Chemistry Xi

Prrofessor Jee Collegiate Chemistry XI

CHAPTER # 1INTRODUCTION TO FUNDAMENTAL CHEMISTRY

CHEMISTRY:

Branch of chemistry which deals with the study of substances, their structure, properties And the reactions that changes them into the substances is called Chemistry.

PHYSICAL CHEMISTRY:

Branch of chemistry which deals with the study of physical properties and the structure of matter and the laws and theories of chemistry is called Physical Chemistry.

ATOM: Greek: Atom → indivisible

OLD DEFINITION:

About 400 B.C, the Greek Philosopher Democritus defined an atom is extremely small particle of a matter which cannot be further substances. Modern Definition: According to the modern ideas an atom is the smallest particle of a matter which takes part of chemical reaction (except the atom of inert gases). It may or may not have Independent existence.

MOLECULE:

A molecule may be defined as the smallest particle of matter formed by the combination of two or more like or unlike atoms which have independent existence in nature of H2O , CO2, N2, O2 etc.

ELEMENT:

A substance cannot be further simplified into more simple substances by any chemical action and possesses identical chemical and physical properties is called Element.

OR

A substance is which all of the atoms have the same atomic number i.e. same number of protons in the nucleus is called Element.

COMPOUND:“A compound is a substance that is composed by two or more different elements united chemically”. E.g. CO2, HCl, NaCl etc.

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MIXTURE:

“A material containing two or more components which have not been chemically united and the components of the mixture retain their original properties is called Mixture”.Mixtures are of two types.

1. HOMOGENEOUS MIXTURE:

Mixture whose composition is uniform throughout is called Homogeneous Mixture.e.g. salt solution, sugar solution, air etc.

2. HETEROGENEOUS MIXTURE:

Mixture which is not uniform in composition throughout is called heterogeneous mixture of salt + sand, oil and water etc.

FORMULA:

A chemical formula is the symbolic representation of a molecule of a compound. It consists of symbols and the number of atoms of each element is a molecule of a compound.Formulas are of two types.1: Simple Formula or Empirical Formula2: Molecular Formula

EMPIRICAL OR SIMPLE FORMULA:

That formula which gives simple ratio or relative number of each kind of atoms present in the simple formula of a compound is called Empirical Formula”

Empirical formula is based on unit or formula. It does not show the actual number of the atoms in a molecule of a compound.

For example the empirical formulas of benzene and glucose are CH and CH2O. These formulas shoes that in benzene carbon and hydrogen are combining in the ratio of 1:1 and in glucose, carbon, hydrogen and oxygen combining in the ratio of 1:2:1

CHARACTERISTICS OF EMPIRICAL FORMULA:

1: Empirical formula of two or more compounds may be the same of benzene and acetylene have same empirical formula CH.

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2: A compound may have same empirical and molecular formulas e.g. methyl alcohol have same empirical and molecular formulas is CH3OH of carbon dioxide is CO2 etc.

MOLECULAR FORMULA:

Formula which gives the total number of each kind of atoms present in the molecule of a compound is called Molecular formula.

For example: Molecular formula of benzene is C6H6 . This shows that a molecule of benzene is composed of six atoms of carbon and six atoms of hydrogen. Molecular formula of a compound may be the same or simple multiple of the empirical formula.

For example: Molecular formula and Empirical formula of water and methane are same is H2O and CH4 respectively. On the other hand molecular formula of glucose is C6H 12O6 but its empirical formula is C H2O.

RELATION BETWEEN EMPIRICAL FORMULA AND MOLECULAR FORMULA:

Molecular formula is either same or simple multiple of empirical formula i.e.

Molecular formula = n × Empirical formula

Where ‘n’ is any simple whole number and is determined by following relationship.

N = Molecular weight (Mass) Empirical formula mass

When the value of n = 1, the molecular and empirical formulas are same.

GRAM ATOM:

Mass of one atom of an element is called Atomic mass v when this mass is expressed in gram is called Gram Atom.

E.g. The mass of one atom of hydrogen is 1.67 × 10-24 g.

ATOMIC MASS OR ATOMIC WEIGHT:

“The weight if one atom of the element as compared with the weight of one atom of carbon -12

the light isotope of carbon taken as 12.0 a.m.u”.

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Thus one atom of hydrogen which weigh 1/12th approximately the weight of light isotope of carbon-12 has atomic weight 1.00 a.m.u.

i.e. (1/12 × 12 = 1.00) a.m.u.

Thus atomic weight is expressed in atomic mass unit (a.m.u.). The atomic mass unit is the average mass of the isotopes as they occur in nature.

When atomic weight expressed in gram, it is called gram atomic weight.e.g. atomic weight(mass) of H = 1.0 a.m.u. Expressed in g = 1.0 g

MOLECULAR MASS (WEIGHT):

The molecular weight of substance is the average weight of its molecules as compared to the weight of one atom of carbon-12 or oxygen-16 of the molecular weight of H2 molecule is 2.0 a.m.u.

2/12 × 12 = 2.00 a.m.u.

The molecular weight of benzene (C6H6) is 78 a.m.u.

FORMULA WEIGHT:

Formula weight is the sum of the atomic (weights) masses of all the atoms present in the simple formula of the compressed.

Formula weight is expressed in atomic mass unit, when it is expressed in gram, it is called

Gram formula weight (Mass)Formula mass of H2O = 18.00 a.m.u.Expressed in gram = 18.00 gFormula mass of NaCl = 58.5 a.m.u.Expressed in gram = 58.5 g

MOLE:

When atomic mass of an element expressed in grams is called gram atomic mass, the molecular mass expressed in gram is called gram molecular mass and when formula weight ‘g’ a compound expressed in gram is called gram formula (Weight) mass. Each of these mass is known as a mole.

“Hence a gram atomic mass, gram molecular mass and formula mass is called a mole”.

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e.g. Atomic mass of H = 1.0 a.m.u. Expressed in gram = 1.0 a.m.u = 1 g mole of H-atomsMolecular mass of CO2 = 44.0 a.m.u.Expressed in gram = 44.0 g = 1 g mole of CO2 moleculesFormula weight of NaCl = 58.0 a.m.u.Expressed in gram = 58.0 g = 1 g mole of NaCl

MODERN DEFINATION:

“A mole is the amount of substances that contains as many number of particles (atoms, molecules or ions) as there are atoms contained in 12.0 gram of pure carbon-12”.

AVOGADRO’S NUMBER: (N OR NA)

One gram mole of any substance contains definite number of atoms, molecules or ions i.e.

6.02 × 1023. 6.02 × 1023 quantity is called Avogadro’s number and is denoted by “NA” or “N”.e.g. One mole of H atoms = 1 gram = 6.02 × 1023 atomsOne mole of H2 molecules = 2.0 g = 6.02 × 1023 moleculesOne mole of g NaCl = 58.5g = 6.02 × 1023 Na+ or = 6.02 × 1023 Cl-

Or = 6.02 × 1023

Or = 12.04 × 1023 ions of Na+Cl-

SIGNIFICANCE OF MOLE:

The significance of mole is that the having an atom of an element is, the more will be the mass of 6.02 × 1023 atoms of that element be. Also the lights an atom of an element is, the less will be the mass ‘g’ 6.02 × 1023 atoms of that element.

SIGNIFICANT FIGURES:

“Significant Figures are the reliable digits in a number that are known with certainly.The last digit of a number is generally considered uncertain by ± 1 in the absence of qualifying information. For example: The weight of an object can be expressed as 0.0112 g or 11.2 mg with changing the uncertainly of the weight or the number of significant figures. The weight is still uncertain by ± 1 in the last digit. This can be expressed is 0.0112 ± 0.0001 g or as 11.2 ± 0.1 mg.

RULES OF DETERMINING SIGNIFICANT FIGURES:

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These are the following rules for determining the significant figures.

1: In a number all the non-zero digits (1 to 9) are significant figures.e.g. 363 has three significant figures. 0.68 has only two significant figures.

2: In a number zero between non zeros are significant and, they are known as Interior Zeros.e.g. 5004 has four significant figures. 202.03 has five significant figures.

3: Zeros locating the decimal point in numbers less than one are not significant.e.g. 0.062 has two significant figures. 0.001 has one significant figures.

4: Final zero to the right of the decimal point are significant.e.g. 2.000 has four significant figures.and 506.40 has five significant figures.

5: Zeros that locate the decimal point in number large than one are not significant.e.g. 40 has one significant figure. 236000 has three significant figures.

RANDOM ERRORS:

“The errors which are unavailable are called random errors. i.e. the errors which remaining variations indicate are called random errors”.

EXPLANATION:

When all the systematic errors have either been eliminated or corrected, we still do not obtain exact or true measurements because there is some uncertainly is every physical measurement.A random error may be positive or negative. This is the only reason that we taken the average of the several measurements, which is more reliable than any individual measurements.

EXPONENTIAL NOTATION:

The powers to the base 10 are known as exponents. Thus all numbers may be expressed as a power of 10 and 102, 103, 105, 10-1, 10-2, 10-5 etc are generally called exponential notation in which the base is 10.

For example: One mole of any element contains 602,300,000,000,000,000,000,000 atoms. This is a very big value. On the other hand the mass of an electron is

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0.0000000000000000000000000000911gram. This is a very small value. Such numbers can in a better way be written in exponential notation.

Thus 1 mole of an element will contains 6.02 × 1023 atoms and mass of an electron is 9.11 × 10-28gm.

PRECISION AND ACCURACY:

A term used to express the closeness between a measured value and the average value for a series of measurements is called Precision.

The departure of the measured values from one another is called deviation.When the measured values are close to one another and they are also close to the actual value, then the result has good precision as well as good accuracy.

STOICHIOMETRY: ( Stoichiometry → Element )

DEFINITION:

The quantitative study of the relationships between the amount of reactants and products of a certain reaction as given by a balanced chemical equation is called Stoichiometry.

EXPLANATION:

The concept of mole finds one of its most useful applications in the calculations of masses and volumes of chemical substances taking part in or produced by chemical reactions.e.g. the following balanced equation indicates.

2H2S + 3O2 → 2SO2 + 2H2O

1: H2S gas and O2 gas react in the ratio of 2 molecules to 3 molecules to produce So2 gas and H2O vapors in the ratio of 2 molecules to 2 molecules.

2: Two moles (68 grams) of H2S reacts with 3 moles (96 grams) of oxygen gas to give 2 moles (36 grams) of water vapors.These amounts of reactants and products are called Stoichiometric amounts.

STOICHIOMETRIC CALCULATIONS:

While considering calculation based on the stoichiometry of chemical reaction. Three important assumptions are considered.1: Reactants are completely converted to products.2: No side reactions to be considered.

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3: A balanced chemical equation is provided.

STOICHIOMETRIC RELATIONSHIP:

There are three relationships involved for the stoichiometric calculations from the balanced chemical equation.

1: Mass – Mass Relationships ( weight-weight relationship)2: Mass - Volume Relationships3: Volume – Volume Relationships

1: MASS – MASS RELATIONSHIPS:

Such relationships are useful in determining an unknown weight of a reactant or product from a given weight of some substances in a chemical reaction. Hence for this calculation the term mass – mass is used.

IMPORTANCE:

The correct result by this relation is obtained if we provided with a complete and balanced chemical equation.

This balanced chemical equation tell us the relationships of the masses or weights of the reactants to one another and to the products.

UNITS:

The units of mass (weight) may be gram, kilogram, tons etc.

2: MASS - VOLUME RELATIONSHIPS:

Gases are most commonly measured by volume. The molar quantities of gases can be expressed in terms of volume as well as mass.

Therefore this relation gives a relation between number of moles of a gas and its volume and allows calculations of quantity of a gas consumed or released in a given chemical change as a volume as well as mass.

CONVERSION UNITS:

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For converting weight into volume or volume into weight (mass). The Avogadro’s Law is applied, according to which

i. One gram mole of gas at S.T.P occupy a volume of 2.241dm3 or liters.ii. 22.41dm3 or liters of a gas weigh equal to 1 gm mole (molecular mass expressed in

gram).

3: VOLUME – VOLUME RELATIONSHIPS:

In this relationship the volume of a substance (gas or liquid) is calculated from the given volume of another substance.

EXPLANATION:

The relation between volumes of the gases in a chemical change is based on Gay-lusse’s Law of combining volumes, which states that:“Gases react or are formed in a chemical reaction in the ratio of small whole numbers by volume measured under the same condition of temperature and pressure.

For Example:2CO + O2 → 2CO2

Both reactants and products are gases.The balanced chemical equation tells us that two volumes of carbon monoxide (CO) react with one volume of oxygen (O2) to yield two volumes of carbon dioxide (CO2). Thus coefficients of the formulas gives us the volume – volume relationship.

UNITS:

The units of volume may be liters dm3 etc, but the properties remains the same.

LIMITING REACTANTS:

DEFINITION:

The reactant which is present in limited amount and that after the consumption of which the chemical reaction is stopped is called Limiting Reactant.

As in irreversible chemical reactants, the reactions go to completion in the direction of the arrow, until one of the reactant which will be the limiting reactant is consumed entirely and the reaction stops.

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CHAPTER # 2THE THREE STATES OF MATTER

(SOLID, LIQUID & GAS)

MATTER:

“Anything which has mass and occupies space is called Matter”.

Properties / States Solid Liquid GasShapes Definite Not definite Not DefiniteVolume Definite Definite Not Definite

Diffusion Extremely slow Slow Very fastCompression No, cause deformity Limited High

Translational Motion No, particles vibrate at their mean position

Yes Yes

Kinetic Energy Negligible Moderate HighForce of attraction High Moderate Negligible

GASEOUS STATE:

The state of matter in which molecules are present wide apart and can move freely in the space provided, it neither has definite volume nor definite shape. The density of gaseous state is less than other states of matter.

KINETIC MOLECULAR THEORY OF GASES:

Postulates of KMT of gases are given below.

1: LARGE DISTANCE BETWEEN MOLECULES:

All gas consists of small particles called ‘molecule’ which are widely separated from each other.

2: NEGLIGIBLE ACTUAL VOLUME:

The actual volume of molecules of a gas is negligible as compared to the volume of the gas.

3: RANDOM ERROR:

The gas molecules are in a state of continuous random motion, travelling in straight paths between collisions, but in random directions.

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4: ELASTIC COLLISIONS:

The gas particles continuously collide with one another and with the walls of container their collision are perfectly elastic (i.e. no loss or gain of energy take place).

5: GAS PRESSURE:

The collision of gas molecules with the walls of container produces gas pressure.

6: NO ATTRACTIVE OR REPULSIVE FORCE:

In an ideal gas there are no attractive or repulsive forces among the molecules. Thus each molecule acts independently.

7: K.Eave α K:

The average kinetic energy of gas molecules is directly proportional to the absolute temperature. K.Eave α K

and at any given temperature the molecules of all gases possess some average kinetic energy.

BEHAVIOUR OF GASES:

DIFFUSIBILITY:

The distribution or spreading of gas molecules throughout the vessels is known as diffusion.

IN TERMS OF KMT:

The molecules of gas are widely separated and there is large empty space between them due to which they are free to move. Due to this free movement, molecules of gases intermix and spread out easily throughout the vessel.

EFFUSION:

“The escape of gas molecules from a very small whole of the vessel is called Effusion”.OR

Gas is forced out through a pore or tiny holes in vessels by applying little pressure.

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COMPRESSIBILITY:

The ability of matter by which they can squeeze their volume due to empty space between molecules by an external pressure.

OR

When external pressure is applied the volume of gas decreases that is known as compression.

IN TERMS OF KMT:

Due to large empty spaces between their molecules gases are easily compressed. When pressure is applied, molecules come closer to each other and hence gas is compressed.

EXPANSIBILITY:

The increase in temperature causes increase in average kinetic energy resulting in expansion and gas occupies greater volume.

