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Chapter 6, CHEMICAL THERMODYNAMICS, What is Thermo dynamics and why is it useful ?, Thermodynamics is the branch of science that describes the behaviour of, matter and the transformation between different forms of energy on a, macroscopic scale. Thermodynamics describes a system in terms, of its bulk properties. Only a few such variable are needed to describe the, system, and the variables are generally directly accessible through, measurements. A thermodynamic description of matter does not make, reference to its structure and behaviour at the microscopic level., , The laws of Thermodynamics :, The law's of thermodynamics is the law of observation. No one has ever, observed that any thing goes in contrary to thermodynamics law. So we elevate, this observation to the status of thermodynamic law. The real justification of, this comes when things we derive using this law turn's out to be true that is, verified by experiments, , Applicat ion of Thermodynamics :, (i) It provides relationship between heat, work and measurable proper ties of, matter., (ii) It predicts direction of natural change - like what circumstances are best for, rusting of iron., (iii) It predicts up to what extent a chemical reaction can proceed in forward, direction., Example : How much ammonia (NH3) can be formed from N2 and H2 in a, closed container., (iv) It help in understanding why different phases of matter exist - and provide, simple relationship between various measurable proper ties of system, (thermodynamical variables), Salient features of Thermodynamics :, During study of this chapter you will observe that mostly you will be dealing, with macroscopic properties (bulk properties) like pressure, volume,, temperature density of system. This is because thermodynamics, is macroscopic science and it do not concern's with detailed microscopic make, up of the system., Limitations :
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(1) It tells us whether a given chemical reaction will take place or not under the, given set of conditions, but doesn't tell us anything about the rate of reaction., (2) It tells us about the initial and final properties of the system but doesn't tell, us anything about the path or mechanism followed by the system., , Important Terms and Definitions, System: Refers to the portion of universe which is under observation., Surroundings: Everything else in the universe except system is called, surroundings., The Universe = The System + The Surroundings., Open System: In a system, when there is exchange of energy and matter, taking place with the surroundings, then it is called an open system., For Example: Presence of reactants in an open beaker is an example of an, open system. Closed System: A system is said to be a closed system when there, is no exchange of matter ‘ but exchange of energy is possible., For example: The presence of reactants in a closed vessel made of conducting, material., Isolated System: In a system, when no exchange of energy or matter takes, place with the surroundings, is called isolated system., For example: The presence of reactants in a thermoflask, or substance in, an insulated closed vessel is an example of isolated system., , Homogeneous System: A system is said to be homogeneous when all the, constituents present is in the same phase and is uniform throughout the, system.
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For example: A- mixture of two miscible liquids., Heterogeneous system: A mixture is said to be heterogeneous when it, consists of two or more phases and the composition is not uniform., For example: A mixture of insoluble solid in water. ’, The state of the system: The state of a thermodynamic system means its, macroscopic or bulk properties which can be described by state variables:, Pressure (P), volume (V), temperature (T) and amount (n) etc., They are also known as state functions., ❖ State variables : To define a thermodynamics states of a system, we, have to specify the values of certain mesurable quantities. These are, called thermodynamic variable or state variable., A system can be completely defined by four variables namely pressure,, temperature, volume and composition. A system is said to be in a certain, definite state when all of its properties have definite, value.. Between two fixed state the change in the value of state function is, same irrespective of the path connection two states., Differential of a state function integerated over a cyclic path returns zero. In, other words summation, of change in state function in a cyclic process is equal to zero., if, , dX = 0 =>X is a state function (property of state function), , note that if X is a state function, dX is called definite quantity, Example : T, V, P and U (internal energy), H (enthalpy) are state variables., , ❖ Path function or path dependent quantities :, The value of path function depends upon path connection two states . There, can be infinite vaules of path function between two states depending upon path, or process., Path functions are also called indefinite quantities since between two fixed, state the value of path function is not fixed. Heat and Work are two important, path dependent quantities with which we deal in this chapter., , ❖ Extensive and Intensive variables :, Properties which depend on the amount of the substance (or substances), present in the system are called extensive propterties., e.g. Mass, volume, heat capacity, internal energy, entropy, Gibb's free, energy (G), surface area etc. These properties will change with change in the, amount of matter present in the system., It is important to note that the total value of an extensive property of a system, is equal to the sum of the values of different parts into which the system is, divided., ❖ Intensive properties : Properties which are independent of the, amount of substance (or substances) present in the system are called, intensive properties,
