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PRINCIPLES OF CHEMISTRY II CHEM 1212 CHAPTER 17 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state university CHAPTER 17 CHEMICAL THERMODYNAMICS CHEMICAL THERMODYNAMICS Used to - Describe heat transfer in chemical systems - Evaluate the heat evolved or absorbed by a reaction - Predict the maximum energy that can be produced by a reaction - Predict whether a proposed reaction is feasible - Predict whether a process is spontaneous or not CHEMICAL THERMODYNAMICS Spontaneous Process - Takes place with no apparent cause Nonspontaneous Process - Requires something to be applied in order for it to occur (usually in the form of energy) ENERGY - The ability to do work or to transfer heat - Energy is necessary for life: humans, plants, animals, cars - Forms of energy are interconvertible - Units: kg∙m2/s2 or Joules (J) SYSTEM AND SURROUNDINGS System - The limited and well-defined portion of the universe under study Surroundings - Everything else in the universe Studying energy changes in a chemical reaction - The reactants and products make up the system - The reaction container and everything else make up the surroundings SYSTEM AND SURROUNDINGS Open System - Matter and energy can be exchanged with the surroundings (water boiling on a stove without a lid) Closed System - Energy but not matter can be exchanged with the surroundings (two reactants in a closed cylinder reacting to produce energy) Isolated System - Neither matter nor energy can be exchanged with the surroundings (insulated flask containing hot tea) WORK (w) - The energy transferred when a force moves an object - The product of force (F) and distance (d) through which the object moves w=Fxd - Units: kg∙m2/s2 or Joules (J) HEAT (q) - The energy transferred between a system and its surroundings due to difference in temperature - A form of energy necessary to raise the temperature of a substance - Units: kg∙m2/s2 or Joules (J) INTERNAL ENERGY (E) - Sum of all potential and kinetic energies of all components - Change in internal energy = final energy minus initial energy E = Efinal - Einitial - Energy can neither be created nor destroyed - Energy is conserved INTERNAL ENERGY E = Efinal - Einitial If Efinal > Einitial E is positive and system has gained energy from its surroundings If Efinal < Einitial E is negative and system has lost energy to its surroundings INTERNAL ENERGY E = q + w q = heat added to or liberated from a system w = work done on or by a system Internal energy of a system increases when - Heat is added to the system from surroundings (positive q) - Work is done on the system by surroundings (positive w) +w +q system INTERNAL ENERGY E = q + w q = heat added to or liberated from a system w = work done on or by a system Internal energy of a system decreases when - Heat is lost by the system to the surroundings (negative q) - Work is done by the system on the surroundings (negative w) -w -q system INTERNAL ENERGY Endothermic Process - Process in which system absorbs heat (endo- means ‘into’) - Heat flows into system from its surroundings (melting of ice - the reason why it feels cold) Exothermic Process - Process in which system loses heat - Heat flows out of the system (exo- means ‘out of’) (combustion of gasoline) INTERNAL ENERGY State Function - Property that depends on initial and final states of the system - Does not depend on path or how a change occurs - Internal Energy depends on initial and final states - Internal Energy is a state function - q and w, on the other hand, are not state functions INTERNAL ENERGY Internal energy is influenced by - Temperature - Pressure - Total quantity of matter - Internal Energy is an extensive property INTERNAL ENERGY Calculate E for a system absorbing 22 kJ of heat from its surroundings while doing 11 kJ of work on the surroundings. State whether it is an endothermic or an exothermic process q = +22 kJ (heat is added to the system from surroundings) w = -11 kJ (work is done by the system on the surroundings) E = q + w E = (+ 22 – 11) kJ = +11 kJ Endothermic PRESSURE-VOLUME WORK A gas against constant pressure w = - PV w = work and P = pressure V = change in volume = Vfinal – Vinitial w is positive when the gas contracts (negative V) w is negative when the gas expands (positive V) Units: L·atm (1 L·atm = 101.3 