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CHEMICAL THERMODYNAMICS, Thermodynamics deal with the heat changes, different forms of energy and their, quantitative relationship between them. Chemical thermodynamics deals with, energy changes in a chemical reaction., Limitations:, 1. These laws do not deal with microscopic systems., 2. These laws do not include time as a variable. It does not deal with the rate of, process and its mechanism., 3. Thermodynamics predicts the feasibility of the process but not its success., Some basic terms, System and Surroundings, The part of the universe which is under experimental consideration is called system, and the part of the universe except system is called surroundings. The real or, imaginary surface which separates the system from the surroundings is called the, boundary., Types of systems, There are three types of systems:, 1. Isolated system:, It is the system which cannot exchange matter and energy with the, surroundings. For example: Tea in a thermos flask., 2. Closed system:, It is a system which can exchange energy but not matter. For example: Hot, water in a closed glass bottle., 3. Open system:, It is a system which can exchange Both energy and matter with the, surroundings. For example: Hot water in an open beaker., On the basis of composition, there are two types of systems:, 1. Homogeneous system: A system is said to be homogeneous when it is, completely uniform throughout. It is made up of only one phase. For, example: aqueous solution of sugar, glucose etc., 2. Heterogeneous system: A system is said to be heterogeneous when it is, not uniform throughout, i.e it consists of two or more phases. For, example: Ice in contact with water.

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Macroscopic properties of system, The properties associated with the bulk quantity of matter are called macroscopic, properties. These properties are of two types:, 1. Intensive properties, The properties whose magnitude does not depend upon the quantity of, matter present in a system are called intensive properties. For example:, Pressure, temperature, density, specific heat, surface tension, etc., 2. Extensive properties, The properties whose magnitude depends upon the quantity of matter, present in the system are called extensive properties. For example: mass,, volume, internal energy, entropy, etc., It is important to note that the quotient obtained by dividing two extensive, properties may give an intensive property, eg, mass and volume are, extensive properties but density (mass per unit volume) is an intensive, property., State functions (State variables) and Path functions, Those macroscopic properties which can change the state of a system are, called state variablesP, V, T, composition, internal energy, entropy are state, functions because we can describe the state of a system by quoting these, variables, Heat(q) and work (w) are not the state functions, The functions which depend on the path followed are path dependant, functions. For eg: Heat (q) and work (w)., Thermodynamic processes, A thermodynamic process is said to occur when the state of a system, changes from one state (initial state) to another state (final state):, 1. Reversible and Irreversible Process, On the basis of nature of change by which a system travels from initial to, final state, process is divided into two types:, Reversible Process, Irreversible Process,  It is an ideal process which takes, infinite time to complete., ,  It is a spontaneous process and, takes finite time to complete.

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 The driving force (P) is, infinitesimally greater than the, opposing force (p).,  It is in equilibrium in all stages., ,  The driving force (P) is much, greater than the, the opposing force (p).,  Equilibrium exists in the initial and, final stages only.,  Work obtained in this process is,  Work obtained in this process is less, maximum., than, reversible process.,  It can be brought back to initial state  It can be brought back to initial state, without, without, It can not be brought, producing any permanent effect in, back to the initial state without, the adjacent surroundings., avoiding a permanent change in the, surroundings., Cyclic Process, A cyclic process is one in which the system after performing a series of operations,, returns to its initial state. A cyclic process may both be reversible and irreversible, depending upon the manner how it is carried out. The net work done in this process, is zero, ie., w = 0., Isothermal process, It is the process carried out at constant temperature dT = 0. For this process, the, system is usually kept in contact with thermostat., Adiabatic Process, It is the process in which heat cannot leave or enter the system dq = 0. For this, process, the system is thermally insulated from the surrounding., Isobaric process, It is the process carried out at constant pressure , dp = 0. All reactions carried out at, atmospheric pressure are examples of isobaric process. However volume change, always takes place in this process., , Isochoric Process, It is the process in which the volume of the system is kept constant. (dV = 0)

