Thermochemistry

Thermochemistry Energy-capacity to do work or to produce heat Law of Conservation of Energy 1st law of thermodynamics Energy can be neither created nor destroyed Energy of universe is constant Potential energy in chemistry, this is usually the energy stored in reactants and products Stored energy; sometimes because of position PE = mgh Kinetic energy Energy involving movement

Energy • System Reactants and products of a reaction • Surroundings Anything other than reactants and products • Temperature Measure of motion of particles (average kinetic energy)

Exothermic reaction • ∆H is negative • energy flows out of the system (heat) (-q ) • temp↑ • work done by system (on the surroundings ( -w ) • ex. compression of gas reactants E products Reaction proceeds 

Endothermic reaction • ∆H is positive • reaction that absorbs energy • energy flow into a system (heat) ( +q ) • temp↓ • work done by surroundings( + w ) • ex. expansion of gas products E reactants Reaction proceeds 

Exothermic ProcessCH4 + 2O2 CO2 + 2H2O + energy System 2 mol O2 1 mol CH4 (Reactants) Energy released to the surroundings as heat Potential energy 2 mol H2O 1 mol CO2 (Products)

Endothermic ProcessN2 + O2 + energy (heat)  2NO 2 mol NO (Product) System Potential Energy Heat absorbed from the surroundings 1 mol N2 1 mol O2 (Reactants)

∆E = q + w Example 1: Calculate the ∆E for a system undergoing an endothermic process in which 15.6 kJ of heat flows and where 1.4 kJ of work is done on the system. ∆E = + 15.6 kJ + + 1.4 kJ = 17.0 kJ q is positive b/c endothermic w is positive b/c work is done on the system

Example 2: q = + 51 kJ and w = - 15 kJ ∆E = 36 kJ q = +47 kJ and w = 88 kJ ∆E = 41 kJ Which of the above has work done on surroundings? -15 kJ

Enthalpy H Heat content at constant pressure ∆H = Hproducts - Hreactants Enthalpy of reaction ∆Hrxn heat absorbed or released by a chemical reaction Enthalpy of combustion ∆Hcomb heat absorbed or released when burning Enthalpy of formation ∆Hf heat absorbed or released when ONE mole of compound is formed from elements in their standard states Enthalpy of fusion ∆Hfus heat absorbed to melt 1 mole of solid to liquid @ MP Enthalpy of vaporization ∆Hvap heat absorbed to change change 1 mole of liquid to gas @ BP

Endothermic reactions +H Exothermic reactions -H Enthalpy can be calculated: Stoichiometrically From tables Using Hess’ Law

S(s) + O2 (g)  SO2 (g) ∆H = -296 kJ/mol • calculate the heat evolved when 275 g S is burned 275 g 1mol S 296 kJ = -2535.8 kJ 32.1 g S 1 mol = 2.54 X 10-3 kJ b) calculate the heat evolved when 150 g SO2 is produced 150 g SO2 1 mol SO2 1 mol S 296 kJ = 693 kJ 64.1 g SO2 1 mol SO2 1 mol S

When 1 mole of methane (CH4) is burned at constant pressure, 890 kJ of energy is released as heat. Calculate H for a process in which a 5.8 g sample of methane is burned at constant pressure. At constant pressure, q = H = -890 kJ/mol

Calorimetry Science of measuring heat Calorimeter constant pressure (coffee cup) or constant volume (bomb) Heat capacity C = heat absorbed increase in temperature • Specific heat capacity- energy required to raise the temperature of one gram of a substance by one degree Celsius q = m c  T • Molar heat capacity-energy required to raise the temperature of one mole of a substance by one degree Celsius

Energy released by a reaction = energy absorbed by the solution • If two reactants at the same temperature are mixed and the resulting solution gets warmer, this means the reaction taking place is exothermic  H = (-) • If the solution gets cooler, the reaction is endothermic  H = (+)

Consider the dissolution of CaCl2: CaCl2(s) Ca2+(aq) + 2Cl-(aq) H = -81.5 kJ An 11.0 g sample of CaCl2 is dissolved in 125 g of water with both substances at 25.0°C. Calculate the final temperature of the solution assuming no heat lost to the surroundings and assuming the solution has a specific heat capacity of 4.18J/goC

A 46.2 g sample of copper is heated to 95.4 °C and then placed in a calorimeter containing 75.0 g water at 19.6 °C. The final temperature of the metal and water is 21.8°C. Calculate the specific heat capacity of copper, assuming that all the heat lost by the copper is gained by the water.