OR

When pressure on a compressed gas is reduced, the gas molecules rush out the gas expands. This property of a gas is called expansibility.

Expansibility is the reverse of compressibility.

PRESSURE:

The force exerted by gas molecules on a unit area is called pressure of a gas.Pressure = Force / Area

At 0ºC at Sea level 1 atom = 760 mm of Hg = 76 cm of Hg = 760 torr = 14.7 Psi (Pound per square inch) = 101300 Nm-2 or Pa

IN TERMS OF KMT:

According to KMT gas molecules are in continuous random motion due to which they collide with one another and with the walls of container. So, gas pressure is the result of collision of gas molecules with walls of container.

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GAS LAWS

BOYLE’S LAW:

“The volume of a given mass of a gas is inversely proportional to the pressure exerted on it at a given temperature”.Mathematically;V α 1/P (at constant temperature)V = K 1/PK = VPPV = KOn the basis of this relation, Boyle’s law can also be stated as,“The product of pressure and volume of a given mass of gas is always constant at constant temperature”.Hence,For initial state of a gas P1V1 = K ----------- (i)For final state of a gas P2V2 = K ------------ (ii)Comparing equation (i) and (ii)P1V1 = P2V2

IN TERMS OF KMT:

If the volume of a gas is decreased at constant temperature, the average velocity of the gas molecules remains constant, so in collide more frequently with the walls of smaller vessel. The most frequent collision cause higher pressure.

GRAPHICAL REPRESENTATION:

(i) A graph of volume v/s 1/P gives straight line.

(ii) A graph between volume and total pressure gives parabolic curve.

LIMITATIONS OF BOYLE’S LAW:

Gases do not obey Boyle’s law at higher pressure and lower temperature.

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CHARLE’S LAW:

“The volume of a given mass of gas is directly proportional to absolute temperature at given pressure”.Mathematically;V α T (Pressure is constant)V = KTK = V / TOn the basis of this relation, Charle’s law can also be stated as,“Under constant pressure, the ratio of volume of a given gas to its absolute temperature is always constant”.Hence,For initial state of a gas V1 / T1 = K -------------- (i)For Final state of a gas V2 / T2 = K -------------- (ii)Comparing equation (i) and (ii)V1 / T1 = V2 / T2

IN TERMS OF KMT:

A decrease in temperature, decrease the kinetic energy of the gas molecules, that is average molecular velocity decreases. At constant pressure the decreased velocity causes the gas to shrink and occupy a smaller volume.

GRAPHICAL REPRESENTATION:

A graph between temperature and volume give straight line.

ABSOLUTE ZERO:

The hypothetical temperature at which the volume of all gases reduced to zero and all the motion cease to exist is called Absolute Zero.It is -273.16ºC or 0 K.

GRAPHICAL EXPLANATION OF ABSOLUTE ZERO:

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The volume of a gas decrease by decrease in temperature and theoretically it becomes zero at -273ºC or 0 K.

But practically this temperature has never been achieved because gases condense to liquids at temperature above this point.

LIMITATION OF CHARLE’S LAW:

This law is not applicable to low temperature and high pressure.

AVOGADRO’S LAW:

“Volume of a gas is directly proportional to the number of moles of gas at constant temperature and pressure.

OR“Equal volume of all the gases contains same number of moles at constant temperature and pressure”.Mathematically;V α n (temperature and pressure are constant)V = KnK = V / nOn the basis of this relation, Avogadro’s law can also be stated as,“At constant temperature and pressure the ratio of volume of gas and its number of moles is constant”.For two different sample of gas at same temperature and pressure a relation can be represented as,V1 / n1 = V2 / n2

GENERAL GAS EQUATION OR IDEAL GAS EQUATIONOR EQUATION OF STATE:

When we combine all gas laws; i.e. Boyle’s, Charle’s, Avogadro’s law. These equation is called General Gas Equation.

DERIVATION:

According to Boyle’s LawV α 1 / P ( T and n constant )According to Charle’s LawV α T ( P and n constant )According to Avogadro’s LawV α n ( T and P constant )

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Combining all three laws:V α 1 / P .T.nV = R.1/P.T.nV = RTn/PPV = nRT ---- General Gas EquationHere ‘R’ is constant of proportionally called General Gas constant.For 1 mole of gas (n=1)PV = nRTORR = PV/TIf P1, V1 and T1 are the initial pressure, volume and temperature of gas, thenR = P1V1 / T1

And similarly P2, V2 and T2 are final pressure, volume and temperature, thenR = P2V2 / T2

Hence,P1V1 / T1 = P2V2 / T2 It is another form of general gas equation.

VALUE OF ‘R’ (GAS CONSTANT):

Value of universal gas constant can be calculated by following formula.R = PV / nTIt depends upon the unit in which P, V and T are expressed.

VALUE Of ‘R’ IN dm 3 atm mol -1 K -1 :

P = 1 atmV = 22.4 dm3

n = 1 mol.T = 273 KR= 1atm × 22.4 dm 3 1 mol × 273 K

R = 0.0821 dm3 atm mol-1 K-1

VALUE OF ‘R’ IN J mol -1 K -1 OR Nm mol -1 K -1 :

P = 101.300 Nm-2

V = 0.0224 m3

T = 273 KR = 101300 Nm -2 × 0.0224 m 3 1 mol × 273 KR = 8.3143 Nm.mol-1.K-1

Since,

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1 Nm = 1 JouleTherefore,R = 8.3143 J.mol-1.K-1

GRAHAM’S LAW OF DIFFUSION:

In 1881, Graham established a relationship between the rate of diffusion of gases and their densities which is known as Graham’s law of diffusion, which is stated as,“Rate of diffusion of gases is inversely proportional to the square root of the density of that gas provided the pressure and temperature constant”.Mathematically;r α 1/ √dr = K / √dwhere,K = rate constantr = rate of diffusiond = density of gasIf rate of diffusion of two gases are compare provided they are at same temperature and then, for one gasr1 = K / √d1 ---------------- (i)the other gas,r2 = K / √d2 -------------- (ii)Comparing equation (i) and (ii)r1 = K / √d1

r2 = K / √d2

r1 = √d2

r2 = √d1

Since, the density of gas is proportional to its molecular mass.r1 = √ M1

r2 M2

Since the rate of diffusion is inversely proportional to the time taken for its diffusion.T1 = √ M1

T2 M2

IN TERMS OF KMT:

According to KMT there are large empty spaces present among gas molecules so of another gas is introduced the molecules of the gas will occupy the empty spaces. Hence gases can diffuse easily.And according to KMT two gases at equal conditions of temperature and pressure have same kinetic energy.(K.E) = (K.E)2

½ m1v12 = ½ m2v2

2

m1 = v22

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m2 v12

v2 = √ m1 This is Graham’s law.v1 m2 As the rate of diffusion is directly proportional to velocity of molecules and density of gas is directly proportional to mass of molecules.

LIMITATIONS OF GRAHAM’S LAW:

Therefore r2/r1 = √α1/α2

Gases do not obey Graham’s law of diffusion at high pressure and low temperature.

APPLICATION OF GRAHAM’S LAW:

(i) To determine molecular weight and density.(ii) In diluting the poisonsous gases and minimizing their toxic effect.(iii) In separating (a) dlf isotopes of gas (b) one gas from another gas

PARTIAL PRESSURE:

When two or more gases which do not react chemically, are mix in the same container each gas will exert the same pressure as it would exert if its alone occupied the volume containing the mixed gases. This portion of total pressure of mixture is known as its partial pressure.

ORThe pressure exerted by an individual gas in a gaseous mixture is called its partial partial pressure.

DALTON’S LAW OF PARTIAL PRESSURE:

In 1801, John Dalton formulated a law about partial pressure of gases, it is stated as,“The total pressure of a mixture of gases is a sum of partial pressure of gases in the mixture”.Mathematically:Consider a mixture of three gases A, B and C AndPA is the partial pressure of gas A.PB is the partial pressure of gas B.PC is the partial pressure of gas C.

Pt is the total pressure of mixture.Then, According to Dalton’s lawPt = PA + PB + Pc

As we know the general gas equation PV= nRT.

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If PT is the total pressure and nT is the total number of moles.Then,PTV = nTRTPT = nTRT/ V -------------(i)Similarly, PA is the partial pressure of gas A and nA is the number of moles of gas A.PA = nART/ V -------------(ii)Comparing equation (i) and (ii)PA / PT = nA / nT

PA = nAPT / nT This equation can be used to find out partial pressure.

IN TERMS OF KMT:

In a non reactive mixture of gases each gas exerts a separate on the container because of collision of its molecules with the walls of container. Thus, the total pressure in the container is caused by the sum of all the collision.

LIMITATION OF DALTON’S LAW:

This law can only be applied for a mixture of gases which do not react chemically.

APPLICATION:

A gas is collected over water also contain some of water vapor. Thus, the pressure exerted by gas would be pressure of pure gas plus the pressure of water vapors’.Pmoist = Pdry + Pwater vapors

Pdry = Pmoist – Pwater vapors

Thus, partial pressure of dry gas can be obtained and its volume at S.T.P can be calculated by following formula.V2 = P1T2V1 / T1P2

IDEAL GAS:

An ideal gas is one whose behavior can be predicted precisely on the basis of kinetic molecular theory and which obeys gas laws at all temperature and pressure.

REAL OR NON-IDEAL GAS:

A gas which does not follow the gas laws at all temperature and pressure and whose behavior cannot be predicted precisely on the basis of KMT is non-ideal or real gas.

DEVIATION FROM IDEAL BEHAVIOR:

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A gas which obeys gas laws & KMT is called an ideal gas but actually no such gas exists. The deviation of real gases from ideal behavior is due to two postulates of kinetic molecular theory which are:

(i) There are no forces of attraction among the molecules of a gas.(ii) The actual volume of gas molecules is negligible as compared to the volume of

vessel.

But at high pressure and low temperature the attractive forces among molecules becomes significant and the actual volume of gas does not remain negligible.

LIQUID STATE:

This state of matter is intermediate between gaseous and solid state. It has definite volume but no definite shape.

KINETIC THEORY FOR LIQUIDS:

1. All liquids consist of molecules e.g. H2O and C6H6.2. The molecules of a liquid are very close together and are arranged in random manner.3. The molecules are free to move. They assume the shape of container. It means liquids

have no definite shape. They can flow and can be poured from one container to another.

4. Close and compact arrangements of molecules enable liquids to possess definite volume.

5. The molecules can be compressed to negligible extent.6. Average kinetic energy of molecules increases by rise of temperature ( Thus, liquids

on heating and contract on cooling).

BEHAVIOR OF LIQUID:

1: DIFFUSIBILITY:

The distribution of liquid molecules throughout the container is called diffusion of liquids and this property is called Diffusibility.

IN TERMS OF KMT:

The freedom of liquid molecules permits diffusion to take place but the cohesive forces among the molecules cause diffusion to be slow.

2: COMPRESSIBILITY:

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The squeezing of liquid molecules under the influence of external pressure is called Compressibility.Liquids are not absolutely incompressible but they may be compressed to a little extend by high pressure.

IN TERMS OF KMT:

The liquid molecules due to their closeness roll over one another. Owing to very little space the liquid molecules cannot be pushed close by pressure. Very high pressure is required for squeezing a liquid.

3: EXPANSION AND CONTRACTION:

A liquid normally expands on heating and contracts on cooling.

IN TERMS OF KMT:

On heating the K.E and velocity of molecules increase, it results in expansion.On cooling the K.E is lowered and negative forces overcome so molecule take less volume and contraction occurs.

VISCOSITY:

“The internal resistance between the molecules of a liquid to flow due to intermolecular attraction is called Viscosity”.

ORIt is the property by virtue of which it tends to oppose the relative motion between two adjacent layers.It is represented by ‘ŋ’ (eta).

UNITS:

It is measured in poise, centipoises and millipoise and S.I unit is Newton second per meter square (Nsm-2).1 poise = 1 Nsm-2

1.C.poise = 10-2 Nsm-2

1 millipoise = 10-3 Nsm-2

FACTORS AFFECTING VISCOSITY:

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1: SIZE OF THE MOLECULES:

It is more difficult for larger molecules to slip over one another than small ones. So viscosity increases with increase in molecular size or weight.For Example: Grease and oils are more viscous than water because of having large molecules.

2: MOLECULAR SHAPE:

The irregular shaped molecules offer more resistance to flow than regular shaped molecules.Example: Honey is very viscous due to irregular shape of its molecules.

3: INTERMOLECULAR FORCES:

The liquids having strong intermolecular forces are more viscous as their molecules are less easily moved about.Example: C2H5-OH is more viscous than C2H5-O-C2H5 as alcohol molecules have hydrogen bond.

4: TEMPERATURE:

With rise in temperature, the K.E increases and intermolecular forces become less effective and hence viscosity decreases.Example:

♦ Viscosity of ethyl alcohol is 1.78 poise at 0ºC and 0.7 poise at 50ºC.♦ Viscosity of water is 1.005 poise at 20ºC while 0.55 at 50ºC.

5: DENSITY:

The greater the density of liquid, the more viscous will be the liquid and vice versa.

SURFACE TENSION:

“The force per unit length acting at right angles of the surface of liquids at a given temperature is called Surface tension”.

OR“Energy per unit area at the surface of liquid”>It is represented by ‘γ’ (gamma).

UNITS:

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It is measured in two units dynes/cm and ergs/cm.

FACTORS AFFECTING SURFACE TENSION:


Surface tension of a liquid is directly proportional to intermolecular forces.Example: Surface tension of water is 72.6 at 20ºC while that of benzene is 28.9 at 20ºC.

2: TEMPERATURE:

If temperature increases the surface tension decreases and Vice Versa.

EFFECTS OF SURFACE TENSION:

1: SPEHRICAL FALLING DROP:

Surface tension tends to reduce the surface area of falling drop of liquid. Since all sphere has least area of a given volume of liquid. That’s why falling drop of liquid are spherical.

2: CAPILLARY ACTION:

The fall or rise of a liquid in capillary tube is called Capillary action.

EXPLANATION:

If a capillary tube is placed in liquid which wets the wall of container the surface area of liquid is decreased by raising the tube because of surface tension. Such liquid (as water) will rise in capillary tube until the upward force due to surface tension is just balanced by downward gravitational pull.

The level of non-wetting liquid in capillary tube will fall below the level of liquid in surrounding space.

3: CONCAVE & CONVEX SURFACE OF LIQUID:

Liquid which have stronger adhesive forces between liquid molecules and surface of glass will form concave meniscus/surface e.g. H2O.While liquids which have stronger cohesive force and do not wet the glass will form convex meniscus/surface e.g. Hg.

VAPOR PRESSURE:

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The vapor pressure of liquid is a measure of tendency of that liquid to evaporate and is defined as:

“The pressure exerted by its vapors when they are in equilibrium with the liquid phase at a given temperature”

It is constant at constant temperature and is independent of the amount of liquid present.

UNITS:

It is measured in following units, mm Hg, atmosphere, N/m2, pascal etc.

FACTORS AFFECTING VAPOR PRESSURE:

1: NATURE OF LIQUID:

Low B.P liquid exerts more vapor pressure and liquid that evaporates readily is said to be volatile.Example:Gasoline exerts more vapor pressure at a given temperature.

2: TEMPERATURE:

Vapor pressure increase with the rise of temperature. This is because average K.E of molecules increase with the increase of temperature and result in escape of vapors’ from the surface.Example:Vapor pressure of H2O is 18 torr at 20ºC while 760 torr at 100ºC.


Liquids having low intermolecular forces have high vapor pressure, while liquidsWith strong force have low vapor pressure.

BOILING POINT:

“The temperature at which vapor pressure becomes equal to the atmosphere pressure is called Boiling Point”.

FACTORS ON WHICH BOILING POINTS DEPENDS:

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1: EXTERNAL PRESSURE:

Boiling point increases with the increase of external pressure and Vice Versa.


