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Class XI Chemistry, Unit 6, THERMODYNAMICS, Topic:-THERMODYNAMIC TERMS, Sub topics:-, The System and the Surroundings, Types of the System, State of the System, State Functions, Adiabatic System, By Vijay Kumar Sethi
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Thermodynamics , The branch of science which deals with the quantitative relationship between heat and other forms of energies is called thermodynamics, For example,, Chemical energy stored by molecules can be released as heat during chemical reactions when a fuel like methane, cooking gas or coal burns in air. , The chemical energy used to do mechanical work when a fuel burns in an engine, The chemical energy used to provide electrical energy through a galvanic cell like dry cell., The study of these energy transformations forms the subject matter of thermodynamics
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The System and the Surroundings, The part of universe in which observations are made or which is under investigation is called system, Surroundings is the part of universe other than system . The surroundings include everything other than the system. , System and the surroundings together constitute the universe ., The universe = The system + The surroundings
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For example, if we are studying the reaction between two substances A and B kept in a beaker, the beaker containing the reaction mixture is the system and the room where the beaker is kept is the surroundings, System and surroundings are separated by a boundary., This boundary may be real or imaginary., Continue…, The System and the Surroundings
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Types of the System, Open System, There is exchange of energy and matter between system and surroundings., For example. the presence of reactants in an open beaker., Here the boundary is an imaginary, Closed System, There is no exchange of matter, but exchange of energy is possible between system and the surroundings., For example, the presence of reactants in a closed vessel made of conducting material., Isolated System, There is no exchange of energy and matter between the system and the surroundings, For example, the presence of reactants in a thermos flask or any other closed insulated vessel
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Continue…, Types of the System
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The State of the System, The state of a thermodynamic system is described by its measurable or macroscopic (bulk) properties. , We can describe the state of a gas by quoting its pressure (p), volume (V), temperature (T ), amount (n) etc.
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State Functions, The macroscopic properties whose values depend only on the state of the system(initial and final) and not on how it is reached are called state functions., For example, P, T, V, Internal energy, Enthalpy, Entropy, Path functions depend on the path taken to reach one state from another.
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Adiabatic System, The system which will not allow exchange of heat between the system and surroundings through its boundary but energy of the system may change by work., Adiabatic process is a process in which there is no transfer of heat between the system and surroundings., The wall separating the system and the surroundings is called the adiabatic wall., Bursting of tyre is an example of adiabatic process., it happens so fast, there is not enough time to exchange heat with surroundings
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Class XI Chemistry, Unit 6, THERMODYNAMICS, Topic:-, Internal Energy (U), Mathematical statement of the first law of thermodynamics, , By Vijay Kumar Sethi
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Internal Energy (U), Total energy of the system is called internal energy of the system., It is the sum of kinetic(translational , vibrational and rotational) potential , chemical, electrical, mechanical energies etc. , Internal energy is a state function-depends on initial and final states of the system, Absolute value of U can not be determined., Change in internal energy (ΔU) is determined. ΔU = U2-U1, It 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.
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Ways of Changing Internal Energy, Work: Consider an adiabatic system, Initial state of the system = A and its temperature =TA. , Internal energy of the system in state A = UA. , We can change the state of the system in two different ways., One way: By mechanical work (1 kJ):- by rotating a set of small paddles, Final State =B state and its temperature=TB., if TB > TA and the change in temperature, ΔT = TB–TA. , The internal energy of the system in state B = UB and , The change in internal energy, ΔU =UB– UA =Wad.
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Continue…, Ways of Changing Internal Energy Work, Second way: We now do an equal amount (i.e., 1kJ) electrical work with the help of an immersion rod and note down the temperature change. , We find that the change in temperature is same as in the earlier case, TB – TA, The change in internal energy, ΔU =UB– UA =Wad., Therefore, internal energy, U, of the system is a state function., J. P. Joule showed that a given amount of work done on the system, no matter how it was done (irrespective of path) produced the same change of state, as measured by the change in the temperature of the system. By conventions:-, Work done on the system = positive work (+W) internal energy of system increases. , Work done by the system = negative work (-W) internal energy of the system decreases.