Camphor (C10H16O) has an energy of combustion of -5903.6 kJ/mol. When a sample of camphor with mass 0.1204 g is burned in a bomb calorimeter, the temperature increases by 2.28°C. Calculate the heat capacity of the calorimeter.

Hess’s Law In going from a particular set of reactants to a particular set of products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps.

Key Concepts • When a reaction is reversed, the sign changes. • When the equation is multiplied by an integer, ΔH must be multiplied by the same integer. • ΔH° for a reaction can be calculated from the ΔH°f of the reactants and products: ΔH°rxn = ∑np ΔH°f(products) − ∑nr ΔH°f(reactants) • Elements in their standard states are not included in the calculations. ΔH°f for an element in its standard state is zero.

N2(g) + O2(g) 2NO(g)H2 = 180 kJ 2NO(g) + O2(g) 2NO2(g)H3 = -112 kJ ___________________________________________ Net reaction: N2(g) + 2O2(g) 2NO2(g) H2 + H3 = 68kJ

2B(s) + 3H2(g) B2H6(g) H 2B(s) + 3/2 O2(g)  B2O3(s) -1273 kJ B2H6(g) + 3O2(g)  B2O3(s) + 3H2O(g) -2035 kJ H2(g) + ½ O2(g)  H2O(l) -286 kJ H2O(l)  H2O(g) 44 kJ

The bombardier beetle uses an explosive discharge as a defensive measure. The chemical reaction involved is the oxidation of hydroquinone by hydrogen peroxide to produce quinone and water. C6H4(OH)2(aq) + H2O2(aq) C6H4O2(aq) + 2H2O(l) Calculate H for this reaction from the following data: C6H4(OH)2(aq)  C6H4O2(aq) + H2(g) DH = +177.4 kJ H2(g) + O2(g)  H2O2(aq)DH = -191.2kJ H2(g) + ½ O2(g)  H2O(g)DH = -241.8 kJ H2O(g)  H2O(l)DH = -43.8 kJ

Standard Enthalpy of Formation Change in enthalpy that accompanies the formation of one mole of a compound from its elements with all substances in their standard states ∆H° indicates that the process has been carried out under standard conditions.

Standard States • Compound • Standard state of gaseous substance is a pressure of exactly 1 atm • For a pure substance in a condensed state (liquid or solid), the standard state is the pure liquid or solid • For a substance present in a solution, the standard state is a concentration of exactly 1 M • Element • The standard state of an element is the form in which the element exists under conditions of 1 atm and 25°C

Enthalpy of Formation Calculate the ∆H°rxn for CaCO3 CaCO3 (s) CaO (s) + CO2 (g) ∆H°f [CaCO3 (s)] = -1206.9 kJ/mol ∆H°f [CaO (s)] = -635.1 kJ/mol ∆H°f [CO2 (s)] = -393.5 kJ/mol

Calculate the change in enthalpy for this reaction: 3 Al(s) + 3NH4ClO4(s) Al2O3(s) + AlCl3(s) + 3NO(g) + 6H2O(g) Appendix 4 Page A21

Write an equation for the combustion of glucose, C6H12O6. Its heat of combustion has been determined in a bomb calorimeter to be -2800 kJ. Calculate the enthalpy of formation of glucose given the following: ΔH°f CO2(g) = -393.5 kJ/mol ΔH°f H2O(l) = -285.8 kJ/mol

Consider the reaction 2ClF3(g) + 2NH3(g) N2(g) + 6HF(g) + Cl2(g) ΔH° = -1196 kJ Calculate ΔH°f for ClF3(g)

Propane, C3H8, is a hydrocarbon that is commonly used as fuel for cooking. • Write a balanced equation for the complete combustion of propane gas, which yields CO2(g) and H2O(l). The heat of combustion of propane is -2220.1 kJ/mol. • Calculate the heat of formation, ΔH°f, of propane given that ΔH°f of H2O(l) = -285.3 kJ/mol and ΔH°f of CO2(g) = -393.5 kJ/mol • Assuming that all of the heat evolved in burning 30.0 grams of propane is transferred to 8.00 kilograms of water (specific heat = 4.18 j/g·K), calculate the increase in temperature of the water.

Calculate the volume of air at 30.0°C and 1.00 atmosphere that is needed to burn completely 10.0 grams of propane. Assume that air is 21.0 percent O2 by volume.

Thermochemistry