matter.
Source: http://www.emsb.qc.ca/laurenhill/science/99scsoln.html#dec
-----------------------------
web sites
Atomic Structure, Isotopes, Periodic Table and Electronic Structure of Atoms
http://www.docbrown.info/page04/4_71atom.htm
Study Guide Ch. 4. PERIODICITY OF PROPERTIES OF ELEMENTS
JEE Syllabus
No specific syllabus on this topic
-----------------
Main Topics in TMH Book Chapter
CLASSIFICATION OF ELEMENTS
MAIN CHARACTERISTICS OF REPRESENTATIVE ELEMENTS
MAIN CHARACTERISTICS OF TRANSIENT ELEMENTS
PERIODICITY IN PROPERTIES
METALLIC AND NONMETALLIC CHARACTER
LANTHANIDE CONTRACTION
-----------------
Periodic table was covered in the X class syllabus. But there are some interesting points about the periodic table that were covered in Redmund’s book in 5th chapter.
We all know that the horizontal rows in the periodic table are called periods, and the vertical columns are called groups. The groups are subdivided into A and B subgroups. The A subgroups, due to their similarities within a group, are often called families. Some of these families are referred to by special names, such as the alkali metals for Group IA, alkaline earth metals for Group IIA, and halogens for Group VIIA. The other A subgroups are sometimes classified according to the first member of the subgroup or family. Thus the IIIA elements are sometimes referred to as the boron family, the IVA elements as the carbon family, the VA elements as the nitrogen family, and the VIA elements as the oxygen family. The last group, called Group zero (or sometimes Group VIIIA), is now called the noble gases.
As well as being classified as metals or nonmetals, the elements are also divided into the following four types: representative elements, noble gases, transition elements and inner transition elements.
-----------------------------
web sites
Introduction to the Periodic Table
http://www.docbrown.info/page03/3_34ptable.htm
No specific syllabus on this topic
-----------------
Main Topics in TMH Book Chapter
CLASSIFICATION OF ELEMENTS
MAIN CHARACTERISTICS OF REPRESENTATIVE ELEMENTS
MAIN CHARACTERISTICS OF TRANSIENT ELEMENTS
PERIODICITY IN PROPERTIES
METALLIC AND NONMETALLIC CHARACTER
LANTHANIDE CONTRACTION
-----------------
Periodic table was covered in the X class syllabus. But there are some interesting points about the periodic table that were covered in Redmund’s book in 5th chapter.
We all know that the horizontal rows in the periodic table are called periods, and the vertical columns are called groups. The groups are subdivided into A and B subgroups. The A subgroups, due to their similarities within a group, are often called families. Some of these families are referred to by special names, such as the alkali metals for Group IA, alkaline earth metals for Group IIA, and halogens for Group VIIA. The other A subgroups are sometimes classified according to the first member of the subgroup or family. Thus the IIIA elements are sometimes referred to as the boron family, the IVA elements as the carbon family, the VA elements as the nitrogen family, and the VIA elements as the oxygen family. The last group, called Group zero (or sometimes Group VIIIA), is now called the noble gases.
As well as being classified as metals or nonmetals, the elements are also divided into the following four types: representative elements, noble gases, transition elements and inner transition elements.
-----------------------------
web sites
Introduction to the Periodic Table
http://www.docbrown.info/page03/3_34ptable.htm
Study Guide Ch.5. BONDING AND MOLECULAR STRUCTURE
See for past JEE questions from this chapter
http://iit-jee-chemistry-ps.blogspot.com/2007/12/past-jee-questions-ch5.html
See for some application questions
http://iit-jee-chemistry-ps.blogspot.com/2007/12/application-questions-ch-5-bonding.html
JEE Syllabus
Orbital overlap and covalent bond;
Hybridisation involving s, p and d orbitals only;
Orbital energy diagrams for homonuclear diatomic species;
Hydrogen bond;
Polarity in molecules, dipole moment (qualitative aspects only);
VSEPR model and shapes of molecules (linear, angular, triangular, square planar, pyramidal, square pyramidal, trigonal bipyramidal, tetrahedral and octahedral).