J) PRESSURE-VOLUME WORK Calculate the work associated with the expansion of a gas from 32 L to 58 L at a constant pressure of 12 atm w = - PV w = - (12 atm)(58 L - 32 L) = - 310 L·atm Gas expands hence work is done by system on surroundings PRESSURE-VOLUME WORK Calculate the work associated with the compression of a gas from 58 L to 32 L at a constant pressure of 12 atm w = - PV w = - (12 atm)(32 L - 58 L) = + 310 L·atm Gas contracts hence work is done on system by surroundings PRESSURE-VOLUME (P-V) WORK Expansion of Volume - V is a positive quantity and w is a negative quantity - Energy leaves the system as work - Work is done by the system on the surroundings Compression of Volume - V is a negative quantity and w is a positive quantity - Energy enters the system as work - Work is done on the system by the surroundings THE FIRST LAW OF THERMODYNAMICS - The law of conservation of energy - Energy can be neither created nor destroyed - Concerned with change in energy In an isolated system - Neither matter nor energy can enter or leave E = 0 In a closed system - Energy can enter or leave in the form of heat and work E = q + w ENTHALPY (H) - Heat flow in processes occurring at constant pressure - Only P-V work are performed H = E + PV H, E, P, and V are all state functions Change in Enthalpy H = (E + PV) ENTHALPY (H) Change in Enthalpy at Constant Pressure H = E + PV E = q + w PV = - w Implies H = (qp + w) - w = qp qp = heat at constant pressure ENTHALPY (H) H = qp Change in enthalpy = heat gained or lost at constant pressure Positive H - System gains heat from the surroundings - Endothermic process Negative H - System releases heat to the surroundings - Exothermic process ENTHALPY (H) H = E + PV For reactions involving solids and liquids V ≈ 0 H ≈ E For gases n = nfinal – ninitial n = total gas moles of products – total gas moles of reactants Implies PV = (n)RT H = E + (n)RT (R = 8.314 J/mol·K) ENTHALPY (H) E = q + w At constant volume V = 0 w = - PV = 0 and E = qv qv = heat gained or lost at constant volume ENTROPY (S) - The amount of disorder in a process - Is a measure of randomness - Many spontaneous reactions are accompanied by release of energy (exothermic processes) - Some endothermic processes, however, are spontaneous (dissolution of some salts such as barium hydroxide) - Disorder plays an important role in predicting the spontaneity of a reaction ENTROPY (S) - Entropy is a state function Change in entropy ΔS = Sfinal – Sinitial ENTROPY (S) Some General Concepts - The entropy of a substance increases as the substance changes from solid to liquid to gas - Due to increase in randomness of the molecules Ssolid < Sliquid < Sgas Generally ΔS from liquid to gas > ΔS from solid to liquid ENTROPY (S) Some General Concepts - The entropy of a substance increases when a molecular solid or liquid dissolves in a solvent Solid → Aqueous - Increase in disorder when a solute dissolves in a solvent - Increase in disorder of the solute is greater than the increase in disorder of the solvent - Net decrease in entropy occurs with some ionic solids (solids with highly charged ions due to hydration) ENTROPY (S) Some General Concepts - The entropy decreases when a gas dissolves in a solvent CO2(g) → CO2(aq) - Solute molecules go from the gas phase to the liquid phase - Solute molecules become less random - Increase in disorder of the solvent is small - Hence a net decrease in disorder ENTROPY (S) Some General Concepts - Entropy increases with increase in temperature - Kinetic energy of particles increase as temperature increases - Disorder increases as kinetic energy (energy of motion) increases ENTROPY (S) Other Concepts - Entropy increases in a chemical reaction when ∆n is positive - Entropy decreases in a chemical reaction when ∆n is negative - Entropy of a system increases with increasing volume - Entropy of a solution increases with dilution - Osmosis (spontaneous process) is driven by positive ΔS THE SECOND LAW OF THERMODYNAMICS - For a spontaneous process there is always an increase in the entropy of the universe ΔSuniv = ΔSsys + ΔSsurr ΔSuniv > 0 for a spontaneous process If ΔSsys is negative, ΔSsurr must be large and positive to make ΔSuniv positive Example - Combustion of hydrocarbons ΔSsys is negative but ΔSuniv is positive THE SECOND LAW OF THERMODYNAMICS If ΔS < 0 - The reverse reaction is