AP Chemistry Thermochemistry IMPROVED

Thermochemistry

Study of energy changes in chemical reactions and physical changes

Energy

Ability to do work or produce heat

Heat (q)

Energy transferred due to temperature difference

System

Part of the universe being studied

Surroundings

Everything outside the system

Exothermic

Heat released from system to surroundings

Endothermic

Heat absorbed by system from surroundings

Heat of reaction (q_rxn)

Heat absorbed or released in a reaction

Enthalpy (H)

Total heat content of a system

Enthalpy change (ΔH)

Heat of reaction at constant pressure

ΔH = q_p

Enthalpy change equals heat at constant pressure

Bond breaking

Requires energy

Bond forming

Releases energy

Overall ΔH

Depends on net energy absorbed or released

Specific heat capacity (c)

Heat to raise 1 g by 1°C

Units of c

J/g·°C

Heat capacity (C)

Heat to raise entire object by 1°C

q = mcΔT

Heat using mass and specific heat

q = CΔT

Heat using heat capacity

ΔT

Tfinal − Tinitial

Calorimetry

Measurement of heat flow

Calorimeter

Device measuring heat

Constant-pressure calorimeter

Measures heat at constant pressure

Constant-volume calorimeter

Measures heat at constant volume

qcal = −qrxn

Heat absorbed by calorimeter equals negative of reaction heat

Heat flow convention

Heat gained positive

Bomb calorimeter

Constant-volume calorimeter for combustion

ΔE_rxn

Internal energy change measured in bomb calorimeter

q = C_calΔT

Heat absorbed by calorimeter

Thermochemical equation

Balanced equation with ΔH

ΔH applies

ΔH valid only as written

Reverse reaction

Sign of ΔH reversed

Multiply reaction

Multiply ΔH by same factor

Add reactions

Add ΔH values

Stoichiometry with ΔH

Heat scales with amount reacting

Limiting reactant

Determines total heat change

Hess’s Law

ΔH depends only on initial and final states

Path independence

Different paths give same ΔH

Using Hess’s Law

Add

Phase change

Physical change with energy transfer

Change of state

Solid-liquid-gas transitions

Heating curve

Graph of temperature vs heat added

Heating curve for water

Solid → liquid → gas

Sloped region

Temperature changes in one phase

Flat region

Temperature constant during phase change

Solid warming

Particles vibrate faster

Liquid warming

Particles move faster but close

Gas warming

Particles move faster and farther apart

Melting plateau

Solid-liquid coexistence

Boiling plateau

Liquid-gas coexistence

Latent heat

Energy during phase change at constant temperature

Energy use

Overcomes intermolecular forces

Melting/fusion

Solid to liquid

Freezing

Liquid to solid

Vaporization

Liquid to gas

Condensation

Gas to liquid

Sublimation

Solid to gas

Deposition

Gas to solid

ΔH_fus

Heat to melt 1 g

ΔH_vap

Heat to vaporize 1 g

ΔHvap > ΔHfus

Vaporization requires more energy

Boiling point

Temp where vapor pressure = external pressure

Condensation point

Same temp as boiling but releases energy

Heat transfer w/o phase change

Changes temperature

Heat transfer w/ phase change

Changes state

Cooling curve

Reverse of heating curve

Energy flow

Heating absorbs

Intermolecular forces

Energy needed to overcome attractions

Stronger forces

Require more energy to change phase

Standard heat of formation (ΔH°_f)

ΔH for forming 1 mol from elements

Standard state

1 atm

ΔH°_f elements

Zero

Formation reaction rule

Must produce 1 mol of product

ΔH°_rxn

ΣΔH°f(products) − ΣΔH°f(reactants)

Coefficient rule

Multiply ΔH°_f by balanced coefficient

Phase matters

ΔH°_f depends on state

Phase change enthalpy

Can combine using Hess’s Law

ΔHsub = ΔHfus + ΔH_vap

Sublimation equals fusion plus vaporization

Methods to find heat

Calorimetry

Units of ΔH

kJ or kJ/mol

Sign of ΔH

Negative exothermic

Magnitude of ΔH

Proportional to amount reacting