The stronger the intermolecular forces the higher the B.P and Vice Versa.

VACCUME DISTILLATION:

Certain liquid tends to decompose at their boiling point. Such liquids are distilled under reduced pressure at low temperature. This process is called Vacuumed Distillation.

By lowering the pressure, boiling point of liquid is reduced and their decomposition is prevented.

PRESSURE COOKER:

Pressure Cooker helps to cook food rapidly even at higher altitude.

EVAPORATION:

“The spontaneous escape of liquid molecules from its surface without boiling is called Evaporation”.

BOILING EVAPORATION(i) Boiling takes place only a

particular temperature at which the vapor pressure of a liquid is equal to the pressure of the atmosphere.

(ii) Boiling involves the formation of bubbles of the vapor throughout the bulk of liquid.

(i) Evaporation occurs spontaneously at all temperature.

(ii) Evaporation takes place only at the surface of liquid.

COHESIVE FORCES:

The forces of attraction between similar particles of liquid are called Cohesive Force.

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ADHESIVE FORCES:

The forces of attraction between a liquid and another surface such as glass are called Adhesive Force.

SOLID STATE:

The state of matter in which atoms, molecules or ions are held together by strong attractive forces and cannot move freely. However, they possess vibration energy. It has definite volume as well as definite shape.

KINETIC THEORY OF SOLIDS:

1: Solids consist of atoms, ions and molecules.2: Particles of solids are closely packed due to intermolecular and interatomic forces. Thus solids have definite shape and volume.3: The particles are unable to move freely, however they may exhibit vibration motion.4: The particles are arranged in fixed patterns in the form of crystals lattice of vibrating masses.5: The average kinetic energy of particles increase by the rise of temperature.

BEHAVIOR OF SOLIDS:

Solids have highest degree of order which accounts different behaviors.

1: COMPRESSIBILITY:

The effect of pressure is negligible on solid and they cannot be compressed instead the mass is deformed.

IN TERMS OF KMT:

Particles of solids are so tightly bound together that only slight unfilled space is left so cannot be compressed.

2: DEFORMITY:

Solids are deformed by high pressure.

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IN TERMS OF KMT:

When some particles are dislocated under high pressure their force of attraction are enough strong that they rearrange well with their new neighbors.

3: DIFFUSIBILITY:

Diffusion in solid is very slow or almost negligible.

IN TERMS OF KMT:

There is no translational movement of molecules in solid but nevertheless particles are vibrating at their mean position and this vibration motion is responsible for diffusion in solid.

4: MELTINA:

When Solid is heated at certain temperature it starts changing into liquid this phenomenon is called melting and this temperature is called melting point of solid.

IN TERMS OF KMT:

When solid is heated, the kinetic energy of molecules increases which increases their vibration, energy and when it overcome the forces of attraction between molecules the solid starts melting.

5: SUBLIMATION:

“The process by which solids directly change to vapor without passing through liquid phase, process is called Sublimation.”Such solid is known as ‘Sublime’.Example: Nepthelene, Camphor, NH4Cl etc.

IN TERMS OF KMT:

The molecular forces of such solid are so weak that when they get energy the molecules get free and change into vapor.

CLASSIFICATION OF SOLIDS:

There are two types of solid.1: Crystalline 2: Amorphous

CRYSTALLINE SOLIDS:

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“Solids which have definite geometrical shape due to highly ordered three dimensional arrangements of particles are called Crystalline Solids.”Example: Sugar, Alum, Metals and Diamonds.

AMORPHOUS SOLIDS:

“Solids which do not have definite shape and non-repetitive three dimensional arrangements.”Example: Glass, Plastic and Rubber.

CRYSTALLINE SOLIDS AMORPHOUS SOLIDS(i) Particles have definite three

dimensional arrangements.(ii) They have sharp melting points.(iii) Regular intermolecular forces

throughout crystal.(iv) Possess repeating unit cell.(v) Breakdown in a fixed cleavage

plane.(vi) They are anisotropy.(vii) They are three solids.(viii) Examples: NaCl, Graphite,

Diamond etc.

(i) Particles are randomly arranged.(ii) They do not have sharp melting

point.(iii) Intermolecular force very form

place to place.(iv) They do not possess unit cell.(v) Do not break in a fixed cleavage

plane.(vi) They are isotropy.(vii) They are pseudo solid and super

cooled liquids.(viii) Examples: Glass, Plastic Rubber

etc.

CLEAVAGE:

The breakage of big crystals into smaller crystals of identical shape is called Cleavage.

CLEAVAGE PLANE:

The crystals split up along particular directions. The plane which contains the direction of cleavage is called Cleavage Plane.

ANISOTROPY:

The behavior of crystals showing variation in property in different direction is known as Anisotropy.Example: Graphite can conduct electricity parallel to its plane of layers but not perpendicular to plane.

ISOTROPY:

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The behavior of crystals showing similar property in all directions is known as isotropy.Example:Electrical conductance, thermal conduction etc.

TYPES OF CRYSTALS:

Crystals are classified into four classes.

ATOMIC CRYSTALS:

The crystals in which atoms are packed or held together by metallic bond,l hence also known as metallic crystals.Example: All metals from ionic crystals.They have following properties.♦ Luster♦ High melting point♦ Good conductor♦ Malleability♦ Ductility

IONIC CRYSTALS:

The crystals which consist of positively and negatively charged ions held together by electrostatic forces of attraction are called Ionic Crystals.Example: NaCl, CaCl2, CuSO4

They have following properties.♦ High melting point♦ Conduct electricity in fused state or in solution form♦ Brittleness and hardness.

3: COVALENT CRYSTALS:

Many non-metals consist of polyatomic molecules and exist in solid and their atoms are held together by covalent bond, hence known as Covalent Crystal.They require large amount of energy to break up.Example: Diamond, Graphite, Silica.They have following properties.♦ High melting point♦ Low density♦ High refractive index

MOLECULAR CRYSTALS:

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The crystals which consist of molecules held together by hydrogen bonding or vander waal’s forces are called molecular crystals.Example: Ice, Solid CO2, Solid I2

They have following properties.♦ Low melting point♦ Bad conductor

ISOMORPHISM:

“When two different substances have same crystal structure, they said to be isomorphous and the phenomenon is called Isomorphism”.Examples:♦ NaF & MgO → Cubic♦ CaCO3 & NaNO3 → Trigonal♦ ZnSO4 & NiSO4 → OrthorhombicThey have following properties.♦ Th ey have different physical and chemical properties.♦ They have same empirical formula.♦ When their solution are mixed, they form mixed crystal.♦ They show property of over growth.♦ Length of its edges denoted by a, b and c.♦ Angle between its edges by α, β and γ.

SYMMETRY:

Repetition of similar edges, corners or faces at a fixed distances along definite directions is called Symmetry.

CRYSTAL LATTICE OR SPACE LATTICE:

If atoms, ions or molecules constituting a crystal are replaced by points and placed at the same place as in a unit cell, then the three dimensional array of points is called Crystal Lattice or Space Lattice.

CRYSTAL SYSTEM:

There are seven types of crystal system on the basis of cell dimensions and shapes.

1: CUBIC SYSTEM:

A crystal system in which all axis of equal length and all angles of 90ºC called Cubic System.

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a = b = cα = β = γ = 90ºCExample: Nacl, ZnCl and diamond.

2: TETRAGONAL SYSTEM:

It has two axis of equal length and is different, while all angles are of 90º.a = b ≠ cα = β = γ = 90ºCExample: SnO2, BaSO4, 4H2O

POLYMORPHISM:

“The substance which can exist in more than one crystalline form is called polymorphous and phenomenon is called Polymorphism”.Example: CaCo3 occurs in two crystalline forms.

(i) Calcite → Trigonal(ii) Argonite → Orthorhombic

ISOMORPHISM POLYMORPHISM(i) The phenomenon is which two

different substance have same crystal structure.

(ii) They have different chemical properties.

(iii) Example: NaF & MgO bith have cubic structure.

(i) The phenomenon in which one substance exists in more than one crystalline from.

(ii) They have similar chemical properties.

(iii) Example: CaCO3 exist in Calcite and Argonite crystalline from.

UNIT CELL:

“The basic structural unit which when repeated in three dimensional produces the crystal structure is called a Unit Cell”.

CELL DIMENSION OR CELL PARAMETER:

The length and angles of a unit cell are collectively known as cell dimensions or cell parameters.

3: ORTHORHOMBIC SYSTEM:

It has all three axis of different length but all angles are of 90º.

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a ≠ b ≠ cα = β = γ > 90ºExample: FeSO4.7H2O , ZnSO4.7H2O

4: TRIGONAL SYSTEM:

It has all three axis of equal length and all three angles are equal but more than 90º and less than 120º.a = b = cα = β = γ > 90ºExample: Ice, Calcite, NaNO3, AgNO3, KNO3.

5: HEXAGONAL SYSTEM:

It has two axis of equal length but third is different, while two angles are of 90º and one is 120ºExample: Graphite, SiO2

6: MONOCLINIC SYSTEM:

It has all three axis of unequal length one of the axis is at right angle to the other two while two angles are of 90º and one is greater than 90º.a ≠ b ≠ cα = γ = 90ºB > 90ºExample: NaCO3, 10H2O

7: TRICLINIC:

It has all axis of different length and all angles are also different.a ≠ b ≠ cα ≠ β ≠ γ ≠ 90ºExample: K2Cr2O, CuSO4.5H2O

MELTING POINT:

“Melting point is the temperature at which equilibrium is established between solid and liquid phase”.

FACTORS AFFECTING MELTING POINT:

1: IMPURITY:

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It lowers the melting point of a substance, so melting point of pure substance is sharp.

2: PRESSURE:

There are two types of effect of melting point.(a) Those substance which expand on melting have a rise in melting point when pressure

is increased.(b) Those substance which contract on melting have a fall in melting point when pressure

is increased.

LATENT HEAT OF FUSION:

“The amount of heat energy required to change 1 g of a solid into liquid at its M.P”.

LATENT HEAT OF VAPORIZATION:

“The amount of heat energy required to change 1 g of a pure liquid into vapors at its B.P”.

CHAPTER # 3CROOK’S TUBE OR DISCHARGE TUBE EXPERIMENT

PASSAGE OF ELECTRICITY THROUGH GASES UNDER LOW

PRESSURE:

INTRODUCTION:

The first of the subatomic particles to be discovered was electron. The knowledge about the electron was derived as a result of the study of the electric discharge in the discharge tube by J.J.Thomson in 1896. This work later extended by W.Crooke.

CONSTRUCTION OF THE TUBE:This discharge tube consists of a glass tube with metal electrodes fused in the walls as shown in the figure.

Through a glass sidearm air can be drawn with a pump. The electrodes are connected to a source of high voltage.

WORKING OF A DISCHARGE TUBE:

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When a very high voltage about 10,000 volts is applied between the two electrodes, no electric discharge occurs until the part of the air has been pumped out of tube. When the pressure of the gas inside the tube is less than 1mm, a dark space appears near the cathode and thread like lines are observed in rest of the tube. On further decrease of pressure the dark space increases in size. At 0.01 mm Hg it fills the whole tube. The electric discharge passes between the electrodes and the residual gas in the tube begins to glow. These rays which proceed from cathode and move away from it at right angle in straight lines are called cathode rays.

PROPERTIES OF CATHODE RAYS:

They travel in straight lines away from cathode and produce shadow of the object placed in their paths.

The rays carry a negative charge. These rays can also be easily deflected by an electrostatic field. The rays can exert mechanical pressure showing that these consist of material particle

which are moving with kinetic energy. The produce fluorescence when they strike the glass wall of the discharge tube. Cathode rays produce x-rays when they strike a metallic plate. These rays consist of material particle whose e/m resembles with electron. These rays emerge normally from the cathode and can be focused by using a concave

cathode.

POSITIVE RAYS:

In 1896, Goldstein used a discharge tube with a hole in cathode. He observed that while cathode rays were emitting away from cathode, there was colored rays produced simultaneously which passed through the perforated cathode and caused a glow on the wall opposite to anode. Thomson studied these rays and observed and showed that they consisted of particles carrying a positive charge. He called them Positive Rays.

PROPERTIES OF POSITIVE RAYS:

These rays travel in straight line in a direction opposite to the cathode. These are deflected by electric as well as magnetic field in the way indicating that

they are positively charged. The charge to mass ratio (e/m) of positive particles varies with the nature of the gas

placed in discharge tube. Positive rays are produced from the ionization of gas and not from anode electrodes. Positive rays are deflected in electric field. This deflection shows that these are

positively charged so these are named as protons.

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THE INFORMATION OBTAINED FROM DISCHARGE TUBE EXPERIMENT:

The negatively charged particles electrons and positively charge particles are the fundamental part of every atom.

RADIOACTIVITY:

In 1895, Henry Beequeral observed that uranium and its compounds spontaneously emitted certain type of radiation which affected a photographic plate in dark and were able to penetrate solid matter. He called these rays as radioactive rays and a substance which possessed the property of emitting these radioactive rays and a substance which possessed the property of emitting these radioactive rays was said to be radioactive element and the phenomenon was called Radioactivity.On further investigation by Mark Curie, It was found that the radiation emitted from the element uranium as well as its salts is independent of temperature and source of mineral but depend upon the quantity of uranium present e.g. pitchblende U3O8 was found to be about four times more radioactive than uranium.

RADIOACTIVE RAYS:

Soon after the discovery of radium, it was suspected that these rays given out by radium and other radioactive substances were not of one kind. Rutherford in 1902 devised an ingenious method for separating these rays from each other by passing them between two oppositely charged plates. It was observed that the radioactive rays were of three kinds, the one bending toward negative plate obviously carrying positive charge were called α-rays and those deflected to the positive plate and carrying negative charge were named β-rays. The third type gamma rays pass unaffected and carry no charge.

PROPERTIES OF α -RAYS:

These rays consist of positively charge. These particles are fast moving helium nuclei. The velocity of α-particles is approximately equal to 1/10th of the velocity of light. Being relatively large in size, the penetrating power of α-rays is very low. They ionize air and their ionization power is high.

PROPERTIES OF β-RAYS: These rays consist of negatively charged particles. These particles are fast moving electrons. The velocity of β-particles is approximately equal to the velocity of light. The penetrating power of β-rays is much greater than α-rays.

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These rays ionize gases to lesser extent.

PROPERTIES OF γ-RAYS:

Gamma rays do not consist of particles. These are electromagnetic radiations. They carry no charge so they are not deflected by electric or magnetic field. Their speed is equal to that of light. These are weal ionizer of gases. Due to high speed and non-material nature they have great power of penitrarion.

CHADWICK EXPERIMENT (DISCOBERY OF NEUTRON):

When a light element is bombarded by α-particles, these α-particles leave the nucleus in an unstable disturbed state which on settling down to stable condition sends out radioactive rays. The phenomenon is known as “Artificial Radioactivity”.

In 1933, Chadwick identified a new particle obtained from the bombardment of beryllium by α-particles. It had a unit mass and carried no charge. It was named ‘Neutron’. The process for this reaction is given below.

4Be9 + 2He4 → 6C12 + 0n1

With the help of artificial radioactivity, the particle neutron was discovered.

SPECTROSCOPIC EXPERIMENT:

After the discovery of fundamental particles which are electron, proton, neutron, the next question concerned with electronic structure of atom.The electronic structure of atom was explained by the spectroscopic studies. In this connection Planck’s Quantum theory has great impact on development of theory of structure of atom.

PLANCK’S QUANTUM THEORY:

In 1900, Max Planck studied the spectral lines obtained from hot body radiations at different temperatures. According to him:

“When atoms or molecules absorb or emit radiant energy, they do so in separate unit of waves called Quanta or Protons”.

Thus light radiation obtained from excited atoms consists of a stream of protons and not continuous waves.