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e.g. pressure, density, temperature, viscosity, surface tension,, refractive index, emf, chemical potential, sp. heat etc, These are intensive, properties., An extensive property can be converted into intensive property by defining it, per unit of another extensive property., E x . Concentration = mole / volume, Density = mass / volume, heat capacity = heat absorbed / rise in temperature, While mole, mass, heat are extensive properties, concentration, density and, heat capacity are intensive properties., , ❖ Thermo dynamic equilibrium :, Thermodynamic generally deals the equilibrium state of the system in which, the state variable are uniform and constant throughout the whole system., The term thermodynamic equilibrium implies the existence of three different, types of equilibria in the system. These are :, ( i ) Mechanical equilibrium :, When there is no macroscopic movement within the system itself or of the, system with respect to surroundings, the system is said to be in a state of, mechanical equilibrium., ( i i ) Chemical equilibrium :, When the system consists of more than one substance and the composition of, the system does not vary with time, the system is said to be in chemical, equilibrium. The chemical compositionof a system at equilibrium must be, uniform and there should be no net chemical reaction taking place., ( i i i ) Thermal equilibrium :, When the temperature throughout the entire system is the same as that of the, surroundings then the system is said to be in thermal equilibrium., , ❖ Equation of state :, An equation that relates the variables T, p, V, and n to each other is called the, "equation of state." The most general form for an equation of state is., f (p, V, T, n) = 0., The ideal gas equation of state :, The best known equation of state for a gas is the "ideal gas equation of state". It, is usually written in the form, pV = nRT, This equation contains a constant, R, called the gas constant., The vander Walls equation of state for real gases :, The vander Walls equation of state is,, , Notice that the vander Walls equation of state differs from the ideal gas by the, addition of two adjustable parameters, a, and b (among other things)., Note : Equation of state for liquid and solids are also defined empirically.
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Process : Anything which changes state of system is called process. Usually as, a result of heat and work, interactions change of state take place. e.g. isothermal process, Path of a process : The exact sequence of steps through which system, changes state is called path of a, process.e.g. reversible or irrervisible path, Isothermal process: When the operation is carried out at constant, temperature, the process is said to be isothermal. For isothermal process, ΔT = 0 Where ΔT is the change in temperature., Adiabatic process: It is a process in which no transfer of heat between, system and surroundings, takes place., Isobaric process: When the process is carried out at constant pressure, it is, said to be isobaric. i.e. ΔP = 0, Isochoric process: A process when carried out at constant volume, it is, known as isochoric in nature., Cyclic process: If a system undergoes a series of changes and finally returns, to its initial state, it is said to be cyclic process., Reversible Process: When in a process, a change is brought in such a way that, the process could, at any moment, be reversed by an infinitesimal change. The, change r is called reversible., Poly tropic processes : It is defined as a process in which PVn=constant =k, All of the above mentioned processes can be performed in two ways ,, reversibly and irreversibly, Reversible process : When the difference between driving force and opposing, force is very small and the process is carried out infinitesimally slowly, then the, process is called reversible process. The reversible process is carried out in, such a manner that at any moment of the change the directon of process can, be reversed by making a small change in driving force.A reversible process is, also called quasi static process., During a reversible process, the system and surrounding remain in equilibrium, through the process. The reversible processes are idealized processes which, cannot be actually carried out, but nevertheless they are very important, because they help in calculation of change in state function in the process. In, other words the reversible processes are hypothetical processes., A quasistatic process is the one in which system remain infinite simally closer, to the state of equilibrium through out t he process., Irreversible process : Any process which does not take place in the above, manner and difference between driving force and the opposing force is quite, large is called irreversible process.All natural processes are example of, irreversible process.
