Thursday, February 4, 2016

Chemistry Knowledge History - February





February 1 - Chemistry Knowledge History

Birthdays

Emilio Segrè (1905): codiscovered technetium (Tc, element 43) and astatine (At, 85); spontaneous fission; antiproton; Nobel Prize (Physics), 1959
http://iit-jee-chemistry.blogspot.in/2007/12/chapter-13a-halogens.html  Astatine is a halogen.

Roger Yonchien Tsien (1952): green fluorescent protein, GFP; Nobel Prize, 2008.

February 2 - Chemistry Knowledge History



1923 Leaded gasoline first marketed in the US in Dayton, OH,
1923. Thomas Midgley, Jr., of General Motors Research labs added tetraethyllead to gasoline.

Birthdays

Jean Baptiste Boussingault 1802: agricultural chemistry; isolated and named sorbitol; role of nitrogen in plant nutrition.

Albert Schatz  1920: discovery of antitubercular agent streptomycin. Schatz's version of the discovery differs from the standard account in which Selman Waksman receives near-exclusive credit. The patent (US 2,449,866).

February 3 - Chemistry Knowledge History

Birthdays
Leonora Neuffer Bilger  1893: asymmetric nitrogen compounds; Received Garvan Medal in 1953
https://en.wikipedia.org/wiki/Leonora_Bilger

February 4  - Chemistry Knowledge History

Joseph Goldberger begins the experiment that demonstrates that pellagra is a dietary disease, 1915.


John Jacob Livingood made radium E (210Bi) by bombarding common bismuth with deuterons, 1936, the first synthetis of a radioactive substance in the US.

Birthdays

Friedrich Hund born 1896: Hund's rules for electron configurations, the first of which predicts maximum multiplicity of spin; molecular-orbital theory (Hund-Mulliken approach).
IIT - JEE  http://iit-jee-chemistry.blogspot.in/2015/05/ch1-atomic-structure-and-chemical.html

February 5  - Chemistry Knowledge History

Birthdays - 5 Feb

John Boyd Dunlop (1840); manufactured pneumatic rubber tires.

Lafayette Benedict Mendel  (1872): modern science of nutrition; codiscovered vitamin A and B complex; linked nutritive value of proteins to their amino acids.

February 6  - Chemistry Knowledge History


William Parry Murphy born 1892: diabetes; pernicious anemia and other blood diseases; Nobel Prize (Medicine), 1934
Clemens Winkler, in the course of analyzing a mineral, discovered element (germanium, Ge, element 32) in 1886, consistent with predictions by J. A. R. Newlands and Dmitrii Mendeleev.
Nikolai Dmitrievich Zelinskii born 1861: catalysis of hydrocarbon disproportionations; bromination of fatty acids (Hell-Volhard-Zelinsky reaction)

February 7  - Chemistry Knowledge History


Ulf Svante von Euler born 1905: identification of noradrenaline (norepinephrine) as a neurotransmitter; son of 1929 Nobel laureate biochemist Hans von Euler-Chelpin; Nobel Prize (medicine), 1970.
John Brown Francis Herreshoff born 1850: manufacture of sulfuric acid

February 8

 - Chemistry Knowledge History


Bernard Courtois born 1777: discovered iodine (I, element 53) from seaweed
Moses Gomberg born 1866: work on triphenylmethyl (first stable organic free radical); tautomerism
Robert Holton announces total synthesis of taxol, an important cancer drug, 1994.
Francis Robert Japp  born 1848: benzil, benzoin, and phenanthraquinone.
Dmitrii Mendeleev born 1834 : periodic law and periodic table.
Friedlieb Runge born 1795: discovered carbolic acid (phenol) and aniline in coal tar; dry distillation

February 9

 - Chemistry Knowledge History


Edward Charles Baly born 1871: showed that organic compounds, including sugars, can be formed photochemically from water, carbon dioxide, and ammonia
Californium (Cf, element 98) discovered by  Kenneth Street, Jr., Stanley G. Thompson, Glenn T. Seaborg, and Albert Ghiorso using ion-exchange chromatography at University of California, Berkeley, 1950.
Lloyd Ferguson born 1918: chemical educator
Norman Bruce Hannay born 1921: materials for solid state electronics

February 10

 - Chemistry Knowledge History


Per Teodor Cleve born 1840: discovered holmium and thulium; suggested "didymium" was not elementary; naphthalene derivatives.
John Franklin Enders born 1897: showed polio virus was not only neurotropic; Nobel Prize (Medicine), 1954.
Ira Remsen born 1846: prominent American organic chemist; founder of American Chemical Journal; first professor of chemistry at Johns Hopkins University; saccharin was discovered in his lab

February 11

 - Chemistry Knowledge History

Fred Basolo born 1920: organometallics.
Thomas Alva Edison born 1847: inventor (incandescent light (US 233,898), phonograph (US 200,521, electrical systems, etc.).
Josiah Willard Gibbs born 1839: thermodynamics and the phase rule; the Gibbs free energy is named after him.
Izaak Kolthoff born 1894: analytical chemistry.
Alwin Mittasch  and Christian Schneider filed US patent application for catalytic production of methanol from carbon monoxide and hydrogen (U.S. patent 1,201,850) in 1914.
William Henry Fox Talbot born 1800: photography pioneer.