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The energy E of a quantum or proton is given by relation.E = hvWhere ‘v’ is the frequency of emitted radiation and ‘h’ the Planck’s constant. The value of h = 6.62 × 10-27 erg.sec

The main point of this theory is that the amount of energy gained or lost is quantized which means that energy change occurs in small packets or multiple of those packets, hv, 2hv, 3hv and so on.

SPECTRA:

A spectrum is an array of waves or particles spread out according to the increasing or decreasing of some property. E.g. when a beam of light is allowed to pass through a prism, it splits into seven colors. This phenomenon is called dispersion and the bend of colors is called spectrum. The spectrum is also known as emission spectrum. Emission spectra are of two types.

Continuous Spectrum Line Spectrum

CONTINUOUS SPECTRUM:

When a beam of white light is passed through a prism, different wave length is refracted through different angles. When received on a screen these form a continuous series of color bands, violet, indigo, blue, green, yellow and red (VIBGYOR). The colors of this spectrum are so mixed up that there is no line of between different colors. This series of bands that form a continuous rainbow of colors is called Continuous Spectrum.

LINE SPECTRUM:

When light emitted from a gas source passes through a prism a different kind of spectrum may be obtained.

If the light emitted from the discharge tube is allowed to pass through a prism some discrete sharp lines on a completely dark background are obtained. Such spectrum is known as line spectrum. In this spectrum, each line corresponds to a definite wave length.

IDENTIFICATION OF ELEMENT BY SPECTRUM:

Each element produces a characteristic set of lines, so line spectra came to serve as ‘finger prints’ for the identification of elements. It is possible because same elements always emit the same wave length of radiation. Under normal condition only certain wave lengths are emitted by an element.

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RUTHERFORD’S ATOMIC MODEL:

EVIDENCE FOR NUCLEUS AND ARRANGEMENTS OF PARTICLES:

Having known that atom contain electrons and a positive ion, Rutherford and Marsden performed their historic “Alpha particle scattering experiment” in 1909 to know how and where these fundamental particles were located in the structure of atom.

Rutherford took a thin sheet of gold with thickness 0.0004cm and bombarded it with α-particles. He observed that most of the α-particles passed straight through golf foil and thus produced a flash on screen behind it. This indicated that gold atoms had a structure with plenty or empty space, but some flashes were also seen on portion of screen. This showed that gold atoms deflected or scattered α-particles through large angles so much so that some of these bounced back to the source.

Based on these observations, Rutherford proposed a model of the atom which is known as Rutherford’s atomic model.

ASSUMPTION DRAWN FROM MODEL:

1: Atom has a tiny dense central core or nucleus which contains practically the entire mass of the atom leaving the rest of atom almost empty.2: the nucleus carries a positive charge +2e.3: The electrons were moving in orbits or closed circular paths around the nucleus like planets around the sun.4: The greater part of atomic volume comprises of empty space in which electrons revolve and spin.

WEAKNESS OF RUTHERFORD ATOMIC MODEL:

According to the classical electromagnetic theory if a charged particle accelerate around on oppositely charge particle it will radiate energy. Its speed will decrease and it will go into spiral motion finally falling into the nucleus. Similarly if an electron moving through Orbital’s of ever decreasing radii would give raise to radiations of all possible frequencies. In other words, it would given rise to a continuous spectrum. In actual practice, atom gives discontinuous spectrum.

X-RAYS AND ATOMIC NUMBER:

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In 1895, W.Roentgen discovered that when high energy electrons from cathode collide with the anode in the Crooke’s tube, very penetrating is produced. These rays were named as X-Rays.

EXPLANATION:

When an electron coming from the cathode strike with anode in the crook’s tube. It can remove an electron from the inner shell of the atom. Due to removal of this electron, the electronic configuration of this ion is unstable and an electron from orbital of higher energy drops into inner orbital by emitting energy in form of photon. This photon corresponds to electromagnetic radiation in x-rays region.

RELATIONSHIP BETWEEN WAVE LENGTH & NUCLEAR CHARGE:

In 1911, Mosley established a relationship between the wave length and nuclear charge. He found that when cathode rays struck elements used as anode targets in the discharge tube, characteristic x-rays were emitted. The wave length of the x-rays emitted decreases regularly with the increase of atomic mass. On careful examination of his data, Mosley found that the number of positive charges on the nucleus increases from atom to atom by single electronic unit. He called the number of positive charges as the atomic number.

BOHR’S THEORY:

Rutherford’s model of atom fails to explain the stability of atom and appearance of the line spectra. Bohr in 1913 was first to present a simple model of atom which explained the appearance of line spectra.

Some of the postulates of the Bohr’s theory are given below:1: An atom has a number of stable orbits or stationary states in which an electron can reside without emission or absorption of energy.

2: An electron may pass from one of these non-radiating states to another of lower energy with the emission of radiations whose energy equals the energy difference between the initial and final states.

3: In any of these states the electrons move in a circular path about the nucleus.

4: the motion of electron in these states is governed by the ordinary laws of mechanics and electrostatic provided its angular momentum is an integral multiple of h/2π.

It can be written as:mvr = nh/2π

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Here mvr becomes the angular momentum of electron. Thus Bohr’s first condition defining the stationary states could be stated as:“Only those orbits were possible in which the angular momentum of the electrons would be an integral multiple of h/2π”.These stationary states correspond to energy levels in atom.

CALCULATION OF RADIUS OF ORBITS:

Consider an electron of charge e revolving.

Let ‘m’ be the mass of electron ‘r’ the radius of orbit and ‘v’ the tangential velocity of revolving electron.

The electrostatic force of attraction between the nucleus and electron according to coulomb’s law.

= Z e x e / r2

The centrifugal force acting on the electron.= mv2 / rBohr assumed that these two opposing faces must be balanced each other exactly to keep the electron in an orbit.Therefore,Zee / r2 = mv2 / rMultiply both side by rr × Ze2 / r2 = r × mv2 / rZe2 / r = mv2

Or r = Ze2 / mv2 ------------ (i)The Bohr’s postulate states that only those orbits are possible in which:mvr = nh / 2πTherefore, V = nh / 2πmrSubstituting the value of V in equation 1r = Ze2 / m (nh / 2πmr)2

Or r = Ze2 × 4π2 mr2 / n2h2

Or 1/r = 4π2mZe2 / n2h2

Or r = n2h2 / 4π2mZe2 --------------- (ii)This equation gives the radii of all possible stationary states. The values of constants present in this equation are as follows.H = 6.625 × 10-27 ergs.sec OR 6.625 × 10-37J.sMe = 9.11 × 10-28 gm OR 9.11 × 10-31 kgE = 4.802 × 10-10 e.s.n OR 1.602 × 10-19 CBy substituting these values we get for first shell of H atom.r = 0.529 × 10-8 m OR 0.529

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The above equation may also be written as:r = n2 (h2 / 4π2mZe2) = n2a0 -------------- (2)For the first orbit n = 1 and r = 0.529. This is the value of the terms in the bracket sometimes written as a0 called ‘Bohr’s radius. For second shell n = 2 and for third orbit n = 3 and so on.

HYDROGEN ATOM SPECTRUM:

BALMER SERIES:

The simplest element is hydrogen which contains only one electron in its valence shell.Ballmer in 1885 studied the spectrum of hydrogen. For this purpose he used hydrogen gas in the discharge tube. Ballmer observed that hydrogen atom spectrum consisted of a series of lines called Ballmer Series. Ballmer determined the wave number of each of the lines in series and found that the series could be derived by a simple formula.V = RH = 1.09677 × 10-5 cm-1

LYMAN SERIES:

Lyman series is obtained when the electron returns to ground. State i.e. n = 1 from higher energy level n2 = 2, 3, 4, 5, 6 etc. This series of lines belongs to the ultraviolet region of spectrum. It is represented by the following equation.V = RH (1 / 12 – 1 / n2

2)Where n2 = 2, 3, 4, 5 etc.

PASCHEN SERIES:

Paschen series is obtained when the electron returns to the 3rd shell i.e. n = 3 from higher energy levels n2 = 4, 5, 6 etc. This series belongs to Infrared region. The equation for paschen series may be written as:V = RH (1 / 32 – 1 / n2

2)

BRACKETT SERIES:

This series is obtained when an electron jumps from higher energy levels to 4th+ energy level. The equation for this series is:V = RH (1 / 42 – 1 / n2

2)

PFUND SERIES:Pfund series is obtained when an electron jumps from higher energy level to 5th energy level.

HEINSBERG UNCERTAINTY PRINCIPLE:

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According to Bohr’s theory an electron was considered to be a particle but electron also behaves as a wave according to de-Broglie.

Due to this dual nature of electron in 1925 Heinsberg gave a principle known as Heinsberg Uncertainty Principle which is stated as:

“It is impossible to calculate the position and momentum of a moving electron simultaneously”.

It means that if one was known exactly it would be impossible to know the other exactly. Therefore if the uncertainty in position is ∆x then according to this principle the product of these two uncertainties may written as:

∆px. ∆x = h

So if one of these uncertainties is known exactly then the uncertainty in its determination is zero and the other uncertainty will become infinite which is according to the principle.

ENERGY LEVELS AND SUB-LEVELS:

According to Bohr’s atomic theory, electrons are revolving around the nucleus in circular orbits which are present at different distance from nucleus. These orbits are associated with definite energy of electron increasing outwards from nucleus, so these orbits are referred as Energy levels or Shells.

These shells or energy levels are designated as 1, 2, 3, 4 etc Or K, L, M, N etc.The spectral lines which correspond to the transition of an electron from one energy level to another consists of several separate close lying lines as doubles, triples, and so on. It indicates that some of electrons of given energy level have different energies or the electron belonging to same energy level may differ in their energy. So the energy levels are accordingly divided into sub-energy levels which denoted by letters s, p, d, f (sharp, principle, diffuse and fundamental).

The number of sub-levels in a given level of energy or shell is equal to its value of n.e.g. in third shell where n=3 three sublevels s, p, d are possible.

QUANTUM NUMBERS:

There are four quantum numbers which describe the electron in an atom.

1: PRINCIPAL QUANTUM NUMBER:

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It is represented by “n”, which describes the size of orbital Or energy level. This energy level K, L, M, N, O correspond to n= 1, 2, 3, 4, 5 etc.If n = 1 then electron is in K shell. n = 2 then electron is in L shell. n = 3 then electron is in M shell.

2: AZIMUTHAL QUANTUM NUMBER:

This quantum number is represented by ‘l’, which describes the shape of the orbit. The value of Azimuthal quantum number may be calculated by a relation.l= 0 → n – 1So for different shell the value of l aren = 1 K shell l = 0n = 2 L shell l = 0, 1n = 3 M shell l = 0, 1, 2n = 4 N shell l = 0, 1, 2, 3when l = 0 the orbit is s.when l = 1 the orbit is p.when l = 2 the orbit is d.when l = 3 the orbit is f.

3: MAGNETIC QUANTUM NUMBER:

It is represented by ‘m’ and explains the magnetic properties of an electron. The value of m depends upon the value of ‘l’. It is given by:m = + l → 0 → -1where l = 1, m has three values (-1, 0, 1) which correspond to p orbital. Similarly,when l = 2, m has five values which correspond to d orbit.

4: SPIN QUANTUM NUMBER:

It is represented by s which represents the spin of a moving electron. This spin may be either clockwise or anticlockwise so the values for s may be +1/2, 0, -1/2.

PAULI’S EXCLUSION PRINCIPLE:

According to this principle,“No tow electrons in the same atom can have the same four quantum number”.Consider an electron is present in 1s orbital. For this electron n = 1, l = 0, m = 0. Suppose the spin of this electron is s = +1/2 which will be indicated by an upward arrow ↑. Now if another electron is put in the same orbital (1s) for that electron n=1, l = 0, m = 0. It can occupy this orbital only if the direction of its spin is opposite to that of the first electron so s= -1/2 which

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is symbolized by downward arrow ↓. From this example, we can observe the application of Pauli’s exclusion principle on electronic structure of atom.

SHAPES OF ORBITS:

The shape of s orbital is spherical so the probability of finding the electron is maximum because the electrons are uniformly distributed around the nucleus.The p orbital are dumbbell shaped. These are oriented in space along the three mutually perpendicular axis (x, y, z) and are called px, py, pz. these are degenerate orbital of equal energy. Each p orbital has two lobes. One of which is labeled (+) and other ( - ).

ELECTRONIC CONFIGURATION:

The distribution of electrons in the available orbital is proceeded according to these rules.1: Pauli exclusion principle2: Aufbau rule3: (n+l) rule4: Hund’s rule

AUFBAU PRINICIPLE:

It is stated as:“The orbital are filled up with electrons in the increasing Order of their energy”.It means that the orbital are filled with the electrons according to their energy level. The orbital of minimum energy are filled up first and after it the orbital of higher energy are filled.

HUNDS RULE:

If orbital of equal are provided to electron then electron will go to different orbital and having their parallel spin.

In other words we can say that electrons are distributed among the orbital of a sub-shell in such a way as to give the maximum number of unpaired electrons and have the same direction of spin. According to this rule the electronic configuration of some elements are given below:N = 1s2, 2s2, 2px

1. 2py1, 2pz

1

O = 1s2, 2s2, 2px2 . 2py

1, 2pz1

(n + 1) RULE:

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According to this rule:“The orbital with the lowest value of (n+1) fills first but when two orbital have the same value of (n+1) the orbital with the lower value of n fills first.For the electronic configuration the order of the orbital is as follows:1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s etc

ATOMIC RADIUS:

For homonuclear diatomic molecules the atomic radii may be defined as:“The half of the distance between the two nuclei present in a homonuclear diatomic molecule is called atomic radii”.

In case of hetronuclear molecule like AB, the bond length is calculated which is (rA + rB) and if radii of any one is known the other can be calculated.

For the element present in periodic table the atomic radius decreases from left to right due to the more attraction on the valence shell but it increases down the group with increases of the number of shells.

IONIC RADIUS:

Ionic radius is defined as:“The distance between nucleus of an ion and the point up to which nucleus has influence of its electron cloud”.When an electron is removed from a neutral atom the atom is left with an excess of positive charge called a cat ion. e.g.Na → Na+ + e-

But when an electron is added in a neutral atom a negative ion or onion is formed.Cl + e- → Clas the atomic radius, the atomic radii are3 known from α-rays analysis. The value of ionic radius depends upon the ions that surround it.Ionic radii if cations have smaller radii then the neutral atom because when an electron is removed. The effective charge on nucleus increases and pulls the remaining electrons with a greater force.Ionic radii of an ions have a larger radii then the neutral atom because an excess of negative charge results in greater electron repulsion.Radius of Na- atom = 1.57Radius of Na+ atom = 0.95 (smaller than neutral atom)Radius of Cl atom = 0.99Radius of Cl- atom = 1.81 (larger than neutral atom)

IONIZATION POTENTIAL:

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DEFINATION:

The amount of energy required to remove most loosely bounded electron from the outermost shell of an atom in its gaseous state is called ionization potential energy.It is represented as:M(gas) → M+

(gas) + e- The energy required to remove first electron is called first I.P. The energy required to remove 2nd or 3rd electron is called 2nd I.P or 3rd I.P.M(gas) → M+

(gas) + e- ∆E = 1st I.PM+

(gas) → M++(gas) + e- ∆E = 2nd I.P

M++(gas) → M+++

(gas) + e- ∆E = 3rd I.P

FACTORS ON WHICH I.P DEPENDS:

SIZE OF THE ATOM:

If the size if an atom is bigger than the I.P of atom is low, but if the size of atom is smaller than I.P will be high, due to fact if we move down the group in periodic table. The I.P value decreases down the group.

MAGNITUDE OF NUCLEAR CHARGE:

If nuclear charge of atom is greater than force of attraction on the valence electron is also greater so the I.P value for atom is high therefore as we move from left to right in periodic table the I.P is increased.