------------------
Main Topics in TMH JEE Book
COVALENT BOND
COORDINATE BOND
IONIC BOND
QUANTUM MECHANICAL EXPLANATION OF COVALENT BOND
HYBRIDIZATION
VALENCE SHELL ELECTRON-PAIR-REPULSION MODEL
HYDROGEN BOND
RESONANCE
MOLECULAR ORBITAL METHOD
--------------------
JEE syllabus and sections of Jauhar's Book (Class XI)
Orbital overlap and covalent bond; 6.5, 6.6, 6.7, 6.15, 6.18
Hybridisation involving s, p and d orbitals only; 6.18 (d orbitals not described)
Orbital energy diagrams for homonuclear diatomic species;
Hydrogen bond; 6.20
Polarity in molecules, dipole moment (qualitative aspects only); 6.14
VSEPR model and shapes of molecules (linear, angular, triangular, square planar, pyramidal, square pyramidal, trigonal bipyramidal, tetrahedral and octahedral). 6.13 and 6.13
Section of Jauhar Chapter 6
1. Cause of chemical combination
2. Lewis symbols
3. Octetl rule andmodes of chemical combination
4. Ionic or electrovalent bond
5. Covalent bond
General properties of covalent bonds
7. Co-ordinate covalent bond
8. Formation of ionic bond
9. Lattice enthalpy of ionic crystals
10. Born-haber Cycle for lattice enthalpies
11. General properties of ionic compounds
12. Geometry of shapes of molecules
13. Valence Shell Electron Pair Repulsion (VSEPR) theory
14. electronegativity - polar and nonpolar character of covalent bonds
15. Valency bond approach of covalent bond
---Orbital overlap concept of covalent bond
16. Bonding parameters
17. Resonance
18. Directional properties of covalent bonds
19. Metallic bonding
20. Hydrogen bonding
Ionic Bond
In an ionic bond negative ions are surrounded by positive ions and positive ions are surrounded by negative ions. NaCl does not mean each Na and Cl ions are bound to each other. It only means the proportion of the ions is 1:1.
Hydrogen Bonding
The attractive force which binds hydrogen atom of one molecule with electronegative atom (F,O or N) of another molecule is known as hydrogen bond or hydrogen bonding.
Hydrogen bonds are formed when hydrogen is bonded to strongly electronegative elements uch as F, O or N.
The size of electronegative atom should be small for formation of Hydrogen bonds. Hydrogen bonds are not formed by Cl because of its bigger size.
Example of Hydrogen bonds
HF
H2O
NH3
Influence of Hydrogen bonding on properties
.1 Association
2. Higher melting and boiling points
3. Influence on the physical state
4. Solubility
----------------------
Examples in the Chapter
Double bond
O2 molecule
CO2 molecule C forms double bond with each O atom.
CS2
Triple bond
N2 molecule
CO molecule
Exceptions to octet rule
Hydrogen molecule only 2 electrons make it stable.
Incomplete octet of central atom
LiCl
BeH2
BeCl2
BH3
BF3
LiCl 4 electrons around central Li-atom
BeCl2 4 electrons around central Be-atom
BF3 6 electrons around central B-atom
Expanded octet of the central atom
PF5 has ten around P
SF6 hs twelve around S
IF7 has fourteen electrons around I
H2SO4 12 electrons around sulphur atoms
Odd elctron molecules
Nitric oxide, NO
Nitrogen 7 shared electrons
Oxygen 8
Nitrogen dioxide, NO2(there is a coordinate bond)
Nitrogen 7
both oxygens 8
compounds of noble gases
XeF2, KrF2, XeOF2, XeOF4, XeF6
Coordinate covalent bonds
ozone
Hydronium ion
ammonia
sulphuric acid
Molecule shapes
Linear
BeF2, BeCl2, ZnCl2, HgCl2
Trigonal planar
BF3, BCl3, AlCl3
Tetrahedral
CH4, SiF4, CCl4, SiH4, NH4^+
Trigonal bipyramidal
PCl5, PF5
Octahedral
SF6
Pentagonal bipyramidal
IF7
Molecules containing lone pairs and their shapes
Three electron pairs that include lone pairs
standard shape Trigonal planar
SO2 angular or V shaped or bent shape
Four electron pairs that include lone pairs
standard shape: tetrahedral geometry
Ammonia molecule pyramidal
PCl3, NF3, H3O^+
H2O molecule
two bond pairs and two lone pairs bent or angular
same shape H2S, F2O, SCl2
Five electron pairs that include lone pairs
standard shape: trigonal bipyramidal
SF4 4 bond pairs and 1 lp distorted tetrahedron or a folded square
Chlorine triflouride 3 bp and 2 lp T shaped
Xennon difluoride XeF2 2 bp and 3 lp linear geometry
six electron pairs that include lone pairs
BrF5 5bp and 1 lp square pyramidal
XeF4 4 bp and 2 lp square planar
Polar covalent molecules
HCl, BrCl, H2O, HF
Dipole moment
HCL 1.03 D
CO2 zero
H2O 1.84 D
NH3 1.49D
BF3 zero
CCl4 zero
NF3 0.24D
Molecules described by resonance property
O3 ozone
CO2
CO
SO2
SO3
Benzene
CO3^2-
NO2 ion
Hybridization
sp
BeCl2
BeF2
BeH2
C2H2
sp^2
BCl3
C2H4
sp^3
CH4
NH3
H2O
Hydrogen bonds
HF
H2O
Ammonia NH3
-------------
VSEPR Model
I found the tutorial in the site given below very useful understand this topic.
Molecular geometry: VSEPR http://winter.group.shef.ac.uk/vsepr/intro.html
----------------------------
Molecular orbital theory
I found material from the site below to be useful to understand this topic.