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Continue…, Ways of Changing Internal Energy, Heat(q) :, If system and surroundings are at different temperatures , internal energy of the system is changed by transfer of heat through thermally conducting walls, If temperature of System = TA and Temperature of surroundings = TB , TB < TA heat is transferred from system to surroundings till TB=TA, TB > TA heat is transferred from surroundings to system till TB=TA, In this case change in internal energy, ΔU= q, when no work is done at constant volume.
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Heat(q) :, By conventions , The q is positive (+q)= heat is transferred from the surroundings to the system and the internal energy of the system increases and , q is negative (-q) = heat is transferred from system to the surroundings resulting in decrease of the internal energy of the system.., Continue…, Ways of Changing Internal Energy
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Mathematical statement of the first law of thermodynamics, ΔU = U2 – U1 = q + w, ΔU = q + w mathematical statement of the first law of thermodynamics, , which states that The energy of an isolated system is constant., It is commonly stated as the law of conservation of energy i.e., energy can neither be created nor be destroyed.
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Continue…, Mathematical statement of the first law of thermodynamics, q and w are path functions but q +w = ΔU will depend only on initial and final state therefore (q + w) is state function, It will be independent of the way the change is carried out. , If there is no transfer of energy as heat or as work (isolated system) i.e., if w = 0 and q = 0, then Δ U = 0.
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Problem 6.1, Express the change in internal energy of a system when, No heat is absorbed by the system from the surroundings, but work (w) is done on the system. What type of wall does the system have ?, Solution:- Δ U = wad, wall is adiabatic, (ii) No work is done on the system, but q amount of heat is taken out from the system and given to the surroundings. What type of wall does the system have?, Solution:- Δ U = – q, thermally conducting walls, (iii) w amount of work is done by the system and q amount of heat is supplied to the system. What type of system would it be?, Solution:- Δ U = q – w, closed system.
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Class XI Chemistry, Unit 6, THERMODYNAMICS, , Topic:-, Pressure-Volume Work, Reversible and Irreversible process , , By Vijay Kumar Sethi
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Pressure-Volume Work, A cylinder contains one mole of an ideal gas fitted with a frictionless piston. , Initial volume the gas is Vi and pressure of the gas inside is p. , pex = external pressure , pex > p , piston is moved inward till pex= p, Final volume = Vf . , During this compression, piston moves a distance, l and is cross-sectional area of the piston is A
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Pressure-Volume Work, The negative sign of this expression is required to obtain conventional sign for w, which will be positive., compression work is done on the system., (Vf – Vi ) will be negative and negative multiplied by negative will be positive., This formula is also applicable for expansion work which is done by the system so it will be negative., In expansion (Vf – Vi ) will be positive
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Reversible process, A process or change is said to be reversible, if a change is brought out in such a way that the process could, at any moment, be reversed by an infinitesimal (इन्फ़िनिˈटे̮सिम्ल् extremely small )change. , A reversible process proceeds infinitely slowly by a series of equilibrium states such that system and the surroundings are always in near equilibrium with each other. , Irreversible processes , Processes other than reversible processes are known as irreversible processes., All natural processes are irreversible.
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Continue…, Pressure-Volume Work, If the compression of gas is carried out in number of finite steps, work done on the gas will be summed over all the steps and will be equal to Ʃ pΔV, Fig :- , pV-plot when pressure is not constant and changes in finite steps during compression from initial volume, Vi to final volume, Vf , Work done on the gas is represented by the shaded area.
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Continue…, Pressure-Volume Work in Reversible process, Fig. pV-plot when pressure is not constant and changes in infinite steps (reversible conditions) during compression from initial volume, Vi to final volume, Vf ., Work done on the gas is represented by the shaded area.
