Week 10 Flashcards
ΔG: the Change in the Gibbs Free Energy
ΔG = ΔH -TΔSsystem = -TΔStotal
Precipitation
a loss of entropy → ΔS < 0
a loss of enthalpy → ΔH < 0
Shifting equilibrium with temperature
- The exo direction of a reaction will be favoured as we lower T
- The endo direction of a reaction will be favoured as we raise T
If the concentration ____, the system acts to ____ some of it
- increases, consumes
- decreases, produces
There are 3 ways in which the pressure of a system can be changed:
- By adding (or removing) a gaseous reactant or product:
- By adding a gas that is not involved in the reaction:
- By changing the volume of the container.
Change in volume
- If the volume decreases, the system acts to decrease the # of moles present
- If the volume increases, the system acts to increase the # of moles to fill it.
We have two approaches to determining the direction of spontaneous change, what is the connection?
ΔGº = – RT lnKeq
If ΔHº < 0 then increasing T
will make Keq smaller
* The equilibrium will shift to favour the reactants and so to reduce the heat released
If ΔHº > 0 then increasing T
will make Keq larger
* The equilibrium will shift to favour the products and absorbing more heat
increase reactant concentrations ⇒
more products produced
Increase pressure / decrease volume ⇒
less gas produced
Increase temperature ⇒
endothermic reaction favoured
Brønsted - Lowry
(H+) + (A-) ⟺ HA
* ACID: proton (H+) donor.
* BASE: proton (H+) acceptor.
Arrhenius
H+(aq) + OH-(aq) ⟺ H2O(l)
* ACID: H+ producer in aqueous solution.
* BASE: OH- producer in aqueous solution
What is H+(aq)?
In water: an acid (e.g., HCl) ionises to produce H+ (aq)
Acids & equilibrium
A STRONG acid has equilibrium to the right
* (HA completely ionised)
A WEAK acid has equilibrium to the left
* (HA partly/mostly intact)
Conjugate acid-base pairs
conjugate base has one less proton than its conjugate acid
H2SO4 is a dibasic or diprotic acid:
H2SO4(aq) + H2O(l) ⟺ H3O+(aq) + HSO4-(aq)
HSO4-(aq) + H2O(l) ⟺ H3O+(aq) + SO42-(aq)
Autoionisation of Water
2H2O(l) ⟺ H3O+(aq) +OH-(aq)
pH
pH = -log[H+]
* Acid : pH < 7
* Neutral: pH = 7
* Basic: pH > 7
Strong acids and bases…
Completely ionise in water:
* e.g. HCl(aq) ⟺H+ (aq) + Cl- (aq)
* Equilibrium lies completely to the right, Ka ≈ ∞
Weak acids
They do not completely ionise in water
* HA(aq) ⟺ H+(aq)+ A-(aq)
Acid dissociation constant:
* Ka= (H+)(A-)/(HA)
pKa = - log(Ka)
The larger the value of Ka…
the stronger the acid and the lower the value of pKa
Weak base
Ionisation of a weak base:
* NH3(aq) + H2O(l) ⟺ NH4+ (aq) + OH-(aq)
Equilibrium constant is called base ionisation constant:
* Kb=(NH4+)([OH-)/(NH3)
