1. Thermochemistry

2. The study of heat changes in chemical reactions Exothermic: reactions that release heat Endothermic: reactions that absorb heat Enthalpy: the heat content of a system at constant pressure

3. Units of Heat 1 Calorie = 1 kilocalorie = 1000 calories 1 J = 0.239 cal 4.186 J = 1 cal Food Calorie

4. Heat Capacity The heat capacity of an object depends on: – mass – chemical composition. The amount of heat needed to increase the temperature of an object exactly 1°C is the heat capacity of that object.

5. Specific Heat q = mcΔT Where: q = heat (energy) m = mass c = specific heat ΔT= temperature change

6. Specific Heat of Substances

7. Why is the sand hot and the water cool at the beach?

8. Jet Stream

9. Enthalpy The heat content of a system at constant pressure is the same as a property called the enthalpy (H) of the system.

10. Calorimetry The measurement of heat changes for physical and chemical processes

11. Calorimetry

12. Two Key Equations 1. q reaction = -q surroundings q r = -q s 2. q = mcΔT = ΔH Where:m = mass c = specific heat ΔT= temperature change

13. Calorimetry Example 1.50 mL of water is placed in a beaker. A piece of calcium is weighed and placed in the cup. The temperature change of the water is measured. Calculate the heat released to the water by the calcium. Note: c = 4.184 J/g°C 2. Calculate the heat released per mole of calcium based on the previous results

14. Thermochemistry Exothermic reaction: Heat is a product C 3 H 8 + 5O 2 → 3CO 2 + 4H 2 O + 2043 kJ ΔH = -2043kJ H Reaction Progress (t)

15. Thermochemistry Endothermic reaction: Heat is a reactant C + H 2 O + 113 kJ → CO + H 2 ΔH = +113kJ H Reaction Progress (t)

16. Enthalpy (ΔH) Heat content of a system at constant pressure ΔH = H products - H reactants ΔH ◦ is the standard enthalpy reported at 25 degrees C and 1 atm. Sign of ΔH ProcessHeat +endothermicabsorbed -exothermicreleased

17. Standard Heats of Formation For a reaction that occurs at standard conditions, you can calculate the heat of reaction by using standard heats of formation.

19. Standard Heats of Formation

20. Entropy and Reaction Tendency There is a tendency in nature to proceed in a direction that increases the randomness of a system. A random system is one that lacks a regular arrangement of its parts. This tendency toward randomness is called entropy. Entropy, S, can be defined in a simple qualitative way as a measure of the degree of randomness of the particles, such as molecules, in a system.

21. Standard Entropy Changes for Some Reactions

22. Entropy In a solid, the particles are in fixed positions, and we can easily determine the locations of the particles. In a liquid, the particles are very close together, but they can move around. Locating an individual particle is more difficult. The system is more random, and the entropy is higher. In a gas, the particles are moving rapidly and are far apart. Locating an individual particle is much more difficult, and the system is much more random. The entropy is even higher.

23. Entropy Absolute entropy, or standard molar entropy, of substances are recorded in tables and reported in units of kJ/(molK). Entropy change, which can also be measured, is defined as the difference between the entropy of the products and the reactants. An increase in entropy is represented by a positive value for ∆ S, and a decrease in entropy is represented by a negative value for ∆S.

24. Free Energy Processes in nature are driven in two directions: toward least enthalpy and toward largest entropy. As a way to predict which factor will dominate for a given system, use this equation: ∆G 0 = ∆H 0 – T ∆ S 0 This combined enthalpy-entropy function is called the free energy, G, of the system; it is also called Gibbs free energy.

25. Relating Enthalpy and Entropy to Spontaneity ∆G 0 = ∆H 0 – T ∆ S 0 A spontaneous reaction is one that goes to completion unaided Example: rusting, ice melting