SCREENING EFFECT:

The shell present between the nucleus and valence electrons also decreases the force of attraction due to which I.P will be low for such elements.

ELECTRON AFFINITY:

DEFINATION:

The amount of energy liberated by an atom when an electron is added in it is called electron affinity.It shows that this process is an exothermic change which is represented as:Cl + e- → Cl- ∆H = =348kj/mole

FACTORS ON WHICH ELECTRON AFFINITY DEPENDS:

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SIZE OF ATOM:

If the size of atom is small, force of attraction from the nucleus on the valence electron will be high and hence the E.A for the element will also be high but if the size of atom is larger, the E.A for these atoms will be low.

MAGNITUDE OF NUCLEAR CHARGE:

Due to greater nuclear charge the force of attraction on the added electron is greater so E.A of the atom is also high.

ELECTRONIC CONFIGURATION:

The atom with stable configuration has no tendency to gain an electron so the E.A of such elements is zero. The stable configuration may exist in following cases.

Inert gas configuration Fully filled orbital Half filled orbital

ELECTRONEGATIVITY:The force of attraction by which an atom attracts a shared pair of electrons is called electro negativity.

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CHAPTER # 4CHEMICAL BOND

INTRODUCTION:

Atom of all the elements exert noble gases have incomplete outermost orbits and tends to complete them chemical combination with the other atoms.

In 1916, W. Kossel described the ionic bond which is formed by the transfer of electron from one atom to another and also in 1916 G.N. Lewis described about the formation of covalent bond which is formed by the mutual sharing between two atoms.

Both these scientists based their ideas on the fact that atoms attain greatest stability when they acquire an inert gas electronic configuration.

CHEMICAL BOND:

When two or more than two atoms are combined with each other in order to complete their link between them is produced which is known as “Chemical Bond”.

ORThe force if attraction which holds atoms together in the molecule of a compound is called “Chemical Bond”.

TYPES OF CHEMICAL BOND:

There are three main types of chemical bond.1: Ionic bond or electrovalent bond2: Covalent bond3: Co-ordinate covalent bond or Dative covalent bond

IONIC BOND OR ELECTROVALENT BOND:

A chemical bond which is formed by the complete shifting of electron between two atoms is called “Ionic bond” OR “Electrovalent bond”.

ORThe electrostatic attraction between positive and negative ions is called “Ionic bond”.

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CONDITIONS FOR THE IONIC BOND FORMATION:

1: ELECTRONEGATIVITY:

Ionic bond is formed between the elements having a difference of electronegative more than 1.7 or equal to 1.7 eV.Therefore ionic bond is generally formed between metals (low electronegative) and non-metal (high electronegative) elements.

2: IONIZATION POTENTIAL:

We know that ionic bond is formed by the transference of electron from one atom to another, so in the formation of ionic bond an element is required which can lose its electron from the outer most shell, it is possible to remove electron from the outer most shell of metals because of their low ionization potential values.

3: ELECTRON AFFINITY:

In the formation if ionic bond an element is also required which can gain an element is also required which can gain electron. Since non-metals can attract electrons with a greater force due to high electro negativity. So a non-metal is also involved in the formation of ionic bond due to high electron affinity.

EXAMPLE OF IONIC BOND:

In order to understand ionic bonds consider the example of NaCl.During the formation of ionic bond between Na and Cl2, Sodium loses one electron to form Na+ ion while chlorine atom gains this electron to form Cl- ion. When Na+ ion and Cl- ion attract to each other NaCl is formed. The stability if NaCl is due to the decrease in the energy. These energy changes which are involved in the formation of ionic bond between Na and Cl are as follows:

i. Sodium has one valence electron. In order to complete its octet Na loses its valence electron. The loss of the valence electron required 495 KJ/mole.

Na → Na+ + e- ∆H = 495 KJ/moleii. Chlorine atom has seven electrons in its valence shell. It requires only one electron to

complete its octet, so chlorine gains this electron of sodium and release -348 KJ/mole energy.

Cl + e- → Cl- ∆H = -348 KJ/moleHere the energy difference is 147 KJ/mole (495 – 348 = 147). This loss of energy is balanced when oppositely charged ions are associated to form a crystal lattice.iii. In third step, positively charged Na+ ion and negatively charged Cl- ion attract to each

other and a crystal lattice is formed with a definite pattern.Na+ + Cl- → Na+Cl- ∆H = -788 KJ/mole

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This energy which is released when one mole of gaseous ions arrange themselves in definite pattern to form lattice is called Lattice energy.From this example: We can conclude that it is essential for the formation of ionic bond that the sum of energies released in the second and third steps must be greater than the energy required for the first step.

CHARACTERISTICS OF IONIC COMPOUNDS:

1: An ionic compound the oppositely charged ions are tightly packed with each other. So, these compounds exist in solid state.2: Due to strong attractive forces between ions a large amount of energy is required to melt or to boil the compound and hence the melting and boiling points of ionic compound are generally high.3: Ionic compounds are soluble in water but insoluble in organic solvents like benzene CCl2

etc.4: In the aqueous solution, the ionic compounds are good electrolytes, because in water the interionic forces are so weakened that the ions are separated and free to move under the influence of electric current. Due to this free movement of ions, the ionic compounds conduct electricity in their solutions.

COVALENT BOND:

A link which is formed by the mutual sharing of electrons between two atoms is called Covalent bond.

Explanation:

In the formation of covalent bond, mutual sharing of electron takes place. This mutual sharing is possible in non-metals, therefore covalent bond is generally formed between the atoms of non-metals.

For example:

In Cl2 molecule, two atoms of chlorine are combined with each other to form Cl2 molecule. Each atom of chlorine having seven electrons in its valence shell. These atoms are united with each other by sharing one of its valence electron as shown:

:Cl Cl: → :Cl :Cl OR Cl - Cl

In this molecule, one shared pair of electrons forms a single covalent bond between two chlorine atoms with the formation of a covalent bond the energy of the system is also decreased.Cl + Cl → Cl – Cl ∆H = -242KJ/moleThis released energy lowered the energy of the stability of the compound is also increased.

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TYPES OF COVALENT BOND:

There are three main types of covalent bond.

1: SINGLE COVALENT BOND:

When a covalent bond is formed by sharing of one electron from each atom, then it is called Single Covalent bond and denoted by ( − ) single line between the two bonded atoms.e.g. Cl – Cl, H – H, H – Br etc.

2: DOUBLE COVALENT BOND:

In a covalent bond, if two electrons are shared from each of the bonded atom then this covalent bond is called Double Covalent bond and denoted by (═) two lines.e.g. O ═ O O: :O

3: TRIPLE COVALENT BOND:

When a covalent bond is formed by sharing of three electrons from each atom then this type of covalent bond is called Triple Covalent bond and denoted by (≡) three lines between the two bonded atoms.e.g. N ≡ N N: :NThe bond distance of multiple bonds is shorter and the bond energies are higher.

CHARACTERISTICS OF COVALENT COMPOUND:

The main characteristic properties of covalent compounds are as follows:1: The covalent compound exists as separate covalent molecules, because the particles are electrically neutral so they possess solid, liquid or gaseous state. This intermolecular force of attraction among the molecules.2: Since the covalent compound exist in all the three states of matter so their melting points and boiling points may be high or low.3: Covalent compounds are non-electrolytes so they do not conduct electricity from their aqueous solution.4: Covalent compounds are generally insoluble in water and similar polar solvent but soluble in the organic solvents.

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CO-ORDINATE OR DATIVE COVALENT BOND:

It is a type of covalent bond in which both the shared electrons are donated only by one atom, this type of bond is called co-ordinate covalent bond.The co-ordinate covalent bond between two atoms is denoted by an arrow (→). The atom which donates an electron pair is called as a donor of electron and the other atom involved in this bond is called acceptor.e.g. A + B → A : B OR A → B

EXAMPLES OF CO-ORDINATE COVALENT BOND:

1: FORMATION OF NH4+ RADICAL:

Electronic configuration of nitrogen is 1s2, 2s2, 2p3. So nitrogen contains five electrons in the valence shell which are shared with three hydrogen atom to form three covalent bond. HH : N : HWhen this molecule of NH3

is reacted with H+ a co-ordinate covalent bond is formed. H H HH : N : + H+ → H : N : H → H – N − H H H H

PHOSPHORUS OXYCHLORIDE POCl3:

Electronic configuration of P is = 1s2, 2s2, 2p6, 3s2, 3p3. Therefore phosphorus atom contains five electrons in the outermost shell.The three unpaired electron of P are shared with three chlorine atoms.:Cl :P. . Cl: :Cl:When an oxygen atom is attached with this molecule the lone pair of electron is provided by the Phosphorus atom and a co-ordinate covalent bond is formed. :O: O:Cl :P. . Cl: :O: → :Cl :P. .Cl: → Cl−P−Cl :Cl: :Cl: Cl

DIPOLE MOMENT:

The product of the charge and the distance present in a polar molecule is called Dipole Moment and represented by μ.

ORThe extent of tendency of a molecule to be oriented under the influence of an electric field is called Dipole Moment.

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MATHEMATICAL REPRESENTATION OF DIPOLE MOMENT:

Suppose the charge present on a polar molecule is denoted by ‘e’ and the separation between the two oppositely charged poles of the molecule is‘d’, and then the produce of these two may be written as:E × = μWhere μ is dipole moment.

DIPOLE MOMENT IN DIATOMIC MOLECULES:

The diatomic molecules which are made up of similar atoms will be non-polar and their dipole moment is zero but the diatomic molecules made up of two different atoms e.g. HCl or HI are polar and have some dipole moment. The value of the dipole moment depends upon the difference of electro negativities of the two bonded atoms. If the difference of electro negativity between the atoms is greater, the polarity and also the dipole moment of the molecule is greater e.g.The dipole moment of HCl = 1.03 debyeWhereas dipole moment of HF = 1.90 debye

DIPOLE MOMENT OF POLY ATOMIC MOLECULES:

In poly atomic molecules, the dipole moment of molecules depends upon the polarity of the bond as well as the geometry of the molecule. Some examples for dipole moment in poly atomic molecule are given below.

1: DIPOLE MOMENT IN CO2:

The molecular structure of Co2 is linear in which the angle between oxygen carbon and oxygen is 180º as shown in figure.O5 C5 =The dipole moment for each C=O bond is 2.3 Debye, but due to the linear structure of the molecule, the dipole moment of the two C = O bonds cancel each other and net dipole moment of the molecule is zero.

2: DIPOLE MOMENT IN CS2:

The molecular structure of CS2 is also linear as shown.S-8 = C+8 = SDue to this linear structure of the dipole moment of C = S bond of one side is cancelled by the dipole moment of other C = S so the dipole moment of the molecule is zero.3: DIPOLE MOMENT OF H2O:

The molecular structure of H2O is angular as shown.

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O-5

H+5 H+5

Due to this angular structure the two O – H bonds do not cancel the dipole moment of each other. Therefore H2O molecule has a net dipole moment.

UNITS OF DIPOLE MOMENT:

In S.I units, dipole moment is measured in coulomb-meter (c – m). However commonly used unit is Debye (D).

IONIC CHARACTER OF COVALENT BOND:

In homonuclear diatomic molecules like Cl2, O2, I2, H2 both the atoms are identical so the shared electrons are equally attracted due to identical electro negatives and hence the molecules are non-polar.

When two dissimilar atoms are linked by a covalent bond the shared electrons are not attracted equally by the two bonded atoms. Due to unsymmetrical distribution of electrons one end of the molecules acquire partial positive charge and the other end acquire a partial negative charge. This character of a covalent bond is called Ionic Character of a covalent bond.

The ionic character of a covalent bond depends upon the difference of electro negativity of the two dissimilar atoms joined with each other in a covalent bond. E.g. the H – F bond is 43% ionic whereas the H – Cl bond is 17% ionic. The ionic character greatly affects the properties of a molecules. E.g. melting points, boiling point of polar molecules are high and they are soluble in polar solvent like H2O. Similarly the presence of partial polar character shortens the covalent bond and increases the bond energies.

BOND ENERGY:

The amount of energy required to break a bond between two atoms in a diatomic molecule is known as Bond Energy.

ORThe energy releasing in forming a bond from the free atoms is also known as Bond Energy.It is expressed in Kilo Joules per mole or KCl/mole.

EXAMPLES:

1: The bond energy for hydrogen molecule is:H(g) + H(g) → 2 H(g) ∆H = -435 KJ/moleORH(g) + H(g) → H – H ∆H = -345 KJ/mole

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Its can be observed from this example that the breaking of bond is endothermic whereas the formation of the bond is exothermic.2: The bond energy for oxygen molecule is:O ═ O(g) → 2O(g) ∆H = 498 KJ/moleO(g) + O(g) → O ═ O ∆H = -498 KJ/moleBond energy of a molecule also measures the strength of the bond. Generally bond energies of polar bond are greater than pure covalent bond.Cl – Cl → 2Cl ∆H = 244 KJ/moleH – Cl → H+ + Cl- ∆H = 431 KJ/moleThe value of bond energy also depends upon the bond and greater would be the bond energy e.g. triple bonds are usually shorter than the double bond therefore the bond energy for triple bond is greater than double bond.

SIGMA AND Pi BOND:

When the two orbital which are involved in a covalent bond are symmetric about an axis then the bond formed between these orbital is called Sigma Bond.

ORA bond which is formed by head to head overlap of atomic orbital is called Sigma Bond.

EXPLANATION:

In the formation of a sigma bond the atomic orbital lies on the same axis and the overlapping of these orbital is maximum therefore, all such bonds in which regions of highest density around the bond axis are termed as sigma bond.

TYPES OF OVERLAPPING IN SIGMA BOND:

There are three types of overlapping in the formation of sigma bond.1: S – S orbital overlapping2: S – P orbital overlapping3: P – P orbital overlapping

In all the three types, when the two atomic orbital are overlapped with each other two molecular orbital are formed, In these two molecular orbital the energy if one orbital is greater than the atomic orbital which is known as Sigma ant bonding orbital while the energy of the other orbital is less than the atomic orbital this orbital of lower energy is called Sigma bonding orbital and the shared electron are always present in the Sigma bonding orbital.

1: S – S ORBITAL OVERLAPPING:

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In order to explain S – S overlapping consider the example of H2 molecule in this molecule 1s orbital of one hydrogen overlaps with 1s orbital of other hydrogen to form sigma bonding orbital. Due to this bonding a single covalent bond is formed between the two hydrogen atoms. The two molecular orbital formed due to this overlapping are shown below.

2: S – P ORBITAL OVERLAPPING:

This type of overlapping takes place in H – Cl molecule. 1s orbital of hydrogen overlaps with 1p orbital of chlorine to form a single covalent bond. In this overlapping two molecular orbital are formed, one of the lower energy while the other orbital is of higher energy. The shapes of these orbital are as follows:

3: P – P ORBITAL OVERLAPPING:This type of overlapping takes place in fluorine molecule. In this mole 1p orbital of a fluorine atom is overlapped with 1p orbital of the other fluorine atom. The molecular orbital formed in this overlapping are given in the figure.

Pi BOND:

When the two atomic orbital involved in a covalent bond are parallel to each other then the bond formed between them is called Pi bond.In this overlapping, two molecular orbital are also formed the lower energy molecular orbital is called π bonding orbital while the higher energy molecular orbital is called π anti bonding orbital. The shape of these molecular orbital are as follows:

HYBRIDIZATION:

The process in which atomic orbital of different energy and shape are mixed together to form new set of equivalent orbital of the same energy and same shape.There are many different types of orbital hybridization but we will discuss. Here only three main types.

