Thermochemistry | General Chemistry 2

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

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:

Internal energy U (in J):

Sum of kinetic and potential energies of the particles of a system:
U_sys = E_k + E_p

Energy change of a system: ΔU = U_f - U_i
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

ΔU_univ = ΔU_sys + ΔU_surr = 0
⇒ ΔU_sys = - ΔU_surr = q + w

Enthalpy H (in J):

Energy transferred between a system and the surroundings under isobaric conditions
H_sys = U_sys + PV

H is a state function:
ΔH_rxn = H_prod – H_react 

The change in enthalpy for a reaction is:
ΔH_rxn = ΔU_rxn + Δ(PV)_rxn 
ΔH_rxn = ΔU_rxn + PΔV_rxn     (P = cst, isobaric conditions)

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

Hess’s Law:

Enthalpy changes for chemical equations are additive

Consequences:

- reaction with 2 steps (or more): ΔH_rxn = ΔH_rxn (step 1) + ΔH_rxn (step 2)
- reverse chemical equation: ΔH_rxn = ΔH_rxn (forward) + ΔH_rxn (reverse) = 0
⇒ ΔH_rxn (reverse) = - ΔH_rxn (forward)
- multiplication of a chemical equation A by a factor n: ΔH_rxn = n x ΔH_rxn (A)

ΔH_sublimation = ΔH_fusion + ΔH_vaporization

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 ΔH⁰_rxn (in J):

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

Standard enthalpy of formation ΔH⁰_f (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
 

ΔH⁰_f of H₂ (g) = 0 J
ΔH⁰_f of O₂ (g) = 0 J
ΔH⁰_f of C (graphite) = 0 J

ΔH⁰_rxn = Σ n ΔH⁰_f (products) - Σ n ΔH⁰_f (reagents)
with n = stoichiometric coefficients in the reaction
 

ΔH⁰_rxn of the combustion of propane:
C₃H₈ (g) + 5 O₂ (g) → 3 CO₂ (g) + 4 H₂O (l)
ΔH⁰_rxn = 3 ΔH⁰_f (CO₂) + 4 ΔH⁰_f (H₂O) - ΔH⁰_f (C₃H₈) - 5 ΔH⁰_f (O₂)

Bond enthalpy ΔH_bond (in J):

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

ΔH_rxn = Σ H_{bond broken} - Σ H_{bond formed} = Σ H_bond (reagents) - Σ H_bond (products)

Heat capacity c_P (in J.K^-1):

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

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

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

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

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