Molecular Orbital Theory http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch8/mo.html
Diamagnetic and Paramagnetic substances
Atoms or molecules in which the electrons are paired are diamagnetic repelled by both poles of a magnetic. Those that have one or more unpaired electrons are paramagnetic attracted to a magnetic field. Liquid oxygen is attracted to a magnetic field and can actually bridge the gap between the poles of a horseshoe magnet. The molecular orbital model of O2 is therefore superior to the valence-bond model, which cannot explain this property of oxygen.
----------------------
web sites
The Structure, Bonding and properties of substances
http://www.docbrown.info/page04/4_72bond.htm
JEE question on dipole moment
Among the following, the molecule with the highest dipole moment is:
(A) CH-3Cl (B) CH-2Cl-2
(C) CHCl-3 (D) CCl-4
answer A
----------------------------
JEE Question 2007 paper I
The percentage of p-character in the orbitals forming P–P bonds in 4 P is
(A) 25
(B) 33
(C) 50
(D) 75
Solution: (D)
Phosphorous will show sp^3 hybridisation having 75% p-character.
-----------------------
JEE Question 2007 paper I
Statement - 1
Boron always forms covalent bond
Because
Statement - 2
The small size of B3^+ favours formation of covalent bond.
(A) Statement – 1 is True, Statement – 2 is True; Statement – 2 is a correct explanation for statement – 1
(B) Statement – 1 is True, Statement – 2 is True; Statement – 2 is Not a correct explanation for Statement – 1.
(C) Statement – 1 is True, Statement – 2 is False
(D) Statement – 1 is False, Statement – 2 is True
Answer: A
-------------------------------
JEE Question paper II 2007
Among the following metal carbonyls, the C — O bond order is lowest in
(A) [Mn(CO)-6]^+
(B) [Fe(CO)-5]
(C) [Cr(CO)-6]
(D) [V(CO)-6]^-
answer: B
-----------------
JEE 2006
If the bond length of CO bond in carbon monoxide is 1.128 A,
then what is the value of CO bond length in Fe(CO)-5?
(A) 1.15 A
(B) 1.128 A
(C) 1.72 A
(D) 1.118 A
Answer: (A)
----------------------
http://iit-jee-chemistry-ps.blogspot.com/2007/12/past-jee-questions-ch5.html
See for some application questions
http://iit-jee-chemistry-ps.blogspot.com/2007/12/application-questions-ch-5-bonding.html
JEE Syllabus
Orbital overlap and covalent bond;
Hybridisation involving s, p and d orbitals only;
Orbital energy diagrams for homonuclear diatomic species;
Hydrogen bond;
Polarity in molecules, dipole moment (qualitative aspects only);
VSEPR model and shapes of molecules (linear, angular, triangular, square planar, pyramidal, square pyramidal, trigonal bipyramidal, tetrahedral and octahedral).
------------------
Main Topics in TMH JEE Book
COVALENT BOND
COORDINATE BOND
IONIC BOND
QUANTUM MECHANICAL EXPLANATION OF COVALENT BOND
HYBRIDIZATION
VALENCE SHELL ELECTRON-PAIR-REPULSION MODEL
HYDROGEN BOND
RESONANCE
MOLECULAR ORBITAL METHOD
--------------------
JEE syllabus and sections of Jauhar's Book (Class XI)
Orbital overlap and covalent bond; 6.5, 6.6, 6.7, 6.15, 6.18
Hybridisation involving s, p and d orbitals only; 6.18 (d orbitals not described)
Orbital energy diagrams for homonuclear diatomic species;
Hydrogen bond; 6.20
Polarity in molecules, dipole moment (qualitative aspects only); 6.14
VSEPR model and shapes of molecules (linear, angular, triangular, square planar, pyramidal, square pyramidal, trigonal bipyramidal, tetrahedral and octahedral). 6.13 and 6.13
Section of Jauhar Chapter 6
1. Cause of chemical combination
2. Lewis symbols
3. Octetl rule andmodes of chemical combination
4. Ionic or electrovalent bond
5. Covalent bond
General properties of covalent bonds
7. Co-ordinate covalent bond
8. Formation of ionic bond
9. Lattice enthalpy of ionic crystals
10. Born-haber Cycle for lattice enthalpies
11. General properties of ionic compounds
12. Geometry of shapes of molecules
13. Valence Shell Electron Pair Repulsion (VSEPR) theory
14. electronegativity - polar and nonpolar character of covalent bonds
15. Valency bond approach of covalent bond
---Orbital overlap concept of covalent bond
16. Bonding parameters
17. Resonance
18. Directional properties of covalent bonds
19. Metallic bonding
20. Hydrogen bonding
Ionic Bond
In an ionic bond negative ions are surrounded by positive ions and positive ions are surrounded by negative ions. NaCl does not mean each Na and Cl ions are bound to each other. It only means the proportion of the ions is 1:1.
Hydrogen Bonding
The attractive force which binds hydrogen atom of one molecule with electronegative atom (F,O or N) of another molecule is known as hydrogen bond or hydrogen bonding.
Hydrogen bonds are formed when hydrogen is bonded to strongly electronegative elements uch as F, O or N.
