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, E k : macroscopic energy associated with an object’s movement –Potential energy, E p : 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 –E tot = E k + E p + 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 = q p (enthalpy change = reaction heat at constant pressure) –An exothermic reaction might have q p = -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) + 2H 2 O(l)  2NaOH(aq) + H 2 (g); ∆H = -368.6 kJ 368.6 kJ of heat is evolved when 2 mol Na react with 2 mol H 2 O to form 2 mol NaOH and 1 mol H 2 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 CH 4 (g) + 2O 2 (g)  CO 2 (g) + 2H 2 O(l); ∆H = -890.3 kJ In excess oxygen, how many kJ of heat can be obtained from burning 36.0 g CH 4 ? 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 = t f - t i 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 q calorimeter = -q reaction –Divide q reaction 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, ∆H f °: enthalpy change for formation of one mole of the substance in its standard state from its elements ∆H ° =  n ∆H f °(products) -  n ∆H f °(reactants)