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2. Chemical Thermodynamics • Chemical reactions obey two fundamental laws: 1. The law of conservation of mass • States that matter can be neither created nor destroyed • Explains why equations must balance and is the basis for stoichiometry and equilibrium calculations • Stoichiometry that allows us to compare apples and oranges • Equilibrium predictions of reversible reactions which leads to • Kinetics allowing us to determine how fast the reaction will occur 2. The law of conservation of energy • States that energy can be neither created nor destroyed • Energy takes various forms that can be converted from one to the other

3. Some Thermodynamic Terms • Thermodynamics - The study of the relationship between heat, work, and other forms of energy of a system at equilibrium. Predicts whether a particular reaction is energetically possible in the direction as written and the composition of the reaction system at equilibrium. Thermodynamics does not say whether an energetically feasible reaction will actually occur as written. Thermodynamics tells nothing about the rate of the reaction or the pathway by which it will occur. • Thermochemistry- A branch of thermodynamics which focuses on the study of heat given off or absorbed in a chemical reaction. • Temperature - An intensive property of matter; a quantitative measurement of the degree to which an object is either "hot" or "cold". There are 3 scales: • Fahrenheit- relative • 32 ◦F is the normal freezing point temperature of water; 212 ◦ F is the normal boiling point temperature of water. • Celsius (centigrade) - relative • 0 ◦ C is the normal freezing point temperature of water; 100 ◦ C is the normal boiling point temperature of water. • Kelvin - absolute • 0 K is the temperature at which the volume and pressure of an ideal gas extrapolate to zero.

4. Standard States and Standard Enthalpy Changes • Thermochemical standard state conditions • The thermochemical standard T = 298.15 K. • The thermochemical standard P = 1.0000 atm. • Be careful not to confuse these values with STP. • Thermochemical standard states of matter • For pure substances in their liquid or solid phase the standard state is the pure liquid or solid. • For gases the standard state is the gas at 1.00 atm of pressure. • For gaseous mixtures the partial pressure must be 1.00 atm. • For aqueous solutions the standard state is 1.00 M concentration.

5. Thermodynamics and Work • A system is that part of the universe in which we are interested (in chemistry this is the reactant side of the chemical equation); the surroundings are everything else—the rest of the universe. • System + surroundings = universe. • A closed systemcan exchange energy but not matter with its surroundings; an open systemexchanges both, and an isolated system exchanges neither. • State function — the property of a system that depends only on the present state of the system and not on its history. • State Functions are independent of pathway: • T (temperature), P (pressure), V (volume), E (change in energy), H (change in enthalpy – the transfer of heat), and S (entropy) • Examples of non-state functions are: • n (moles), q (heat), w (work) • A change in state function depends only on the difference between the initial and final states, not on the pathway used to go from one to the other. • Thermodynamics is concerned with state functions and does not deal with how the change between the initial and final state occurs.

6. Some Thermodynamic Terms • Latent Heat versus Sensible Heat • Sensible heat- Heat that can be detected by a change in the temperature of a • system. • Latent heat- Heat that cannot be detected because there is no change in • temperature of the system. e.g.  The heat that is used to melt ice or to evaporate water is latent heat. • There are two forms of latent heat: • Heat of fusion- The heat that must be absorbed to melt a mole of a solid. • e.g.  melting ice to liquid water • Heat of vaporization- The heat that must be absorbed to boil a mole of a liquid. • e.g.  boiling liquid water to steam

8. Thermochemical equationsare a balanced chemical reaction plus the H value for the reaction. For example, this is an exothermic thermochemical equation. The stoichiometric coefficients in thermochemical equations must be interpreted as numbers of moles. 1 mol of C5H12 reacts with 8 mol of O2 to produce 5 mol of CO2, 6 mol of H2O, and releasing 3523 kJ is referred to as one mole of reactions. Thermochemical Equations • this is an endothermic thermochemical equation. H2O(S) + 6.02 kJ →H2O(l)

