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;
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).
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Showing posts with label Bonding. Show all posts
Showing posts with label Bonding. Show all posts
Sunday, December 28, 2008
Monday, February 4, 2008
IIT JEE Revision Ch.5. BONDING AND MOLECULAR STRUCTURE -Core Points
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).
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Covalent bonds are formed as a result of the sharing of one or more pairs of bonding electrons. Each atom donates half of the electrons to be shared.
This sharing of electrons is as a result of the electronegativity(electron attracting ability) of the two bonded atoms are either equal or the difference is no greater than 1.7.
Atomic Orbital Approach to Bonding: The basic premise of this theory is that bonds are formed when atoms get close enough so that atomic orbitals on the individual atoms will be able to overlap so that the three dimensional probability regions share a common volume.
hybridisation
In terms of chemistry we refer to the hybridisation of atomic orbitals to explain the change that seems to happen between the atomic orbitals in an uncombined atom and the orbitals used by the same atom when bonding.
The classic molecule to consider is methane CH4. In this molecule the bond angles indicate that the shape of the molecule is a perfect tetrahedron with bond angles of 109º 28' (approximately 109.5º)
It seems that the orbitals used for bonding in CH4 are arranged as far apart as possible suggesting that they have the same energy (degenerate). We know that the orbitals on the carbon atom do not have the same energy - the 2s orbital is of lower energy that the three 2p orbitals. Hybridisation is a model that allows us to combine the atomic orbitals and then produce four degenerate orbitals to be used for bonding.
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.
Polarity
Bonds in which the electron density is symmertically distributed between the nuclei are called non-polar bonds, while those in which the electron density is unsymmetrically distributed are called polar bonds.
VSEPR
The bonding pairs and lone pairs around any particular atom in a molecule adopt positions in which their mutual interactions are minimized. The logic here is simple. Electron pairs are negatively charged and will get as far apart from each other as possible.
Shapes of Molecules
linear - 2 electron pairs
triangular,- 3 electron pairs
tetrahedral - 4 electron pairs
trigonal bipyramidal 5 bond pairs,
octahedral – 6 bond pairs.
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).
------------
Covalent bonds are formed as a result of the sharing of one or more pairs of bonding electrons. Each atom donates half of the electrons to be shared.
This sharing of electrons is as a result of the electronegativity(electron attracting ability) of the two bonded atoms are either equal or the difference is no greater than 1.7.
Atomic Orbital Approach to Bonding: The basic premise of this theory is that bonds are formed when atoms get close enough so that atomic orbitals on the individual atoms will be able to overlap so that the three dimensional probability regions share a common volume.
hybridisation
In terms of chemistry we refer to the hybridisation of atomic orbitals to explain the change that seems to happen between the atomic orbitals in an uncombined atom and the orbitals used by the same atom when bonding.
The classic molecule to consider is methane CH4. In this molecule the bond angles indicate that the shape of the molecule is a perfect tetrahedron with bond angles of 109º 28' (approximately 109.5º)
It seems that the orbitals used for bonding in CH4 are arranged as far apart as possible suggesting that they have the same energy (degenerate). We know that the orbitals on the carbon atom do not have the same energy - the 2s orbital is of lower energy that the three 2p orbitals. Hybridisation is a model that allows us to combine the atomic orbitals and then produce four degenerate orbitals to be used for bonding.
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.
Polarity
Bonds in which the electron density is symmertically distributed between the nuclei are called non-polar bonds, while those in which the electron density is unsymmetrically distributed are called polar bonds.
VSEPR
The bonding pairs and lone pairs around any particular atom in a molecule adopt positions in which their mutual interactions are minimized. The logic here is simple. Electron pairs are negatively charged and will get as far apart from each other as possible.
Shapes of Molecules
linear - 2 electron pairs
triangular,- 3 electron pairs
tetrahedral - 4 electron pairs
trigonal bipyramidal 5 bond pairs,
octahedral – 6 bond pairs.