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• Internal Energy, It is the sum of all the forms of energies that a system can possess., In thermodynamics, it is denoted by AM which may change, when, — Heat passes into or out of the system, — Work is done on or by the system, — Matter enters or leaves the system., Change in Internal Energy by Doing Work, Let us bring the change in the internal energy by doing work., Let the initial state of the system is state A and Temp. TA Internal energy = uA, On doing’some mechanical work the new state is called state B and the temp., TB. It is found to be, TB > TA, uB is the internal energy after change., ∴ Δu = uB – uA, Change in Internal Energy by Transfer of Heat, Internal energy of a system can be changed by the transfer of heat from the, surroundings to the system without doing work., Δu = q, Where q is the heat absorbed by the system. It can be measured in terms of, temperature difference., q is +ve when heat is transferred from the surroundings to the system. q is -ve, when heat is transferred from system to surroundings., When change of state is done both by doing work and transfer of heat., Δu = q + w
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First law of thermodynamics (Law of Conservation of Energy)., It states that, energy can neither be created nor be destroyed. The energy of an, isolated system is constant., Δu = q + w., The first law of thermodynamics is based on experience that energy can be, neither created nor destroyed, if both the system and the surroundings are, taken into account.Suppose a blocks of mass 'M' is moving in, gravitational field with velocity v. The total energy of blocks (in ear th frame of, reference) is given as :, E = K + V + U : (K = kinetic energy, V = potential energy, U = internal energy), A thermodynamic system is studies generally at rest so K = 0. If effect of, gravitation field (or other fields are ignored) is also ignored then we left with E =, U. Thus U (internal energy) is energy of system., If a system is present in particular thermodynamic state say 'A' it has fixed, amount of internal energy UA., Suppose by a process the system is taken from state A to state B. In the, process 'q' heat is absorbed by system and w work is done on the system. Thus, in the state 'B' total internal energy of system become, UB = UA + q + w., UB – UA = q + w, , This is mathematical statement of first law., First law of thermodynamics states that energy is conserved. Direct, consequence of this statement is U is state function. This implies between any, two fixed state, there can be infinite process or path, but, ΔU in all process will remain the same.(Property of a function of state), The total of all these forms of energy for the system of interest is given the, symbol U and is called, the internal energy., Hence total internal energy of a system can be written as, , U = Utranslational+Urotational+Uvibrational+Uintermolecular+Uelectronic+Urelativistic, of these Urelativistic and Ueletronic is unaffected by ordinary heating. So, basically the kinetic energy terms and Uintermolecur accommodate heat provided, to the system . Hence heat capacity of a sample depends upon these four, terms., For an ideal gas, Uintermolecular is equal to zero, because of absence of, intermolecular force of attraction, in ideal gas. Uintermolecular have large and negative value in solids and liquids., For an ideal gas U is only function of temperature e.g. U=F(T) +Constant
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Due to absence of pressure or volume terms in ideal gas internal energy, U is, independent of pressure and volume of theoretical ideal gas., ❖ Enthalpy (H):, Chemical reactions are generally carried out at constant pressure (atmospheric, pressure) so it has been, found useful to define a new state function Enthalpy (H) as :, H = U + PV (By definition), , Heat capacity, The heat capacity of a system may be defined as the amount of heat required to, raise the temperature of the system by one degree., If dq is the small quantity of heat added to the system, let the temperature of, the system rises by dT, then heat capacity C of the system is given by, , The increase in temperature is proportional to the heat transferred., q = coeff. x ΔT, q = CΔT, Where, coefficient C is called the heat capacity., C is directly proportional to the amount of substance., Cm = C/n, It is the heat capacity for 1 mole of the substance., • Molar heat capacity, It is defined as the quantity of heat required to raise the temperature of a, substance by 1° (kelvin or Celsius)., Heat capacity at constant volume( Cv) :, It may be defined as the rate of change of internal energy with temperature at, constant volume.
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Molar heat capacity at constant volume is defined by the relation, , Heat capacity at constant pressure(Cp) :, It is the rate of change of enthalpy with temperature at constant pressure., , When pressure is maintained constant, the relation is given, , Hence heat capacity of a system at constant volume Cv is equal to the increase, in internal energy of the system per degree rise of temperature at constant, volume. Similarly heat capacity at constant pressure Cp is numerically equally, to the increase in enthalpy of the system per degree rise of temperature., For 1 mole of an ideal gas, heat capacity at constant pressure i.e. Cp is greater, than the heat capacity at constant volume i.e., Cv, , Cp > Cv, These are called molar heat capacities, • Specific Heat Capacity, It is defined as the heat required to raise the temperature of one unit mass of a, substance by 1° (kelvin or Celsius)., q = C x m x ΔT, where m = mass of the substance, ΔT = rise in temperature., • Relation Between Cp and Cv for an Ideal Gas, At constant volume heat capacity = Cv, At constant pressure heat capacity = Cp, At constant volume qv= CvΔT = ΔU, At constant pressure qp = Cp ΔT = ΔH, For one mole of an ideal gas, ΔH = ΔU + Δ (PV) = ΔU + Δ (RT), ΔH = ΔU + RΔT, On substituting the values of ΔH and Δu, the equation is modified as, Cp ΔT = CvΔT + RΔT, or Cp-Cv = R, ❖
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❖ Measurement of ΔU and ΔH—Calorimetry, Determination of ΔU: ΔU is measured in a special type of calorimeter,, called, bomb, calorimeter., , Working with calorimeter. The calorimeter consists of a strong vessel, called (bomb) which can withstand very high pressure. It is surrounded, by a water bath to ensure that no heat is lost to the surroundings., Procedure: A known mass of the combustible substance is burnt in the, pressure of pure dioxygen in the steel bomb. Heat evolved during the, reaction is transferred to the water and its temperature is monitored.