February 12

 - Chemistry Knowledge History


Pierre-Louis Dulong born 1785: discovered nitrogen trichloride; refractive indices and specific heats of gases; law of Dulong and Petit (specific heat times atomic weight is the same for many elements); suggested that acids were compounds of hydrogen; formula for heat content of fuel (Dulong formula)

Moritz Traube born 1826: physiological chemist; semipermeable membranes, sugars, respiration, fermentation, putrefaction, oxidation, protoplasm, and muscle

February 13

 - Chemistry Knowledge History


Heinrich Caro born 1834: Caro's acid (H2SO5), dye chemistry.
Étienne-François Geoffroy born 1672: chemical affinities; displacement reactions in salt
Henry Clemens Pearson born 1858: rubber scientist and editor; see his books, Crude rubber and compounding ingredients and The rubber country of the Amazon

February 14

 - Chemistry Knowledge History



Herbert Aaron Hauptman born 1917: mathematical methods for crystal structures; Nobel Prize, 1985.
Lawrencium (Lr, element 103) was produced in 1961 by Torbjorn Sikkeland, Albert Ghiorso, and Almon Larsh and Robert Latimer, at University of California, Berkeley.
Julius Nieuwland born 1878: synthetic rubber pioneer (US patent 1,811,959); acetylene chemistry. .
Agnes Pockels born 1862: liquid surfaces: surface tension and films; invention of the slide trough and surface film balance. Read her article on surface tension.
Dennis Searle and E. M. Skillings found borax and other soluble salts near San Bernardino, CA, 1873.

February 15

 - Chemistry Knowledge History


Synthesis of diamond by Francis Bundy, H. Tracy Hall, Herbert Strong, & Robert H. Wentoff, Jr., at General Electric Research Laboratories announced in 1955.

Hans K. A. S. von Euler-Chelpin born 1873: enzymes and fermentation; father of 1970 Nobel laureate Ulf Svante von Euler; Nobel Prize, 1929

George Johnstone Stoney born 1826: suggested that electrical charge came in discrete units; coined term electron for "atom of electricity".

February 16

 - Chemistry Knowledge History


Julius Thomsen born 1826: heats of reaction, relative strength of acids, manufacture of soda from cryolite
John Rex Whinfield born 1901: terephthalic acid polyester fibers (terylene).

Robert Williams born 1886: isolation, synthesis, and manufacture of Vitamin B1 (thiamine).

February 17

 - Chemistry Knowledge History


Friedrich Konrad Beilstein born 1838: his standard reference work on organic chemistry was first published in 1880-83 and has been updated ever since

Wallace Henry Coulter born 1913: instrument maker; developed instrumentation to characterize particles.
Dmitrii Mendeleev sketched his first draft periodic table, 1869.

Otto Stern born 1888: quantization of angular momentum (Stern-Gerlach experiment); Nobel Prize (physics), 1943.


February 18 - Chemistry Knowledge History

Harry Brearley born 1871: development of stainless steel
John Sinfelt born 1931: platinum-iridium catalysts in petroleum refining. Read a book chapter by Sinfelt on materials and catalysis.
Frederick Soddy introduced the term "isotopic" (meaning "same place") for elements which share the same place in the periodic table in 1913.

Alessandro Volta born 1745: invented the voltaic pile, the first electric battery; discovered and isolated methane. The unit of electric potential, the volt, is named in his honor.
(18 February 2015 - Google carries doodle in the honour of Volta)




http://electronics.howstuffworks.com/everyday-tech/battery4.htm

The Voltaic Pile
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Science Online

February 19

 - Chemistry Knowledge History


Svante Arrhenius born 1859: electrolytic dissociation, viscosity, reaction rates, and even the greenhouse effect; Nobel prize, 1903

Louis-Georges Gouy born 1854: interfacial electrical double layer.
Gottlieb Sigismund Kirchhof born 1764: catalytically produced glucose from starch.
Roderick MacKinnon born 1956: structural and mechanistic studies of ion channels; Nobel Prize, 2003
Ernest Marsden born 1889: scattering of alpha particles (work with Hans Geiger in Ernest Rutherford's lab), contributing to the development of the nuclear model of the atom.
One atom of mendelevium (Md, element 101) was produced by Gregory R. Choppin, Glenn Seaborg, Bernard G. Harvey, and Albert Ghiorso in 1955 by bombarding a billion atoms of 253Es with helium.
Ferdinand Reich born 1799: codiscovered indium (In, element 49)

February 20

 - Chemistry Knowledge History


Isaac Adams, Jr. born 1836: pioneer in nickel plating.
Ludwig Boltzmann born 1844: statistical mechanics; thermodynamics, especially the second law; Maxwell-Boltzmann distribution of molecular speeds; Stefan-Boltzmann law of blackbody radiation; Boltzmann constant is named after him

Henry Eyring born 1901: chemical kinetics (transition-state theory, Eyring equation)
Helen Murray Free born 1923: diagnostic chemistry: reagents and instrumentation for clinical diagnosis in blood and urine chemistry, histology, and cytology
Robert Huber born 1937: three-dimensional structure of proteins involved in photosynthesis; Nobel Prize, 1988

February 21

 - Chemistry Knowledge History


Carl Henrik Dam born 1895: vitamin K as a dietary factor in blood clotting; Nobel Prize (medicine), 1943.
Humphry Davy reads paper introducing the name chlorine (to replace oxymuriatic acid) and asserting its elementary nature, 1811.

Oliver Wolcott Gibbs born 1822: early American inorganic and analytical chemist (Harvard); founding member of US National Academy of Sciences
Edwin Land demonstrates Polaroid camera to optical society meeting, 1947.
John Mercer born 1791: treated cotton with caustic soda (mercerized cotton); discovered some calico dyes
Dorothy Virginia Nightingale born 1902: synthetic organic chemistry, particularly reactions of alkylbenzenes in the presence of aluminum chloride; Garvan Medal, 1959

February 22

 - Chemistry Knowledge History


Johannes Nicolaus Brønsted born 1879: acid-base theory and properties of ions; kinetics and catalysis; nitramide

Heinrich Hertz born 1857: discovered electromagnetic waves and the photoelectric effect.
Pierre Jules Cesar Janssen born 1824: astronomical spectroscopy and photography, particularly of the Sun; found a line in the solar spectrum subsequently identified with helium.
Fritz Strassmann born 1902: nuclear fission.
Friedrich Wöhler wrote a letter to J. J. Berzelius stating that he had synthesized urea, an early synthesis of an organic compound from inorganic materials, 1828.