The size of electronegative atom should be small for formation of Hydrogen bonds. Hydrogen bonds are not formed by Cl because of its bigger size.
Example of Hydrogen bonds
HF
H2O
NH3
Influence of Hydrogen bonding on properties
.1 Association
2. Higher melting and boiling points
3. Influence on the physical state
4. Solubility
----------------------
Examples in the Chapter
Double bond
O2 molecule
CO2 molecule C forms double bond with each O atom.
CS2
Triple bond
N2 molecule
CO molecule
Exceptions to octet rule
Hydrogen molecule only 2 electrons make it stable.
Incomplete octet of central atom
LiCl
BeH2
BeCl2
BH3
BF3
LiCl 4 electrons around central Li-atom
BeCl2 4 electrons around central Be-atom
BF3 6 electrons around central B-atom
Expanded octet of the central atom
PF5 has ten around P
SF6 hs twelve around S
IF7 has fourteen electrons around I
H2SO4 12 electrons around sulphur atoms
Odd elctron molecules
Nitric oxide, NO
Nitrogen 7 shared electrons
Oxygen 8
Nitrogen dioxide, NO2(there is a coordinate bond)
Nitrogen 7
both oxygens 8
compounds of noble gases
XeF2, KrF2, XeOF2, XeOF4, XeF6
Coordinate covalent bonds
ozone
Hydronium ion
ammonia
sulphuric acid
Molecule shapes
Linear
BeF2, BeCl2, ZnCl2, HgCl2
Trigonal planar
BF3, BCl3, AlCl3
Tetrahedral
CH4, SiF4, CCl4, SiH4, NH4^+
Trigonal bipyramidal
PCl5, PF5
Octahedral
SF6
Pentagonal bipyramidal
IF7
Molecules containing lone pairs and their shapes
Three electron pairs that include lone pairs
standard shape Trigonal planar
SO2 angular or V shaped or bent shape
Four electron pairs that include lone pairs
standard shape: tetrahedral geometry
Ammonia molecule pyramidal
PCl3, NF3, H3O^+
H2O molecule
two bond pairs and two lone pairs bent or angular
same shape H2S, F2O, SCl2
Five electron pairs that include lone pairs
standard shape: trigonal bipyramidal
SF4 4 bond pairs and 1 lp distorted tetrahedron or a folded square
Chlorine triflouride 3 bp and 2 lp T shaped
Xennon difluoride XeF2 2 bp and 3 lp linear geometry
six electron pairs that include lone pairs
BrF5 5bp and 1 lp square pyramidal
XeF4 4 bp and 2 lp square planar
Polar covalent molecules
HCl, BrCl, H2O, HF
Dipole moment
HCL 1.03 D
CO2 zero
H2O 1.84 D
NH3 1.49D
BF3 zero
CCl4 zero
NF3 0.24D
Molecules described by resonance property
O3 ozone
CO2
CO
SO2
SO3
Benzene
CO3^2-
NO2 ion
Hybridization
sp
BeCl2
BeF2
BeH2
C2H2
sp^2
BCl3
C2H4
sp^3
CH4
NH3
H2O
Hydrogen bonds
HF
H2O
Ammonia NH3
-------------
VSEPR Model
I found the tutorial in the site given below very useful understand this topic.
Molecular geometry: VSEPR http://winter.group.shef.ac.uk/vsepr/intro.html
----------------------------
Molecular orbital theory
I found material from the site below to be useful to understand this topic.
Molecular Orbital Theory http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch8/mo.html
Diamagnetic and Paramagnetic substances
Atoms or molecules in which the electrons are paired are diamagnetic repelled by both poles of a magnetic. Those that have one or more unpaired electrons are paramagnetic attracted to a magnetic field. Liquid oxygen is attracted to a magnetic field and can actually bridge the gap between the poles of a horseshoe magnet. The molecular orbital model of O2 is therefore superior to the valence-bond model, which cannot explain this property of oxygen.
----------------------
web sites
The Structure, Bonding and properties of substances
http://www.docbrown.info/page04/4_72bond.htm
JEE question on dipole moment
Among the following, the molecule with the highest dipole moment is:
(A) CH-3Cl (B) CH-2Cl-2
(C) CHCl-3 (D) CCl-4
answer A
----------------------------
JEE Question 2007 paper I
The percentage of p-character in the orbitals forming P–P bonds in 4 P is
(A) 25
(B) 33
(C) 50
(D) 75
Solution: (D)
Phosphorous will show sp^3 hybridisation having 75% p-character.
-----------------------
JEE Question 2007 paper I
Statement - 1
Boron always forms covalent bond
Because
Statement - 2
The small size of B3^+ favours formation of covalent bond.
(A) Statement – 1 is True, Statement – 2 is True; Statement – 2 is a correct explanation for statement – 1
(B) Statement – 1 is True, Statement – 2 is True; Statement – 2 is Not a correct explanation for Statement – 1.