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Enthalpy, H, The heat absorbed at constant volume is equal to change in the Internal Energy i.e., ΔU = qv, But most of chemical reactions are carried out not at constant volume, but in flasks or test tubes under constant atmospheric pressure., Another thermodynamic function, the enthalpy H [Greek word enthalpien, to warm or heat content] as : H = U + pV, Enthalpy is the sum of internal energy and product of pressure-volume., Enthalpy is a state function-depends on initial and final states of the system, Absolute value of H can not be determined., Change in enthalpy (ΔH) is determined. ΔH = H2-H1
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An extensive property is a property whose value depends on the quantity or size of matter present in the system. , For example, mass, volume, internal energy, enthalpy, heat capacity, etc. are extensive properties., Those properties which do not depend on the quantity or size of matter present are known as intensive properties., For example temperature, density, pressure, molar heat capacity etc. are intensive properties., Extensive and Intensive Properties
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Continue…, Extensive and Intensive Properties, considering a gas enclosed in a container of volume V and at temperature T ., Let us make a partition such that volume is halved, each part now has one half of the original volume,V/2 , but the temperature will still remain the same i.e., T. , It is clear that volume is an extensive property and temperature is an intensive property.
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MEASUREMENT OF ΔU AND ΔH:, CALORIMETRY:-, Experimental technique used to measure energy changes associated with chemical or physical processes is called calorimetry, A calorimeter is a device used to measure the amount of heat changed in a chemical or physical process by measuring temperature changes., Measurements are made under two different conditions:, i) at constant volume, qV, ii) at constant pressure, qp
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ΔU Measurements, For chemical reactions, heat absorbed at constant volume, is measured in a bomb calorimeter . , Here, a steel vessel (the bomb) is immersed in a water bath. The whole device is called calorimeter., The steel vessel is immersed in water bath to ensure that no heat is lost to the surroundings. , A combustible substance is burnt in pure dioxygen supplied in the steel bomb
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Heat evolved during the reaction is transferred to the water around the bomb and its temperature is monitored. , Since the bomb calorimeter is sealed, its volume does not change i.e., the energy changes associated with reactions are measured at constant volume. , Under these conditions, no work is done as the reaction is carried out at constant volume in the bomb calorimeter , Temperature change of the calorimeter produced by the completed reaction is then converted to qV, Continue…, ΔU Measurements
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ΔH Measurements, Measurement of heat change at constant pressure (generally under atmospheric pressure) can be done in a simple calorimeter. ΔΗ = qp (at constant p) , Heat absorbed or evolved, qp at constant pressure is also called the heat of reaction or enthalpy of reaction, ΔrH., In an exothermic reaction, heat is evolved, and system loses heat to the surroundings. Therefore, qp will be negative and ΔrH will also be negative. , Similarly in an endothermic reaction, heat is absorbed, qp is positive and ΔrH will be positive.
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Standard Enthalpy of Reactions, ΔrHƟ, The standard enthalpy of reaction is the enthalpy change for a reaction when all the participating substances are in their standard states., The standard state of a substance at a specified temperature is its pure form at 1 bar. For example, the standard state of liquid ethanol at 298 K is pure liquid ethanol at 1 bar; standard state of solid iron at 500 K is pure iron at 1 bar. Usually data are taken at 298 K., Standard conditions are denoted by adding the superscript Ɵ to the symbol ΔH, e.g., ΔHƟ
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Enthalpy Changes during Phase Transformations, Standard enthalpy of fusion, The enthalpy change that accompanies melting of one mole of a solid substance in standard state is called standard enthalpy of fusion or molar enthalpy of fusion, ΔfusHƟ., Melting of a solid is endothermic, so all enthalpies of fusion are positive, , Standard enthalpy of vaporization, Amount of heat required to vaporize one mole of a liquid at constant temperature and under standard pressure (1bar) is called its standard enthalpy of vaporization or molar enthalpy of vaporization, ΔvapHƟ.