9. Some Thermodynamic Terms • Heat (q)- A form of energy associated with the random motion of the elementary particles in matter. • Heat capacity- The amount of heat needed to raise the temperature of a defined • amount of a pure substance by one degree. • Specific heat- The amount of heat needed to raise the temperature of one gram of a • substance by 1 C (or 1 K) • SI unit for specific heat is joules per gram-1 Kelvin-1 (J/g-K) • Calorie- The specific heat of water = 4.184 J/g-K • Molar heat capacity- The amount of heat required to raise the temperature of one • moleof a substance by 1 C (or 1 K) • SI unit for molar heat capacity is joules per mole-1 Kelvin-1(J/mol-K) • Btu (British thermal unit)- The amount of heat needed to raise the temperature of • 1 lb water by 1 F. • NOTE:  The specific heat of water (4.184 J/g-K) is very large relative to other • substances.  The oceans (which cover over 70% of the earth) act as a giant "heat sink," • moderating drastic changes in temperature.  • Our body temperatures are also controlled by water and its high specific heat.  • Perspiration is a form of evaporative cooling which keeps our body temperatures from • getting too high.

10. Some Thermodynamic Theories • Caloric Theory of Heat • Served as the basis of thermodynamics. • Is now known to be obsolete • Based on the following assumptions • Heat is a fluid that flows from hot to cold substances. • Heat has a strong attraction to matter which can hold a lot of heat. • Heat is conserved. • Sensible heat causes an increase in the temperature of an object when it flows into the object. • Latent heat combines with particles in matter (causing substances to melt or boil) • Heat is weightless. • The only valid part of the caloric theory is that heat is weightless. • Heat is NOT a fluid, and it is NOT conserved.

11. Some Thermodynamic Theories Kinetic Theory of Heat • Divides the universe into two parts: • System. - The substances involved in the chemical and physical changes under investigation: In chemistry lab, the system is the REACTANTS inside the beaker. • Surroundings- Everything not included in the system, i.e. the rest of the universe. • A BOUNDARY separates the system and the surroundings from each other and can be tangible or imaginary. • Heat is something that is transferred back and forth across boundary between a system and its surroundings • Heat is NOT conserved.

12. Some Thermodynamic Theories • The kinetic theory of heat is based upon the last postulate in the kinetic molecular theory which states that the average kinetic energy of a collection of gas particles is dependent only upon the temperature of the gas.  • where R is the ideal gas constant (0.08206 L-atm/mol-K) and T is temperature (Kelvin) The kinetic theory of heat can be summarized as follows: • The set of conditions that specify all of the properties of the system is called the thermodynamic state of a system. • For example the thermodynamic state could include: • The number of moles and identity of each substance. • The physical states of each substance. • The temperature of the system. • The pressure of the system.

13. The Three Laws of Thermodynamics • There are two basic ideas of importance for thermodynamic systems. • Chemical systems tend toward a state of minimum potential energy. • Chemical systems tend toward a state of maximum disorder. • When: • H > 0 disorder increases (which favors spontaneity). • H < 0 disorder decreases (does not favor spontaneity). • When: • S > 0 disorder increases (which favors spontaneity). • S < 0 disorder decreases (does not favor spontaneity).

14. The First Law of Thermodynamics • The first law is also known as the Law of Conservation of Energy. Energy is neither created nor destroyed in chemical reactions and physical changes. • The energy of the universe does not change. • The energy in a system may change, but it must be complemented by a change in the energy of its surroundings to balance the change in energy. • The term internal energy is often used synonymously with the energy of a system.  It is the sum of the kinetic and potential energies of the particles that form the system.  The change in energy of the system is identical in magnitude but opposite in sign to the change in energy of the surroundings.