February 23

 - Chemistry Knowledge History


First organizational meeting of the Chemical Society of London, 1841. (The Royal Society of Chemistry is its successor organization.)

Casimir Funk born 1884: discovered vitamins and named them (vitamines)
Charles Martin Hall first produced electrolytic aluminum in 1886 (US patent 400,766).
Thomas Midgley, Jr., received US patent 1,573,846 for tetraethyllead as an anti-knock agent in gasoline, 1926.
Glenn Theodore Seaborg and coworkers chemically identified plutonium (Pu, element 94) at University of California, Berkeley, 1941.

February 24

 - Chemistry Knowledge History


First atom of element 107, eventually named Bohrium (Bh) was observed at GSI Laboratories, Darmstadt, Germany in 1981.

John Gorham born 1783: wrote Elements of Chemical Science, an early American chemistry text.
Karl Graebe born 1841: organic synthesis (alizarin) and nomenclature (ortho, meta, para prefixes).
Eugène Melchior Peligot born 1811: isolated uranium metal; identified glucose in diabetics' urine.
William Summer Johnson born 1913: synthesis of complex molecules

February 25

 - Chemistry Knowledge History


Ruth Erica (Leroi) Benesch born 1925: oxygen-carrying capacity of hemoglobin; role of sulfur in proteins
Arthur Becket Lamb born 1880: editor of the Journal of the American Chemical Society, 1917-1949.
Phoebus Aaron Theodor Levene born (as Fishel Aaronovich Lenin) 1869: biochemistry, hexosamines, and stereochemistry.
Ida Eva Noddack born  1896: co-discoverer of rhenium (Re, element 75) with husband Walter and Otto Berg; suggested (correctly) that nuclear fission rather than transuranic elements explained results reported by Enrico Fermi.
Mary Locke Petermann born 1908: ribosomes and protein synthesis

February 26

 - Chemistry Knowledge History


Marjorie Beckett Caserio born 1929: physical organic chemistry: kinetics and mechanisms; chemical education: Basic Principles of Organic Chemistry
Benoit Paul Emile Clapeyron born 1799: relationship between temperature, volume, and heat of vaporization (Clapeyron and Clausius-Clapeyron equations).
Herbert Henry Dow born 1866: electrolytic production of bromine; founder of Dow Chemical.
Giulio Natta born 1903: polymer chemistry including polymer stereochemistry; Nobel Prize, 1963
William Joseph Sparks born 1905: advances in synthetic rubber.
Ahmed Zewail born 1946: "femtochemistry" (dynamics on a sub-picosecond time scale); Nobel Prize, 1999

February 27

 - Chemistry Knowledge History


James Chadwick's note announcing the possible discovery of the neutron is published in Nature, 1932.

Robert Grubbs born 1942: metathesis reactions and catalysts; Nobel Prize, 2005.
Alice Hamilton born 1869: occupational medicine; hazards of carbon monoxide, mercury, tetraethyllead, benzene, and others; first woman professor at Harvard.
Felix Hoffmann received US patent 644,077 for acetyl salicylic acid (better known as aspirin), 1900.
Karl Friedrich Wenzel died 1793 (birth date unknown c. 1740): stoichiometry; concentration determines the speed of chemical reactions.

February 28

 - Chemistry Knowledge History


Edward Goodrich Acheson received US patent number 492,767 for production of artificial silicon carbide ("Carborundum"), 1893.

Steven Chu: laser cooling and trapping of atoms; US Secretary of Energy; Nobel (physics), 1997.
Edmond Fremy born 1814: plumbates, stannates, and ferrates; preparation of anhydrous hydrogen fluoride; coloring of flowers and saponification of fats

Philip Showalter Hench born 1896: hormones of the adrenal cortex; Nobel Prize (Medicine), 1950

Linus Carl Pauling born 1901: molecular structure, bonding (hybrid orbitals), electronegativity, and resonance (The Nature of the Chemical Bond); Nobel Prize, 1954; Nobel Peace Prize, 1962
February 29

Heike Kamerlingh Onnes announced solidification of helium, 1908.

National Science Day of India


Chemistry History
http://web.lemoyne.edu/~giunta/February.html

Wednesday, February 3, 2016

6. Chemical Kinetics - JEE Main - CBSE XII - Core Revision Points

Importance of  Core Revision Points: Core Revision Points are important because if you remember them strongly, many more points related to them will come out of your memory and help you to answer question and problems. Read them many times and make sure you remember them very strongly.


Sections in the Chapter Jauhar

6.1 Rate of a chemical reaction
6.2 Experimental measurement of reaction rate
6.3 Factors which influence rates of chemical reactions
6.4 Dependence of reaction rates on concentration
6.5 Order of a reaction
6.6 Integrated rate expansion
6.7 Experimental determination of order of a reaction
6.8 Half life period of a reaction
6.9 Collision theory: Energy and orientation barriers to reactions
6.10 Dependence of reaction rates on temperature
6.11 Concept of activation energy and activated complex theory
6.12 Arrhenius equation and calculation of activation energy
6.13 Effect of radiations on reaction rates: photochemical reactions
6.14 Mechanism of a reaction
6.15 Fast reactions

Revision of the Chapter Video - Hindi
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SardanaTutorials

Revision Points


The topic "Chemical kinetics" consists of reaction rate and reaction mechanism.

The branch of chemistry which deals with the rates of chemical reactions and the mechanism by which they occur, is called chemical kinetics.