1: Sp 3 HYBRIDIZATION:

The mixing of one S and three p orbital to form four equivalents Sp3 hybrid orbital is called Sp3 hybridization. These Sp3 orbital are directed from the center of a regular tetrahedron to its four corners. The angles between tetrahedrally arranged orbital are 109.5º.Consider the example of CH4 molecule. The electronic configuration of C atom is1s2, 2s2, 2p1x, 2p1y, 2pzIt has two partially filled 2p orbital which indicate that it is divalent but carbon behaves as tetravalent in most of its compounds. It is only possible if one electron from 2s orbital is promoted to an empty 2pz orbital to get four equivalent Sp3 hybridized orbital.

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The four Sp3 hybrid orbital of the carbon atom overlap 1s-orbital of four hydrogen atoms to form a methane CH4 molecule.

2: Sp 2 HYBRIDIZATION:

The mixing of one S and two p orbital to form three orbital of equal energy is called Sp 2 or 3Sp2 hybridization. Each Sp2 orbital consists of S and P in the ratio of 1:2. These three orbital are co-planner and at 120ºangle as shown.

A typical example of this type of hybridization is of ethylene molecule. In ethylene, two Sp2

hybrid orbital of each carbon atom share and overlap with 1s orbital of two hydrogen atoms to form two σ bonds. While the remaining Sp2 orbital on each carbon atom overlaps to form a σ bond. The remaining two unhybridized p orbital (one of each) are parallel and perpendicular to the axis joining the two carbon nuclei. These generate a parallel overlap and results in the formation of 2π orbital. Thus a molecule of ethylene contain five σ bonds and one π bond.

3: Sp HYBRIDIZATION:

When one S and one p orbital combine to give two hybrids orbital the process is called Sp hybridization. The Sp hybrid orbital has two lobes one with greater extension in shape than the other and the lobes are at an angle of 180º from each other. Its means that the axis of the two orbital form a single straight line as shown.Now consider the formation of acetylene molecule HC ≡ Ch. The two C – H σ bonds are formed due to Sp – S overlap and a triple bond between two carbon atoms consists of an σ bond and two π bonds. The sigma bond is due to Sp – Sp overlap whereas π bonds are formed as a result of parallel overlap between the unhybridized four 2p orbital of the two carbon.

VALENCE SHELL ELECTRON PAIR REPULSION THEORY:

The covalent bonds are directed in space to give definite shapes to the molecules. The electron pairs forming the bonds are distributed in space around the central atom along definite directions. The shared electron pairs as well as the lone pair of electrons are responsible for the shape of molecules.

Sidwick and Powell in 1940 pointed out that the shapes of the molecules could be explained on the basis of electron pairs present in the outermost shell of the central atom. Pairs of electrons around the central atom are arranged in space in such a way so that the distances between them are maximum and coulumbic repulsion of electronic cloud are minimized.The known geometrics of many molecules based upon measurement of bond angles shows that lone pairs of electrons occupy more space then bonding pairs. The repulsion between electronic pairs in valence shell decreases in the following order.

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Lone pair – Lone pair > Lone pair – Bond pair > Bond pair – Bond pairWhen we apply this theory we can see the variation of angle in the molecular structures.Consider the molecular structures of NH3, CH4 and H2O.

Variations from ideal bond angles are caused by multiple covalent bonds and lone electron pairs both of which require space than single covalent bonds are therefore cause compression of surrounding bond angles.

Thus the number if pairs of electrons in the valence shell determine the overall molecular shape.

STRUCTURE OF BeCl2:

The two bond pairs of electrons is BeCl2 arrange themselves as far apart as possible in order to minimize the repulsion between them.The only arrangement which can given to satisfy this condition is linear.

STRUCTURE OF BF3 OR BCl3:

In this molecule three bond pairs are present around boron, to arrange themselves as far apart as possible a trigonal structure is fromed

It should therefore be a planar triangular molecule.

HYDROGEN BOND:

When hydrogen is bonded with a highly electro negative element such as nitrogen oxygen, fluorine. The molecule will be polarized and dipole is produced. The slightly positive hydrogen atom is attracted by the slightly negatively charged electro negative atom. An electrostatic attraction between the neighboring molecules is setup when the positive pole of one molecule attracts the negative pole of the neighboring molecule. This type of attractive force which involved hydrogen is known as Hydrogen bonding.

Hydrogen bonding causes the association of molecules with each other.Example of hydrogen bonding are H2O and HF.

The hydrogen bonding is the strongest of the secondary bond, but it is much weak than a normal covalent bond.The hydrogen bonding greatly affects the physical properties of the molecules. Due to this bond the melting point, boiling point, density and viscosity the compounds are increased.

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CHAPTER # 5ENERGETICS OF CHEMICAL REACTION

THERMODYNAMICS:

It is a branch of chemistry which deals with the heat energy change during a chemical reaction.

TYPES OF THERMODYNAMICAL REACTION:

Thermo-chemical reactions are of two types.1: Exothermic Reactions2: Endothermic Reactions

1: EXOTHERMIC REACTION:

A chemical reaction in which heat energy is evolved with the formation of product is known as Exothermic reaction.An exothermic process is generally represented as:Reactants → Products + Heat

2: ENDOTHERMIC REACTION:

A chemical reaction in which heat energy is absorbed during the formation of product is known as Endothermic reaction.Endothermic reaction is generally represented as:Reactants + Heat → Products

THERMODYNAMIC TERMS:

1: SYSTEM:

Any real or imaginary portion of the universe which is under consideration is called System.

2: SURROUNDINGS:

All the remaining portion of the universe which is present around a system is called Surroundings.3: STATE:

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The state of a system is described by the properties such as temperature pressure and volume when a system undergoes a change of state, it means that the final description of the system is different from the initial description of temperature, pressure or volume.

PROPERTIES OF SYSTEM:

The properties of a system may be divided into two main types.

1: INTENSIVE PROPERTIES:

These properties which are independent of the quantity of matter are called intensive properties.e.g. melting point, boiling point, viscosity, surface lension, refractive index etc.

2: EXTENSIVE PROPERTIES:

Those properties which depend upon the quantity of matter are called extensive properties.e.g. mass, volume, enthalpy, entropy etc.

FIRST LAW OF THE THERMODYNAMICS:

This law was given by Helmholtz 1917. According to this law:“Energy can neither be created nor destroyed but it can be changed from one form to another”.In other words the total energy of a system and surroundings must remain constant.

MATHEMATICAL DERIVATION OF FIRST LAW OF THERMODYNAMICS:

Consider a gas is present in a cylinder which contains a frictionless piston as shown.Let a quantity of heat q is a provided to the system from the surrounding. Suppose the internal energy of the system is E and after the absorption of q amount of heat it changes to E2. Due to the increase of this internal energy the collisions offered by the molecules also increases or in other words the internal pressure of the system is increased after the addition of q amount of heat with the increase of internal pressure the piston of the cylinder moves in the upward direction to maintain the pressure constant so a work is also done by the system.Therefore if we apply first law of thermodynamic on this system we can write.

q = E2 – E1 + Wq = ∆E + WOR∆E = q – w

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This is the mathematical representation of first law of thermodynamics.

PRESSURE – VOLUME WORK:

Consider a cylinder if a gas which contain a frictionless and weightless piston, as shown above. Let the area of cross-section of the piston = aPressure on the piston = PThe initial volume of the gases = V1

And the final volume of the gases = V2

The distance through which piston moves = dSo the change in volume = ∆V = V2 – V1

OR∆V = A × dThe work done by the system.W = Force × distance

But pressure =

Or

P =

F = PAPut in equation 1W = P.A × dHereW = pressure × (length × width × height)Pressure change of volumeSoW = P × ∆VIt is called pressure work volume.By substituting the value of work. The first law of thermodynamic may be written as:∆E = q – w∆E = q - P∆VQ = ∆E + P∆VThe absorption or evolution of heat during chemical reaction may take place in two ways.1: Process at constant volume (Isochoric process).2: Process at constant pressure (Isoboric process).

ISOCHORIC PROCESS:

A process in which volume remains constant is called Isochoric process.Let qv be the amount of heat absorbed at constant volume.

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According to first law of thermodynamic.Q = ∆E + P∆VBut volume is constant.v1 = v2 = vSoqv = ∆E + P∆Vqv = E2 – E1 + P(V – V )qv = ∆E + P (O)qv = ∆EAs constant volume just internal energy of the system will be change when heat supplied by the system.

ISOBORIC PROCESS:

A process in which pressure will remain constant is called Isoboric process.Let qp is the amount of heat energy provided to system at constant pressure.Due to this addition of heat the internal energy of the gas is increase from E1 to E2 and volume is changed from V1 to V2.So,According to first law of thermodynamic.qp = ∆E + P∆VBut pressure is constantP = PSoqp = E2 – E1 + P (V2 – V1)qp = E2 – E1 + PV2 – PV1

qp = E2 _ PV2 – E1 – PV1

qp =( E2 + PV2) – (E1 + PV1) --------- (2)HereH = E + PVThe sum of the product of PV and energy is called Enthalpy and it is denoted by H.SoH1 = E1 + PVH2 = E2 + PV2

Put it in equation 2qp = H2 – H1

qp = ∆HThe heat at constant pressure is called Enthalpy.

SIGN OF ∆H:

All represent the change of enthalpy. It is a characteristic property of a system which depends upon the initial and final state of the system.

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For all exothermic process ∆H is negative and for all endothermic reactions. All is positive.

THERMOCHEMISTRY:

It is a branch of chemistry which deals with the measurement of heat evolved or absorbed during a chemical reaction.The unit of heat energy which are generally used are Calorie and Kilo Calorie or Joules and Kilo Joules.1Cal = 4.184J OR 1Joule = 0.239 Cal

HESS’S LAW OF CONSTANT HEAT SUMMATION:

STATEMENT:

If a chemical reaction is completed in a single step or in several steps the total enthalpy change for the reaction is always constant.ORThe amount of heat absorbed or evolved during a chemical reaction must be independent of the particular manner in which the reaction takes place.

EXPLANATION:

Suppose in a chemical reactant A changes to the product D in a single step with enthalpy change ∆H.

This reaction may proceed through different intermediate stages i.e. A first changes to B with enthalpy change ∆H1 then B changes to C with enthalpy change ∆H2 and finally C changes to D with enthalpy changes ∆H3.According to Hess’s law:∆H = ∆H1 + ∆H2 + ∆H3

VERIFICATION OF HESS’S LAW:

When CO2 reactants with excess of NaOH sodium carbonate is formed with the enthalpy change of 90KJ/mole. This reaction may take place in two step via Sodium bicarbonate as shown in figure.In the first step for the formation of NaHCO3 the enthalpy change is -49KJ/mole and in the second step the enthalpy change is =41KJ/mole.

CO2 + 2NaOH Na2CO3 + NaOH

∆H = -49KJ/moleNaHCo3 + NaOH → Na2Co3

∆H = -41KJ/mole

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According to Hess’s law:∆H = ∆H1 + ∆H2

∆H = -41 – 49 = -90KJ/moleThe total energy change when the reaction is completed in a single step is -90KJ/mole which is equal to the enthalpy change when the reaction is completed into two steps. Thus the Hess’s law is verified from this example.

CHAPTER # 6CHEMICAL EQUILIBRIUM

REVERSIBLE REACTIONS:

“Those chemical reactions which take place in both the directions and never proceed to completion are called Reversible reactions”.For these types of reaction both the forward and reverse reaction occurs at the same time so these reaction are generally represented as:Reactant □ ProductThe double arrow □ indicates that the reaction is reversible and that both the forward and reverse reaction can occur simultaneously.Some examples of reversible reactions are given below.1: 2HI □ H2 + I2

2: N2 + 3H2 □ 2NH3

IRREVERSIBLE REACTIONS:

“Those reactions in which reactants are completely converted into product are called Irreversible reaction”.These reactions proceed only in one direction. Examples of such type of reaction are given below.1: NaCl + AgNO3 → AgCl + NaNO3

2: Cu + H2SO4 → CuSo4 + H2

EQUILIBRIUM STATE:

“The state which the rate of forward reaction becomes equal to the rate of reverse reaction is called Equilibrium State”.

EXPLANATION:

Consider the following reaction.A + B □ C + D

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It is a reversible reaction. In this reaction both the changes (i.e. forward and backward) occur simultaneously. At initial stage reactant A and B are separated from each other therefore the concentration of C and D is zero.

When the reaction is started and the molecules of A and B react with each other. The concentration of reactant is decreased while the concentration of product is increased with the formation of product, the reverse process of the reaction starts. Due to the decrease of the concentration of reactant, the rate of forward reaction decreases with time but the rate of reverse reaction is increased with the formation of product C and D.Ultimately a stage reaches when the number of reacting molecules in the forward reaction equalizes the number of reacting molecules in the reverse direction. So this state at which the rate of forward reaction becomes equal to the rate of reverse reaction is called Equilibrium State.

LAW OF MASS ACTION:

STATEMENT:

“The rate of which a substance reacts is proportional to its active mass and the rate of a chemical reaction is proportional to the product of the active masses of the reactant”.The term “active mass” means the concentration in terms of mole/dm3.

DERIVATION OF EQUILIBRIUM CONSTANT EXPRESSION:

Consider in a reversible reaction “m” mole of A and “n” mole of B reacts to give “x” mole of C and “y” mole of D as shown in equation.mA + nB □ xC +yDIn this process:The rate of forward reaction = α [A]m [B]n

Or The rate of forward reaction = Kf [A]m [B]n

And The rate of reverse reaction = α [C]x [D]y

Or The rate of reverse reaction = Kf [C]x [D]y

But at equilibrium state:Rate of forward reaction = Rate of reverse reactionTherefore:Kf [A]m [B]n = Kf [C]x [D]y

Or

=

Kc =

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This is the expression for equilibrium constant which is denoted by Kc and defined as:The ratio of multiplication of active masses of the products to the product of active masses of reactant is called Equilibrium Constant.

EQUILIBRIUM CONSTANT FOR A GASEOUS SYSTEM:

Consider in a reversible process. The reactants and products are gases as shown:A(g) + B(g) □ C(g) + D(g)

When the reactants and products are in gaseous state, their partial pressures are used instead of their concentration, so according to law of mass action.

Kp =

Where PA, PB, Pc, PD are the partial pressures of reactant and product.

RELATIONSHIP BETWEEN Kc AND Kp:

There are three possibilities in the relationship between Kc and Kp which are as follows:1: Kc = Kp

In a chemical reaction if there is no change in number of moles before and after the reaction then Kc = Kp.e.g. H2 + I2 □ 2HI2: When reaction occur with increase in volume such as:2NH3 □ N2 + 3H2

For such reaction Kp > Kc

3: When there is decrease in volume on the product side e.g.O2 + 2So2 □ 2So3

In this case Kp , Kc.

DETERMINATION OF EQUILIBRIUM CONSTANT:

The value of equilibrium constant Kc does not depend upon the initial concentration of reactants. In order to find out the value of Kc, we have to find out the equilibrium concentration of reactant and product.

1: ETHYL ACETATE EQUILIBRIUM:

Acetic acid reacts with equal alcohol to form ethyl acetate and water as shown:CH3COOH + C2H5OH → CH3COOC2H5 + H2OSuppose ‘a’ moles of acetic acid and ‘b’ moles of alcohol are mixed in this reaction. After some time when the state of equilibrium is established. Suppose ‘x’ moles of H2O and ‘x’

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moles of ethyl acetate are formed while the number of moles of acetic acid and alcohol are a – x and b – z respectively at equilibrium so we can write as:CH3COOH + C2H5OH → CH3COOC2H5 + H2OInitial moles a b zero zeroMoles at a – x b – x x xEquilibriumConcentration a – x/v b – x/v x/v x/vAt equilibrium:According to law of mass action:Kc = [CH3COOC2H5] [H2O] / [CH3COOH] [C2H5OH]

Kc = [x/v] [x/v] / [a – x/v] [b – x/v]

Kc = (x) (x) / (a – x) (b – x)

2: HYDROGEN IODIDE EQUILIBRIUM:

For the reaction between hydrogen and iodine suppose ‘a’ mole of hydrogen and ‘b’ mole of iodine are mixed in a sealed bulb at 444ºC in the boiling sulphur for some time. The equilibrium mixture is then cooled and the bulbs are opened in the solution of NaOH. Let the amount of hydrogen consumed at equilibrium be ‘x’ moles which means that the amount of hydrogen left at equilibrium is a –x moles. Since 1 mole of hydrogen reacts with 1 mole of iodine to form two moles of hydrogen iodide hence the amount of iodine used is also x moles so its moles at equilibrium are b – x and the moles of hydrogen iodide at equilibrium are 2x so we can write as:H2 + I2 □ 2HIInitial molesa b zeroMoles at equilibriuma – x b – x 2xConcentration equilibriumA – x/v b – x/v 2x/vAccording to law of mass action:Kc = [III]2 / [H2] [I2]Kc = [2x/v]2 / [a – x/v] [b – x/v]

APPLICATIONS OF LAW OF MASS ACTION:

There are two important applications of equilibrium constant.1: It is used to predict the direction of reaction.2: Kc is also used to predict the extent of reaction.

