1. Thermochemistry “The Quick and Dirty”

2.  Energy changes accompany every chemical and physical change.  In chemistry heat energy is the form of energy that we are most often interested in.  Kinetic energy (energy of motion)  Potential energy (stored energy)  Chemical bond energy is the major form of potential energy we are concerned about in chemistry.  Heat transfer is always from the warmer object to the colder object.

3.  The standard unit of heat energy is joule (J).  Kelvin to 0 C + 273  Open/closed systems: Closed systems can exchange energy but not matter.  Law of conservation of Energy = 1 st Law of Thermodynamics

4. Phase Change Graph For Water

5.  Exothermic reaction – molar enthalpy ( Δ H) is lost by conversion to heat or light. - energy is lost to the surroundings - energy of the system decreases  Endothermic reaction – energy in the surroundings is absorbed and converted to molar enthalpy.  Calorimeter – thermally insulated container in which the exchange between the system and its surroundings can be measured

6.  q = cm Δ T  m= q Δ T = q c Δ T cm  Energy changes during a state change: q = n Δ H phase n- number of moles H f – heat of fusion H v – heat of vapourization

7.  Exothermic Reactions - net release of energy  Energy term is on the product side  Fe 2 O 3 + 2Al  2Al 2 O 3 + 2Fe + 847.6 KJ  Endothermic Reactions – net input of energy  Energy term is on the reactants side.  2 SO 3 + 198 KJ  2 SO 2 + O 2  Chemical Reactions occur spontaneously for two reasons: 1. The products of the reaction have less energy than the reactants (burning a match) always exothermic.

8. 2. Products are more random than the reactants. This is the entropy (S). Changes that produce substances with greater randomness (+ Δ S) are favoured in nature and drive the reaction to occur. More Randomness: 1. Solid state  liquid 2. Liquid state  gas 3. Solid  gas 4. Formation of a mixture 5. Increase in volume of a gas

9.  Gas Highest S Aqueous Liquid Solid Lowest S Enthalpy: Δ H = H final – H initial OR Δ H = H products – H reactants Exothermic: - Δ H Endothermic: + Δ H

10. The reactants have less potential energy than do the products. Energy must be input in order to raise the particles up to the higher energy level. Energy + A + B --> AB

11. The reactants have more potential energy than the products have. The extra energy is released to the surroundings. A + B --> AB + Energy

12.  Writing Equations: 1. Δ H notation for 1 mol of CO exothermic Fe2O3(s) + 3CO(g)  3CO2(g) + 2 Fe (s) + 25 kJ 1/3 Fe 2 O 3(s) + CO (g)  CO 2 (g) + 2/3 Fe (s) Δ H = -8.3 kJ 2. Using Energy as a term: endothermic 3FeCl 3(s)  3FeCl 2(s) + 3/2 Cl 2 Δ H = + 173 kJ 6 FeCl 3(s) + 346 kJ  6 FeCl 2(s) + 3Cl 2(g)

13.  Calculating Heat:  How much heat is produced when 95 g of methane is burned in oxygen? CH 4 (g) + 2 O 2(g)  CO 2(g) + 2H 2 O (g) Δ H = -891kJ 95g x 1mol = 5.9 mol 16g Δ H = 891 kJ/mol 1 mol so 891 kJ q= n Δ H = 5.9 mol x 891 kJ/mol = 5300 kJ

14.  Calorimetry  Δ H substance = mc Δ T n m -mass of water (usually taking in heat) C -specific heat of water 4.18 J/g o C Δ T – change in temperature n – moles of the substance that you are calculating the Δ H of

15.  Hess’s Law  Based on 1 mole.  Reaction occurs in a series of steps in which the intermediates are cancelled. OR Δ H reaction = Σ H f products – Σ H f products Standard Heat of Formations from elements( you will need to write a balanced equation for the formation of the substance). Bond Energies – uses Lewis Structures to draw structural formulas (energy values come from table) Δ H = Σ Reactants – Σ products Δ G: - Δ G reaction is spontaneous