JEE Revision - Covalent Bond
The Nature of the Co-Valent Bond
Covalent bonds are formed as a result of the sharing of one or more pairs of bonding electrons. Each atom donates half of the electrons to be shared.
This sharing of electrons is as a result of the electronegativity(electron attracting ability) of the two bonded atoms are either equal or the difference is no greater than 1.7.
If the electronegativity difference is greater than 1.7 then the higher electronegative atom has an electron attracting ability large enough to force the transfer of electrons from the less electronegative atom. This would be an ionic bond.
As long as the electronegativity difference is no greater than 1.7 the atoms can only share the bonding electrons.
Single Co-valent Bond
A single co-valent bond would be the sharing of two electrons between the two bonded atoms. Examples are:
H-H
H-Cl
F-F
Double Co-valent Bond
A double co-valent bond is two pairs of electrons being shared. Examples are:
O=O
C=O
C=C
Triple Co-valent Bond
A triple co-valent bond is the sharing of three pairs of electrons. Examples are:
triple bond between two Nitrogen atoms
Triple bond between two carbon atoms
Triple bond between a Carbon and a Nitrogen
The Polarity of the Co-valent Bond
Two atoms with the same electronegativity will share the bonding electron pairs equally. As a result the bonding electrons will be evenly distributed between the bonded atoms. There will be no accumulation of bonding electrons on any one atom and the bond dipole moment will be zero.
Such a co-valent bond will be called a "non-polar" bond.
The bond between two Hydrogens as in H2 or two Oxygens as in O2 or two Nitrogens like N2 will all be non-polar bonds.
On the other hand, if the two bonded atoms have a different electronegativity then the bonding pairs of electrons will be shared unequally. The atom with the higher electronegativity will attract the bonding electrons closer to itself. As a result the electron distribution will be unequal and a bond dipole moment will be formed.
For example the single bond between a Hydrogen and a Chlorine as in H-Cl will have the bonding pair closer to the higher electronegative atom (Chlorine). As a result the Chlorine end will be partially negative since the electrons are closer to the Chlorine. The Hydrogen end will be partially positive since the bonding pair is farther from the Hydrogen.
This two pole condition is called a dipole and it generates a dipole moment that is a vector force directed toward the higher electronegative atom in the bond. Such a bond is referred to as a "polar bond".
The greater the difference in the electronegativity between the two bonded atoms the more polar the bond.
Covalent bonds are formed as a result of the sharing of one or more pairs of bonding electrons. Each atom donates half of the electrons to be shared.
This sharing of electrons is as a result of the electronegativity(electron attracting ability) of the two bonded atoms are either equal or the difference is no greater than 1.7.
If the electronegativity difference is greater than 1.7 then the higher electronegative atom has an electron attracting ability large enough to force the transfer of electrons from the less electronegative atom. This would be an ionic bond.
As long as the electronegativity difference is no greater than 1.7 the atoms can only share the bonding electrons.
Single Co-valent Bond
A single co-valent bond would be the sharing of two electrons between the two bonded atoms. Examples are:
H-H
H-Cl
F-F
Double Co-valent Bond
A double co-valent bond is two pairs of electrons being shared. Examples are:
O=O
C=O
C=C
Triple Co-valent Bond
A triple co-valent bond is the sharing of three pairs of electrons. Examples are:
triple bond between two Nitrogen atoms
Triple bond between two carbon atoms
Triple bond between a Carbon and a Nitrogen
The Polarity of the Co-valent Bond
Two atoms with the same electronegativity will share the bonding electron pairs equally. As a result the bonding electrons will be evenly distributed between the bonded atoms. There will be no accumulation of bonding electrons on any one atom and the bond dipole moment will be zero.
Such a co-valent bond will be called a "non-polar" bond.
The bond between two Hydrogens as in H2 or two Oxygens as in O2 or two Nitrogens like N2 will all be non-polar bonds.
On the other hand, if the two bonded atoms have a different electronegativity then the bonding pairs of electrons will be shared unequally. The atom with the higher electronegativity will attract the bonding electrons closer to itself. As a result the electron distribution will be unequal and a bond dipole moment will be formed.