15. The First Law of Thermodynamics If a system is more complex than an ideal gas, then the internal energy must be measured indirectly by observing any changes in the temperature of the system.  The change in the internal energy of a system is equal to the difference between the final and initial energies of the system: The equation for the first law of thermodynamics can be rearranged to show the energy of a system in terms of the energy of its surroundings. This equation indicates that the energy lost by one must equal the energy gained by the other: • Esys = KEsys + PEsys • KE – kinetic energy: translational, rotational, vibrational • PE – energy stored in bonds (Bond energy)

16. The First Law of Thermodynamics • The energy of a system can change by the transfer of work and or heat between the system and its surroundings.  Any heat that is taken, given off, or lost must be complemented by an input of work to make up for the loss of heat.  Conversely, a system can be used to do any amount of work as long as there is an input of heat to make up for the work done. • This equation can be used to explain the two types of heat that can be added to a system: • Heat can increase the temperature of a system.  This is sensible heat. • Heat that does ONLY WORK on a system is latent heat. • Any machine that converts energy to work is designed to want to maximize the amount of work obtained and to minimize the amount of energy released to the environment as heat

17. The First Law of Thermodynamics Most chemical reactions occur at constant P, so Heat transferred at constant P = qp qp = ∆H where H = enthalpy and so ∆E = ∆H + w (and w is usually small) ∆H = heat transferred at constant P ≈ ∆E ∆H = change in heat content of the system ∆H = Hfinal - Hinitial ∆Horxn =  ∆Hfo(prod) -  ∆Hfo(react) C3H8(g) + 5 O2(g) )  3 CO2(g) + 4 H2O(g) -103.8 kJmol-1 0 kJmol-1 -393.5 kJmol-1 -241.8 kJmol-1

18. The First Law of Thermodynamics • Exchange of heat (q) Endothermic and exothermic • Work is performed (w) • DE = q + w Gases Why only gases? Because changes in volume results in work w = Fd F = Pressure x Area d = Dh W = P (A Dh) = DV Solids, Liquids, Solutions Changes in volume are negligible Therefore w is effectively zero DE = q + 0 = DH DH is change in enthalpy which is the transfer of heat and is measured experimentally by determining changes in temperature.

19. The First Law of Thermodynamics heat transfer in (endothermic), +q heat transfer out (exothermic), -q w transfer in (+w) Compression of system w transfer out (-w) Expansion of system SYSTEM ∆E = q + w By convention except for some engineers whose frame of reference is the work done on the surroundings. hi hf A(hf-hi)>0 DV w = -PDV DE =DH + w = DH – PDV = DH – D(PV) A(hf-hi)<0 DV hi hf

20. DE = DH – D(PV) Constant Volume w = -PDV DV = 0 DE = q + 0 = DH Check the temperature change Constant Pressure Apply some stoichiometry And the Ideal Gas Law PV=nRT D(PV)=D(nRT) Hold Temperature constant k1 D(PV)=D(nRk1) Combine constants and multiply through by -1 -D(PV) = -R1Dn w = -PDV = -R1Dn DE = DH + w = DH - R1Dn

21. Thermochemical Equations Write the thermochemical equation for CuSO4(aq) + 2NaOH(aq) Cu(OH)2(s) + Na2SO4(aq) 50.0mL of 0.400 M CuSO4 at 23.35 oC 50.0mL of 0.600 M NaOH at 23.35 oC Tfinal 25.23oC CH2O = 4.184 J/goC Density final solution = 1.02 g/mL

22. The Second Law of Thermodynamics Enthalpy changes are not the only factors that determine whether a process is spontaneous. Most spontaneous reactions are exothermic, but there are many that are not exothermic, however, reactions can be both spontaneous and highly endothermic. • The second law of thermodynamics states, “In spontaneous changes the universe tends towards a state of greater disorder.” • Spontaneous processes have two requirements: • The free energy change of the system must be negative. • The entropy of universe must increase. • Fundamentally, the system must be capable of doing useful work on surroundings for a spontaneous process to occur. Changes in S are usually quite small compared to E and H. Notice that S has units of only a fraction of a kJ while E and H values are much larger numbers of kJ.