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Thermodynamic equilibrium, A system is in thermodynamic equilibrium if its properties do not change with, time. It is of three types:, 1. Mechanical equilibrium, Ie., position or velocity of a system does not change with time or when no, work is done on the system or done by the system., 2. Thermal equilibrium, Ie,. temperature remains constant throughout the system including the, surroundings., 3. Chemical equilibrium, Ie,. composition of system remains constant and definite and does not, change with time., Thermodynamic Quantities, 1. Internal energy (U or E), The total energy including all forms of kinetic, potential and chemical, energies contained in a thermodynamic system is called internal energy (U, or E). We can not define how much energy is associated with the energy in, all the above form of energy but can only predict and measure the change in, internal energy. It is given as ∆U = Ufinal – Uinitial. The SI unit is Joule., 2. Work, Work can be defined as the displacement of system against some force, originating from the surroundings and acting on the boundary of the system ., Work may be in the form of expansion or compression of system known as, the mechanical work, ie., pressure – volume work. The volume change in, expansion of solids and liquids is small, so the work done is small. Internal, change can be brought out by doing adiabatic work on the system., Therefore,, ∆U = U2 – U1 = Wadiabatic

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Heat: We can change the internal energy of a system by transfer of heat from the, surroundings to the system or viceversa without expenditure of work. This, exchange of energy which is result of temperature difference is called heat (q). The, change in internal energy is equal to the heat exchanged when no work is done by, the system. Therefore, ∆U = q. The q is positive when heat is transferred from the, surrounding to the system and q is negative when heat is tranferred from the, system to the surroundings. The quantity of heat is expressed in calories., First law of Thermodynamics, It is an application of the law of conservation of energy which means that energy, can neither be created nor be destroyed that is the total energy of the system, remains constant though it may change from one from form to another., Derivation, Consider a system whose internal energy is U1. If the system is supplied ‘q’, amount of heat the internal energy of the system will become U1 + q. Now if work, is done on the system, the final internal energy is U2. Thus, U2 = U1 + q + w, U2 - U1 = q + w, ∆U = q + w, If ‘q’ is the heat absorbed and ‘W’ is the workdone by the system, then relationship, becomes, ∆U = q + (-w) = q – w, Limitations of first law of thermodynamics, 1. The law does not give any information about the direction of flow of energy, or heat., 2. It does not explain the irreversible nature of natural spontaneous processes., 3. This law does not explain why chemical reactions does not proceed to, completion., Expression for first law for different thermodynamic processes

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Enthalpy:, Most of the chemical reactions in gaseous state proceed under constant, atmospheric pressure, so thermodynamic changes need to be evaluated at constant, pressure conditions. Consider a change in the system from initial state (1) to final, state (2) carried out at constant pressure. The amount of heat absorbed during the, process is q and the volume changes from V1 to V2. The according to first law of, thermodynamics, we have, ∆U = q + w = qp – P(V2 – V1), qp = (U2 – U1) + P(V2 –V1), qp = (U2 + PV2) – (U1 + PV1), qp = H2 – H1, qp = ∆H, The quantity U + PV is known as heat content or enthalpy (H) of the system and, ∆H is the enthalpy change for the process. It is a state function and is an extensive, property.

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∆H = (U2 – U1) + (P2V2 – P1V1), ∆H = ∆U + P∆V, At constant pressure conditions,, ∆H = ∆U + P∆V, At constant volume, ∆V = 0, the above expression becomes,, ∆H = ∆U = qv, In chemical reaction involving gases, we have, P∆V = ∆ngRT, Or ∆ng = nB – nA. Therefore, ∆H = ∆U + P∆V, ∆H = ∆U + ∆ngRT, qp = qv + ∆ngRT, or where ∆ng is the difference between stoichiometric coefficients of all gaseous, products and all gaseous reactants. Conditions under which ∆H = ∆U or qp = qv., i), ii), iii), , When the reaction is carried out in closed vessel, ie., ∆V = 0, When reaction involves only solids or liquids or solutions, When np = nr, , Standard Enthalpy Change (∆H˚), The enthalpy change of a reaction when all the reactants and the products are in, their standard states (ie., temperature, 298K and pressure, 1 atm or 1 bar) is known, as the standard enthalpy change (∆H˚)., Spontaneous and non spontaneous process, A process which can take place either of its own or under some initiation is called, spontaneous process or reaction. Eg., i), ii), iii), , Dissolution of common salt, sugars, salts etc., in water., Flow of heat from hot body to cold body., Burning of coal etc.

