THERMOCHEMISTRY

ENERGY 1. Categories: Potential Stored energy, energy of position Gravitational Chemical Nuclear Kinetic Energy of motion Mechanical Electrical Heat Sound Radiant: Electromagnetic radiation

ENERGY Law of Conservation of Energy In any chemical or physical process, energy is neither created nor destroyed. BUT energy CAN be transformed from one form to another. Chemical → Light + Heat

HEAT vs TEMPERATURE Heat vs. Temperature – NOT the same a. Temperature 1. A measure of the average kinetic energy of the particles (atoms/molecules) 2. Determines the DIRECTION of heat flow b. Heat - flows from the hotter object to the colder object.

Heat Vs. Temperature Heat Temperature Heat can travel Temp Can NOT Measured in Joules Measured in Degrees Measured with calorimeter Measured with Thermometer Heat is E it can do work Temp is a man made scale

HEAT 1. Factors affecting heat: Mass Specific heat Temperature

HEAT the amount of energy needed to raise the temperature of one 2. Measured by means of difference Q = m * c * Δt m = mass c = specific heat Δt = temperature change UNITS of energy: calorie (cal) the amount of energy needed to raise the temperature of one gram of water 1ºC. kilocalorie (kcal) = 1000 cal or 1 Calorie Joule (J) 1 cal = 4.186 J 1 kcal = 4.186 kJ Kilowatt-hour 1 kwh = 860 kcal

But wait…. 1. food Calorie = 1,000 calories or 1 kilocalorie Food Calories listed are actually kilocalories! Note: That Big Mac (576 Calories) can actually raise _____________ grams of water 1°C. 576,000

HEAT 3. Specific heat (capacity) Energy needed to raise 1 gram of a substance 1oC Unit: cal or _J_ goC goC

HEAT FLOW ENDOTHERMIC Energy is absorbed ∆ H is + a. Heat flows into the system b. Ex. CHEMICAL CHANGES 1. PHOTOSYNTHESIS 6 CO2(g) + 6 H2O(g) → C6H12O6(s) + 6 O2(g) 2. ELECTROLYSIS 2 H2O(l) → 2 H2(g) + O2(g) c. Ex. PHYSICAL CHANGES 1. BOILING: H2O(l) → H2O(g)

HEAT FLOW 2. MELTING: H2O(s) → H2O(l)

HEAT FLOW EXOTHERMIC Energy is released ∆ H is – a. Heat flows out of the system b. Ex. CHEMICAL CHANGES 1. BURNING C(s) + O2(g) → CO2(g) 2. FORMATION OF A COMPOUND FROM ITS ELEMENTS 2 H2(g) + O2(g) → 2 H2O(l)

HEAT FLOW PHYSICAL CHANGES 1. FREEZING 2. CONDENSATION H2O(l) → H2O(s) H2O(g) → H2O(l) IF A GIVEN PROCESS IS EXOTHERMIC, THE REVERSE PROCESS IS ENDOTHERMIC. Most Synthesis reactions are EXOTHERMIC. Most Decomposition reactions are ENDOTHERMIC.

H (ENTHALPY) H (Enthalpy) – HEAT CONTENT of a system at constant pressure ∆ H (change in heat content) = H products – H reactants

Exothermic Heat Diagram

Endothermic Heat Diagram

HEATS OF REACTION Heat of formation: ∆Hf The change in heat content when 1 mole of a compound is formed from its elements a. ∆H is usually _____ Exothermic b. Ex. The heat of formation of CaO is -151.9 kcal/mol. Write the thermochemical equation. 1. Write the balanced equation. 2. Convert to show 1 mole product. 3. Show the physical states and give the value for ∆H.

