1. Unit 1: Energy Changes in Chemical Reactions Not All Reactions Go Off with a Bang!

2. Energy The capacity to do work or to produce heat The capacity to do work or to produce heat Forms of Energy: 1. Kinetic Energy– The energy of motion

3. 2. Potential Energy – The energy of position (gravitational potential energy) eg. Diefenbaker Dam – The energy stored in chemicals because of their composition eg. Chocolate bars (Yum!)

4. Measuring Energy ☼ Common Units → calorie (cal) - food ≈ 2000 – 3000 kcal/day 1000 cal = 1 kcal (kilocal) 1000 cal = 1 Calorie → British Thermal Unit (BTU) - fuels fuel = 20 900 BTU/pound coal = 9 600 BTU/pound plastic bags = 18 700 BTU/pound

5. ☼ Scientific Unit → joule (J) → joule (J) 1 J = 1 newton ● metre (N ● m) i.e. The work required to move an object one metre using a force of one newton (N = kg ● m/s 2 ). 4.184 J = 1 cal

6. ☼ Most chemical reactions involve energy or changes in energy. ☼ In a Chemical Reaction (rxn)… - existing bonds are broken (requires E) - atoms are rearranged - new bonds are formed (releases E) Thus, almost all chem. Rxns either absorb or release energy. This results in an exchange of energy (aka HEAT)

7. Heat-- The energy that is transferred from one object to another due to a difference in temperature (flow of energy is usually from hot to cold until equilibrium is reached) Thermochemistry-- The study of the changes in heat in chem. rxns (part of thermodynamics-- energy/work) → thermes = greek for heat → thermes = greek for heat → we look at systems during changes of heat (i.e. A beaker or flask) → we also look at surroundings into which heat may be lost or gained → we also look at surroundings into which heat may be lost or gained

8. → A system is part of the universe on which we focus our attention, the surroundings include everything else in the universe → A system is part of the universe on which we focus our attention, the surroundings include everything else in the universe → energy is neither created or destroyed when heat is transferred (1 st Law of Thermodynamics) → energy is neither created or destroyed when heat is transferred (1 st Law of Thermodynamics) → closed vs open system → closed vs open system → we can’t tell how much energy is in something until it is released → we can’t tell how much energy is in something until it is released

9. Temperature = degree of hotness or coldness of an object, which is a measure of average kinetic energy of the molecules Heat = the energy transferred from one body to another because of the temp. difference Heat is a form of energy; temp. in NOT! There is more heat in a large iceberg than in a cup of boiling water! This is b/c heat is trapped inside as opposed to being released so the ice berg doesn’t feel warm.

10. Exothermic Reactions → rxns that release heat into the surroundings → PE is converted to heat energy; temp. ↑ Endothermic Reactions → rxns that absorbs heat from the surroundings → KE decreases; temp. ↓

11. Examples: 1. Combustion of propane C 3 H 8(g) + 5O 2 → 3CO 2(g) + 4H 2 O (g) + 2043 kJ 1 mole of propane produces 2043 kJ of heat

12. Energy released as new bonds are formed in the products is greater than the energy required to break the old bonds in the reactants. All combustion rxns are exothermic!

13. 2. Water Gas Reaction C (s) + H 2 O (g) + 113 kJ → CO (g) + H 2(g) C (s) + H 2 O (g) + 113 kJ → CO (g) + H 2(g) 1 mole of solid carbon requires 113 kJ of heat to produce 1 mole of carbon monoxide. This is an endothermic rxn. Chemical equations with heat values incorporated into them are known as thermochemical equations.

14. Enthalpy (H) is… … the heat content or the amount of heat a substance has at a given temp. and pressure (the total energy stored by a substance is the sum of KE & PE. The enthalpy is this energy plus a small added term that takes into account the pressure and volume.) ∆H = Change in enthalpy ∆H = H products – H reactants ∆H = H products – H reactants ∆H = + (pos.) = heat absorbed = endo ∆H = + (pos.) = heat absorbed = endo ∆H =  (neg.) = heat released = exo ∆H =  (neg.) = heat released = exo

15. Energy Diagrams

16. For reporting enthalpy changes, chemists use 1 atm = 101.3 kPa and 25°C (298K) as standard enthalpy changes ∆H°. 1 atm = 101.3 kPa and 25°C (298K) as standard enthalpy changes ∆H°. CH 4(g) + 2O 2(g) → CO 2(g) + 2H 2 O (l) ∆H° =  890.3 kJ ½H 2(g) + ½I 2(g) → HI (s) ∆H° = + 26.5kJ Equations are stoichiometrically correct! SATP

17. Your Turn! 1. Draw the Energy diagram for… a) the combustion of propane b) the formation of CO (g) 2. Write a balanced equation and state the ∆H° value (i.e. don’t include the value in the eqn but state it separately as pos. or neg.) ∆H° value (i.e. don’t include the value in the eqn but state it separately as pos. or neg.)

18. Summary Exothermic (  ∆H°) → energy is released → energy appears as a product → surroundings increase in temp. (warmer) Endothermic (+∆H°) → energy is absorbed → energy appears as a reactant → surroundings decrease in temp. (cooler) ☼ Equations are stoichiometrically correct!

19. Calculation Questions: 1. How much heat is transferred when 9.22 g of glucose (C 6 H 12 O 6 ) in your body reacts with O 2 according to the following eqn.? C 6 H 12 O 6(s) + 6O 2(g) → 6CO 2(g) + 6H 2 O (l) ∆H° =  2803 kJ 9.22g C 6 H 12 O 6 x 1 mol x 2803 kJ 180.0 g 1 mol C 6 H 12 O 6 180.0 g 1 mol C 6 H 12 O 6 = 143.57589 kJ = 144 kJ released

20. 2. How much heat is transferred when 147g of NO 2(g) is dissolved in an excess of H 2 O? 3NO 2(g) + H 2 O (l) → 2HNO 3(aq) + NO (g) ∆H° =  138 kJ ∆H° =  138 kJ 147g NO 2 x 1 mol NO 2 x 138 kJ 46.0 g 3 mol NO 2 46.0 g 3 mol NO 2 = 147 kJ released Assignment : Pg. 187, #10 & 11 Pg. 191 – 195, #30 – 33 (Addison-Wesley Text)