Reaction rate is the speed with which a reaction takes place. This shows the rate or speed at which the reactants are consumed and products are formed.

Reaction mechanism is the path by which a reaction takes place.


6.1 Rate of a chemical reaction


Reaction rate is the speed with which a reaction takes place. This shows the rate or speed at which the reactants are consumed and products are formed.

Reaction mechanism is the path by which a reaction takes place.


Rate of reaction

The rate of reaction is a quantity that tells how the concentration of reactants or product changes with time.

So this can be expressed as Δ concentration/Δ time. That is change in concenation divided by time taken for the change.

Molar concentration i.e., moles per liter (M), is used in these equations.

The brackets, [ ] are always used to to indicate molar concentrations.


6.2 Experimental measurement of reaction rate


6.3 Factors which influence rates of chemical reactions


Temperature

As temperature increases, the average kinetic energy increases. So there are more molecules with activation energy and hence reaction rate increases.

As a general approximation, the rate roughly doubles for each 10°C rise in temperature.


6.4 Dependence of reaction rates on concentration


Law of Mass Action

In 1867, Cato Guldberg, and Peter Waage, proposed this law. According to this law, for the rate determining step in a reaction, the rate of reaction is proportional to the product of the concentrations of the reactants, each raised to the power of its coefficient in the balanced equation.

For the reaction aA + bB → cC (when it is a rate determining step)

Rate of reaction is proportional to [A]^a[B]^b

The above proportionality can be written as an equation, by putting in a proportionality constant k.

Rate = k *[A]^a[B]^b

K is called the specific rate constant

Rate law

The rate for a reaction is a mathematical expression that relates the rate of reaction to the concentrations of the reactants.

For the reaction aA + bB → products

The rate law is expressed as, rate of reaction is proportional to [A]^x[B]^y.
x and y are determined experimentally. These values can be whole or fractional numbers or zero.

Rate = k[A]^x[B]^y

k = the rate constant.
[A] and [B] are molar concentrations of reactants mol/litre

Units of rate: Rate is the change in concentration with time.

If the concentrations are expressed in moles/litre and time in seconds, then the units for rate of reaction are mol litre-1 s-1 or mol L-1s-1

Units of rate constant

Units of rate constant are different for different orders of reaction.

For zero order reactions units of rate constant are mol L-1s-1

For first order reactions units of rate constant are s-1

For second order reactions units of rate constant are L mol-1s-1

Basically units of rate constant are changing to give the rate of reaction in required units mol L-1s-1 (change in concentration with time).

In case of gases, the concentrations are expressed in terms of pressure in the units of atmosphere. Therefore the rate of reaction has the units of atm per second.


Integrated Rate Expressions

For zero order reactions

k0 = {[A]0 - [A]}/t

Where k0 = rate constant in the case of zero order reactions
[A]0 = Initial concentration of reactant A

[A] = concentration of reactant A at time t.
t = time

This can be alternatively expressed.

a = Initial concentration of reactant A (in moles per litre)
x = moles reactants that changed into products in time t
a-x = concentration of reactant A after time t

k0 = x/t

Where k0 = rate constant in the case of zero order reactions
x = moles reactants that changed into products in time t

For first order reactions

k1 = (2.303/t)log{[A]0/[A]}

Where k1 = rate constant for first order equations

Alternative expression



a = Initial concentration of reactant A (in moles per litre)
x = moles reactants that changed into products in time t
a-x = concentration of reactant A after time t

k1 = (2.303/t)log{a/(a-x)}

6.5 Order of a reaction

Order of a reaction

The sum of the powers to which the concentration terms are raised in the rate law expression.

For the expression Rate = k[A]^x[B]^y, the order of the reaction is x+y. The order of the reaction is represented by n.

When n = 1, the reaction said to be first order reaction.
n = 2 second order reaction etc.

There are number of reactions where rate of reaction is independent of concentration of reactants. The order of reaction is zero.





6.6 Integrated rate expansion
6.7 Experimental determination of order of a reaction
6.8 Half life period of a reaction
6.9 Collision theory: Energy and orientation barriers to reactions
6.10 Dependence of reaction rates on temperature
6.11 Concept of activation energy and activated complex theory

6.12 Arrhenius equation and calculation of activation energy


Reaction Rate Depends on Temperature

Temperature has influence on reaction rates. In general, an increase in temperature increases the rate of almost all reactions.

A general approximate rule is that the rate of a reaction becomes almost double for every 10° rise in temperature.

Activation Energy

For many reactions some extra energy is to be supplied to the reactants to initiate the reaction. This excess energy is required to bring the energy of reactants to the energy that is required to start the reaction. The energy of the reactants at which the reaction starts is called threshold energy.

Activation energy is the extra energy supplied to initiate the reaction. Thus activation energy is equal to the difference between the threshold energy and the average kinetic energy of the reacting molecules at the the given temperature (Note as activation energy is being given the temperature of the reactants increases)

Arrhenius Equation

Arrhenius proposed a quantitative relationship between rate constant and temperature

k = Ae(–Ea/RT)

where k = rate constant
A is a constant known as frequency factor. In a JEE problem it was termed as preexponential factor
–Ea is the activation energy
Both A and –Ea0 are characteristic of the equation
T is the absolute temperature and R is the gas constant

In log form the equation becomes

log k = log A - (Ea)/2.303 RT

As the activation energy –Ea0 increases, the value of k decreases and therefore, the reaction rate decreases.

Find the value of –Ea0

If log k is plotted against 1/T (both found through experiments), the intercept of the line will be equal to - (Ea)/2.303 R. Hence from the slope found from the graph - (Ea) can be found out as -2.303 R multiplied by slope.