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A process which has no natural tendency or urge to occur is said to be a non –, spontaneous process, eg.,, i), ii), iii), , Flow of water from ground to overhead tank, Diffusion of gas from a low pressure to high pressure., Flow of heat from a cold body to hot body., , Driving force for a spontaneous process, The force which is responsible for the spontaneity of a process is called driving, force. This driving force is due to tendency of a system (i) to acquire minimum, energy (∆H) and (ii) to have maximum randomness or disorderness. The above two, factors determine the spontaneity of a process (or reaction) together., Entropy (S), Entropy is a measure of randomness or disorderness of the molecules of the, system., Some facts about entropy, 1. The value of entropy depends on the mass of the system. Hence it is an, extensive property., 2. The magnitude of entropy changes on changing the physical states of a, system, ie.,, Solid < liquid < Gas, 3. Entropy increases with increase in temperature. Thus heat (q) has, randomising influence on the system., Entropy Change (∆S), It is equal to the heat absorbed isothermally and reversibly divided by the, temperature at which heat is absorbed., Mathematically,, ∆S =, , 𝑞𝑟𝑒𝑣, 𝑇, , It has been found that heat added to a system at lower temperature causes greater, randomness than when the same quantity of heat is added to the system at higher, temperature. This shows that entropy change (∆S) is inversely proportional to the, temperature (T). Entropy is a state function and it depends on the initial and final, states of the system so that

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∆S = Sfinal - Sinitial, For a chemical reaction,, ∆S = S(products) – S(reactants), Unit of entropy and entropy change, Entropy is an extensive property. It is expressed as (cal K-1mol-1) or joules deg1, mol-1, Entropy change is expressed in terms of cal K-1 or JK-1(SI unit), Spontaneity in terms of entropy change, 1. For a spontaneous process in an isolated system such as mixing of gases, the, change in entropy is positive, ie., ∆S > 0, 2. If a system is not isolated, both system and surroundings together constitute, the isolated system. Thus, ∆S(total) = Ssystem + Ssurrounding, For a spontaneous process, ∆S(total) = Ssystem + Ssurrounding > 0, ∆Stotal also known as ∆Suniverse. For natural processes, entropy of universe is, increasing., 3. During a spontaneous process, the entropy of the system goes on increasing, till the system attains the equilibrium, ie., it becomes maximum. Therefore,, ∆S = 0 (at equilibrium for an isolated system), If ∆Stotal is negative, the direct process is non-spontaneous whereas the, reverse process is spontaneous.

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4. Entropy change in reversible and irreversible process:, , Second Law of the thermodynamics, 1. Claussius statement: It is impossible to make heat flow from a body at a, lower temperature to a body at a higher temperature without doing external, work on the working substance., 2. Kelvin –Planck statement: No thermodynamic process is possible in which, the only event is the conversion of heat into work., 3. All spontaneous processes or naturally ocurring processes are, thermodynamically irreversible. Without the help of any external agency, a, spontaneous process cannot be reversed., 4. The entropy of the universe increases in a spontaneous process and remains, unchanged in an equilibrium process. Mathematically, the second law of, thermodynamic can be expressed as:, At spontaneous process:, ∆Suniverse = Ssystem + Ssurroundings > 0, At equilibrium process:, ∆Suniverse = Ssystem + Ssurroundings = 0, Thus, the entropy of universe is continuously increasing and tends to be, maximum., Therefore, the mathematical expression for second law of thermodynamics, would be:, dS ≥, , 𝑑𝑞, 𝑇, , Gibb’s energy (G), Change in free energy (∆G) and spontaneity