(C) Statement – 1 is True, Statement – 2 is False
(D) Statement – 1 is False, Statement – 2 is True
Answer: A
-------------------------------
JEE Question paper II 2007
Among the following metal carbonyls, the C — O bond order is lowest in
(A) [Mn(CO)-6]^+
(B) [Fe(CO)-5]
(C) [Cr(CO)-6]
(D) [V(CO)-6]^-
answer: B
-----------------
JEE 2006
If the bond length of CO bond in carbon monoxide is 1.128 A,
then what is the value of CO bond length in Fe(CO)-5?
(A) 1.15 A
(B) 1.128 A
(C) 1.72 A
(D) 1.118 A
Answer: (A)
----------------------
Study Guide Ch. 6 ENERGETICS
JEE Syllabus
Energetics:
First law of thermodynamics;
Internal energy, work and heat,
pressure-volume work;
Enthalpy,
Hess's law;
Heat of reaction, fusion and vapourization;
Second law of thermodynamics;
Entropy;
Free energy;
Criterion of spontaneity.
------------------
Main Topics in TMH Book Chapter
FIRST LAW THERMODYNAMICS
INTERNAL ENERGY AND ENTHALPY
ENTHALPY CHANGE OF A CHEMICAL EQUATION
MOLAR ENHTALPIES OF FORMATIONS
HESS'S LAW OF CONSTANT HEAT SUMMATION
TYPES OF REACTIONS
RELATION BETWEEN DELTA H AND DELTA U OF A CHEMICAL EQUATION
--------------------------
Enthalpy Change, ∆H, With Temperature and State Change
Enthalpy, represented by the symbol H, is essentially a chemistry term for heat, and a term for total kinetic energy of particle motion in a sample. If the reaction is exothermic, the energy contained in the substances is reduced so ∆H has a negative value. On the other hand, if the reaction is endothermic, the substances absorb energy and ∆H is positive.
Temperature is a measure of the average kinetic energy of the particles in a sample.
The faster the molecules in a sample of water the higher the temperature.
Changing temperature of a sample requires a change in enthalpy
It makes sense that you have to add heat to increase temperature, and as a sample cools it gives off heat. The amount of heat required to change the temperature of a sample depends on two factors:
# Mass. The bigger the sample the more heat needed to change its temperature.
# Specific heat capacity of the substance it is made of. This is the energy needed to change the temperature of a 1 gram sample 1 Celsius and it is different for different materials. Its symbol is Cp.
When you go to the beach on a hot summer day, you may have noticed that dry sand can get really hot, while wet sand stays cooler, though the air over them is obviously the same temperature and they are getting the same sunlight. How about dry pavement versus wet pavement? Metal versus tile?
The difference is the specific heat capacity of the different materials. Substances with low specific heat capacities will get very hot very fast, while those with large specific heat capacities warm up much slower since more energy is required to change their temperature. In fact, liquid water has about the highest specific heat of any common substance, so it tends to "heat up" and 'cool down' very slowly compared to most other substances. Ice and steam have lower heat capacities than liquid water.
Some people who prize their outdoor plants set gallon jugs of water around their plants. How does this help the plants? During the heat of the day, that water absorbs energy from the air around it, leaving the air cooler. During the cold night, that water releases that energy as it slowly cools, warming the air around it. This is known as a heat sink.
To find the energy required to change the temperature of a sample use Changing temperature of a sample requires a change in enthalpy
It makes sense that you have to add heat to increase temperature, and as a sample cools it gives off heat. The amount of heat required to change the temperature of a sample depends on two factors:
? Mass. The bigger the sample the more heat needed to change its temperature.
? Specific heat capacity of the substance it is made of. This is the energy needed to change the temperature of a 1 gram sample 1? Celsius and it is different for different materials. For some reason its symbol is Cp. Sorry.
When you go to the beach on a hot summer day, you may have noticed that dry sand can get really hot, while wet sand stays cooler, though the air over them is obviously the same temperature and they are getting the same sunlight. How about dry pavement versus wet pavement? Metal versus tile?
The difference is the specific heat capacity of the different materials. Substances with low specific heat capacities will get very hot very fast, while those with large specific heat capacities warm up much slower since more energy is required to change their temperature. In fact, liquid water has about the highest specific heat of any common substance, so it tends to "heat up" and 'cool down' very slowly compared to most other substances. Ice and steam have lower heat capacities than liquid water.
Some people who prize their outdoor plants set gallon jugs of water around their plants. How does this help the plants? During the heat of the day, that water absorbs energy from the air around it, leaving the air cooler. During the cold night, that water releases that energy as it slowly cools, warming the air around it. This is known as a heat sink.
In Chicago, which sits on the shores of Lake Michigan, the weather near the lake is warmer in winter and cooler in summer than the weather farther inland. This is known as the lake effect. The lake functions as a giant heat sink.
To find the energy required to change the temperature of a sample use ∆H = (m)(∆T)(Cp)
Stoichiometry with Energy
Enthalpy values can be included as part of balanced chemical equations. In exothermic reactions, ∆H is negative, but energy is listed as a positive value on the product side of the equation. In endothermic reactions, ∆H is positive and energy is required and is listed on the reactant side of the equation.