For example the single bond between a Hydrogen and a Chlorine as in H-Cl will have the bonding pair closer to the higher electronegative atom (Chlorine). As a result the Chlorine end will be partially negative since the electrons are closer to the Chlorine. The Hydrogen end will be partially positive since the bonding pair is farther from the Hydrogen.
This two pole condition is called a dipole and it generates a dipole moment that is a vector force directed toward the higher electronegative atom in the bond. Such a bond is referred to as a "polar bond".
The greater the difference in the electronegativity between the two bonded atoms the more polar the bond.
JEE Revision - Orbital Overlap
Atomic Orbital Approach To Bonding
This second approach is also called the Atomic Orbital Approach to Bonding. The basic premise of this theory is that bonds are formed when atoms get close enough so that atomic orbitals on the individual atoms will be able to overlap so that the three dimensional probability regions share a common volume. This effectively increases the probability of finding bonding electrons between the two atoms. It also effectively results in the lowering of the energy state of the molecular system making the molecule more stable as a result of the overlap. The greater the overlap, the greater is the strength of the bond.
Pure Atomic Orbital Overlap
The simplest of these bonds involve the overlap of two "s" orbitals as in the example when two Hydrogen atoms get close enough to bond. The "s" orbitals overlap to form a "sigma" bond between the two "s" orbitals.
A second type of overlap is between two "p" orbitals to form a sigma bond between two "p" orbitals. An example is the p-p overlap between two Chlorine atoms.
A third type of sigma overlap is the overlap between an "s" orbital and a "p" orbital such as when a Hydrogen atom's "s" orbital overlaps with a "p" orbital of another atom like a Chlorine atom.
This second approach is also called the Atomic Orbital Approach to Bonding. The basic premise of this theory is that bonds are formed when atoms get close enough so that atomic orbitals on the individual atoms will be able to overlap so that the three dimensional probability regions share a common volume. This effectively increases the probability of finding bonding electrons between the two atoms. It also effectively results in the lowering of the energy state of the molecular system making the molecule more stable as a result of the overlap. The greater the overlap, the greater is the strength of the bond.
Pure Atomic Orbital Overlap
The simplest of these bonds involve the overlap of two "s" orbitals as in the example when two Hydrogen atoms get close enough to bond. The "s" orbitals overlap to form a "sigma" bond between the two "s" orbitals.
A second type of overlap is between two "p" orbitals to form a sigma bond between two "p" orbitals. An example is the p-p overlap between two Chlorine atoms.
A third type of sigma overlap is the overlap between an "s" orbital and a "p" orbital such as when a Hydrogen atom's "s" orbital overlaps with a "p" orbital of another atom like a Chlorine atom.
JEE Revision - Hybridisation
Hybridisation involving s, p and d orbitals only
Hybridisation
Hybridisation means making something new from an amalgamation or combination of other parts. A hybrid plant is one made from two different plants blended together. The hybrid shows the characteristics of both plants.
In terms of chemistry we refer to the hybridisation of atomic orbitals to explain the change that seems to happen between the atomic orbitals in an uncombined atom and the orbitals used by the same atom when bonding.
The 's' orbital is spherical about the nucleus and the 'p' orbitals are like double headed balloons arranged along the axis of (imaginary) three dimensional coordinates.
However, it is apparent that the shapes of these orbitals are inadequate to explain the orientation of the bonds produced in molecules. The 'p' orbitals are oriented at 90º to one another and yet there are few molecules that show a bond angle of 90º (in fact the bond angle 90º does appear in some of the larger moolecules but that is due to different reasons).
The classic molecule to consider is methane CH4. In this molecule the bond angles indicate that the shape of the molecule is a perfect tetrahedron with bond angles of 109º 28' (approximately 109.5º)
It seems that the orbitals used for bonding are arranged as far apart as possible suggesting that they have the same energy (degenerate). We know that the orbitals on the carbon atom do not have the same energy - the 2s orbital is of lower energy that the three 2p orbitals. Hybridisation is a model that allows us to combine the atomic orbitals and then produce four degenerate orbitals to be used for bonding.