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Standard enthalpy of sublimation, Sublimation is direct conversion of a solid into its vapour., ΔsubHƟ is the change in enthalpy when one mole of a solid substance sublimes at a constant temperature and under standard pressure (1bar)., Solid CO2 or ‘dry ice’ sublimes at 195K with ΔsubHƟ=25.2 kJ mol–1;, Naphthalene sublimes slowly and for this ΔsubHƟ = 73.0 kJ mol–1, The magnitude of the enthalpy change depends on the strength of the intermolecular interactions in the substance undergoing the phase transformations., Continue…, Enthalpy Changes during Phase Transformations
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Standard Enthalpy of Formation, The standard enthalpy change for the formation of one mole of a compound from its elements in their most stable states of aggregation (also known as reference states) is called Standard Molar Enthalpy of Formation. Its symbol is ΔfHƟ , The reference state of an element is its most stable state of aggregation at 25°C and 1 bar pressure., For example, the reference state of dihydrogen is H2 gas and those of dioxygen, carbon and sulphur are O2 gas, Cgraphite and Srhombic respectively.
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Continue…, Standard Enthalpy of Formation, standard molar enthalpy of formation, ΔfHƟ, is just a special case of ΔrHƟ, where one mole of a compound is formed from its constituent elements., It is not an enthalpy of formation of calcium carbonate, since calcium carbonate has been formed from other compounds, and not from its constituent elements.
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Continue…, Standard Enthalpy of Formation, By convention, standard enthalpy for formation, ΔfHƟ, of an element in reference state, i.e., its most stable state of aggregation is taken as zero., How much heat is required to decompose calcium carbonate?, decomposition of CaCO3 (s) is an endothermic process
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Thermochemical Equations, A balanced chemical equation together with the value of its ΔrH is called a thermochemical equation. , We specify the physical state (alongwith allotropic state) of the substance in an equation., , The above equation describes the combustion of liquid ethanol at constant temperature and pressure. , The negative sign of enthalpy change indicates that this is an exothermic reaction.
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Continue…, Thermochemical Equations, remember the following conventions regarding thermochemical equations., The coefficients in a balanced thermochemical equation refer to the number of moles (never molecules) of reactants and products involved in the reaction., The numerical value of ΔrHƟ refers to the number of moles of substances specified by an equation. Standard enthalpy change ΔrHƟ will have units as kJ mol–1.
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Continue…, Thermochemical Equations, The unit for ΔrHƟ is kJ mol–1, which means per mole of reaction., Once we balance the chemical equation in a particular way, this defines the mole of reaction.
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Continue…, Thermochemical Equations, When a chemical equation is reversed, the value of ΔrHƟ is reversed in sign. For example
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Hess’s Law of Constant Heat Summation, Enthalpy is a state function , Enthalpy change for a reaction is the same whether it occurs in one step or in a series of steps., Statement:, If a reaction takes place in several steps then its standard reaction enthalpy is the sum of the standard enthalpies of the intermediate reactions into which the overall reaction may be divided at the same temperature.
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Continue…, Hess’s Law of Constant Heat Summation
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Standard Enthalpy of Combustion (symbol : ΔCHƟ), Combustion reactions are exothermic in nature., Standard enthalpy of combustion is defined as the enthalpy change per mole (or per unit amount) of a substance, when it undergoes combustion and all the reactants and products being in their standard states at the specified temperature., Cooking gas in cylinders contains mostly (C4H10).
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Enthalpy of Atomization (symbol: ΔaHƟ), It is the enthalpy change on breaking one mole of bonds completely to obtain atoms in the gas phase., , In case of diatomic molecules, like dihydrogen , the enthalpy of atomization is also the bond dissociation enthalpy., Other examples, In this case, the enthalpy of atomization is, same as the enthalpy of sublimation
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Bond Enthalpy (symbol: ΔbondHƟ), (i) Bond dissociation enthalpy, (ii) Mean bond enthalpy, (i) Bond dissociation enthalpy, For diatomic molecules the bond dissociation enthalpy is the change in enthalpy when one mole of covalent bonds of a gaseous covalent compound is broken to form products in the gas phase., it is the same as the enthalpy of atomization
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(ii) Mean bond enthalpy, Used for Polyatomic Molecules, In methane, all the four C – H bonds are identical in bond length and energy. However, the energies required to break the individual C – H bonds in each successive step differ :
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Lattice Enthalpy, The lattice enthalpy of an ionic compound is the enthalpy change which occurs when one mole of an ionic compound dissociates into its ions in gaseous state., , , Since it is impossible to determine lattice enthalpies directly by experiment, we use an indirect method where we construct an enthalpy diagram called a Born-Haber Cycle
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The importance of the cycle is that, the sum of the enthalpy changes round a cycle is zero.
