Unit 6: Thermochemistry

6.0 Introduction, 6.1 Endothermic and Exothermic Processes, 6.2 Energy Diagrams, 6.3 Heat Transfer and Thermal Equilibrium, 6.4 Heat Capacity and Calorimetry, 6.5 Energy of Phase Changes, 6.6 Introduction to Enthalpy of Reaction, 6.7 Bond Enthalpies, 6.8 Enthalpy of Formation, 6.9 Hess's Law

types of energy

kinetic and potential energy

thermal energy

type of kinetic energy involving heat; molecule vibrations

chemical energy

type of potential energy involving stored energy in bonds or IMFs

unit for energy

joules

enthalpy

denoted H, total energy/heat content of system

enthalpy change

denoted ΔH, amount of heat transferred during a reaction

heat

denoted q, form of energy directly measured by changes in temperature

equation for heat

q = mcΔT

assume that q =

ΔH

in calorimetry, often assume

solution has same density (1g/ml) and specific heat capacity (4.18J/g°C) as water

units for q

J

units for ΔH

kJ/mol

+ΔH means reaction is

endothermic

-ΔH means reaction is

exothermic

endothermic reaction

products have more PE than reactants; more energy was absorbed to break bonds than what was released to form new bonds

exothermic reaction

reactants have more PE than products; less energy was absorbed to break bonds than what was released to form new bonds

what type of reactions are always exothermic

combustion, condensing, freezing, deposition (molecules release energy to form IMFs)

what type of reactions are always endothermic

melting, evaporation, sublimation (molecules absorb energy to weaken/overcome IMFs)

dissolution in thermodynamics

the ionic solid and water must undergo an endothermic reaction to break apart, then come together and mix in an exothermic reaction; enthalpy of solution is the difference between these processes

how to represent endothermic reactions on a diagram

products have higher energy than reactants

how to represent exothermic reactions on a diagram

reactants have higher energy than products

temperature

measure of average kinetic energy

thermal equilibrium

two substances reach same temperature and average KE

which way does energy flow

from hot to cold

how is KE transferred

through collisions

heat capacity

denoted C, quantity of energy needed to change temperature

specific heat capacity

denoted c, quantity of heat required to raise the temperature of 1g of a specific substance by 1°C

specific heat capacity of water

4.18 J/g°C, characteristically high

specific heat of water vs. metals

metals have lower specific heat than water—easier to change temperature

why do we set q = -q

heat gained = heat lost

what happens during phase changes (assume substance is being heated, so melting/boiling)

PE increases, IMFs are weakened/broken, temperature does not change

what happens in between phase changes as a substance is heated

temperature increases, KE increases

what equation to use when substance is heating/cooling

q = mcΔT

what equation to use when substance is changing phases

q = ΔHm, where ΔH is given in J/g or kJ/mol, m = mass or moles depending on unit of ΔH

standard conditions

room temperature, 1 atm, solutions 1M, substances in pure form

bond enthalpy

enthalpy change required to break/form given bond

enthalpy of bond formation

always exothermic, energy released to form 1 mole of covalent bonds

enthalpy of bond breaking

always endothermic, energy required to break 1 mole of covalent bonds

how to use bond enthalpy to calculate ΔH

1. given enthalpy of bond breaking

2. count up number of each specific type of bond in reactants and products

3. multiply by given values for each type of bond

4. for products (bond formation) take the negative of given values

5. add all values together

enthalpy of formation

energy change when 1 mole of substance formed (must maintain 1 mole of product); enthalpy of an element in pure form is 0 kJ/mol

how to use enthalpy of formation to calculate ΔH

1. given enthalpies of each molecule

2. use given values to sum up and calculate enthalpy of reactants and products

3. remember, elements in pure form (e.g. O₂) have 0 enthalpy

4. ΔH = ∑ΔH_products - ∑ΔH_reactants

how to use Hess’s Law to calculate ΔH

1. given separate rxns and ΔH_rxn of each

2. manipulate equations to create original equation (multiply to get desired coefficient, multiply by -1 to reverse rxn):

3. start with the substances in the original equation that show up once in the list of equations

4. manipulate the rest so they cancel each other out

5. add up resulting enthalpies to get ΔH