23. The Second Law of Thermodynamics Entropy (S)– is the measure of the disorder in a system. Entropy of the universe is unchanged in reversible processes and constitutes part of the second law ofthermodynamics: the entropy of the universe remains constant in a reversible process, whereas the entropy of the universe increases in an irreversible (spontaneous) process.  Entropy is a state function described by the equation: where k is a proportionality constant equal to the ideal gas constant (R) divided by Avogadro's number (6.022 x 10-23) and lnW is the natural log of W, the number of equivalent ways of describing the state of a system. In this reaction, the number of ways of describing a system is directly proportional to the entropy of the system.

24. The Second Law of Thermodynamics Number of Equivalent Combinations for Various Types of Poker Hands

25. The Second Law of Thermodynamics • Entropy of Reaction (DS) • The difference between the sum of the entropies of the products and the sum of the entropies of the reactants: • In the above reaction, n and m are the coefficients of the products and the reactants in the balanced equation. • As with H, entropies have been measured and tabulated. • When: • S > 0 disorder increases (which favors spontaneity). • S < 0 disorder decreases (does not favor spontaneity).

26. The Second Law of Thermodynamics Natural processes that occur in an isolated system are spontaneous when they lead to an increase in the disorder, or entropy, of the system. Isolated system - System in which neither heat nor work can be transferred between it and its surroundings.  This makes it possible to ignore whether a reaction is exothermic or endothermic. If DSsys > 0, the system becomes more disordered through the course of the reaction If DSsys < 0, the system becomes less disordered (or more ordered) through the course of the reaction. • There are a few basic principles that should be remembered to help determine whether a system is increasing or decreasing in entropy. • Liquids are more disordered than solids. • WHY? - Solids have a more regular structure than liquids. • Gases are more disordered than their respective liquids. • WHY? - Gases particles are in a state of  constant random motion. • Any process in which the number of particles in the system increases • consequently results in an increase in disorder. • In general for substances with the same molar mass and number of atoms in its three states of matter, Sº values fall in the : • Sgas> Sliquid > Ssolid

27. The Second Law of Thermodynamics • Does the entropy increase or decrease for the following reactions? • CaCO3(s) → CaO(s) + CO2(g) • N2(g) + 3H2(g) 2NH3(g) • NH4NO3(s) → NH4+(aq) + NO3-(aq) • H2O(g) H2O(l) → → → →

28. Entropy, S • The Third Law of Thermodynamics states, “The entropy of a pure, perfect, crystalline solid at 0 K is zero.” • This law permits us to measure the absolute values of the entropy for substances. • To get the actual value of S, cool a substance to 0 K, or as close as possible, then measure the entropy increase as the substance heats from 0 to higher temperatures.The coldest place in nature is the depths of outer space. There it is 3 degrees above Absolute Zero. • Notice that Appendix L has values of S not S. Predicted 1924... ...Created 1995 Bose-Einstein Condensation in a gas: a new form of matter at the coldest temperatures in the universe... Cornell and Wieman cooled a small sample of atoms down to only a few billionths (0.000,000,001) of a degree above Absolute Zero A. Einstein S. Bose

29. Entropy, S BEC When a gas expands into a vacuum, its entropy increases because the increased volume allows for greater atomic or molecular disorder; the greater the number of atoms or molecules in the gas, the greater the disorder. The magnitude of the entropy of a system depends on the number of microscopic states, or microstates, associated with it; the greater the number of microstates, the greater the entropy.

30. S increases slightly with T Entropy and Temperature S increases a large amount with phase changes

31. Calculating S from Standard Molar Entropy Values • Similar molecular structures have similar Sº values • Those with the lowest entropies tend to be rigid crystals composed of small atoms linked by strong, highly directional bonds • Those with higher entropies are soft crystalline substances that contain larger atoms and increased molecular motion and disorder • Absolute entropy of a substance tends to increase with increasing molecular complexity because the number of available microstates increases with molecular complexity • Substances with strong hydrogen bonds have lower values of Sº, reflecting a more ordered structure • To calculate Sº for a chemical reaction from standard molar entropies, the “products minus reactants” rule is used; here the absolute entropy of each reactant and product is multiplied by its stoichiometric coefficient in the balanced chemical equation

32. Entropy, S • Entropy changes for reactions can be determined similarly to H for reactions.This is only true, i.e. conserved, for the system. This is not included for the surroundings.

