Complete vs Equilibrium Reactions
Complete Equilibrium
Open system Closed system
If endothermic, forward reaction then Ea is too high Rate of reverse = rate of forward
for reverse reaction to occur - reaction will likely go
to completion
Written with single arrow Written with double arrows
No further macroscopic changes occur
(e.g. constant colour, temp, pressure)
Constant concentrations of reactants
and products
Continual but equal interchange
between reactants and products on
microscopic level
Open vs Closed Systems
Open System Closed System
Open to atmosphere Not open to atmosphere
Can gain or lose energy across boundary Can gain or lose energy across boundary
Can gain or lose mass across boundary Cannot gain or lose mass across boundary
Static vs Dynamic Equilibrium
Static Dynamic
Irreversible reaction goes to completion Reversible reaction goes to eqm
Rate of reverse = rate of forward = 0 Rate of reverse = rate of forward
May be considered to be irreversible reactions Amount of products formed = amount of
products reacted
Constant concentrations Constant concentrations (not necessarily
equal)
Reaches completion Never reaches completion
Single arrow Double arrow
No interaction between products and reactants Constant interaction between products and
reactants
No microscopic/macroscopic changes Only microscopic changes
, Concentration vs Time Graphs
Change Concentration Volume Temperature
Increase Sharp spike in Sharp Gradual
one species spike in change/no
all species sharp spike
Amount vs Time Graphs
Change Concentration Volume Temperature
Increase Sharp spike in one Gradual change/no sharp Gradual change/no sharp
species spikes spike
Non-Equilibrium Systems
Enthalpy
● Enthalpy (ΔH): amount of heat produced in a reaction carried out at constant pressure
● Energy changes in chemical reactions are caused by bond breaking (endothermic) and
bond forming (exothermic)
● Bond energies (bond enthalpies) can be used to estimate the enthalpy
● ∆H = (energy to break reactant bonds) – (energy released from product bonds)
● ΔH = -ve: exothermic
● ΔH = +ve: endothermic
Entropy
● Entropy (ΔS): measure of disorder in a system
● ΔS = ∑S(products) – ∑S(reactants)
● ΔS = +ve: disordered
● ΔS = -ve: ordered
● For gaseous chemical reactions, the more particles of gas there are in the system, the
higher the entropy of the system
Gibbs Free Energy
● ∆G = ∆H - TΔS
● REMEMBER: ΔH is in kJ but ΔS is normally in J
● Unit: kJ/mol
● ΔS = -ve: spontaneous
● ΔS = 0: eqm
● ΔS = +ve: not spontaneous
Combustion Reactions
● Irreversible, exothermic reaction
Complete Equilibrium
Open system Closed system
If endothermic, forward reaction then Ea is too high Rate of reverse = rate of forward
for reverse reaction to occur - reaction will likely go
to completion
Written with single arrow Written with double arrows
No further macroscopic changes occur
(e.g. constant colour, temp, pressure)
Constant concentrations of reactants
and products
Continual but equal interchange
between reactants and products on
microscopic level
Open vs Closed Systems
Open System Closed System
Open to atmosphere Not open to atmosphere
Can gain or lose energy across boundary Can gain or lose energy across boundary
Can gain or lose mass across boundary Cannot gain or lose mass across boundary
Static vs Dynamic Equilibrium
Static Dynamic
Irreversible reaction goes to completion Reversible reaction goes to eqm
Rate of reverse = rate of forward = 0 Rate of reverse = rate of forward
May be considered to be irreversible reactions Amount of products formed = amount of
products reacted
Constant concentrations Constant concentrations (not necessarily
equal)
Reaches completion Never reaches completion
Single arrow Double arrow
No interaction between products and reactants Constant interaction between products and
reactants
No microscopic/macroscopic changes Only microscopic changes
, Concentration vs Time Graphs
Change Concentration Volume Temperature
Increase Sharp spike in Sharp Gradual
one species spike in change/no
all species sharp spike
Amount vs Time Graphs
Change Concentration Volume Temperature
Increase Sharp spike in one Gradual change/no sharp Gradual change/no sharp
species spikes spike
Non-Equilibrium Systems
Enthalpy
● Enthalpy (ΔH): amount of heat produced in a reaction carried out at constant pressure
● Energy changes in chemical reactions are caused by bond breaking (endothermic) and
bond forming (exothermic)
● Bond energies (bond enthalpies) can be used to estimate the enthalpy
● ∆H = (energy to break reactant bonds) – (energy released from product bonds)
● ΔH = -ve: exothermic
● ΔH = +ve: endothermic
Entropy
● Entropy (ΔS): measure of disorder in a system
● ΔS = ∑S(products) – ∑S(reactants)
● ΔS = +ve: disordered
● ΔS = -ve: ordered
● For gaseous chemical reactions, the more particles of gas there are in the system, the
higher the entropy of the system
Gibbs Free Energy
● ∆G = ∆H - TΔS
● REMEMBER: ΔH is in kJ but ΔS is normally in J
● Unit: kJ/mol
● ΔS = -ve: spontaneous
● ΔS = 0: eqm
● ΔS = +ve: not spontaneous
Combustion Reactions
● Irreversible, exothermic reaction
