1. Thermodynamics

2. study of energy changes that accompany physical and chemical processes. Thermochemistry is one component of thermodynamics which focuses on how energy changes are measured and predicted.

3. Energy ability to do work or transfer heat. 1. kinetic energy-energy of motion (example: electrical) KE = 1/2mv 2 2. potential energy-stored energy due to composition (example: chemical)

4. Types of Changes Endothermic- a reaction in which energy is absorbed from the surroundings.

5. Exothermic-a reaction in which energy is released.

6. First Law of Thermodynamics the total amount of energy in the universe is constant. Also known as the Law of Conservation of Energy which states that energy cannot be created or destroyed, but can only change form.

7. Thermodynamic Terms System-the substances being studied. Surroundings- everything in the system’s environment Universe-the system plus the surroundings. State Function-property whose value depends only on the state of the system-not on the pathway it took to get there.

8. Examples of State Functions Pressure, volume, and temperature are all examples of state functions. Example: T i = 30 o C and T f = 22 o C ∆ T = -8 o C How the temperature change occurred does not matter ∆ X = X f – X i If there is an increase in X, ∆X > 0. If there is a decrease in X, ∆X < 0.

9. Enthalpy Enthalpy change (∆H) is the quantity of heat (q) transferred in or out of a system. ∆H = H final – H initial ∆H = H products - H reactants

10. Exothermic Reactions ∆H = energy of products – energy of reactants If the reaction is exothermic, ∆H < 0.

11. Endothermic Reactions ∆H = energy of products – energy of reactants If the reaction is endothermic, ∆H > 0.

12. Measuring Energy Changes Calorimetry-process of measuring energy changes Calorimeter-device used to measure heat Heat released by reaction = heat gained by calorimeter + heat gained by the solution q = m(∆T)C

13. Internal Energy Internal Energy, E, is all of the energy contained within a substance. 1) kinetic energy of the molecules 2) energy of attraction and repulsion between subatomic particles, molecules, ions, etc. 3) other forms of energy Internal energy is a state function

14. ∆E = E final – E initial = E products – E reactants = q + w q represents heat and w represents work ∆E = heat absorbed by system + work done on the system

15. The following conventions apply to the signs of q and w q is positive: heat is absorbed by the system from the surroundings (endothermic) q is negative: heat is released by the system to the surroundings (exothermic) w is positive: work is done on the system by the surroundings w is negative: work is done by the system on the surroundings.

16. Writing Equations When ∆E < 0, energy is released by the system and can be written as a product. When ∆E > 0, energy is absorbed by the system and can be written as a reactant.

17. Effect of Pressure on Work Work done on a system = -P∆V If volume decreases (could be due to a decrease in the number of moles of gas), work is done on the system so the sign of w is positive. If volume increases (could be due to an increase in the number of moles of gas), work is done by the system so the sign of w is negative. In constant volume reactions, no work is done so E = q

18. Relationship between ∆H and ∆E. ∆H = ∆E + P∆V (constant T and P) *use with physical changes ∆H = q p (constant T and P) ∆H= ∆E + (∆n)RT or ∆E = ∆H – (∆n)RT (constant T and P) *use with chemical changes

19. Hess’s Law Since enthalpy is a state function, the change in enthalpy in going from some initial state to some final state is independent of the pathway. The change in enthalpy in going from a particular set of reactants to a particular set of products is the same whether the reaction takes place in one step or a series of steps.

20. Characteristics of Enthalpy Changes If the reaction is reversed, the sign of ∆H is also reversed. The magnitude of ∆H is directly proportional to the quantities of reactants and products. If the coefficients are multiplied by an integer, the value of ∆H is multiplied by the same number.

21. Standard Enthalpies of Formation For a reaction under standard conditions of constant pressure, enthalpy changes can be measured using a calorimeter. Because some reactions proceed too slowly, a process is needed that allows the enthalpy change to be calculated.

22. Standard Enthalpies of Formation The standard enthalpy of formation (∆H f o ) is defined as the change in enthalpy that accompanies the formation of one mole of a compound from its elements with all substances in their standard states. The standard state is a precisely defined reference state. See page 246 for definitions of standard states.

