1. Unit 12

2.  Tuesday, May 8 – EOC Field Testing for Chemistry  Wednesday, May 9 – Unit 12 Test Review  Thursday, May 10 – Unit 12 Test

3.  The Flow of Energy  Energy – the capacity to do work or supply heat  Chemical Potential Energy – energy stored within the bonds of chemical compounds  Activation Energy – the minimum energy colliding particles must have in order to react

4.  Thermodynamics – the study of energy in chemical reactions; literally means “changes in heat.”  All chemical reactions either release or absorb energy when they occur. Another way to say this is that reactions either “give off” or “take in” energy.  The Law of Conservation of Energy – 1 st Law of Thermodynamics; states energy can neither be created nor destroyed, only transformed  ALL energy is either work performed, stored potential, or heat lost.

5.  System – the chemical reaction under study  Surroundings – every place in the universe except the system  Universe – the system and the surroundings

7.  Exothermic Reaction – a reaction in which heat is released by the system to the surroundings; from the perspective of the system, “heat is given off”

8.  Endothermic Reaction - is a reaction in which heat is absorbed by the system from the surroundings. From the perspective of the system, “heat is taken in.”

9. Energy level diagram for an exothermic chemical reaction without showing the activation energy; it could also be seen as quite exothermic with a highly unlikely zero activation energy, but reactions between two ions of opposite charge usually has a very low activation energy. Very endothermic reaction with a large activation energy. Moderately exothermic reaction with a moderately high activation energy. A small activation energy reaction with no net energy change; this is theoretically possible if the total energy absorbed by the reactants in bond breaking equals the energy released by bonds forming in the products. Very exothermic reaction with a small activation energy. Energy level diagram for an endothermic chemical reaction without showing the activation energy; it could also be seen as quite endothermic with zero activation energy. Moderately endothermic reaction with a moderately high activation energy.

10. Potential Energy – Read from the x-axis to the reaction line EX: b is the potential energy of ….. (answer to question 2) A+B and C+D are NOT answers on the questions. They represent the reactants (A+B) and the products (C+D). It might help to draw dashed lines so that you can visualize the potential energy all the way across the graph Reactants Products

11. 1. Is the above reaction endothermic or exothermic? 2. What letter represents the potential energy of the reactants? 3. What letter represents the potential energy of the products? 4. What letter represents the change in energy for the reaction? b f d

12. 5. What letter represents the activation energy of the forward reaction? 6. What letter represents the activation energy of the reverse reaction (read the chart backwards)? 7. What letter represents the potential energy of the activated complex? 8. Is the reverse reaction endo or exothermic? 9. If a catalyst were added, what letter(s) would change? a e c a and c

13. Day 2

14.  Measuring and Expressing Heat Changes  Calorimetry – the accurate and precise measurement of heat change for chemical and physical processes  Calorimeter – the insulated device used to measure the absorption or release of heat in chemical or physical processes  Enthalpy (H) – heat energy content of a system at constant pressure  Enthalpy is the heat absorbed or released by a system when pressure is constant.  It is impossible to record enthalpy directly, but change in enthalpy (ΔH ) can be measured.  Units of heat energy: calorie (cal), joules (J), or kilojoules (kJ).

15.  For an exothermic reaction, the sign of ΔH is negative.  When a reaction is exothermic (ΔH is negative), that is a favorable condition. Enthalpy is just one of the variables involved when predicting whether or not a reaction will occur, but, in general, reactions which release heat are more likely to occur than ones in which heat is required.

16.  For an endothermic reaction, the sign of ΔH is positive.  When a reaction is endothermic (ΔH is positive), that is an unfavorable condition. Enthalpy is just one of the variables involved when predicting whether or not a reaction will occur, but reactions which absorb heat are less likely to occur than ones in which heat is released, all things being equal.

17.  Thermochemical Equation – an equation that includes the heat change  Ex. CaO(s) + H2O(l) → Ca(OH)2(s) + 65.2kJ  Heat of Reaction – the heat of change for the equation exactly as it is written

18. 1. 2NO (g) + O 2(g)  2NO 2(g) + 113.04 kJ endothermicorexothermic 2. 2H 2(g) + O 2(g)  2H 2 O (l) ; ΔH = -571.6 kJ endothermicorexothermic 3. 4NO (g) + 6H 2 O (l)  4NH 3(g) + 5O 2(g) ; ΔH = +1170 kJ endothermic or exothermic 4. SO 2 (g) +296 kJ  S (s) + O 2 (g) endothermic orexothermic

20. Day 3

21.  Calculating Heat Changes  Standard Heat of Formation (ΔH f 0 ) – the change in enthalpy that accompanies the formation of one mole of a compound from its elements  Heat of Reaction (ΔH 0 )– the heat released or absorbed during a chemical reaction, or Enthalpy  ΔH 0 = ΔH f 0 (products) - ΔH f 0 (reactants)  READ AS: Heat of Rxn EQUALS the SUM of Heat of Formation of the Products minus the SUM of Heat of Formation of the Reactants  Standard Heats of Formation have been determined for many common pure substances, both elements and compounds. Elements in their natural state are understood to have a ΔH f 0 = 0

22.  In order to calculate ΔH 0, the standard heats of formation of the reactants and products must be known; they can be found on the following table:

23.  Steps to Calculate Heat of Reaction: 1. Find the balanced chemical equation for the reaction; must have states of matter 2. Find the sum of the Heats of Formation for the reactants a. Multiply the ΔH f 0 for each reactant by its corresponding number of moles (coefficient) from the balance equation b. Sum the reactants 3. Find the sum of the Heats of Formation for the products a. Multiply the ΔH f 0 for each product by its corresponding number of moles (coefficient) from the balance equation b. Sum the products 4. Subtract the ΔH f 0 (reactants) from the ΔH f 0 (products)

24. 1. Find the balanced chemical equation for the reaction; must have states of matter 2. Find the sum of the Heats of Formation for the reactants a. Multiply the ΔH f 0 for each reactant by its corresponding number of moles (coefficient) from the balance equation b. Sum the reactants 3. Find the sum of the Heats of Formation for the products a. Multiply the ΔH f 0 for each product by its corresponding number of moles (coefficient) from the balance equation b. Sum the products REMEMBER! Elements in their natural state are understood to have a ΔH f 0 = 0 OTHERWISE, use the table provided to lookup individual ΔH f 0

25. 4. Subtract the ΔH f 0 (reactants) from the ΔH f 0 (products) ΔH 0 = ΔH f 0 (products) - ΔH f 0 (reactants) (show work here)