In order for the electrons to be ready for this process one of them must be promoted from the 2s orbital to the 2pz orbital.
It should be emphasised at this point that this is the norm rather than the exception. It seems that all elements undergo this hybridisation process (or a similar one) when bonding. Logically, the 2s orbital is in no position to overlap directly with another orbital from another atom without interfering with the p orbitals. This hybridisation process allows the 2s electrons to be involved in bonding.
If we study the shape of the water molecule we find that the four electron pairs around the oxygen are tetrahedrally arranged. They have hybridised and the sp3 orbitals so formed overlap directly with the 1s orbitals of the two hydrogens. The two lone pairs on the oxygen remain in the sp3 orbitals that are not used in bonding.
Other types of hybridisation
Carbon can also bond to three other atoms instead of four (as in methane) and it seems that it hybridised its orbitals using only the 2s and two of the 2p orbitals to do this.
There is the formation of a double bond in molecules such as ethene. The three sp2 hybrid orbitals are degenerate (same energy) and consequently arrange as far apart as possible in space i.e. at 120º to each other. This creates a trigonal shape that is planar leaving the remaining 2pz orbital to orientate itself above and below the plane of the other orbitals. This 2p orbital can then laterally overlap with adjacent singly occupied 'p' orbitals on adjacent atoms.
In sp hybridisation, carbon bonds to two other atoms by hybridising the 2s and only one of the 2p orbitals to produce two sp orbitals arranged at 180º to one another. The remaining two 2p prbitals can overlap with suitable orbitals on adjacent atoms to produce pi systems. Examples include ethyne, the nitrogen molecule, hydrogen cyanide, and any other triple bond systems.
Other forms of hybridisation
It should be mentioned that this hybridisation process can be extended to allow atoms to bond with more than four other atoms (octet expansion). In this case the hybridisation invariably involves one or more of the 'd' orbitals. Sulphur hexafluoride forms six attacments to the six fluorines and consequently needs six available orbitals. It gets these by promoting one electron from the 3s and 3px orbitals into two of the 3d orbitals. It can then hybridise the 3s, 3px, 3py, 3pz, and 3dxy, 3dxz orbitals into an octagonal arrangement each with one electron.
Summary
Atoms rearrange their atomic orbitals when bonding to produce orbitals with shapes more suitable for the bonding process. This is called hybridisation.
It is performed by almost all atoms when bonding although carbon provided the easiest examples to show.
It is easy to recognise the hybridisation used by simply observing the double or triple bonds.
Only single bonds = sp3 hybridisation
1 double bond = sp2 hybridisation
1 triple bond =sp hybridisation
Hybridisation
Hybridisation means making something new from an amalgamation or combination of other parts. A hybrid plant is one made from two different plants blended together. The hybrid shows the characteristics of both plants.
In terms of chemistry we refer to the hybridisation of atomic orbitals to explain the change that seems to happen between the atomic orbitals in an uncombined atom and the orbitals used by the same atom when bonding.
The 's' orbital is spherical about the nucleus and the 'p' orbitals are like double headed balloons arranged along the axis of (imaginary) three dimensional coordinates.
However, it is apparent that the shapes of these orbitals are inadequate to explain the orientation of the bonds produced in molecules. The 'p' orbitals are oriented at 90º to one another and yet there are few molecules that show a bond angle of 90º (in fact the bond angle 90º does appear in some of the larger moolecules but that is due to different reasons).
The classic molecule to consider is methane CH4. In this molecule the bond angles indicate that the shape of the molecule is a perfect tetrahedron with bond angles of 109º 28' (approximately 109.5º)
It seems that the orbitals used for bonding are arranged as far apart as possible suggesting that they have the same energy (degenerate). We know that the orbitals on the carbon atom do not have the same energy - the 2s orbital is of lower energy that the three 2p orbitals. Hybridisation is a model that allows us to combine the atomic orbitals and then produce four degenerate orbitals to be used for bonding.