The number part of the energy value (as opposed to the unit) can act as a coefficient of a mole ratio, just as any other coefficient would. This way you can convert energy information to information about any substance in a balanced equation and vice versa.
2H-2 + O-2 → 2 H-2O + 561.6 kJ
Enthalpy changes of different types have different names.
The enthalpy change when something dissolves is heat of solution (∆H sol) The enthalpy change during a chemical reaction is heat of reaction (∆Hrxn). The enthalpy change during a comustion reaction is heat of combustion (∆Hcomb). The enthalpy change during a reaction in which a compound is formed from its elements is heat of formation (∆Hf).
Heat of Formation is an important concept.
Heat of Formation Problems.
The enthalpy change can be calculated by taking the total enthalpy of the products - the total enthalpy of the reactants. The formula is...
DHrxn = (the sum of ∆Hf products ) - (the sum of the ∆Hf reactants )
To do this, multiply the number of moles, or coefficient from the balanced equation, of each substance by its heat of formation), and add them up for the products, then do the same for the reactants. Then subtract. Heats of formation are given in kJ / mol, but ∆H is in kJ, since the moles cancel out.
Hess's Law
Hess's Law states that the enthalpy change for a reaction that occurs in many steps is the same as if it occurred in one step. Another way to put this is if several reactions add up to some total reaction, then their enthalpy changes will add up to the enthalpy change for the total reaction.
DHtotal = ∆Hrxn 1 + ∆Hrxn 2 + ∆Hrxn 3 + etc.
Hess's law problems usually give you two or three reactions with their enthalpy change information, then ask you to find the enthalpy change for some target reaction. You must figure out how to make the given reactions add up to the target. This can mean reversing the reactions (and reversing the sign on the enthalpy change), or using them multiple times, or both.
Entropy
Entropy (S) is a measure of the amount of disorder in a substance Gases with their rapid random motion are high in entropy, and solids with their ordered crystalline lattice are low in entropy.
The change in entropy (∆S) is determined just like a heat of formation problem, only use entropy values instead.
∆Srxn = (the sum of ∆Sproducts ) - (the sum of the ∆Sreactants )
Note that the units for entropy are given in J / mol K. ∆S is in J / K, since the moles cancel out.
K stands for Kelvins, the temperature unit on the absolute scale. Also called the Kelvin scale, it is named for Lord Kelvin, who developed it, so the units should be capitalized.
Entropy is temperature affected. It is large at high temperatures, and small at low temperatures. Enthalpy is not temperature affected.
How Enthalpy and Entropy Drive Change
If a reaction will occur spontaneously, it will occur so the products are said to be favored. If a reaction will not occur spontaneously, then the reverse reaction will(in a reversible reaction), so the reactants are said to be favored.
Exothermic changes, those with negative enthalpy changes, are favored to occur spontaneously. So if ∆H is negative, a reaction is more likely than if it is not.
On the other hand, changes that involve an increase in disorder, or entropy, are favored to occur spontaneously. So if ∆S is positive, a reaction is more likely than if it is not.
If ∆H is negative and ∆S is positive, a reaction will certainly occur, no matter what the temperature. If ∆H is positive and ∆S is negative, there will not be a reaction at any temperature, since both indicators say it wont.
The trouble is that often the two indicators disagree. When they do, the enthalpy tends to win out at low temperatures, and the entropy (since it is temperature affected) tends to win out at high temperatures.
For example, if the enthalpy and entropy change values are both negative, the enthalpy indicates the reaction will occur, and it will at low temperatures. The entropy indicates that there will be no reaction, and at high temperatures there wont. The reverse is true for positive enthalpy and entropy changes. These reactions are more likely to occur at high, but not at low temperatures.
Gibbs Free Energy (∆G) determines for sure whether a reaction will be favored to occur. It is simply a formula that compares ∆H to ∆S in a special way.
∆G = ∆H - T∆S
Temperature must be in Kelvins. If ∆G has a negative value, the reaction will occur spontaneously. If ∆G has a positive value, it will not occur
One complication in calculating the free energy change is that the enthalpy values are typically given in kilojoules (kJ), while entropy values are given in J / K, so you must convert so that both use Joules, or both use kilojoules. It doesn't matter which.
Sometimes you may be asked to find the temperature above which a reaction will or wont occur. This is the temperature at which ∆G is between negative and positive, or when it equals zero.
∆G = 0 so 0 = ∆H - T∆S
Rearrange the equation to solve for T, and you will find that...
T = ∆H/∆S
Above that temperature entropy change determines whether a reaction occurs, and below that temperature, enthalpy change determines whether a reaction occurs.
∆G = 0 means there will be equilibrium in a reversible reaction.
