Chapter 6: Thermochemistry

1. Chapter 6: Thermochemistry Chemistry 1061: Principles of Chemistry I Andy Aspaas, Instructor

2. Thermochemistry • Thermodynamics: relationships between heat and other forms of energy • Thermochemistry: an area of thermodynamics that involves heat transferred due to a chemical reaction • Energy: potential or capacity to move matter • Kinetic energy, Ek: macroscopic energy associated with an object’s movement • Potential energy, Ep: macroscopic energy assiciated with an object’s position in a field of force (only relative values) • Internal energy, U: microscopic sum of energy contained in a substance’s particles • Etot = Ek + Ep+ U

3. Heat of reaction • System: collection of substances in which the thermodynamic change is happening • Surroundings: everything outside the system, includes the flask, the room, and the universe • Heat, q: energy that flows into or out of a system because of a difference in temperature • Thermal equilibrium: heat flows from areas of high temperature to areas of low temperature until the temperatures are equal

4. Energy added or subtracted from a system • The sign of q is viewed from the perspective of the system, not the surroundings • Exothermic process, the reaction vessel warms, so energy must have left the system: q is – • Endothermic process, the reaction vessel cools, so energy must have been added to the system: q is +

5. Enthalpy • Enthalpy, H: an extensive property (depends on quantity) which is related to the amount of heat that can be absorbed or evolved in a chemical reaction • H is a state function (only depends on present state, independent of any history) • ∆H = change in enthalpy = H(products) - H(reactants) • ∆H = qp(enthalpy change = reaction heat at constant pressure) • An exothermic reaction might have qp= -400 kJ • So, ∆H = -400 kJ

6. Thermochemical equations • Thermochemical equation: balanced molar chemical equation, with enthalpy of reaction given after the equation 2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g); ∆H = -368.6 kJ • 368.6 kJ of heat is evolved when 2 mol Na react with 2 mol H2O to form 2 mol NaOH and 1 mol H2 • Phase labels are important, ∆H may be different depending on phase of product • If a thermochemical equation is multiplied by anything, ∆H is also multiplied • If a thermochemical equation is reversed, the sign of ∆H is also reversed

7. Stoichiometry and heats of reaction CH4(g) + 2O2(g)  CO2(g) + 2H2O(l); ∆H = -890.3 kJ • In excess oxygen, how many kJ of heat can be obtained from burning 36.0 g CH4? • Use ∆H as a conversion factor from mol to kJ

8. Heat capacity • Heat capacity, C: quantity of heat required to raise temperature of a substance one degree Celsius q = C∆t where ∆t = tf - ti • Specific heat, s: quantity of heat required to raise temperature of one gram of a substance by one degree Celsius q = sm∆t • Water has a very high specific heat: 4.18 J / (g · °C)

9. Measuring heat of reaction • Calorimeter: device used to measure heat of reaction • Insulated reaction vessel, with thermometer to record temperature change • Two nested styrofoam cups works well • First find amount of heat absorbed by calorimeter q = sm∆t and q = C∆t • Heat absorbed by surroundings is opposite of heat given off by system qcalorimeter = -qreaction • Divide qreaction by correct number of moles of limiting reactant to set up a thermodynamic equation

10. Hess’s law • Adding reactions together will also add ∆H for each reaction • So, ∆H can be found for reactions where it would be difficult to measure experimentally • Remember, multiplying a reaction by a constant also multiplies the ∆H by that constant, and reversing a reaction reverses the sign of ∆H

11. Standard enthalpies of formation • Standard enthalpy of reaction, ∆H °: ∆H at 25° C and 1 atm • Standard enthalpy of formation, ∆Hf°: enthalpy change for formation of one mole of the substance in its standard state from its elements ∆H ° =  n ∆Hf°(products) -  n ∆Hf°(reactants)