1. Chapter 6 Thermochemistry

2. 6.1: I. Nature of Energy A. Energy (E): capacity for work or producing heat B. Law of Conservation of Energy: can’t be created or destroyed only changed to different forms C. Potential Energy (PE): due to position or composition, Ex. Ball held in the air D. Kinetic Energy (KE): due to motion E. Activation Energy (E a ): The energy that must be overcome for a chemical rxn to occur

3. II. Energy Transfer A. Frictional heating: loss of energy as heat by friction against surface B. Temperature: property of random motion of particles in a substance C. Heat (q): transfer of energy due to temp. difference between objects D. Work (w): force acting over a distance Ex. Lifting a dumbbell *** Energy can be transferred as heat or work***

4. III. Chemical Energy A. For energy transfer we need to identify parts B. “System”: part we focus on C. “Surroundings”: everything around system D. This reaction between Na and H 2 O would be your system E. The rest of the universe is the surroundings

5. IV. Endothermic C. The reaction has absorbed energy as reactants become products A. Endothermic: the system absorbs energy B. (“Endo” = in, “thermic” = heat)‏

6. V. Exothermic A. Exothermic: when system loses energy (“Exo”= Exit, “Thermic” = heat)‏ B. The reaction has released energy as reactants become products

7. VI. 1 st Law of Thermodynamics A. Energy of the “universe” is constant B. Internal energy of a system (E) is the sum of the kinetic and potential energy C. ***∆ Before anything means how much that value is changing*** D. ∆E = q + w (q= heat, w= work)‏ E. ∆E = - for an exothermic process, energy lost F. ∆E = + for endothermic, energy gained

8. VII. Gas Energy A. Related to compression/expansion of gas B. Gas Work (W) = Force (F) x distance (∆h) C. F = Pressure (P) x Area (A) D. W = P x A x ∆h E. A x ∆h = ∆V (∆V + expanding, - compressed) F. W is released (-) for expanding gas, work added (+) to compress a gas G. Since W and ∆V have opposite signs, must be a “–” sign in front of pressure H. W = - P ∆V

9. 6.2: I. Enthalpy and Calorimetry A. Enthalpy (H): E + PV B. Enthalpy change is equal to heat (q) at constant pressure C. ∆H = - for exothermic reaction D. ∆H = + for endothermic reaction

10. II. Calorimetry A. Measure of heat change in a system B. Heat capacity (C): heat absorbed/increase in temperature C. Specific heat capacity (C): heat capacity per gram of substance (J/gºC) or (J/gK) D. Molar heat capacity (C): (J/mole ºC) or (J/mole K)

11. III. Constant Pressure Calorimetry A. Allows us to measure heat change or enthalpy change B. ∆Q or ∆H = C x m x ∆T Constant Volume Calorimetry C. Bomb calorimeter: container where substance is ignited to measure energy of combustion D. ∆E = q + w, w = P ∆V, ∆V = 0 E. ∆E = q

12. 6.3: I. Hess’s Law A. States that the heat (or enthalpy) for reactions yielding a net reaction can be added to yield the heat or enthalpy of the net reaction B. Ex.  H for 1 mole H 2 O(s) to H 2 O(g)‏ H 2 O(s)  H 2 O(l)  H = 6.01KJ/mole H 2 O(l)  H 2 O(g)  H = 40.7 KJ/mole

13. Add Reactions when Done Multiply heat by coefficient of reaction Reverse Rxn = Flip Sign II. Characteristics of Enthalpy Change A. If a rxn reversed, sign of q or ∆H switches (because direction of heat flow reverses)‏ B. Magnitude of q or ∆H is directly proportional to the quantities of reactants and products in a reaction (if double rxn, then double heat)‏

14. 6.4: I. Standard Enthalpy of Formation A. ΔH for formation of 1 mole of a compound from its elements in their standard states B. Std. conditions are 1 atm, 1 M, and the states found at 25 ºC C. ∆H = Sum ∆H f (products) - Sum ∆H f (reactants)‏