In order for the electrons to be ready for this process one of them must be promoted from the 2s orbital to the 2pz orbital.
It should be emphasised at this point that this is the norm rather than the exception. It seems that all elements undergo this hybridisation process (or a similar one) when bonding. Logically, the 2s orbital is in no position to overlap directly with another orbital from another atom without interfering with the p orbitals. This hybridisation process allows the 2s electrons to be involved in bonding.
If we study the shape of the water molecule we find that the four electron pairs around the oxygen are tetrahedrally arranged. They have hybridised and the sp3 orbitals so formed overlap directly with the 1s orbitals of the two hydrogens. The two lone pairs on the oxygen remain in the sp3 orbitals that are not used in bonding.
Other types of hybridisation
Carbon can also bond to three other atoms instead of four (as in methane) and it seems that it hybridised its orbitals using only the 2s and two of the 2p orbitals to do this.
There is the formation of a double bond in molecules such as ethene. The three sp2 hybrid orbitals are degenerate (same energy) and consequently arrange as far apart as possible in space i.e. at 120º to each other. This creates a trigonal shape that is planar leaving the remaining 2pz orbital to orientate itself above and below the plane of the other orbitals. This 2p orbital can then laterally overlap with adjacent singly occupied 'p' orbitals on adjacent atoms.
In sp hybridisation, carbon bonds to two other atoms by hybridising the 2s and only one of the 2p orbitals to produce two sp orbitals arranged at 180º to one another. The remaining two 2p prbitals can overlap with suitable orbitals on adjacent atoms to produce pi systems. Examples include ethyne, the nitrogen molecule, hydrogen cyanide, and any other triple bond systems.
Other forms of hybridisation
It should be mentioned that this hybridisation process can be extended to allow atoms to bond with more than four other atoms (octet expansion). In this case the hybridisation invariably involves one or more of the 'd' orbitals. Sulphur hexafluoride forms six attacments to the six fluorines and consequently needs six available orbitals. It gets these by promoting one electron from the 3s and 3px orbitals into two of the 3d orbitals. It can then hybridise the 3s, 3px, 3py, 3pz, and 3dxy, 3dxz orbitals into an octagonal arrangement each with one electron.
Summary
Atoms rearrange their atomic orbitals when bonding to produce orbitals with shapes more suitable for the bonding process. This is called hybridisation.
It is performed by almost all atoms when bonding although carbon provided the easiest examples to show.
It is easy to recognise the hybridisation used by simply observing the double or triple bonds.
Only single bonds = sp3 hybridisation
1 double bond = sp2 hybridisation
1 triple bond =sp hybridisation
JEE Revision - Hydrogen Bond
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
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
JEE Revision - Polarity in Molecules
Good tutorial
http://www.ausetute.com.au/molpolar.html
Polarity
Covalent bonds are formed when two atoms share one or more pairs of electrons.
We now want to examine the nature of that sharing in greater detail.
First, we will consider the question of whether two atoms share a pair of electrons equally, i.e. is the electron density in the inter-nuclear region symmetrically distributed or not?
Bonds in which the electron density is symmertically distributed between the nuclei are called non-polar bonds, while those in which the electron density is unsymmetrically distributed are called polar bonds.
Non-Polar Bonds
Any bond between two identical atoms is non-polar since the electronegativities of the two atoms is identical. The simplest examples are the diatomic molecules such as H2, N2, and F2. The C-C bond in ethane, H3C-CH3, is also non-polar.
Polar Bonds
Any bond between two non-identical atoms is polar.
The bond in HF is polar.
So are the C-H bonds in CH4.
Both C-C bonds in propane CH3CH2CH3 are polar.
This is because the terminal carbon atoms and the central carbon are not identical. The terminal carbon atoms are both bonded to three hydrogen atoms and the central carbon atom. The central carbon, however, is bonded to two hydrogen atoms and two carbon atoms.