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Gibb’s energy is a measure of the capacity of a system to do maximum useful, work. Decrease in Gibb’s free energy is equal to the useful work done by the, system during the process. For example: the electrical energy supplied to, water decomposes it into H2 and O2. The volume of the system increases, because the products of electrolysis are gases. The work done by the system, is the mechanical work, ie., work of expansion. The useful work that has, been obtained is the decomposition of water into H2 and O2, ie., nonmechanical useful work. Gibb’s energy is a state function and is an extensive, property. It is given by the equation, G = H – TS, For a change and constant temperature, the above equation may be written as:, ∆G = ∆H - T∆S, This equation is called Gibbs Helmholtz equation. ∆G is a measure of spontaneity, of a chemical reaction., Free energy change and non-mechanical work, From first law of thermodynamics, we know, q = ∆U + w, or, , q = ∆U + wexpansion + wnon-expansion, , At constant pressure, the expansion work in P∆V and electrical work is nonexpansion work., q = ∆U + P∆V + wnon-expansion, q = ∆H + wnon-expansion, For a reversible change taking place at a constant temperature, ∆S =, , 𝑞𝑟𝑒𝑣, 𝑇, , qrev = T∆S, Substituting the value of eq (i), we get, T∆S = ∆H + wnon-expansion, Or,, , ∆H - T∆S = -wnon-expansion, , For a change taking place under condition of constant temperature and pressure,

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∆G = ∆H - T∆S, Substituting this value in eq(iii), we get, -∆G = wnon-expansion = wuseful, -∆G = wmax, , Or,, , For a spontaneous process, there is decrease in free energy, ie., ∆G is negative, i), ii), iii), , If ∆G < 0, the process is spontaneous., If ∆G = 0, the process is in equilibrium., If ∆G > 0, the process is non-spontaneous., , Standard free energy of formation (∆G˚), It is defined as the free energy change for a process at 298K and, 1 bar pressure in, which the reactants in their standard state are converted into products in their, standard state. It is denoted by ∆G˚ and it may be expressed as, ∆G˚ = ∆H˚ - T∆S˚, For a chemical reaction,, ∆G˚ = Σ ∆fG˚(Products) – Σ ∆fG˚(Reactants), Standard free energy of formation (∆fG˚), The standard free energy of formation is equal to the change in free energy when 1, mole of the compound is formed from its constituent elements in their standard, states. The standard free energy of formation of an element in its standard state is, zero., Relation between ∆G˚ and Equilibrium constant (Kc)

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∆G˚ = -2.303 RT log Kc, Helmholtz Free Energy or Work Function (A), The fraction of internal energy which is isothermally available is called the work, function (A). Mathematically,, A = U – TS, For a change in work function,, ∆A = U - T∆S, Relationship between ∆G and ∆A, Gibbs – Helmholtz equation is written as:, ∆G = ∆H - T∆S, , ------(i), , The enthalpy change for a system at constant pressure is given by, ∆H = ∆U + P∆V, , ------(ii), , Comparing equation (i) and (ii), we have, ∆G = ∆U + P∆V - T∆S, , ------(iii), , Now change in Helmholtz free energy is given by, ∆A = ∆U - T∆S, , ------(iv), , From equation (iii) and (iv), we get, ∆G = ∆A + P∆V, If volume remains constant, ie., ∆V = 0, ∆G = ∆A, ∆G as a driving force is preferred over ∆A because when ∆G is used to predict the, direction of a process, it is essential to keep temperature (T) and pressure (P), constant as (∆G)T,P < 0 for a spontaneous process., But if A is used to decide the direction of a process, temperature (T) and volume, (V) are kept constant as (∆A)T,V < 0 for a spontaneous process. As we know that, the experiments are carried out in laboratory at constant T and P, the use of ∆G is, more preferred., Thermal Death