Second Method

Measure rate constant at two temperatures k1 and k2 at T1 and T2

log (k2/k1) = (Ea/2.303 R)[ (1/T1) - (1/T2)]

JEE 2009 problem


For a first order reaction A→P, the temperature (T) dependent rate constant(k) was found to follow the equation logk = – (2000)(1/T) + 6.0
The pre-exponential factor A and the activation energy Ea, respectively, are -

Answer:
1.0 × 1066 s-1 and 38.3 kJ mol-1



6.13 Effect of radiations on reaction rates: photochemical reactions
6.14 Mechanism of a reaction
6.15 Fast reactions

Chemical Kinetics - 28 Videos Playlist - Examfearvideos
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Updated 3 Feb 2016, 22 May 2015

5. Electrochemistry - JEE Main - CBSE Class XII - Core Revision Points

Importance of  Core Revision Points: Core Revision Points are important because if you remember them strongly, many more points related to them will come out of your memory and help you to answer question and problems. Read them many times and make sure you remember them very strongly.


Sections in the Chapter


5.1 Electrochemical changes: electrolytic and galvanic cells
5.2 Electrolysis and laws of electrolysis
5.3 Metallic and electrolytic conductance
5.4 Electrolytic conduction
5.5 Factors for the variation of molar conductance
5.6 Kohlrausch’s law
5.7 Electrochemical cell or galvanic cell
5.8 Representation of an electrochemical cell
5.9 Electrode potential and E.M.F. of a galvanic cell
5.10 Standard Electrode Potential
5.11 Electrochemical series
5.12 Differences between Galvanic cell and electrolytic cell
5.13 Dependence of electrode and cell potentials on concentration: Nernst Equations
5.14 Equilibrium constant form Nernst Equation
5.15 Electrochemical cell and free energy
5.16 Some commercial cells
5.17 Electrode potential electrolysis and criteria for product formation
5.18 Corrosion
5.19 Commercial production of chemicals
5.20 Manufacture of some important metals an chemical compounds



Sections in the Chapter


5.1 Electrochemical changes: electrolytic and galvanic cells

5.2 Electrolysis and laws of electrolysis


Faraday's laws of electrolysis:
---------------------------------------
Quantitative Relationships in Electrolytic Cells

Determining the amount of electrical energy necessary for accumulating a given amount material from the electrolytic cell.


First law: It states that the amount of any substance that is liberated at an electrode during electrolysis is directly proportional to the quantity of electricity passed through the electrolyte.

W α Q (w = weight of substance deposited and Q is charge = ampere * time)

Second law: It states tht when the same quantity of electricity is passed through different electrolytes amount of different substances liberated or deposited at the different electrodes are directly proportional to the chemical equivalents9i.e., equivalent weight) of substances.

One faraday (F) is the amount of electrical energy required for flow of 1 mole of electrons.

To three significant digits, 1 faraday equals 96,500 coulombs(coul).

Current flow is measured in amperes (A)which is coulombs/seconds or coul/s,

5.3 Metallic and electrolytic conductance

Electrolytic conductance, specific, equivalent and molar conductance,



Electrolytic conductance


The flow of electric current through an electrolytic solution is known as electrolytic conduction.

Electrolytic conduction also follows Ohm's law.

V = I/R

R = ρ* l/a

ρ is called specific resistance.
The reciprocal of specific resistance is termed specific conductance. It may be defined as the conductance of a solution of 1 cm length and having 1 sq.cm as the area of cross section.

Specific conductance is the conductance of one centimetre cube of a solution of an electrolyte. It is denoted by k (kappa)

κ = 1/ρ

The equivalent conductivity of an electrolyte may be defined as the conductance of a volume of solution containing one equivalent mass of a dissolved substance when placed between two parallel electrodes which are at a unit distance apart, and large enough to contain between them the whole solution.

The molar conductivity of a solution gives the conducting power of ions produced by one molar mass of an electrolyte at any particular concentration.

It is denoted by Λm (Lambda).

Λm = κ/M

where M is the molar concentration





5.4 Electrolytic conduction
5.5 Factors for the variation of molar conductance

5.6 Kohlrausch’s law


Kohlrausch's Law on the independence of migrating ions: The molar conductivity of an electrolyte equals the sum of the molar conductivities of the cations and the anions; n = number of anions or cations.

Λ = v+Λ+ + vˉΛˉ

According to this law, the molar conductance of infinite dilution for a given salt can be expressed as the sum of the contributions from each ion of the electrolyte. If molar conductivity of the cation is denoted by Λˉ and anion by Λ+,and vˉ and v+ are number of cations and anions respectively, total molar conductance will be given by Λ.

Revision

1. Calculation of molar conductance at infinite dilution for weak electrolytes
2. Calculation of degree of dissociation of weak electrolytes

5.7 Electrochemical cell or galvanic cell




5.8 Representation of an electrochemical cell

5.9 Electrode potential and E.M.F. of a galvanic cell


The difference between the electrode potentials of the two electrodes constituting an electrochemical cell is known as electromotive force or cell potential of a cell.


5.10 Standard Electrode Potential

5.11 Electrochemical series


The electrochemical series is built up by arranging various redox equilibria in order of their standard electrode potentials (redox potentials). The most negative E° values are placed at the top of the electrochemical series, and the most positive at the bottom.



The electrochemical series

equilibrium E° (volts)
Li-3.03
K -2.92
Ca -2.87
Na -2.71
Mg -2.37
Al -1.66
Zn -0.76
Fe-0.44
Pb -0.13
H 0
Cu +0.34
Ag+0.80
Au +1.50

Remember that in terms of electrons:

OIL RIG

Oxidation is loss Reduction is gain


Reducing agents and oxidising agents

reducing agent reduces something else. That must mean that it gives electrons to it.

