1. THERMOCHEMISTRY Courtesy of lab-initio.com

2. Definitions #1 Energy: The capacity to do work or produce heat. Potential Energy: Energy due to position or composition. Kinetic Energy: Energy due to the motion of the object and depends on its mass and velocity.

3. Definitions #2 Law of Conservation of Energy: Energy can neither be created nor destroyed, but can be converted between forms. The First Law of Thermodynamics: The total energy content of the universe is constant.

4. State Functions State Functions refers to a property of the system that depends only on its present state. ENERGY IS A STATE FUNCTION A person standing at the top of Mt. Everest has the same potential energy whether they got there by hiking up, or by falling down from a plane WORK and HEAT ARE NOT A STATE FUNCTIONS WHY NOT???

5. Work: defined as force acting over a distance. Heat: involves the transfer of energy between two objects due to a temperature difference. Remember heat and temperature are different. Temperature is a measure of the hotness or coldness of something and is proportional to the average molecular kinetic energy of the atoms, molecules, or ions present.

6. System: The portion of the universe under study. Surroundings: Everything else besides the system. Interaction: Exchange of energy and or matter between the system and its surroundings. Systems: Open System: Exchanges both matter and energy with its surroundings. Closed System: Exchanges only energy with its surroundings. Isolated System: Exchanges neither energy nor matter with its surroundings.

7. Internal Energy (E) of a system = sum of the PE + KE of all “particles” of the system. The joule (J) is the SI unit for energy. J = kgm 2 /s 2 Can be changed by a flow of work, heat, or both. ΔE = q + w ΔE = change in system’s internal energy q = heat w = work

8. Thermodynamic quantities = two parts: a number, giving the magnitude of the change, and a sign, indicating the direction of flow. q = heat flowing into or out of the system -q if energy is leaving to the surroundings +q if energy is entering from the surroundings

9. w = work done by, or on, the system -w if work is done by the system on the surroundings +w if work is done on the system by the surroundings

10. A common type of work associated with chemical processes is work done by a gas (through expansion) or work done to a gas (through compression). For example, in an automobile engine, the heat from the combustion of the gasoline expands the gases in the cylinder to push back the piston, and this motion is then translated into the motion of the car. The gases are expanding so the volume is increasing therefore, the change in volume (ΔV) is positive. Work and volume must have opposite signs.

11. Work, Pressure, and Volume Expansion Compression +  V (increase)-  V (decrease) -w results + w results E system decreases Work has been done by the system on the surroundings E system increases Work has been done on the system by the surroundings

12. Exothermic: energy flows out of the system. Which has lower potential energy? Reactants or products?

13. CH 4 (g) + 2O 2 (g) → CO 2 (g) + 2H 2 O(g) + energy (heat) Total energy is conserved = energy gained by the surroundings must be equal to the energy lost by the system. In the case above, the heat flow into the surroundings results from a lowering of the PE of the reaction system.

14. In an exothermic process, the bonds in the products are stronger (on average) than those of the reactants. More energy is released by forming the new bonds than is consumed in breaking the bonds in the reactants. In any exothermic reaction, some of the PE stored in the chemical bonds is being converted to thermal energy via heat.

15. Endothermic process = reverse situation Products have higher PE (weaker bonds on average) than the reactants. Energy flows into system as heat to increase the PE of the system.