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 V

Pext = constant external pressure
ΔV = Vf - Vi

 


Internal energy U (in J):

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

The 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)


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)
- reversing a chemical equation: ΔHrxn = ΔHrxn (forward) + ΔHrxn (reverse) = 0
⇒ ΔHrxn (forward) = - ΔHrxn (reverse)
- multiplication of a chemical equation A by a factor of n: ΔHrxn = n x ΔHrxn (A)

 

ΔHsublimation = ΔHfusion + ΔHvaporization

Enthalpy of Formation

Standard states: reactant 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 reactants in their standard states change to product 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 (reactants)
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 (reactants) - 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:

cPqPT

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 cPn

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 = cPm

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