HEATS OF REACTION c. Compounds with very negative heats of formation – VERY STABLE ex. -235.8 kcal/mol d. Compounds with positive or low negative heats of formation – GENERALLY UNSTABLE ex. -4.8 kcal/mol or + 6.3 kcal/mol

HEATS OF REACTION HEAT OF COMBUSTION c. Compounds with high positive heats of formation VERY UNSTABLE ex. HgC2N2O2 ∆H = + 64 kcal (mercury fulminate) HEAT OF COMBUSTION Change in heat content when 1 mole of a substance burns a. ∆H is ________ EXOTHERMIC b. General format for a hydrocarbon combustion: CxHy + O2 → CO2 + H2O

HEATS OF REACTION c. Construct a heat content diagram for the combustion of CO(g). See table 4.2 for ∆H.

USES OF ∆H 1. If ∆H for the forward reaction is positive, then ∆ H for the reverse reaction is negative. 2. ∆H is directly proportional to the amount of reactants or products in a process. Ex. C(s) + O2(g) → CO2(g) ∆H = -94.1 kcal a. Calculate the heat when 2.80 mol CO2 forms.

USES OF ∆H Ex. C(s) + O2(g) → CO2(g) ∆H = -94.1 kcal b. Calculate the heat when 1.00 g of Carbon burns.

USES OF ∆H Ex. C(s) + O2(g) → CO2(g) ∆H = -94.1 kcal c. What mass of Carbon is needed to generate 10.0 kilocalories of heat?

CHANGES OF STATE HEAT OF FUSION ∆Hfus Change in heat content when 1 mole of a substance changes from a solid to a liquid (melts) a. The heat of fusion of water is 80.0 cal/g. Express as kcal/mol. b. Write a thermochemical equation for the melting of ice.

CHANGES OF STATE HEAT OF VAPORIZATION: ∆Hvap Change in heat content when 1 mole of a substance changes from a liquid to a gas a. The heat of vaporization is 540. cal/g. Express in kcal/mol.

CALORIMETRY CALORIMETRY – MEASURING ∆H When using a calorimeter, Q = ∆H, and the equation becomes: ∆H = m c ∆T m = mass in g or kg c = specific heat (cal/goC or kcal/kgoC) ∆T = temperature change cH2O = 1 cal/goC or 1 kcal/kgoC) KEY: PAY ATTENTION TO THE UNITS!

CALORIMETRY 1. A reaction takes place in a calorimeter during which 40.0 g of water is heated from 24.0oC to 50.0oC. Find the heat of reaction (∆H).

CALORIMETRY 2. The temperature of a piece of copper with a mass of 95.4 g increases from 25.0oC to 48.0oC when the metal absorbs 205.3 calories of heat. What is the specific heat of the copper?

HESS’ LAW If Equation (C) is the sum of equations (A) and (B), then: ∆H for (C) = ∆H for (A) + ∆H for (B) Ex. 1: Given: C(s) + ½ O2(g) → CO(g) ∆H = -26.4 kcal CO(g) + ½ O2(g) → CO2(g) ∆H = -67.7 kcal Find: ∆H for: C(s) + O2(g) → CO2(g)

HESS’ LAW Ex. 2 Calculate ∆H for: NO(g) + ½ O2(g)  NO2(g)

HESS’ LAW Ex. 3 Given: 1) Sn(s) + ½ O2(g)  SnO(s) ∆H = -68 kcal 2) SnO2(s)  SnO(s) + ½ O2(g) ∆H = 70 kcal   Calculate the heat of formation of SnO2.  

HESS’ LAW Ex. 4: Find ∆H for: CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g) See Table 4.3 and H2O(l)  H2O(g) ∆H = 9.72 kcal

HESS’ LAW Ex. 5 Calculate ∆H for: 2 C(s) + H2(g)  C2H2(g) Given: C2H2(g) + 5/2 O2(g)  2 CO2(g) + H2O(l) ∆H = -310.6 kcal C(s) + O2 (g)  CO2(g) ∆H = -94.1 kcal H2(g) + ½ O2(g)  H2O(l) ∆H = -68.3 kcal  Calculate ∆H for: 2 C(s) + H2(g)  C2H2(g)