An alternative way to think about this is to identify the groups that are attached to each carbon in propane. The left-hand carbon is attached to three hydrogen atoms and a CH2CH3 (ethyl) group. The central carbon is attached to two hydrogen atoms and two CH3 (methyl) groups. Since the right-hand carbon is attached to three hydrogen atoms and an ethyl group, it is identical to the left-hand carbon. Looking at propane in this way allows us to introduce the idea of group electronegativities.
Group Electronegativities
In the same way that the electronegativity of an atom is a measure of the tendency of that atom to attract electrons, group electronegativity is a measure of the tendency of a polyatomic group to attract electrons.
Propane contains two CH3 groups and one CH2 group (methylene group). Since these two groups are not identical, they have different group electronegativities. Even though the H3C-CH2 bond is a bond between two carbon atoms, the carbons are not identical because they have different atoms attached to them.
As far as the pair of electrons that the two carbons share goes, they experience a different Coulombic attraction from the CH3 group than they do from the CH2 group.
One way to investigate group electronegativities experimentally involves nuclear magnetic resonance (NMR) spectroscopy.
Polar and Non-Polar Molecules
The C-H bonds in methane are polar. However, a molecule of methane is non-polar. Specifically, the dipole moment of methane is zero.
A dipole moment of zero means that the "center of negative charge" in the molecule corresponds to the "center of positive charge". In the case of methane, the "center of positive charge" and the "center of negative charge" are focused on the carbon atom. Think of the "center of charge", whether positive or negative, in the same way that you think of the "center of mass". From that perspective, a molecule with a dipole moment of zero is like a perectly balanced see-saw.
http://www.ausetute.com.au/molpolar.html
Polarity
Covalent bonds are formed when two atoms share one or more pairs of electrons.
We now want to examine the nature of that sharing in greater detail.
First, we will consider the question of whether two atoms share a pair of electrons equally, i.e. is the electron density in the inter-nuclear region symmetrically distributed or not?
Bonds in which the electron density is symmertically distributed between the nuclei are called non-polar bonds, while those in which the electron density is unsymmetrically distributed are called polar bonds.
Non-Polar Bonds
Any bond between two identical atoms is non-polar since the electronegativities of the two atoms is identical. The simplest examples are the diatomic molecules such as H2, N2, and F2. The C-C bond in ethane, H3C-CH3, is also non-polar.
Polar Bonds
Any bond between two non-identical atoms is polar.
The bond in HF is polar.
So are the C-H bonds in CH4.
Both C-C bonds in propane CH3CH2CH3 are polar.
This is because the terminal carbon atoms and the central carbon are not identical. The terminal carbon atoms are both bonded to three hydrogen atoms and the central carbon atom. The central carbon, however, is bonded to two hydrogen atoms and two carbon atoms.
An alternative way to think about this is to identify the groups that are attached to each carbon in propane. The left-hand carbon is attached to three hydrogen atoms and a CH2CH3 (ethyl) group. The central carbon is attached to two hydrogen atoms and two CH3 (methyl) groups. Since the right-hand carbon is attached to three hydrogen atoms and an ethyl group, it is identical to the left-hand carbon. Looking at propane in this way allows us to introduce the idea of group electronegativities.
Group Electronegativities
In the same way that the electronegativity of an atom is a measure of the tendency of that atom to attract electrons, group electronegativity is a measure of the tendency of a polyatomic group to attract electrons.
Propane contains two CH3 groups and one CH2 group (methylene group). Since these two groups are not identical, they have different group electronegativities. Even though the H3C-CH2 bond is a bond between two carbon atoms, the carbons are not identical because they have different atoms attached to them.
As far as the pair of electrons that the two carbons share goes, they experience a different Coulombic attraction from the CH3 group than they do from the CH2 group.
One way to investigate group electronegativities experimentally involves nuclear magnetic resonance (NMR) spectroscopy.
Polar and Non-Polar Molecules
The C-H bonds in methane are polar. However, a molecule of methane is non-polar. Specifically, the dipole moment of methane is zero.
A dipole moment of zero means that the "center of negative charge" in the molecule corresponds to the "center of positive charge". In the case of methane, the "center of positive charge" and the "center of negative charge" are focused on the carbon atom. Think of the "center of charge", whether positive or negative, in the same way that you think of the "center of mass". From that perspective, a molecule with a dipole moment of zero is like a perectly balanced see-saw.