Magnesium is good at giving away electrons to form its ions. Magnesium must be a good reducing agent.

An oxidising agent oxidises something else. That must mean that it takes electrons from it.

Copper doesn't form its ions very readily, and its ions easily pick up electrons from somewhere to revert to metallic copper. Copper(II) ions must be good oxidising agents.

5.12 Differences between Galvanic cell and electrolytic cell


5.13 Dependence of electrode and cell potentials on concentration: Nernst Equations



Nernst Equation: The cell potential of a half cell (as well as that of a complete cell) depends upon the concentrations of involved ions, pressure of the gaseous species (if involved) and the temperature. The relation connecting them is given by the Nernst equation.

It is expressed as

E = E° - (RT/nF)ln Q°

Q° = Product of concentration (or pressure) of products each raised to the corresponding stochiometric number/Product of concentration (or pressure) of reactants each raised to the corresponding stochiometric number

n = number of electrons involved in the hall cell reaction

5.14 Equilibrium constant form Nernst Equation
5.15 Electrochemical cell and free energy
5.16 Some commercial cells
5.17 Electrode potential electrolysis and criteria for product formation
5.18 Corrosion
5.19 Commercial production of chemicals
5.20 Manufacture of some important metals an chemical compounds


ElectroChemistry - 35 Video Playlist

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ExamFearVideos

Updated 3 Feb 2016,  22 May 2015



4. Chemical Thermodynamics - JEE Main - Class XII - Core Revision Points

Importance of  Core Revision Points: Core Revision Points are important because if you remember them strongly, many more points related to them will come out of your memory and help you to answer question and problems. Read them many times and make sure you remember them very strongly.


Sections in the chapter (Jauhar’s Book CBSE)

4.1 Some basic terms and concepts
4.2 Internal energy, heat and work
4.3 First law of thermodynamics
4.4 Enthalpy and Enthalpy change
4.5 Limitations of first law of thermodynamics: Need of second law
4.6 Randomness and spontaneity
4.7 Entropy
4.8 Second law of thermodynamics
4.9 Entropy change during phase transition
4.10 Gibbs Free energy and free energy change
4.11 Free energy change for predicting feasibility of a reaction
4.12 Standard free energy change
4.13 Standard free energy change and equilibrium constant
4.14 Gibbs free energy change and nonmechanical work
4.15 Absolute entropies and third law of thermodynamics


Sections in the chapter (Jauhar’s Book CBSE)

4.1 Some basic terms and concepts
4.2 Internal energy, heat and work
4.3 First law of thermodynamics
4.4 Enthalpy and Enthalpy change
4.5 Limitations of first law of thermodynamics: Need of second law
4.6 Randomness and spontaneity
4.7 Entropy
4.8 Second law of thermodynamics
4.9 Entropy change during phase transition
4.10 Gibbs Free energy and free energy change
4.11 Free energy change for predicting feasibility of a reaction
4.12 Standard free energy change
4.13 Standard free energy change and equilibrium constant
4.14 Gibbs free energy change and nonmechanical work
4.15 Absolute entropies and third law of thermodynamics


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4. Chemical Thermodynamics - JEE - CBSE Class XII - Revision Questions


Updated  3 Feb 2016,  22 May 2015


3. Solutions - JEE Main - CBSE Class XII - Core Revision Points

Importance of  Core Revision Points: Core Revision Points are important because if you remember them strongly, many more points related to them will come out of your memory and help you to answer question and problems. Read them many times and make sure you remember them very strongly.


Sections in the Chapter

3.1 Types of solutions
3.2 Methods for expressing the concentration of a solution: units of solution
3.3 Solubilioty of gases and solids in liquids
3.4 Vapour pressure of solutions
3.5 Ideal and Nonideal solutions
3.6 Colligative properties
3.7 Relative lowering of vapour pressure
3.8 Elevation in boiling point
3.9 Depression in freezing point
3.10 Osmosis and osmotic pressure
3.11 Electrolytic solutions – Abnormal molar masses



Sections in the Chapter

3.1 Types of solutions



a solution is a homogeneous mixture of two or more substances whose composition can be varied within certain limits.

In a solution, the component which is in excess is called solvent.

The component, that has lesser quantity is called solute.

In a solution particles are of molecular size (about 1000 pm) and the different components cannot be separated by any of the physical methods such as filtration, settling etc.


The concentration of a solution may be defined as the amount of solute present in the given quantity of the solution.

volume percent
Vol% = 100*volume of solute/(volume of solute + volume of solvent)

mass percent
mass % = 100*(mass of solute)/(mass of solute + mass of solvent)

parts per million
ppm = 10^6*(mass of solute)/(mass of solute + mass of solvent)

molality
m = number of moles of solute/kilograms of solvent

molar concentration or moles per liter or Molarity
M = number of moles of solute/liters solution

Mole fraction
Xy (y is subscript) = number of moles of y in mixture/totals moles in mixture

Normality
It is the number of gram equivalents of the solute dissolved per litre of the solution. It is denoted by N.

Normality (N) =

[Number of gram equivalents of solute]/[Volume of solution in litres]
Units of normality are gm equivalent per litre.

Formality
It is the number of formula masses of the solute dissolved per litre of the solution. It is represented by F.

Formality = [Number of formula masses of solute]/[Volume of the solution in litre]

Formality is used to express the concentrations of ionic substances like NaCl, CuSO4 etc. in solutions. They do not exist in solutions as discrete molecules. In these solutions, the sum of the atomic masses of various atoms constituting the formula of the compound (ionic) is called gram formula mass instead of molar mass.




3.2 Methods for expressing the concentration of a solution: units of solution


3.3 Solubilioty of gases and solids in liquids


Solubility of Gases in Liquids

Gases dissolve in liquids to form homogeneous solutions. The solubility of different gases in the same solvent varies. Gases which react with the solvent will be most soluble. The solubility of a gas decreases with temperature and increases with with increase of pressure over the solution at a given temperature.