TO PREDICT THE DIRECTION OF REACTION:

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The value of equilibrium constant Kc is used to predict the direction of reaction for a reversible process.Reactant □ ProductWith respect to the ratio of initial concentration of the reagent.[Product]initial / [Reactant]initial There are three possibilities for the value of K.

1: It is greater than Kc.2: It is less than Kc.3: It is equal to Kc.

CASE 1:

If [Reactant]initial [Product]initial > Kc the reaction will shift towards the reverse direction.

CASE 2:

If [Reactant]initial [Product]initial < Kc the reaction will shift towards the forward direction.

CASE 3:

If [Reactant]initial [Product]initial = Kc this is equilibrium state for the reaction.

TO PREDICT THE EXTENT OF REACTION:

From the value of Kc we can predict the extent of the reaction.If the value of Kc is very large e.g.For 203 □ 302 Kc = 1055

From this large value of Kc it is predicted that the forward reaction is almost complete.When the value of Kc very low e.g.2HF □ H2 + F2 Kc = 10-13

From this value it is predicted that the forward reaction proceeds with negligible speed.But if the value of Kc is moderate, the reaction occurs in both direction and equilibrium will be attained after certain period of time e.g. Kc forN2 + 3H2 □ 2NH3 is 10So the reaction occurs in both the direction.

LE CHATELIER’S PRINCIPLE:

STATEMENT:

“When a stress is applied to a system at equilibrium, the equilibrium position changes so as to minimize the effect of applied stress”

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The equilibrium state of a chemical reaction is altered by changing concentration pressure or temperature. The effect of these changes is explained by Le Chatelier.

EFFECT OF CONCENTRATION:

By changing the concentration of any substance present in the equilibrium mixture. The balance of chemical equilibrium is distributed. For the reaction,A + B □ C + D

Kc =

If the concentration of a reactant A or B is increased the equilibrium state shifts to right and yield of products increases.But if the concentration of C or D is increased then the reactions proceed in the backward direction with a greater rate or more A and B are formed.

EFFECT OF TEMPERATURE:

The effect of temperature is different for different type of reaction.For an endothermic reaction heat is absorbed for the conversion of reactant into product so if temperature during the reaction is increased then the reaction will proceed with a greater rate in forward reaction.Temperature increase → More products are formed

ENDOTHERMIC REACTION

Temperature decrease → More reactants are formedTemperature increase → More reactants are formed

EXOTHERMIC REACTION

Temperature decrease → More products are formed

EFFECT OF PRESSURE:

The state of equilibrium of gaseous reaction is disturbed by the change of pressure. There are three types of reactions which show the effect of pressure change.

1: When the number of moles of product are greater:

In a reaction such as:

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PCl5 → PCl3 + Cl2

The increase of pressure shifts the equilibrium towards reactant side.

1: When the number of moles of reactant are greater:

In a reaction such as:N2 + 3H2 → 2NH3

The increase of pressure shifts the equilibrium towards product side because the number of moles of product are less than the number of moles of reactant.

1: When the number of moles of reactants and products are equal:

In these reactions where the numbers of moles of reactant are equal to the number of moles of product the change of pressure does not change the equilibrium state.e.g. H2 + I2 □ 2HISince the number of moles of reactants and products are equal in this reaction so the increase of pressure does not affect the yield of HI.

IMPORTANT INDUSTRIAL APPLICATION OF LE CHATELIER’S PRINCIPLE:

HABER’S PROCESS:

This process is used for the production of NH3 by the reaction of nitrogen and hydrogen. In this process 1 volume of hydrogen at 500ºC and 200 to 1000 atm pressure in presence of a catalyst.N2 + 3H2 □ 2NH3 ∆H = -4.62KJ/mole

1: EFFECT OF CONCENTRATION:

The value of Kc for this reaction is:Kc = [NH3]2 / [N2] [H2]3

Increase in concentration of reactants which are nitrogen and hydrogen, the equilibrium of the process shifts towards the right so as to keep the value of K c= constant. Hence the formation of NH3 increases with the increase of the concentration of N2 or hydrogen.

2: EFFECT OF TEMPERATURE:

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It is an exothermic process, so heat is liberated with the formation of product. Therefore, according to Le Chatelier’s Principle at low temperature the equilibrium shifts towards right to balance the equilibrium state so low temperature favour the formation of NH3.

3: EFFECT OF PRESSURE:

The formation of NH3 proceeds with the decrease in volume. Therefore, the reaction is carried out under high pressure or in other words high pressure is favourable for the production of NH3.

CONTACT PROCESS:

The process is used to manufacture H2SO4 on large scale. In this process the most important step is the oxidation of SO2 to SO3 in presence of a catalyst Vanadium pentoxide.2So2 + O2 □ 2SO3 ∆H = -395KJ/mole

1: EFFECT OF CONCENTRATION:

The value of Kc for this reaction is:Kc = [SO3]2 / [SO2]2 [O2]Increase in concentration of SO2 or O2 shifts the equilibrium towards the right and more SO3

is formed.

2: EFFECT OF TEMPERATURE:

Since the process is exothermic, so low temperature will favour the formation SO3. The optimum temperature for this reaction is 400 to 450ºC.

3: EFFECT OF PRESSURE:

In this reaction decrease in volume takes place so high pressure is favourable for the formation of SO3.

COMMON ION EFFECT:

STATEMENT:

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“The process in which precipitation of an electrolyte is caused by lowering the degree of ionization of a weak electrolyte when a common ion is added is known as Common ion effect”.

EXPLANATION:

In the solution of an electrolyte in water, there exist equilibrium between the ions and the undissociated molecules to which the law of mass action can be applied.Considering the dissociation of an electrolyte AB we have:AB □ A+ + B-

And[A+] [B-] / [AB] = K (dissociation constant)If now another electrolyte yielding A+ or B+ ions be added to the above solution, it will result in the increase of concentration of the ions A+ or B- and in order that K may remain the same, the concentration AB must evidently increase. In other words “The degree of dissociation of an electrolyte containing a common ion”. This phenomenon is known as Common ion effect.

APPLICATION OF COMMON ION EFFECT IN SALT ANALYSIS:

An electrolyte is precipitated from its solution only when the concentration of its ions exceed from the solubility product. The precipitates are obtained when the concentration of any one ion is increased. Thus by adding the common ion, the solubility product can be exceeded.

In this solution Cu(OH)2 is a weak base while H2SO4 is a strong acid so the PH of the solution is changed towards acidic medium.

When Na2CO3 is dissolved in water, it reacts with water such as:Na2Co3 + 2H2O □ 2NaOH + H2Co3

In this solution H2Co3 which is weak acid and NaOH which is a strong base are formed. Due to presence of strong base the medium is changed towards basic nature.

SOLUBILITY PRODUCT:

When a slightly soluble ionic solid such as silver chloride is dissolved in water. It decomposition into its ions.AgCL □ Ag+ + Cl-

These Ag+ and Cl- ions from solid phase pass into solution till the solution becomes saturated. Now there exists an equilibrium between the ions present in the saturated solution and the ions present in the solid phase, thus:AgCL □ Ag+ + Cl-

Undissolved dissolvedApplying the law of mass action:Kc = [Ag+] [Cl-] / K’

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Or Kc × K’ = [Ag+] [Cl-]Or Ks.p = [Ag+] [Cl-]Where Ks.p is known as solubility product and defined as:“The product of the concentration of ions in the saturated solution of a sparingly soluble salt is called Solubility product”.The value of solubility product is constant for a given temperature.

CALCULATION OF SOLUBILITY PRODUCT FROM SOLUBILITY:

The mass of a soluble present in a saturated solution with a fixed volume of solvent is called Solubility, which is generally represented in the unit of gm/dm3 with the help of solubility we can calculate the solubility product of a substance e.g. the solubility of Mg(OH)2 at 25ºC is 0.00764gm/dm3. To calculate the Ks.p of Mg(OH)2. First of all we will calculate the concentration of Mg(OH)2 present in the solution.Mass of Mg(OH)2 = 0.00764gm/dm3

Mass of Mg(OH)2 = 0.00764/58 mole/dm3

= 1.31 × 10-4 mole/dm3

The ionization of Mg(OH)2 in the solution is as follows:Mg(OH)2 □ Mg+2 + 2OH-

And the solubility product for Mg(OH)2 solution one mole of MG++ ions are present while two moles of OH- ions are present, therefore in 1.31× 10-4 mole/dm3 solution of Mg(OH)2. The concentration of Mg+ is 1.31 × 10-4 mole/dm3 while the concentration of OH- is 2.62 × 10-8 mole/dm3. By substituting these value.Ks.p = [Mg+2] [OH-]2

= [1.31 × 10-4] [2.62 × 10-4]2

= 9.0 × 10-12mole3/dm9

So in this way the solubility product of a substance may be calculated with the help of solubility.

CALCULATION OF SOLUBILITY FROM SOLUBILITY PRODUCT:

If we know the value of solubility product, we can calculate the solubility of the salt.For example: The solubility product of PbCrO4 at 25ºC is 2.8 × 10-13moles/dm3.

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CHAPTER # 7HYDRATION

SOLUTIONS AND ELECTROLYTES

SOLUTION:

A homogeneous mixture or a single phase mixture of two or more substances is called solution.

SOLUTE:

The component of solution present in smaller amount is called solute.

SOLVENT:

The component of solution present in larger amount is called solvent

CONCENTRATION OF SOLUTION:

The amount (No. of Moles, No. of Atoms or No. of Molecules) of a solute present in a given amount of solvent or solution is called concentration of solution.

MOLAR CONCENTRATION (MORALITY):

Definition:

The number of moles of solute present per dm3 of solution is called molar concentration or morality of the solution.

The morality is denoted by “M” and is calculated as:

M = No. of Moles of Solute dm3 of solution

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EXAMPLE:

When 1 mole of HCl (i.e. 36.5 gm HCl) is present in 1 dm3 solution the morality of solution is 1M similarly if 0.5 moles of HCl (i.e. 18.25 gm HCl) is present in 1 dm 3 of solution the morality of solution is 0.5 M

MOLAL CONCENTRATION (MOLALITY):

Definition:

The number of moles of solute present per 1000 gm (or 1 kg) of solvent is called molality of the solution.

Molality of solution is denoted by “m” which is calculated as:

m = No. of moles of solute mass of solvent in kg

EXAMPLE:

When one mole of glucose (180 gm glucose) is dissolved in 1 kg of H2O the molality of solution is 1m similarly if 2 moles of glucose (360 gm glucose) is dissolved in 1 kg H2O the solution is said to be 2 molal.

HYDRATION:

Addition of water or association of water molecules with a substance without dissociation is called Hydration.Water is a good solvent and its polar nature plays very important part in dissolving substances. It dissolves ionic compounds readily.When an ionic compound is dissolved in water, the partial negatively charged oxygen of water molecule is attracted towards the cation ion similarly the partial positively charged hydrogen of water molecule is attracted towards the anions so hydrated ions are formed as shown in figure.

In solution, the number of water molecules which surround the ions is indefinite, but when an aqueous solution of a salt is evaporated, the salt crystallizes with a definite number of water molecules which is called as water of Crystallization. E.g. when CuSO4 recrystallized from its solution the crystallized salt has the composition CuSo4.5H2O. Similarly when magnesium chloride is re-crystallized from the solution, it has the composition MgCl2.6H2O. This

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composition indicates that each magnesium ion in the crystal is surrounded by size water molecules. This type of salts is called hydrated salts.It is observed experimentally that the oxygen atom of water molecule is attached with the cation of salt through co-ordinate covalent bond so it is more better to write the molecular formulas of the hydrated salts as given below.

[Cu(H2O)5]SO4 [Mg(H2O)6]Cl2

It is also observed that these compound exist with a definite geometrical structure e.g. the structure of [Mg(H2O6)Cl2 is octahedral and [Cu(H2O)4]+2 is a square planar as shown.

HYDROLYSIS:

DEFINITION:

“The reaction of cation or anion (or both) of a salt with water so as to change its pH is known Hydrolysis”.

SALIENT FEATURES:

1- Due to hydrolysis new compounds are formed.2- H – O bond of H2O is broken.

3- It is a chemical change.

4- Solution becomes either acidic or basic.

EXAMPLES:

1- When NH4Cl is added in water it forms a strong acid and a weak base. So the resultant solution becomes acidic.

NH4Cl + H – OH → NH4OH + HCl

2- When Na2CO3 id dissolved in water it forms a strong base and a weak acid so the solution has higher conc. of OH- and the pH becomes basic.

Na2CO3 + H – OH → NaOH + H2CO3

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THEORY OF IONIZATION:

INTRODUCTION:

This theory was put forward by a Swedish chemist Arrbenius in 1887. It accounts for electrical conductivity of electrolyte solution and it also describes the effect of ionization on colligative properties of a substance.

POSTULATES OF THEORY:

1: ELECTROLYTES:

Electrolytes are substances that contain charged particles called ions except H+, all positive ions are derived from metals and negative ions are derived from non-metals.

2: DISSOCIATION OF ELECTROLYTES:

Particles of electrolytes (acid, base, salt) dissociate into opposite charged ions on dissolution in water.e.g.H2SO4 → 2H+ + SO4

--

HCl → H+ + Cl-

NaOH → Na+ + OH-

3: EQUILIBRIUM BETWEEN IONIZED AND UNIONIZED MOLECULE:

The ionization is a reversible process. Ions recombine to form unionized molecules. So the solution contains ions of electrolyte together with unionized molecules.

H2SO4 → 2H+ + SO4--

DEGREE OF IONIZATION:

The external of dissociation of an electrolyte is called degree of ionization or dissociation. Greater the dissociation, greater will be conduction of electricity through solution.

Degree of dissociation =

It depends upon

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i. Nature of electrolyteii. Degree of dilutioniii. Temperature

i: NATURE OF ELECTROLYTE:

Strong electrolytes (HCl, NaCl) ionize almost completely, whereas weak electrolytes (CH3COOh, NH4OH) ionize partially.

ii: DEGREE OF DILUTION:

Ionization increases with increase in dilution.

iii: TEMPERATURE:

Ionization increases with increase in dilution.

5: NUMBER OF IONS AND SEED OF IONS:

Greater the number of ion and greater the speed of ions, more will be the electrical conductance.

6: EFFECT OF IONIZATION OF COLIGATIVE PROPERTIES:

Ionization is responsible for deviation in some colligative properties such as:a) Elecvation of boiling pointb) Depression of freezing pointc) Osmotic processd) Lowering of vapor pressure

For example:

1 mole of NaCl produce q mole of Na+ and 1 mole of Cl-. So the number of ions are twice the number of NaCl. Hence atomic pressure must also increase.This is valid when there is no interionic attraction and no hydration.

DRAW BACKS:

This theory cannot explain.i) Conductance of electrolytes in fused state.ii) Why conductance of strong electrolytes does not increase with dilution.