-----------------------------
JEE Question 2007 paper II
For the process (water becoming steam) H-2O(l)(1 bar, 373 K) --> H-2O(g)(1 bar, 373 K), the correct set of thermodynamic parameters is
(A) ∆G=0,∆S= + ve
(B)∆G=0,∆S= -ve
(C)∆G=+ve,∆S=0
(D)∆G=-ve, ∆S= + ve
Solution: A
The answer is A because, because at 100 degree C, the steam and water mixture is at equilibrium. Hence ΔG = 0(G = H - TS), and ΔS is positive. Why? when liquid becomes gas, there is more disorder. More entropy.
--------------
Energetics:
First law of thermodynamics;
Internal energy, work and heat,
pressure-volume work;
Enthalpy,
Hess's law;
Heat of reaction, fusion and vapourization;
Second law of thermodynamics;
Entropy;
Free energy;
Criterion of spontaneity.
------------------
Main Topics in TMH Book Chapter
FIRST LAW THERMODYNAMICS
INTERNAL ENERGY AND ENTHALPY
ENTHALPY CHANGE OF A CHEMICAL EQUATION
MOLAR ENHTALPIES OF FORMATIONS
HESS'S LAW OF CONSTANT HEAT SUMMATION
TYPES OF REACTIONS
RELATION BETWEEN DELTA H AND DELTA U OF A CHEMICAL EQUATION
--------------------------
Enthalpy Change, ∆H, With Temperature and State Change
Enthalpy, represented by the symbol H, is essentially a chemistry term for heat, and a term for total kinetic energy of particle motion in a sample. If the reaction is exothermic, the energy contained in the substances is reduced so ∆H has a negative value. On the other hand, if the reaction is endothermic, the substances absorb energy and ∆H is positive.
Temperature is a measure of the average kinetic energy of the particles in a sample.
The faster the molecules in a sample of water the higher the temperature.
Changing temperature of a sample requires a change in enthalpy
It makes sense that you have to add heat to increase temperature, and as a sample cools it gives off heat. The amount of heat required to change the temperature of a sample depends on two factors:
# Mass. The bigger the sample the more heat needed to change its temperature.
# Specific heat capacity of the substance it is made of. This is the energy needed to change the temperature of a 1 gram sample 1 Celsius and it is different for different materials. Its symbol is Cp.
When you go to the beach on a hot summer day, you may have noticed that dry sand can get really hot, while wet sand stays cooler, though the air over them is obviously the same temperature and they are getting the same sunlight. How about dry pavement versus wet pavement? Metal versus tile?
The difference is the specific heat capacity of the different materials. Substances with low specific heat capacities will get very hot very fast, while those with large specific heat capacities warm up much slower since more energy is required to change their temperature. In fact, liquid water has about the highest specific heat of any common substance, so it tends to "heat up" and 'cool down' very slowly compared to most other substances. Ice and steam have lower heat capacities than liquid water.
Some people who prize their outdoor plants set gallon jugs of water around their plants. How does this help the plants? During the heat of the day, that water absorbs energy from the air around it, leaving the air cooler. During the cold night, that water releases that energy as it slowly cools, warming the air around it. This is known as a heat sink.
To find the energy required to change the temperature of a sample use Changing temperature of a sample requires a change in enthalpy
It makes sense that you have to add heat to increase temperature, and as a sample cools it gives off heat. The amount of heat required to change the temperature of a sample depends on two factors:
? Mass. The bigger the sample the more heat needed to change its temperature.
? Specific heat capacity of the substance it is made of. This is the energy needed to change the temperature of a 1 gram sample 1? Celsius and it is different for different materials. For some reason its symbol is Cp. Sorry.
When you go to the beach on a hot summer day, you may have noticed that dry sand can get really hot, while wet sand stays cooler, though the air over them is obviously the same temperature and they are getting the same sunlight. How about dry pavement versus wet pavement? Metal versus tile?
The difference is the specific heat capacity of the different materials. Substances with low specific heat capacities will get very hot very fast, while those with large specific heat capacities warm up much slower since more energy is required to change their temperature. In fact, liquid water has about the highest specific heat of any common substance, so it tends to "heat up" and 'cool down' very slowly compared to most other substances. Ice and steam have lower heat capacities than liquid water.
Some people who prize their outdoor plants set gallon jugs of water around their plants. How does this help the plants? During the heat of the day, that water absorbs energy from the air around it, leaving the air cooler. During the cold night, that water releases that energy as it slowly cools, warming the air around it. This is known as a heat sink.
In Chicago, which sits on the shores of Lake Michigan, the weather near the lake is warmer in winter and cooler in summer than the weather farther inland. This is known as the lake effect. The lake functions as a giant heat sink.
To find the energy required to change the temperature of a sample use ∆H = (m)(∆T)(Cp)
Stoichiometry with Energy
Enthalpy values can be included as part of balanced chemical equations. In exothermic reactions, ∆H is negative, but energy is listed as a positive value on the product side of the equation. In endothermic reactions, ∆H is positive and energy is required and is listed on the reactant side of the equation.
The number part of the energy value (as opposed to the unit) can act as a coefficient of a mole ratio, just as any other coefficient would. This way you can convert energy information to information about any substance in a balanced equation and vice versa.
2H-2 + O-2 → 2 H-2O + 561.6 kJ
Enthalpy changes of different types have different names.