spontaneous If ΔS = 0 - The system is at equilibrium - The process is not spontaneous in either direction - The process has no tendency to occur THE THIRD LAW OF THERMODYNAMICS - The entropy of any pure crystalline substance at a temperature of 0 K is zero - Absolute value of S can be measured - As the temperature of a substance is increased from 0 K, the motion of the particles increases and entropy increases THE THIRD LAW OF THERMODYNAMICS - Change in entropy is proportional to the added energy - The transfer of a given amount of energy as heat has greater impact on entropy at lower temperatures than at higher temperatures - ΔS is directly proportional to quantity of heat transferred (q) and inversely proportional to temperature (T) Units: J/K q ΔS T THE THIRD LAW OF THERMODYNAMICS - The entropy change of a chemical reaction ∆Srxn = ΣnSo[products] - ΣmSo[reactants] So = standard molar entropy of a substance (at 298 K) - n and m are the number of moles of products and reactants - The standard molar entropy of an element in its standard state is not zero THE THIRD LAW OF THERMODYNAMICS Calculate ∆So at 25 oC for the reaction 2NiS(s) + 3O2(g) → 2SO2(g) + 2NiO(s) Obtain So values from appendix ∆So = [(2 mol)(248.11 J/mol·K) + (2 mol)(37.99 J/mol·K)] – [(2 mol)(52.99 J/mol·K) + (3 mol)(205.03 J/mol·K)] = -148.87 J/K (∆So is negative as ∆n is negative) GIBBS FREE ENERGY (G) G = H – TS - Change in Gibbs free energy at constant temperature and pressure ∆G = ∆H – T∆S - The absolute value of G cannot be measured but ∆G can be measured - ∆G is a state function GIBBS FREE ENERGY (G) If ∆G < 0 - Forward reaction is spontaneous If ∆G = 0 - System is at equilibrium If ∆G > 0 - Reverse reaction is spontaneous GIBBS FREE ENERGY (G) Standard Gibbs Free energy of formation (∆Gfo) - The Gibbs free energy change during the formation of one mole of a substance in its standard state from its constituent elements in their standard states ∆Gfo = ∆Hfo – T∆Sfo GIBBS FREE ENERGY (G) Change in Gibbs free energy of a chemical reaction ∆Gorxn = Σn∆Gfo[products] - Σm∆Gfo[reactants] - n and m are the number of moles of products and reactants - The standard Gibbs free energy of an element in its standard state is zero GIBBS FREE ENERGY (G) Influence of Temperature ∆G = ∆H - T∆S From the equation above - Decrease in ∆H (more negative) favor spontaneous change - Increase in ∆S (more positive) favor spontaneous change - ∆G is strongly influenced by temperature through the T∆S term GIBBS FREE ENERGY (G) Influence of Temperature At Low Temperatures - The sign of ∆H determines the sign of ∆G At High Temperatures - The sign of ∆S determines the sign of ∆G - When ∆H and ∆S have opposite signs temperature change does not influence the direction of spontaneity - It is assumed that the numerical values of ∆H and ∆S are not affected by temperature change GIBBS FREE ENERGY (G) Equilibrium Constant ∆G = ∆Go + RTlnQ ∆G = Gibbs free-energy change at non-standard-state conditons ∆Go = standard Gibbs free-energy change R = ideal gas constant (8.314 J/mol·K) T = temperature (K) Q = reaction quotient GIBBS FREE ENERGY (G) Equilibrium Constant ∆G = ∆Go + RTlnQ At equilibrium ∆G = 0 and Q = Keq Implies ∆Go = - RTlnKeq or K eq e ΔG o /RT GIBBS FREE ENERGY (G) In Summary - Reaction is spontaneous in the forward direction if ∆G is negative - Reaction is spontaneous in the reverse direction if ∆G is positive - Reaction is spontaneous in the forward direction if Q < Keq - Reaction is spontaneous in the reverse direction if Q > Keq - At equilibrium ∆G approaches zero and Q approaches Keq GIBBS FREE ENERGY (G) Temperature and Equilibrium Constant ∆Go = - RTlnKeq Combining and ∆Go = ∆Ho - T∆So We obtain ∆Ho - T∆So = - RTlnKeq and lnK eq ΔS o ΔH o R RT GIBBS FREE ENERGY (G) Temperature and Equilibrium Constant - Equilibrium constant changes with temperature - If K1 and K2 are the equilibrium constant values at temperatures T1 and T2, respectively K1 ΔH o 1 1 ln K2 R T2 T1 - Known as the Clausius-Clapeyron equation GIBBS FREE ENERGY (G) Useful Work - For a spontaneous process at constant temperature and pressure - The maximum useful work that can be performed equals the change in Gibbs free energy wmax = ∆G - The energy that is free to perform useful work - The equation gives the minimum amount of work required to cause a change when ∆G is positive