JEE Revision - VSEPR Model
VSEPR: Valence Shell Electron Pair Repulsion.
Assumptions about the nature of the bonding
The underlying assumptions made by the VSEPR method are the following.
Atoms in a molecule are bound together by electron pairs. These are called bonding pairs. More than one set of bonding pairs of electrons may bind any two atoms together (multiple bonding).
Some atoms in a molecule may also possess pairs of electrons not involved in bonding. These are called lone pairs or non-bonded pairs.
The bonding pairs and lone pairs around any particular atom in a molecule adopt positions in which their mutual interactions are minimized. The logic here is simple. Electron pairs are negatively charged and will get as far apart from each other as possible.
Lone pairs occupy more space than bonding electron pairs.
Double bonds occupy more space than single bonds.
Assumptions about the nature of the bonding
The underlying assumptions made by the VSEPR method are the following.
Atoms in a molecule are bound together by electron pairs. These are called bonding pairs. More than one set of bonding pairs of electrons may bind any two atoms together (multiple bonding).
Some atoms in a molecule may also possess pairs of electrons not involved in bonding. These are called lone pairs or non-bonded pairs.
The bonding pairs and lone pairs around any particular atom in a molecule adopt positions in which their mutual interactions are minimized. The logic here is simple. Electron pairs are negatively charged and will get as far apart from each other as possible.
Lone pairs occupy more space than bonding electron pairs.
Double bonds occupy more space than single bonds.
Sunday, February 3, 2008
JEE Revision - Shapes of Molecules
linear - 2 electron pairs
triangular,- 3 electron pairs
tetrahedral - 4 electron pairs
trigonal bipyramidal 5 bond pairs,
octahedral – 6 bond pairs.
trigonal pyramidal - NH3 3 bp and one lp
square planar,
pyramidal,
square pyramidal,
angular,
To predict the shape of a molecule:
(1) Write down the Lewis dot structure for the molecule.
(2) Count the number of bond pairs and lone pairs around the central atom.
(3) Decide on the electron pair orientation based on the total number of electron pairs (4 = tetrahedral, 5 = trigonal bipyramidal).
(4) Consider the placement of lone pairs and any distortions from "regular" shapes.
(5) Name the shape based on the location of atoms (nuclei).
triangular,- 3 electron pairs
tetrahedral - 4 electron pairs
trigonal bipyramidal 5 bond pairs,
octahedral – 6 bond pairs.
trigonal pyramidal - NH3 3 bp and one lp
square planar,
pyramidal,
square pyramidal,
angular,
To predict the shape of a molecule:
(1) Write down the Lewis dot structure for the molecule.
(2) Count the number of bond pairs and lone pairs around the central atom.
(3) Decide on the electron pair orientation based on the total number of electron pairs (4 = tetrahedral, 5 = trigonal bipyramidal).
(4) Consider the placement of lone pairs and any distortions from "regular" shapes.
(5) Name the shape based on the location of atoms (nuclei).
Saturday, January 19, 2008
IIT JEE Ch.5. BONDING AND MOLECULAR STRUCTURE -Core Points
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).
-------------------
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
------------------
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
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.
VSEPR model and shapes of molecules
linear - 2 electron pairs
angular,
triangular,- 3 electron pairs
square planar,
pyramidal,
square pyramidal,
trigonal bipyramidal,
tetrahedral - 4 electron pairs
octahedral).
The relation between number of electron pairs around the central atom and shape of molecule to be filled
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).
-------------------
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
------------------
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
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.
VSEPR model and shapes of molecules
linear - 2 electron pairs
angular,
triangular,- 3 electron pairs
square planar,
pyramidal,
square pyramidal,
trigonal bipyramidal,
tetrahedral - 4 electron pairs
octahedral).
The relation between number of electron pairs around the central atom and shape of molecule to be filled
Sunday, October 21, 2007
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)
----------------------
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