# Thermochemistry | General Chemistry 2

Thermochemistry is studied in this chapter: internal energy, heat and work, enthalpy, Hess’s law, enthalpy of formation, bond enthalpies, heat capacity

## Important Terms in Thermochemistry

Thermochemistry: study of the energy change associated with chemical reactions

System: object being investigated (= reaction mixture in thermochemistry)
Surroundings: everything else
Surroundings can exchange energy and matter if the system is open or only energy if the system is closed. There is no exchange with an isolated system

Process: path of a system from an initial state to a final state
The final state is generally an equilibrium ⇒ all the properties of the system do not change with time

Isobaric = constant pressure
Isochoric = constant volume
Isothermal process = constant temperature
Adiabatic = no matter or heat transfer from the system to its surroundings

## Internal Energy, Heat and Work

Heat q (in J):

Transfer of energy based on a temperature difference (from high temperature region to low temperature region)
q < 0: reaction releases energy as heat ⇒ exothermic process
q > 0: reaction absorbs heat energy ⇒ endothermic process

Work w (in J):

Force x distance
w < 0: work is done by the system ⇒ expansion
w > 0: work is done on the system ⇒ compression

Gases can do work:

wgas = -

Pext = constant external pressure
ΔV =

Internal energy U (in J):

Sum of kinetic and potential energies of the particles of a system:
Usys = Ek + Ep

Energy change of a system: ΔU = Uf - Ui
ΔU < 0: energy flows from the system into the surroundings
ΔU > 0: energy flows from the surroundings into the system

First Law of Thermodynamics:

The energy of the universe is constant ⇒ energy can neither be created nor destroyed

ΔUuniv = ΔUsys + ΔUsurr = 0
⇒ ΔUsys = - ΔUsurr = q + w

## Enthalpy

Enthalpy H (in J):

Energy transferred between a system and the surroundings under isobaric conditions
Hsys = Usys + PV

H is a state function:
ΔHrxn = Hprod – Hreact

The change in enthalpy for a reaction is:
ΔHrxn = ΔUrxn + Δ(PV)rxn
ΔHrxn = ΔUrxn + PΔVrxn     (P = cst, isobaric conditions)

ΔH < 0 ⇒ exothermic process
ΔH > 0 ⇒ endothermic process

## Enthalpy Relationships

Hess’s Law:

Enthalpy changes for chemical equations are additive

Consequences:

• reaction with 2 steps (or more): ΔHrxn = ΔHrxn (step 1) + ΔHrxn (step 2)
• reverse chemical equation: ΔHrxn = ΔHrxn (forward) + ΔHrxn (reverse) = 0

⇒ ΔHrxn (reverse) = - ΔHrxn (forward)

• multiplication of a chemical equation A by a factor n: ΔHrxn = n x ΔHrxn (A)

ΔHsublimation = ΔHfusion + ΔHvaporization

## Enthalpy of Formation

Standard states: reagents and products in their pure forms at T = 298 K and P = 1 bar
(or at a concentration of 1 mol.L-1 for a liquid solution)

Standard reaction enthalpy ΔH0rxn (in J):

Enthalpy change when reagents in their standard states change to products in their standard states

Standard enthalpy of formation ΔH0f (in J):

Standard reaction enthalpy for the formation of the compound from its elements in their standard states. The standard enthalpy of formation of the most stable form of an element is 0

ΔH0f of H2 (g) = 0 J
ΔH0f of O2 (g) = 0 J
ΔH0f of C (graphite) = 0 J

ΔH0rxn = Σ n ΔH0f (products) - Σ n ΔH0f (reagents)
with n = stoichiometric coefficients in the reaction

ΔH0rxn of the combustion of propane:
C3H8 (g) + 5 O2 (g) → 3 CO2 (g) + 4 H2O (l)
ΔH0rxn = 3 ΔH0f (CO2) + 4 ΔH0f (H2O) - ΔH0f (C3H8) - 5 ΔH0f (O2)

## Bond Enthalpies

Bond enthalpy ΔHbond (in J):

Energy associated with the strength of a chemical bond
ΔHbond > 0 because heat must be supplied to break a bond

ΔHrxn = Σ Hbond broken - Σ Hbond formed = Σ Hbond (reagents) - Σ Hbond (products)

## Heat Capacity

Heat capacity cP (in J.K-1):

Amount of heat needed to raise the temperature of the sample by one degree
At constant pressure:

cP$\frac{{\mathrm{q}}_{\mathrm{P}}}{∆\mathrm{T}}$

qP = energy added as heat (in J)
ΔT = Tf – Ti (in K)

Molar heat capacity CP (in J.K-1.mol-1):

Amount of heat needed to raise the temperature of one mole of a sample by one degree

CP $\frac{{\mathrm{c}}_{\mathrm{P}}}{\mathrm{n}}$

cP = heat capacity (in J.K-1)
n = number of mole (in mol)

Specific heat capacity cS (in J.K-1.g-1):

Amount of heat needed to raise the temperature of one gram of a sample by one degree

cS = $\frac{{\mathrm{c}}_{\mathrm{P}}}{\mathrm{m}}$

cP = heat capacity (in J.K-1)
m = mass (in g)