The enthalpy change when something dissolves is heat of solution (∆H sol) The enthalpy change during a chemical reaction is heat of reaction (∆Hrxn). The enthalpy change during a comustion reaction is heat of combustion (∆Hcomb). The enthalpy change during a reaction in which a compound is formed from its elements is heat of formation (∆Hf).
Heat of Formation is an important concept.
Heat of Formation Problems.
The enthalpy change can be calculated by taking the total enthalpy of the products - the total enthalpy of the reactants. The formula is...
DHrxn = (the sum of ∆Hf products ) - (the sum of the ∆Hf reactants )
To do this, multiply the number of moles, or coefficient from the balanced equation, of each substance by its heat of formation), and add them up for the products, then do the same for the reactants. Then subtract. Heats of formation are given in kJ / mol, but ∆H is in kJ, since the moles cancel out.
Hess's Law
Hess's Law states that the enthalpy change for a reaction that occurs in many steps is the same as if it occurred in one step. Another way to put this is if several reactions add up to some total reaction, then their enthalpy changes will add up to the enthalpy change for the total reaction.
DHtotal = ∆Hrxn 1 + ∆Hrxn 2 + ∆Hrxn 3 + etc.
Hess's law problems usually give you two or three reactions with their enthalpy change information, then ask you to find the enthalpy change for some target reaction. You must figure out how to make the given reactions add up to the target. This can mean reversing the reactions (and reversing the sign on the enthalpy change), or using them multiple times, or both.
Entropy
Entropy (S) is a measure of the amount of disorder in a substance Gases with their rapid random motion are high in entropy, and solids with their ordered crystalline lattice are low in entropy.
The change in entropy (∆S) is determined just like a heat of formation problem, only use entropy values instead.
∆Srxn = (the sum of ∆Sproducts ) - (the sum of the ∆Sreactants )
Note that the units for entropy are given in J / mol K. ∆S is in J / K, since the moles cancel out.
K stands for Kelvins, the temperature unit on the absolute scale. Also called the Kelvin scale, it is named for Lord Kelvin, who developed it, so the units should be capitalized.
Entropy is temperature affected. It is large at high temperatures, and small at low temperatures. Enthalpy is not temperature affected.
How Enthalpy and Entropy Drive Change
If a reaction will occur spontaneously, it will occur so the products are said to be favored. If a reaction will not occur spontaneously, then the reverse reaction will(in a reversible reaction), so the reactants are said to be favored.
Exothermic changes, those with negative enthalpy changes, are favored to occur spontaneously. So if ∆H is negative, a reaction is more likely than if it is not.
On the other hand, changes that involve an increase in disorder, or entropy, are favored to occur spontaneously. So if ∆S is positive, a reaction is more likely than if it is not.
If ∆H is negative and ∆S is positive, a reaction will certainly occur, no matter what the temperature. If ∆H is positive and ∆S is negative, there will not be a reaction at any temperature, since both indicators say it wont.
The trouble is that often the two indicators disagree. When they do, the enthalpy tends to win out at low temperatures, and the entropy (since it is temperature affected) tends to win out at high temperatures.
For example, if the enthalpy and entropy change values are both negative, the enthalpy indicates the reaction will occur, and it will at low temperatures. The entropy indicates that there will be no reaction, and at high temperatures there wont. The reverse is true for positive enthalpy and entropy changes. These reactions are more likely to occur at high, but not at low temperatures.
Gibbs Free Energy (∆G) determines for sure whether a reaction will be favored to occur. It is simply a formula that compares ∆H to ∆S in a special way.
∆G = ∆H - T∆S
Temperature must be in Kelvins. If ∆G has a negative value, the reaction will occur spontaneously. If ∆G has a positive value, it will not occur
One complication in calculating the free energy change is that the enthalpy values are typically given in kilojoules (kJ), while entropy values are given in J / K, so you must convert so that both use Joules, or both use kilojoules. It doesn't matter which.
Sometimes you may be asked to find the temperature above which a reaction will or wont occur. This is the temperature at which ∆G is between negative and positive, or when it equals zero.
∆G = 0 so 0 = ∆H - T∆S
Rearrange the equation to solve for T, and you will find that...
T = ∆H/∆S
Above that temperature entropy change determines whether a reaction occurs, and below that temperature, enthalpy change determines whether a reaction occurs.
∆G = 0 means there will be equilibrium in a reversible reaction.
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JEE Question 2007 paper II
For the process (water becoming steam) H-2O(l)(1 bar, 373 K) --> H-2O(g)(1 bar, 373 K), the correct set of thermodynamic parameters is
(A) ∆G=0,∆S= + ve
(B)∆G=0,∆S= -ve
(C)∆G=+ve,∆S=0
(D)∆G=-ve, ∆S= + ve
Solution: A
The answer is A because, because at 100 degree C, the steam and water mixture is at equilibrium. Hence ΔG = 0(G = H - TS), and ΔS is positive. Why? when liquid becomes gas, there is more disorder. More entropy.
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