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Enthalpy of Solution (symbol : ΔsolHƟ ), Enthalpy of solution of a substance is the enthalpy change when one mole of it dissolves in a specified amount of solvent. , The enthalpy of solution at infinite dilution is the enthalpy change observed on dissolving the substance in an infinite amount of solvent when the interactions between the ions (or solute molecules) are negligible.
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For most of the ionic compounds, Δsol HƟ is positive and the dissolution process is endothermic. , Therefore the solubility of most salts in water increases with rise of temperature. , If the lattice enthalpy is very high, the dissolution of the compound may not take place at all., Continue…, Enthalpy of Solution (symbol : ΔsolHƟ )
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Enthalpy of Dilution, It is the heat change when additional solvent is added to the solution. , The enthalpy of dilution of a solution is dependent on the original concentration of the solution and the amount of solvent added., Enthalpy change for dissolving one mole of gaseous hydrogen chloride in 10 mol of water (10 aq.), The values of enthalpy of solution depend on amount of solvent., Enthalpy of solution approaches a limiting value, i.e, the value in infinitely dilute solution(S-3)
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Subtract equation S-1 from the equation S-2 we obtain–, Continue…, Enthalpy of Dilution, This value (–0.76kJ/mol) of ΔH is enthalpy of dilution. , It is the heat withdrawn from the surroundings when additional solvent is added to the solution.
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Spontaneous Process , A spontaneous process is one that occurs on its own or once started proceeds, without the external input of energy. , It cannot be reversed on its own. It is unidirectional , All natural processes are spontaneous., For example, , a gas expanding to fill the available volume, , burning carbon in dioxygen giving carbon dioxide, Rolling of ball downhill, Rusting of iron, Freezing of water at temperatures below 0oC, Melting of ice
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Criterion for Spontaneity-Driving force for spontaneous process, Decrease in Enthalpy, Spontaneous physical process, flow of water down hill , fall of a stone on to the ground, there is a net decrease in potential energy in the direction of change., Exothermic chemical reactions are spontaneous because decrease in energy has taken place, Driving force for a spontaneous process is ‘decrease in energy (enthalpy)’ i.e. - ΔH
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Now examine the following reactions:, Continue…, Criterion for Spontaneity-Driving force for spontaneous process, Decrease in Enthalpy, These reactions though endothermic (ΔH is positive), are spontaneous., Therefore, it is concluded that decrease in enthalpy may be a contributory factor for spontaneity, but it is not true for all cases.
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Continue…, Criterion for Spontaneity-Driving force for spontaneous process, Another criteria, examine such a case in which ΔH = 0 i.e., there is no change in enthalpy, but still the process is spontaneous., Consider diffusion of two gases into each other in a closed container which is isolated from the surroundings
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Continue…, Criterion for Spontaneity-Driving force for spontaneous process, Another criteria, We may now formulate another postulate:, in an isolated system, there is always a tendency for the systems’ energy to become more disordered or chaotic and this could be a criterion for spontaneous change !, To express this disorder of system, another thermodynamic term is used which is called Entropy
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Entropy (S), Entropy is a measure of the degree of randomness or disorder in the system., The greater the disorder in an isolated system, the higher is the entropy., Qualitatively, Order of entropy :-Solid(lowest entropy) < Liquid < gas (highest entropy), Entropy is a state function and ΔS is independent of path., Whenever heat is added to the system, it increases molecular motions causing increased randomness in the system., A system at higher temperature has greater randomness in it than one at lower temperature., Heat added to a system at lower temperature causes greater randomness than when the same quantity of heat is added to it at higher temperature. , This suggests that the entropy change is inversely proportional to the temperature.