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Most of the natural processes are irreversible and spontaneous in nature. These, processes involve the increase in the entropy of the system. Therefore entropy of, the universe is continuously increasing. It means there is decrease in ordered, energy which is available for useful work. Since the entropy of the universe is, increase day by day, a stage will come when there will be no energy available for, useful work. This is referred to as thermal death of the inhabitants of the universe., At this stage all production of useful work would cease and life would come to an, end., Third law of thermodynamics, Third law of thermodynamics was given by Nernst. According to this law, “ At, absolute zero, the entropy of a perfectly crystalline solid is taken as zero”., Thermochemistry, , , Thermochemistry is a branch of Thermodynamics, a scientific field that comprises the, relationships between heat, work, and other forms of energy resulting from different, chemical and physical processes. Thermochemistry itself is defined as energy changes, when a chemical reaction takes place. Energy is the capacity to supply heat or do work., Heat of reaction or enthalpy of reaction, Enthalpy of reaction is equal to the change in enthalpy when the number of moles of reactants as, represented by a balanced chemical equation react completely., ∆rH = ΣaiHproducts - ΣbiHreactants, Σ represents summation, ai and bi are stoichiometric coefficients of the products and reactants, completely. Eg., For the reaction, CH4(g) + 2O2(g) → CO2(g) + 2H2O(l), ∆rH = [HCO2(g) + 2 × HH2O(l)] – [HCH4(g) + 2 × HO2(g)], If HR > HP then ∆rH = -ive, Exothermic, HR < HP then ∆rH = +ive, Endothermic, Standard enthalpy of reaction (∆rH˚), It is the enthalpy change when all the reactants and products are in their standard states., A substance is said to be in the standard state when it is present in pure and most stable form at 1 bar, pressure and at a specified temperature eg., the standard state of solid iron at 500 K is pure iron at 1, bar.

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Types of enthalpies in chemical reactions

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Enthalpy of neutralisation, It is the change in enthalpy when one gram equivalent of an acid is neutralised by one gram equivalent, of a base in dilute solution or vice-versa., The enthalpy of neutralisation for strong acid and strong base remains constant and is equal to the heat, released in the formation of one of water which is equal to -57.1 KJ., For example: HCl(aq) + NaOH(aq) → NaCl + H2O(l) , ∆cH

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Enthalpy of solution, Enthalpy of solution is the change in enthalpy when one mole of substance is dissolved in a given, amount of the given solvent. If the volume of the solution is so large that further dilution of the solution, does not produce any heat change, then it is called enthalpy of solution at infinite dilution. Enthalpy of, solution may be positive or negative.

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Enthalpy of dilution:, It is equal to the change in enthalpy when a solution containing one mole of the solute is diluted from, one concentration to other concentration by adding solvent. Its value depends on the original, concentration of the solution and the volume of solvent added., Bond Enthalpy (∆bondH), Enthalpy change is related to the types of bonds formed or broken. It may be defined in two terms:

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i), , Bond dissociation enthalpy: The enthalpy change involving the breaking of one mole of, covalent bonds of a gaseous covalent compounds to give products in gaseous state, is called, the bond dissociation enthalpy, H2(g) → 2H(g) , ∆H-HH˚ = 435.0 KJ/mol, It is same as that of enthalpy of atomisation, , ii), , Mean (Average) bond enthalpy, In case of polyatomic molecules, the mean of bond dissociation enthalpies of all the bonds, present in the compound is taken. It is called as the mean bond enthalpy. Eg.,, CH4(g) → C(g) + 4H(g) ∆aH˚ = 1665 KJ/mol, Thus in CH4, ∆C-HH˚ is calculated as, ∆C-HH˚ = 1/4 (∆aH˚), = 1/4 × 1665 = 416 KJ/mol, , Hess’s law of constant heat summation, According to this law, “ If a chemical reaction can be made to take place in number of ways in one or in, several steps, the total enthalpy change is independant of intermediate steps involved in the change.”, According to Hess’s law, total enthalpy change,, ∆H1 = ∆H2 + ∆H3 + ∆H4, This law is used for the determination of heats of formation of various compounds, bond energies,, resonance energies etc.