Thermodynamics - Principles of Chemistry II - Lecture Slides, Slides of Chemistry

Download Thermodynamics - Principles of Chemistry II - Lecture Slides and more Slides Chemistry in PDF only on Docsity! Chemical Thermodynamics Spontaneous Processes Reversible Processes Review First Law Irreversible Processes Second Law Third Law Entropy Temperature Dependence Gibbs Free Energy Equilibrium Constant 9/7/2013 • Thermodynamics is concerned with the question: can a reaction occur? • First Law of Thermodynamics: energy is conserved. • Any process that occurs without outside intervention is spontaneous. • When two eggs are dropped they spontaneously break. • The reverse reaction is not spontaneous. • We can conclude that a spontaneous process has a direction. Spontaneous Processes Reversible and Irreversible Processes • Chemical systems in equilibrium are reversible. • In any spontaneous process, the path between reactants and products is irreversible. • Thermodynamics gives us the direction of a process. It cannot predict the speed at which the process will occur. Spontaneous Processes The Spontaneous Expansion of a Gas • Why do spontaneous processes occur? • Consider an initial state: two flasks connected by a closed stopcock. One flask is evacuated and the other contains 1 atm of gas. • The final state: two flasks connected by an open stopcock. Each flask contains gas at 0.5 atm. • The expansion of the gas is isothermal (i.e. constant temperature). Therefore the gas does no work and heat is not transferred. Entropy and the Second Law of Thermodynamics The Spontaneous Expansion of a Gas • Why does the gas expand? Entropy and the Second Law of Thermodynamics Entropy • Generally, when an increase in entropy in one process is associated with a decrease in entropy in another, the increase in entropy dominates. • Entropy is a state function. • For a system, S = Sfinal - Sinitial. • If S > 0 the randomness increases, if S < 0 the order increases. Entropy and the Second Law of Thermodynamics Entropy • Suppose a system changes reversibly between state 1 and state 2. Then, the change in entropy is given by – at constant T where qrev is the amount of heat added reversibly to the system. (Example: a phase change occurs at constant T with the reversible addition of heat.) Entropy and the Second Law of Thermodynamics )(constant revsys T T q S  The element mercury, Hg, is a silvery liquid at room temperature. The normal freezing point of mercury is -38.9oC, and its molar enthalpy of fusion is 2.29 kJ/mol. What is the entropy change of the system when 50.0 g of Hg(l) freezes at the normal freezing point? MW(Hg) = 200.59 g/mol The element mercury, Hg, is a silvery liquid at room temperature. The normal freezing point of mercury is -38.9oC, and its molar enthalpy of fusion is 2.29 kJ/mol. What is the entropy change of the system when 50.0 g of Hg(l) freezes at the normal freezing point? MW(Hg) = 200.59 g/mol -2.44 J K-1 Calculate ΔSsys , ΔSsurr , ΔSuniv for the reversible melting of 1 mole of Ice in a large, isothermal water bath at 0oC and 1 atm. Heat of Fusion for water is 6.01 kJ/mol. Ssys = +22.0 J mol -1 K-1 • A gas is less ordered than a liquid that is less ordered than a solid. • Any process that increases the number of gas molecules leads to an increase in entropy. • When NO(g) reacts with O2(g) to form NO2(g), the total number of gas molecules decreases, and the entropy decreases. The Molecular Interpretation of Entropy HyperChem The Molecular Interpretation of Entropy • Boiling corresponds to a much greater change in entropy than melting. • Entropy will increase when – liquids or solutions are formed from solids, – gases are formed from solids or liquids, – the number of gas molecules increase, – the temperature is increased. The Molecular Interpretation of Entropy ENTROPY OF THE UNIVERSE Die Enegie der Welt ist Konstant; die entropie der Welt einen Maximum Zu Entropy is the Arrow of Time A Brief History of Time • Absolute entropy can be determined from complicated measurements. • Standard molar entropy, S: entropy of a substance in its standard state. Similar in concept to H. • Units: J/mol-K. Note units of H: kJ/mol. • Standard molar entropies of elements are not zero. • For a chemical reaction which produces n moles of products from m moles of reactants: Entropy Changes in Chemical Reactions      reactantsproducts mSnSS • For a spontaneous reaction the entropy of the universe must increase. • Reactions with large negative H values are spontaneous. • How do we balance S and H to predict whether a reaction is spontaneous? • Gibbs free energy, G, of a state is • For a process occurring at constant temperature Gibbs Free Energy TSHG  STHG  • There are three important conditions: – If G < 0 , forward reaction is spontaneous. – If G = 0 , reaction is at equilibrium. – If G > 0 , then forward reaction is not spontaneous. If G > 0, work must be supplied from the surroundings to drive the reaction. • For a reaction the free energy of the reactants decreases to a minimum (equilibrium) and then increases to the free energy of the products. Gibbs Free Energy Potential energy —> Equilibrium position in valley Position Free energy —> Reactants Equilibrium mixture Products Course of reaction  CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g) ΔGof (kJ/mol): -50.8 0 -394.4 -228.6 -800.8 kJ/mol TABLE 19.4 Effect of Temperature on the Spontaneity of Reactions AH AS —TAS AG = AH — TAS Reaction Characteristics Example = + = Always negative Spontaneous at all temperatures 203(g¢) —— 302(g) + - + Always positive Nonspontaneous at all temperatures; 302(¢) ——> 203(g) reverse reaction always spontaneous - = + Negative at low T; Spontaneous at low T; becomes H,O(/) —> H,O(s) positive at high T nonspontaneous at high T + + - Positive at low T; Nonspontaneous at low T; becomes H,O(s) —> H,O(/) negative at high T spontaneous at high T  • At equilibrium, G = 0 : • From the above we can conclude: – If G < 0, then K > 1. – If G = 0, then K = 1. – If G > 0, then K < 1. Free Energy and The Equilibrium Constant eqKRTG ln Given ΔGo = -33.3 kJ/mol for the ammonia formation from nitrogen and hydrogen, calculate Keq at 25.0 oC.