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ΔS is related with q and T for a reversible reaction as, The total entropy change ( ΔStotal) for the system and surroundings of a spontaneous process is given by, , When a system is in equilibrium, the entropy is maximum, and the change in entropy, ΔS = 0., Entropy for a spontaneous process increases till it reaches maximum and at equilibrium the change in entropy is zero., Since entropy is a state property, we can calculate the change in entropy of a reversible process by, Continue…, Entropy (S)
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Continue…, Entropy (S), Both for reversible and irreversible expansion for an ideal gas, under isothermal conditions, ΔU = 0, but ΔStotal i.e., ΔSsys + ΔSsurr is not zero for irreversible process., Thus, ΔU does not discriminate between reversible and irreversible process, whereas ΔS does., ΔStotal > 0 Process is spontaneous, ΔStotal < 0 Process is Non-spontaneous, ΔStotal = 0 System is in equilibrium, ΔSsys alone can not decide the spontaneity of the process for closed and open system., Conclusion:-neither decrease in enthalpy nor increase in entropy alone can determine the direction of spontaneous change for closed and open systems
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Gibbs Energy or Gibbs Function, G, G = H – TS Gibbs function, G is an extensive property and a state function., Its absolute value also can not be determined., The change in Gibbs energy for the system, ΔGsys can be written as, ΔG has units of energy i.e J, ΔH is the enthalpy change of a reaction, TΔS is the energy which is not available to do useful work. , So ΔG is the net energy available to do useful work and is thus a measure of the ‘free energy’., For this reason, it is also known as the free energy of the reaction.
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Gibbs Energy and Spontaneity, If the system is in thermal equilibrium with the surrounding, then the temperature of the surrounding is same as that of the system., Also, increase in enthalpy of the surrounding is equal to decrease in the enthalpy of the system., Therefore, entropy change of surroundings,
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Continue…, Gibbs Energy and Spontaneity, ΔStotal > 0 Process is spontaneous ΔG < 0 (Negative) , ΔStotal < 0 Process is Non-spontaneous ΔG > 0 (Positive) , ΔStotal = 0 System is in equilibrium ΔG = 0 (Equilibrium)
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Effect of Temperature on Spontaneity of Reactions
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Entropy and Second Law of Thermodynamics, Entropy is the loss of energy available to do work. , Another form of the second law of thermodynamics states that the total entropy of a system either increases or remains constant; it never decreases. , Entropy is zero in a reversible process; it increases in an irreversible process., Entropy of any isolated system always increases., In exothermic reactions heat released by the reaction increases the disorder of the surroundings and overall entropy change is positive which makes the reaction spontaneous.
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Absolute Entropy and Third Law of Thermodynamics, The entropy of any pure crystalline substance approaches zero as the temperature approaches absolute zero. This is called third law of thermodynamics., The importance of the third law is that absolute values of entropy of pure substance can be determined.
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GIBBS ENERGY CHANGE AND EQUILIBRIUM
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Continue…, GIBBS ENERGY CHANGE AND EQUILIBRIUM, Gibbs energy change for a reaction in which all reactants and products are in standard state, is known as Standard Gibbs energy change ΔrGƟ, It is related to the equilibrium constant of the reaction as follows:, ΔrG = ΔrGƟ + RT ln K at equilibrium ΔrG =0 so, ΔrGƟ= – RT ln K, or ΔrGƟ = – 2.303 RT log K where K = equilibrium constant, ΔrGƟ = ΔrHƟ - T ΔrSƟ = – 2.303 RT log K , For strongly endothermic reactions, the ΔrHƟ large and positive, value of K << 1 and the reaction is unlikely to form much product. , In case of exothermic reactions, ΔrHƟ large and negative, and ΔrGƟ large and negative, K >> 1 hence can go to near completion.