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• Enthalpy Changes During Phase Transformation, 1) Enthalpy of fusion: Enthalpy of fusion is the heat energy or change, in enthalpy when one mole of a solid at its melting point is converted, into liquid state., , 2) Enthalpy of vaporisation: It is defined as the heat energy or change, in enthalpy when one mole of a liquid at its boiling point changes to, gaseous state., , 3) Enthalpy of Sublimation: Enthalpy of sublimation is defined as the, change in heat energy or change in enthalpy when one mole of solid, directly changes into gaseous state at a temperature below its melting, point., , 4) Standard Enthalpy of Formation:, Enthalpy of formation is defined as the change in enthalpy in the, formation of 1 mole of a substance from its constituting elements under, standard conditions of temperature at 298K and 1 atm pressure.
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5) Enthalpy of Combustion: It is defined as the heat energy or change, in enthalpy that accompanies the combustion of 1 mole of a substance, in excess of air or oxygen., , • Thermochemical Equation, A balanced chemical equation together with the value of ΔrH and the, physical state of reactants and products is known as thermochemical, equation., , Hess’s Law of Constant Heat Summation:, The total amount of heat evolved or absorbed in a reaction is same, whether the reaction takes place in one step or in number of steps., , • Born-Haber Cycle, It is not possible to determine the Lattice enthalpy of ionic compound by, direct experiment. Thus, it can be calculated by following steps. The, diagrams which show these steps is known as Born-Haber Cycle.
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Spontaneity, Spontaneous Process: A process which can take place by itself or has a, tendency to take place is called spontaneous process., Spontaneous process need not be instantaneous. Its actual speed can vary, from very slow to quite fast., A few examples of spontaneous process are:, (i) Common salt dissolves in water of its own., , (ii) Carbon monoxide is oxidised to carbon dioxide of its own.
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• Entropy (S), The entropy is a measure of degree of randomness or disorder of a system., Entropy of a substance is minimum in solid state while it is maximum in, gaseous state., , The change in entropy in a spontaneous process is expressed as ΔS, Entropy change of system and surrounding, in reversible and irrversible process, , • Gibbs Energy( G) and Spontaneity, A new thermodynamic function, the Gibbs energy or Gibbs function G, can be, defined as G = H-TS, ΔG = ΔH – TΔS, Gibbs energy change = enthalpy change – temperature x entropy change ΔG, , gives a criteria for spontaneity at constant pressure and temperature,, i), (ii), , If ΔG is negative (< 0) the process is spontaneous., If ΔG is positive (> 0) the process is non-spontaneous.
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• Free Energy Change in Reversible Reaction, , The second law of thermodynamics :, , The second law of thermodynamics predict's direction of natural change. It do, so with the help of state, function 'S' - called entropy of system. But for predicting direction of natural, change another quantity Ssurrounding is also needed. Ssurrounding which is called, entropy of surrounding is a path dependent quantity., , Since Ssystem is state function - If a system make transition from state A to, state B - by infinite paths in few of them may be reversible and other may be, irreversible.
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Third law of thermo dynamics :, Third law of thermodynamics helps in determining absolute entropy of, substances. It is based on an assumption that entropy of every perfectly, crystalline substance is zero at zero Kelvin. This is justified because , at, absolute zero every substance is in state of lowest energy and position, of every atom or molecule is defined in solid. Hence at T=0 S(T=0)=0, The third law of thermodynamics modifies this observation and sets, S (T = 0) = 0, for all elements and compounds in their most stable and perfect crystalline, state at absolute zero and one atmosphere pressure. (All except for helium,, which is a liquid at the lowest observable temperatures at one atmosphere.), The advantage of this law is that it allows us to use experimental data to, compute the absolute entropy of a substance.