23. Bond Energy and Enthalpy Bond energy values can be used to calculate approximate energies for reactions also. ∆H = sum of energies required to break old bonds (positive sign/endothermic) plus the sum of energies released in formation of new bonds (negative sign/exothermic)

24. Spontaneous Processes A process is spontaneous if it occurs without outside intervention. Spontaneous processes may be either fast or slow. Thermodynamics lets us predict whether a process will occur but gives no information about the amount of time required for the process.

25. Entropy Entropy (S) is a measure of the molecular randomness or disorder. The driving force of spontaneous processes is an increase in entropy of the universe. The natural progression of things is from order to disorder, from lower entropy to higher entropy.

26. Entropy (continued) Entropy describes the number of arrangements (positions/energy levels) that are available to a system existing in a given state. Nature spontaneously proceeds toward the states that have the highest probabilities of existing. The states with the highest probabilities of existing is that which has the greatest disorder. (S gas > S liquid > S solid ; in the gaseous state, molecules have many more positions available to them and are therefore more disordered).

27. Predicting Entropy Changes ∆S = S final - S initial If the entropy increases, ∆S is > 0 If the entropy decreases, ∆S is < 0

28. Which of the following has the greatest entropy? 10 1.1 mole of solid CO 2 2.1 mole of gaseous CO 2 3.Both are equal 1234567891011121314151617181920 21222324252627282930

29. Which has the greatest entropy? 1.1 mole of N 2 gas at 1 atm 2.1 mole of N 2 gas at 0.01 atm 3.Both have the same entropy 10 123456789 11121314151617181920 21222324252627282930

30. What is the sign for the entropy change when solid sugar is added to water to form a solution? 1.Positive 2.Negative 10 123456789 11121314151617181920 21222324252627282930

31. What is the sign for the entropy change when iodine vapor condenses on a cold surface to form crystals? 1.Positive 2.Negative 10 123456789 11121314151617181920 21222324252627282930

32. Second Law of Thermodynamics The second law of thermodynamics states that in any spontaneous process, there is always an increase in the entropy of the universe. In other words, the entropy of the universe is increasing and is not conserved. ∆S univ = ∆S sys + ∆S surr If ∆S univ > 0, the process is spontaneous as written. If ∆S univ < 0, the process is spontaneous in the opposite direction.

33. The Effect of Temperature on Spontaneity An exothermic process in the system causes heat to flow to the surroundings, increasing the random motions and entropy of the surroundings. ∆S surr >0 The opposite is true for endothermic processes. As a result, nature tends to seek the lowest possible energy.

34. Effect of temp (continued) The magnitude of ∆S surr depends on the temperature. ∆S surr depends directly on the quantity of heat transferred and inversely on the temperature. ∆S surr = -∆H T

35. Predict the sign of ∆S surr for the following process: H 2 O(l)  H 2 O (g) 1.Positive 2.Negative 10 123456789 11121314151617181920 21222324252627282930

36. Predict the sign of ∆S surr for the following process: CO 2 (g)  CO 2 (s) 1.Positive 2.Negative 10 123456789 11121314151617181920 21222324252627282930

37. Third Law of Thermodynamics The Third Law of Thermodynamics states that the entropy of a perfect crystal at 0 K is zero. The standard entropy values S o of many common substances at 298 K and 1 atm are listed in Appendix 4. Because entropy is a state function, ∆S o rxn = Σn p S o products – Σn r S o reactants Generally, the more complex the molecule, the higher the standard entropy value.

38. Free Energy Free energy (G) is the energy that is available to do work. ∆G = ∆ H – T ∆ S where H is enthalpy, T is Kelvin temp, and S is entropy. A process is spontaneous (at constant T and P) in the direction in which the free energy decreases (- ∆G means + ∆S univ ) See the table on page 761 At the melting point and boiling point, ∆G = 0.

39. Free Energy and Chemical Reactions For chemical reactions, we are often interested in the standard free energy change (∆G o ), the change in free energy that will occur if the reactants in their standard states are converted to the products in their standard states.

40. Calculating Free Energy ∆G o cannot be measured directly, but can be calculated from other quantities. (∆G o = ∆H o - T ∆S o ) Free energy is a state function and can be determined using similar procedures as those for finding ∆H using Hess’s law. Free energy can also be calculated using standard free energies of formation (the free energy that accompanies the formation of 1 mole of that substance from its constituent elements with all the reactants and products in their standard states. (∆G o rxn = Σn p G f o products – Σn r G f o reactants )