33. Entropy, S 240 Jmol-1K-1 304.2 Jmol-1K-1 • Calculate the entropy change for the following reactions at 25oC. 210.6 Jmol-1K-1 219.7 Jmol-1K-1 240 Jmol-1K-1 270.2 Jmol-1K-1 0 Jmol-1K-1 197.6 Jmol-1K-1 188.7 Jmol-1K-1 C3H8(g)+ 5O2(g)  3CO2(g) + 4H2O(g)

34. The Third Law of Thermodynamics The entropy of any perfectly ordered, crystalline substance at absolute zero is zero. • Absolute zero is an ideal temperature that is unobtainable, and a perfect single crystal is an ideal that cannot be achieved, however, the combination of these two ideals constitutes the basis for the third law ofthermodynamics. • • The third law of thermodynamics has two important consequences: • It defines as positive the sign of the entropy of any substance at temperatures above absolute zero • It provides a fixed reference point that allows the measurement of the absolute entropy of any substance at any temperature

35. G and Spontaneity • In the mid 1800’s J. Willard Gibbs determined the relationship of enthalpy, H, and entropy, S, and temperature, T, that best describes the maximum useful energy obtainable in the form of work from a process at constant temperature and pressure. G is the difference between the heat released during a process (via a reversible or an irreversible path) and the heat released for the same process occurring in a reversible manner • The relationship also describes the spontaneity of a system. • The relationship is a new state function, G, the Gibbs Free Energy. H = q whether a process is reversible or irreversible TS = qrev(DSuniv = 0) G = q – qrev DG = DH-TDS at constant T and P

36. Free Energy Change, G, and Spontaneity • The change in the Gibbs Free Energy, G, is a reliable indicator of spontaneity of a physical process or chemical reaction. • G does not tell us how quickly the process occurs. • Chemical kinetics, the subject of Chapter 16, indicates the rate of a reaction. • Sign conventions for G. • G > 0 reaction is nonspontaneous • G = 0 system is at equilibrium • G < 0 reaction is spontaneous

37. The Temperature Dependence of Spontaneity • Free energy has the relationship G = H -TS. • Because 0 ≤ H ≥ 0 and 0 ≤ S ≥ 0, there are four possibilities for G. Forward reaction HSG spontaneity < 0 > 0 < 0 at all T’s. < 0 < 0 T dependent at low T’s. > 0 > 0 T dependent at high T’s. > 0 < 0 > 0 Nonspontaneous at all T’s.

38. DG = 0 Equilibrium DG < 0 Spontaneous DG > 0 Non Spontaneous Spontaneity is favored when DH < 0 and/or DS > 0 G = H -TS . . . . . . . . . . . . . . ) ) ) ) ) ) ) . . . . . . . . . . . . . . ) ) ) ) ) )

39. G and Spontaneity • Changes in free energy obey the same type of relationship we have described for enthalpy and entropy changes. Calculate Go298 for the reaction in -23.56 kJmol-1 0 kJmol-1 -394.4 kJmol-1 -237.2 kJmol-1

40. The Temperature Dependence of Spontaneity • Determine the temperature at which the following system is at equilibrium, spontaneous, non-spontaneous. C3H8(g) + 5 O2(g) )  3 CO2(g) + 4 H2O(g) • We know that So298= -1077.4 kJmol-1K-1, • We know that Ho298= -2219.9 kJmol-1, • and that Go298= -2108.5 kJmol-1.