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ELECTRODE POTENTIAL:

DEFINATION:

“The potential difference between a metal and solution of its salt is called electrode potential”.

OR

“It is the measure of tendency of an electrode to loose e- (oxidation) or gain e- (reduction), when it is dipped in its salt solution”.

UNIT:

The unit of electrode potential is volt.

REPRESENTATION:

It is represented by ‘E’.

TYPES OF ELECTRODE POTENTIAL:

There are two types:i. Oxidationii. Reduction Potential

i: OXIDATION POTENTIAL:

DEFINATION:

“It is the measure of tendency of an electrode to loose e- or get oxidized when dipped in the solution of its salt”.It is denoted by Eºox

ii: REDUCTION POTENTIAL:

DEFINATION:

“It is the measure of tendency of an electrode gain e- or get reduced when dipped in its salt solution”.

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It is denoted by EºRED

EXAMPLE:

Eºox Zn / Zn++ = + 0.76ERED Zn++ / Z = - 0.76

FACTORS AFFECTING ELECTRODE POTENTIAL:

The magnitude of electrode potential depends on:i) Nature of electrodeii) Conc. of its ions in solutioniii) Pressure of gas (in case if gas electrodes)iv) Temperature

STANDARD ELECTRODE POTENTIAL:

DEFINATION:

“It is the potential of an electrode (or help cell) measured against standard hydrogen electrode at 25ºC, when ionic concentration of solution is 1M”.

SYMBOL:

It is represented by Eº.

STANDARD ELECTRODE:

Defination:

“A half cell used for measuring electrode potential of another half cell is called standard electrode”.

STANDARD HYDROGEN ELECTRODE (S.H.E):

DEFINITION:

“A half cell of hydrogen electrode used for measuring electrode potential of another half cell is called standard hydrogen electrode (S.H.E).

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CONSTRUCTION OF S.H.E:

A hydrogen electrode consists of a platinum plate immersed in 1 molar solution of H+. A current pure H2 is passed continuously through the solution under 1 atm. The platinum absorbs the gas on its surface and this hydrogen coated platinum acts as the hydrogen electrode.

POTENTIAL OF S.H.E:

Potential of S.H.E udner 1 atm, 25ºC and 1 molar H+ solution is assumed to be zero i.e.

i: FOR OXIDATION:

H2 → 2H+ + 2e- Eºox = 0.0 volt

ii: FOR REDUCTION:

2H+ + 2e- → H2 EºRED = 0.0 volt

i.e.We can say that,

Eº H2 / H+ = Eº H+ / H2 = 0.00 volt

IONIZATION OF STANDARD POTENTIAL OF ZINC:

The standard potential of zinc (Zn) can be determined by making following voltaic cell.

CONSTRUCTION OF CELL:

The cell is constructed as follow:i. One half cells is made of Zn plate immersed in 1M Zn++ solution at 25ºC.ii. The other half cell is S.H.E.iii. Electrodes of both half cells are externally connected by a wire, thorugh a volt

meter.iv. Solutions of both half cells are joined by KCl salt bridge.

OBSERVATION:

Spontaneous flow of electrons (e-) is observed from Zn electrode to hydrogen electrode. As a result voltmeter shows a reading of 0.76 volt.

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CONCLUSION:

It is concluded that:i. Zn is losing e-

, i.e. it oxidized or it is ANODE.ii. As e- are accumulated on zinc so its negatively charged.iii. Standard potential of hydrogen electrode is 0.00 volt hence the meter reading i.e.

0.76 volt is standard potential of Zn.

Standard Ox. Potential of Zn

OR

Eº Zn / Zn++ = 0.76 volt

CELL REACTIONS:

(i): AT ANODE (OXIDATION)

Zn(s) → Zn(aq)++ + 2e- Eºox = +0.76 volt

(ii): AT CATHODE (REDUCTION)

2H+(aq) + 2e- → H2(g) EºRED = 0.0 volt

(iii): COMPLETE CELL REACTION:

Zn(s) + 2H+ → Zn++ + H2

And

ºEcell = ºEox + ºERED

= 0.76v + 0.00v = 0.76v

CELL NOTATION:

Zn(s) ; Zn++ H+(aq) (1M) ; H2(g) (1atm) ; Pt

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ELECTRO-CHEMICAL SERIES: (E.C.S)

The list of elements or ions arranged in the decreasing order of their standard reduction potential values is called as Electro Chemical Series.

SOME USEFUL FACTS ABOUT (E.C.S):

1. Metals placed in the series above hydrogen act as anode and they are oxidized.2. Metals placed in the series below hydrogen act as cathode and they are reduced.3. If the sign of electrode potential is positive then the reaction is spontaneous and the

electrode act the anode.4. If the sign of electrode potential is negative then the reaction is non-spontaneous and

the electrode act as cathode.

OXIDATION NUMBER:

The formal charge (i.e. not real) on the atom in the compound or ion under consideration is known as Oxidation number or state.

OXIDATION:

Oxidation is a chemical change in which electrons are lost by an atom or group of atoms.

EXAMPLE:

Feº → Fe2+ + 2e-

In this reaction neutral iron atom has lost 2 electrons and has changed to ferrous ion so it is oxidation.OROxidation is a process in which oxidation number if an element is increased.

EXAMPLE:

Cº + Oº2 → C+4O2(-2)

Is an oxidation of carbon, since its oxidation number increases from zero to +4.

REDUCTION:

Reduction is a chemical change in which electrons are gained by an atom or group of atoms. Consider the following examples:

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EXAMPLE:

Cl2 + 2e- → 2Cl-

In this reaction, Cl2 gains two electrons and changes to Cl- ions it is therefore reduction.

OR

Reduction is a process in which oxidation number of an element is decreased.

EXAMPLE:

Hº2 + Brº2 → 2H+Br-

In the reduction of bromine as its oxidation number decreases from 0 to 1.

INDICATES:

Organic substance that change color in solution as the PH of the solution change are called Indicates.ORIndicates are complex molecules that are themselves weak acids (Phenolphthalein) or weak base (methyl orange).

USE:

Indicates are used to find the end point of an acid – base titration.

COLORS OF SOME INDICATORS IN ACID< BASE AND NEUTRAL CONDITION:

Indicators Color in Acid Color in Base Color

Litmus Red Blue PurpleMethyl Orange Red Yellow OrangePhenolphthalein Colorless Pink Colorless

Universal indicator Red Purple Green

CHOICE OF INDICATOR FOR DIFFERENT TYPES OF TITRATION:

1. Weak acid strong base Phenolphthalein2. Strong acid weak base Methyl Orange3. Strong acid strong base Any one

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EXPLANATION OF INDICATOR’S BEHAVIOR:

Indicator dissociates in a solution and gives an ion which has different color from the undissociated molecules of indicator.

TYPES OF INDICATORS:

There are three types of Indicators.1. Acid base indicator2. Redox indicator3. Precipitation indicator

1: ACID BASE INDICATORS:

Indicator used in acid base titration are called acid base indicator.

EXAMPLE:

Litmus, methyl orange, phenolphthalein and universal indicator.

2: REDOX INDICATOR:

Indicators used in redox reaction are called redox indicators

EXAMPLE:

Potassium Permanganate

3: PRECIPITATION INDICATOR:

An indicator, which reacts with reactant at the end point and form a precipitate is called Precipitation indicator.EXAMPLE:

Potassium Chromate

pH:

Logarithm of the reciprocal of hydrogen ion concentration is called pH.ORNegative logarithm of hydrogen ion concentration is called pH.

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MATHEMATICAL REPRESENTATION:

pH = log

pH = -log [H+]

EXPLANATION:

Pure water ionizes very slightly and is therefore a weak electrolyte.H2O → H+

+ OHThus equilibrium expression on the above reaction is given by:

Kc =

Kc [H2O] = [H+] [OH-]Kc [H2O] = KwWhere Kw is called the ionic product of water which is equal to 1.0 × 10-14 moles/dm3.Thus[H+] [OH-] = 1.0 × 10-14

[H+] = 10-7

[OH-] = 10-7

pH = - log [H+]pH = - log 10-7

pH = 7 log 10pH = 7(1)pH = 7pH = 7 for neutral solution

METHODS OF DETERMINATION OF PH OF A SOLUTION:

Ph of a solution can be measured by two methods.

1: BY USING UNIVERSAL INDICATOR:

By mixing together various indicators, a universal indicator has been obtained. This indicator indicates different colors of solution at different acidities. Universal indicator colors related to pH value are as follows:

RED ORANGE YELLOW GREEN BLUE PURPLE0 3 6 7 8 14

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pH of a solution is determined by adding a few drops of universal indicator solution. The pH can be obtained by comparing the color of solution from the color of the universal indicator as given above.

2: BY USING pH METER:

A pH meter has a glass electrode which when dipped into solution, will give direct reading of the pH value on the meter.

pOH:

Logarithm of the reciprocal of hydroxyl [OH-] ion concentration is called pOH.

OR

Negative logarithm of hydroxyl ion concentration is called pOH.

MATHEMATICAL REPRESENTATION:

pOH = log

pOH = - log [OH-]

BUFFER SOLUTION:

A solution in which tends to resist change in pH is called a Buffer Solution.

PRINCIPLE OF PREPRATION OF BUFFER SOLUTION:

Buffer solutions of specified pH values are prepared easily by half neutralizing solution of a weak acid with a strong base or a weak base with a strong acid.

1: PREPARATION OF BUFFER SOLUTION OF pH LESS THAN 7:

Such solutions can be prepared by mixing weak acids and their salts with strong bases.

EXAMPLES OF BUFFER ACTION:

A solution of CH3COOH and CH3COONa.

2: PREPARATION OF BUFFER SOLUTION OF pH GREATER THAN 7

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Such solution can be prepared by mixing weak bases and their salts with strong acids.

EXAMPLES OF BUFFER ACTION:

A solution of NH4OH and Nh4Cl.

APPLICATION OF BUFFER:

The use of buffer is an important part of many industrial processes. Examples are electroplating and the manufacture of leather, photographic materials & dyes.

The biological research culture media are generally buffered to maintain a constant pH is required for the growth of the bacteria being studied.

Buffers are used extensively in analytical chemistry and are used to calibrate pH. Human blood is buffered to pH of 7.4 by means of bicarbonates, phosphates and

complex protein systems.

BUFFER ACTION:

The property of a solution due to which it tends to maintain its pH when an acid or an alkali is added to it is called buffer action.

BUFFER CAPACITY:

The capacity of a solution to maintain definite pH is called buffer capacity.

NEUTRALIZATION:

Reaction in which are cid reacts with a base to form salt and water is called Neutralization.

EXAMPLE:

Acid + Base → Water + SaltHCl + NaOH → H2O + NaClORReaction of cation of an acid with anion of base to form water is called Neutralization.EXAMPLE:

H+ + OH- → H2O.

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CHAPTER # 8CHEMICAL KINETICS

INTRODUCTION:

The branch of physical chemistry which deals with the speed or rate at which a reaction occurs is called chemical kinetics.

The study of chemical kinetics therefore includes the rate of a chemical reaction and also the rate of a chemical reaction and also the factors which influence its rate.

SLOW AND FAST REACTION:

Those reactions for which short time is required to convert a reactant into product are called fast reaction but if more time is required for the formation of a product them the reactions are called slow reaction usually ionic reactions which involve oppositely charged ions in aqueous medium are very fast for example reaction between aqueous solution of NaCl and AgNo3

gives white precipitates of AgCl instantaneously

AgNo3 + NaCl AgCl + NaNo3

Such reactions are very fast these are completed in fractions of seconds but those reactions which involve covalent take place very slowly for example conversion of So2 into So3

2So2 + O2 2So3

It is a slow reaction and required more time for the formation of a product

RATE OR VELOCITY OF A REACTION:

DEFINITION:

It is the change in concentration of a reactant or product per unit time

Mathematically it is represented as

Rate of reaction = change in concentration of reactant/product Time taken for the change

The determination of the rate of a reaction is not so simple because the rate of a given reaction is never uniform it falls of gradually with time as the reactants are used up hence we cannot get the velocity or rate of reaction simply by dividing the amount of substance

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transformed by the time taken for such transformation for this reason we take a very small interval of time “dt” during which it is assumed that velocity of reaction remains constant if “dx” is the amount of substance transformed that small interval of time “dt” then the velocity of reaction is expressed as.

Velocity of a reaction = dx / dt

Thus with the velocity of a chemical reaction use mean the velocity at the given moment or given instant.

THE RATE CONSTANT:

DEFINITION:

The proportionality constant present in the rate equation is called rate constant according to law of mass action we know that the rate of chemical reaction is directly proportional to the molar concentration of the reactions for example

R P(Reactant) (Product)

The rate of reaction α [R]Or dx/dt =K [R]

Where K is known as rate constant

SPECIFIC RATE CONSTANT:

When the concentration and temperature both are specified the rate constant is known as specific rate constant. When the concentration of each reactant is mode per dm3 at given temperature the specific rate constant numerically equals to the velocity of the reaction

dx/dt = V = K [R]or K V/[R]

when R = 1 mole/dm3

K = V

But when different reactant are reacting with different number of moles then the value of K may be calculated as

2So2 + O2 2So3

V = dx/dt = K [So2]2 [O2]

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K = V/[So3]2 [O2]

DETERMINATION OF RATE OF REACTION:

There are two methods for the determination of rate of chemical reaction

(1) PHYSICAL METHOD:

When the rate of a chemical reaction is determined by using properties such as colour change state change the method known as physical method

(2) CHEMICAL METHOD:

In the method the change in concentration of reactant or product is noted and with the help of this change rate of reaction is determined e.g.

For the reaction R P

Velocity of reaction = -d [R]/dt = +d [P]/dt

The negative sign indicates decrease in concentration of the reactant while positive sign indicates an increase in the concentration of product

ENERGY OF ACTIVATION:

DEFINITION:

The minimum energy required for a reactant to reach to an excited state or activated state is called energy of activation

OR

The minimum energy required for a collusion to be effective is known as energy of activation

For a chemical reaction the reacting molecules are collided with each other and after collusion if their energy is increased such that they can reach to the activated state then they are able to from product every reaction has its own energy of activation energy takes place at low temperature and a reaction with higher energy of activation will take place at high temperature similarly those reactions which are required low energy of activation are proceeded with a greater rate but those reaction in which high energy of activation is required are generally slow.

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The rate of a chemical reaction is usually increased by adding a catalyst which is actually decreased the energy at activation due to which rate of reaction is increased

FACTORS AFFECTING RATE OF REACTION:

The factors which influence the rate of a chemical reaction are as follows

SURFACE AREA OF REACTANTS:

The surface area of a reactant influence the rate of chemical reaction e.g. powdered zinc reacts more rapidly with water than the piece of zinc because powdered zinc is provided more surface area for the reaction hence we can say that if the surface area of the reactant is greater then the rate of reaction is also greater.

AFFECT OF CONCENTRATION:

As we know that the rate of chemical reaction is directly proportional to the concentration of the reactants are evil able for the reaction the rate of reaction will be greater by increasing the concentration of reactants the collusion frequency increases so the number of effective collusion increases so more product is formed.

TEMPERATURE:

With the increase of temperature the average kinetic energy of the molecules is increased therefore the chance of effective collusion is greater at high temperature so with the increase of temperature the rate of a chemical reaction is generally increased. It is observed that a rise of 10oC in temperature approximately doubles the rate of reaction.

CATALYST:

DEFINITION:

A substance which alters the rate of a chemical reaction its self remaining chemically unchanged at the end of the reaction is called Catalyst.

A catalyst which increases the rate of a chemical reaction by lowering energy of activation is called positive catalyst where as a catalyst which decreases the rate of a chemical reaction by increasing energy of activation is called negative catalyst or inhibitor the effect of a positive catalyst is shown in the figure.

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Chemistry Xi

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