Henry's Law
The mass of a gas dissolved per unit volume of the solvent at a given temperature is proportional to the pressure of the gas in equilibrium with the solution.

m is proportional to p
where m = mass of the gas dissolved in a unit volume of the solvent
p = pressure of the gas in equilibrium.

If pressure is more, more mass of gas is dissolved

3.4 Vapour pressure of solutions


When a liquid is placed in a vessel and is covered with jar, from the liquid evaporation takes place and the vapour of the liquid or molecules of the liquid in gap form fill the available space. As the evaporation takes place over a period of time, the number of gaseous molecules goes up. As evaporation is taking place some molecules in the gaseous phase collide with the surface of the liquid and become liquid molecules. Thus both evaporation and condensation take place simultaneously. But initially there is more evaporation and less condensation. At the some stage, rate of evaporation equals rate of condensation and equilibrium is established between gas and liquid phases. The pressure exerted by the vapours at the equilibrium stage is called vapour pressure.

Definition
The pressure exerted by the vapours above the liquid surface (in a closed vessel) in equilibrium with the liquid at a given temperature is called vapour pressure.

Vapour pressure changes from liquid to liquid. It depends on intermolecular forces. if the forces in a liquid are weak, there is more gas formation and hence more vapour pressure.

A higher temperature there is more gas formation and hence for the same liquid vapour pressures increase with temperature.


Raolt's Law

In the case of a solution of two liquids, A and B, the total vapor pressure Ptot(P total) above the solution is equal to the sum of the vapor pressures of the two components, PA and PB and

PA = PA° * Am
PB = PB° * Bm

Where
PA° = vapour pressure created by 1 mol of liquid A
Am = mole fraction of liquid A in the solution
PB° = vapour pressure created by 1 mol of liquid A
Bm = mole fraction of liquid A in the solution

3.5 Ideal and Nonideal solutions

3.6 Colligative properties

3.7 Relative lowering of vapour pressure


Molecular weight determination from lowering of vapor pressure

Molar mass of a solute can be found from the property of lowering of vapor pressure of a solution.

Mb = (Wb*Ma)/[Wa*(Pa°-Pa)/Pa°]

Wb = weight of solute particles, Wa= weight of solvent
(Pa°-Pa)/Pa° = decrease in vapour pressure of solution
Ma = Molar mass of solvent

3.8 Elevation in boiling point

3.9 Depression in freezing point


Molecular weight determination from depression of freezing point.


The freezing point is the temperature at which the solid and liquid states the substance have the same vapour pressure.

When a non-volatile solute is added to a solvent, the freezing point of the solution is always lower than that of the pure solvent.

The depression in freezing temperature is proportional to the molal concentration of the solution (m).
ΔTf α m Or ΔTf = Kf*m

ΔTf = depression in freezing point.

Kf is the molal depression constant. also called molal cryoscopic constant. It is defined as the depression in freezing point for 1 molal solution i.e., a solution containing 1 gram mole of solute dissolved in 1000 g of solvent.
When m =1; ΔTf = Kf

Depression in freezing point is a colligatvie property as it is directly proportional to the molar concentration of the solute.


To find the molar mass of an unknown substance (nonvolatile compound), a known mass of it is dissolved in a known mass of a solvent and depression in its freezing point (ΔTf)is measured.

weight of solute be Wb g
weight of the solvent be Wa g
Molar mass of the solute be Mb

Molality of the solution, m = Wb*1000/Mb*Wa

Substitute the value of m in ΔTf = Kf*m = Kf*Wb*1000/Mb*Wa

From the above equation Mb can be calculated.

Mb = Kf*Wb*1000/Wa*ΔTf

Example:

Addition of 0.643 g of a compound to 50 ml of benzene (density 0.879 g/ml) lowers the freezing point from 5.51°C to 5.03°C. If Kf for benzene is 5.12 K kg molˉ¹, calculate the molar mass of the compound. (IIT 1992)

The formula of Mb is available above.

weight of solute be Wb g = 0.643 g

weight of the solvent be Wa g = 50*0.879 = 43.95 g
Change in freezing point = 5.51 - 5.03 = 0.48°C

Mb = (5.12 * 0.643 * 1000)/(43.95*0.48)






Mb = [Kf*Wb*1000]/[ΔTf * Wa]

3.10 Osmosis and osmotic pressure
3.11 Electrolytic solutions – Abnormal molar masses


Glossary


Solute: the substance dissolved; the substance present in a solution in the lesser amount.

Solvent: the dissolving medium; the substance in a solution in the greater amount.

Solution: a homogeneous mixture of two or more substances.

volume percent
Vol% = 100*volume of solute/(volume of solute + volume of solvent)

mass percent
mass % = 100*(mass of solute)/(mass of solute + mass of solvent)

parts per million
ppm = 10^6*(mass of solute)/(mass of solute + mass of solvent)

molality
m = number of moles of solute/kilograms of solvent

molar concentration or moles per liter or Molarity
M = number of moles of solute/liters solution

Mole fraction
Xy (y is subscript) = number of moles of y in mixture/totals moles in mixture

Solutions Part 1
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Revision Questions


3. Solutions - JEE - CBSE Class XII - Revision Questions

IIT JEE Chemistry - Study Guide - 3. Solutions

Sections in the Chapter

3.1 Types of solutions
3.2 Methods for expressing the concentration of a solution: units of solution
P.P. 3.1 to 3.20
3.3 Solubilioty of gases and solids in liquids
3.4 Vapour pressure of solutions
3.5 Ideal and Nonideal solutions
P.P. 3.21 to 3.27
3.6 Colligative properties
3.7 Relative lowering of vapour pressure
P.P. 3.28 to 3.34
3.8 Elevation in boiling point
P.P. 3.35 to 3.40
3.9 Depression in freezing point
3.41 to 3.48
3.10 Osmosis and osmotic pressure
P.P. 3.49 to 3.57
3.11 Electrolytic solutions – Abnormal molar masses
P.P. 3.58 to 3.68