41. – Study of reaction rates, or the changes in the concentrations of reactants and products with time • – By studying kinetics, insights are gained into how to control reaction conditions to achieve a desired outcome, its mechanism • – Reaction rate = change in concentration of a reactant or product with time. Three “types” of rates • initial rate • average rate • instantaneous rate Chemical Kinetics • – Chemical kinetics of a reaction depend on various factors • Physical states and surface areas of reactants • Reactant concentrations • Temperature • Solvent and catalyst properties

42. The Rate of Reaction • Consider the hypothetical reaction, A(g) B(g) • equimolar amounts of reactant A will be consumed while product B will be formed as indicated in this graph:

43. The Rate of Reaction • Mathematically, the rate of a reaction can be written as: aA(g) + bB(g) cC(g) + dD(g) • 1. Differential rate law • – Expresses the rate of a reaction in terms of changes in the concentration of one or more reactants, [R], over a specific time interval, t • – Describes what is occurring on a molecular level during a reaction 2. Integrated rate law – Describes the rate of a reaction in terms of the initial concentration, [R]0, and the measured concentration of one or more reactants, [R], after a given amount of time, t – Used for determining the reaction order and the value of the rate constant from experimental measurements

44. The Rate of Reaction • In general, for • a A + b B→ x Xwith a catalyst C • Rate = k [A]m[B]n[C]p • The exponents m, n, and p • • are the order of reactant • • The overall order of reaction is the sum of the order of reactants • • can be 0, 1, 2 or fractions • • must be determined by experiment • Rate law must provide a rate with the units M/s • The proportionality constant, k, is called the rate constant. • 1. Value is characteristic of the reaction and reaction conditions • 2. A given reaction has a particular value of the rate constant under a given set of conditions, such as temperature, pressure, and solvent

45. The rate of a simple one-step reaction is directly proportional to the concentration of the reacting substance. [A] is the concentration of A in molarity or moles/L. k is the specific rate constant. k is an important quantity in this chapter. The Rate of Reaction

46. Zeroth-Order Reactions – Reaction whose rate is independent of concentration – Its differential rate law is rate = k – One can write their rate in a form such that the exponent of the reactant in the rate law is 0 rate = – [A] = k[reactant]0 = k(1) = k t – Since rate is independent of reactant concentration, a graph of the concentration of any reactant as a function of time is a straight line with a slope of –k (concentration decreases with time); a graph of the concentration of any product as a function of time is a straight line with a slope of +k

47. First-Order Reactions – Reaction rate is directly proportional to the concentration of one of the reactants – Have the general form A  products – Differential rate for a first-order reaction is rate = – [A] = k[A] t – If the concentration of A is doubled, the rate of the reaction doubles; – If the concentration of A is increased by a factor of 10, the rate increases by a factor of 10 – Units of a first-order rate constant are inverse seconds, s–1 – First-order reactions are very common The order of a reaction can beexpressed in terms of either each reactant in the reaction or the overall reaction. • For example:

48. Second-Order Reactions • Two kinds of second-order reactions 1. The simplest kind of second-order reaction is one whose rate is proportional to the square of the concentration of the reactant and has the form 2A  products. – Differential rate law is rate = – [A] 2t – Doubling the concentration of A quadruples the rate of the reaction – If the [A] is halved the rate of the reaction will decrease by a factor of 4. (½)2 = ¼ – Units of rate constant is M–1s–1 or L/mols – Concentration of the reactant at a given time is described by the following integrated rate law: • 2. The second kind has a rate that is proportional to the product of the concentrations of two reactants and has the form A + B  products. • – Reaction is first order in A and first order in B • – Differential rate law for the reaction is • rate = – [A] = – [B] = k[A] [B] • t t • – Reaction is first order both in A and in B and has an overall reaction order of 2

49. Factors That Affect Reaction Rates • There are several factors that can influence the rate of a reaction: • The nature of the reactants. • The concentration of the reactants. • The temperature of the reaction. • The presence of a catalyst. • We will look at each factor individually.

50. Phase and Surface Area Effects • If reactants are uniformly dispersed in a single homogeneous solution, the number of collisions per unit time depends on concentration and temperature. • If the reaction is heterogeneous, the reactants are in two different phases, and collisions between the reactants canoccur only at interfaces between phases; therefore, the number of collisions between the reactants per unit time is reduced, as is the reaction rate. The rate of a heterogeneous reaction depends on the surface area of the more condensed phase.