Additional Numerical Problems for Practice: 21

Conceptual Questions with Answers: 29
Key facts to remember
Additional Numerical Problems for Practice: 8
Revision Exercises
Very Short Answer questions 21
Short Answer Questions 66
Long Answer Questions 21

Competition File
Numerical Problems 29
Multiple choice questions: 40
Fill in the blanks: 10
True or False: 10


Study Plan

Day 1

3.1 Types of solutions
3.2 Methods for expressing the concentration of a solution: units of solution

Day 2

P.P. 3.1 to 3.20

Day 3
3.3 Solubility of gases and solids in liquids
3.4 Vapour pressure of solutions
3.5 Ideal and Nonideal solutions

Day 4

P.P. 3.21 to 3.27
3.6 Colligative properties
3.7 Relative lowering of vapour pressure
P.P. 3.28 to 3.34

Day 5

3.8 Elevation in boiling point
P.P. 3.35 to 3.40

Day 6
3.9 Depression in freezing point
3.41 to 3.48

Day 7
3.10 Osmosis and osmotic pressure
P.P. 3.49 to 3.57

Day 8
3.11 Electrolytic solutions – Abnormal molar masses
P.P. 3.58 to 3.68

Day 9

Additional Numerical Problems for Practice: 21


Day 10
Conceptual Questions with Answers: 29

Day 11
Key facts to remember
Additional Numerical Problems for Practice: 8

Day 12
Revision Exercises: Very Short Answer questions 21

Day 13
Revision Exercises: Short Answer Questions 1 to 3366

Day 14
Revision Exercises: Short Answer Questions 34 to 66

Day 15
Competition File: Numerical Problems 1to 20

Revision Period
Day 16
Competition File: Numerical Problems 21to 29

Day 17
Multiple choice questions: 1 to 20

Day 18
Multiple choice questions: 21 to 40

Day 19
Fill in the blanks: 10

Day 20
True or False: 10

Day 21 to 30
Concept revision, formula revision, Test paper problem solving



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35 parts parts are there more can be there.


Playlist 35 parts
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Updated 3 Jan 2016, 11 March 2009

IIT JEE Chemistry - Study Guide - 4. Chemical Thermodynamics (XII Portion)

Sections in the chapter (Jauhar’s Book CBSE)

4.1 Some basic terms and concepts
4.2 Internal energy, heat and work
4.3 First law of thermodynamics
Practice Problems 4.1 to 4.6
4.4 Enthalpy and Enthalpy change
P.P 4.7 to 4.10
4.5 Limitations of first law of thermodynamics: Need of second law
4.6 Randomness and spontaneity
4.7 Entropy
4.8 Second law of thermodynamics
4.9 Entropy change during phase transition
P.P. 4.11 to 4.17
4.10 Gibbs Free energy and free energy change
4.11 Free energy change for predicting feasibility of a reaction
P.P. 4.18 to 4.23
4.12 Standard free energy change
4.13 Standard free energy change and equilibrium constant
P.P. 4.24 to 4.31
4.14 Gibbs free energy change and nonmechanical work
4.15 Absolute entropies and third law of thermodynamics
P.P 4.35 to 4.36


Additional numerical problems for practice 1 to 10
Conceptual Questions with Answers: 12

Revision Exercises
Very Short Answer questions 25
Short Answer Questions 48
Long Answer Questions 10

Competition File
Numerical problems for competitive examinations 17
Objective Questions: 36
Fill in the blanks: 10
True or False: 10


Study Plan

Day 1

4.1 Some basic terms and concepts
4.2 Internal energy, heat and work
4.3 First law of thermodynamics

Day 2

Practice Problems 4.1 to 4.6
4.4 Enthalpy and Enthalpy change
P.P 4.7 to 4.10

Day 3
4.5 Limitations of first law of thermodynamics: Need of second law
4.6 Randomness and spontaneity
4.7 Entropy
4.8 Second law of thermodynamics
4.9 Entropy change during phase transition

Day 4
P.P. 4.11 to 4.17
4.10 Gibbs Free energy and free energy change
4.11 Free energy change for predicting feasibility of a reaction

Day 5
P.P. 4.18 to 4.23
4.12 Standard free energy change
4.13 Standard free energy change and equilibrium constant
P.P. 4.24 to 4.31

Day 6

4.14 Gibbs free energy change and nonmechanical work
4.15 Absolute entropies and third law of thermodynamics
P.P 4.35 to 4.36

Day 7

Additional numerical problems for practice 1 to 10
Conceptual Questions with Answers: 12


Day 8
Revision Exercises: Very Short Answer questions 25

Day 9
Revision Exercises: Short Answer Questions 1 t o 24
Day 10
Revision Exercises: Short Answer Questions 25 to 48

Day 11
Competition File: Numerical problems for competitive examinations 17

Day 12

Competition File: Objective Questions: 1 to 18

Day 13
Competition File: Objective Questions: 19to 36

Day 14

Competition File: Fill in the blanks: 10
Competition File: True or False: 10

Day 15
Concept Revision
Fromula Revision

Day 16 to 30

Revision of the chapter/Extra problems from other books like R.C. Mukherjee, Test Paper books etc..
Use all available time productively to improve your understanding, ability to solve problems and recollection of principles.


11 Videos Playlist
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Updated 3 Feb 2016, 11 March 2009