lecture 2: chemical equilibria

reversible reaction

reactants and products are of similar stability

• goes from products to reactants and vice versa

reverse reaction

read from right to left

forward reaction

read from left to right

state of chemical equilibrium

both forward and reverse reactions occur until the concentration of reactants and products undergo no further change

• reaction has reached chemical equilibrium

• all substances present are being made and unmade at the same rate, so their concentrations are constant at equilibrium, even if they’re not equal

equilibrium constant K; much smaller than 0.001

only reactants are present at equilibrium; essentially no reaction occurs

equilibrium constant K; between 0.001 and 1

more reactants than products are present at equilibrium (reverse reaction is faster)

equilibrium constant K; between 1 and 1000

more products than reactants are present at equilibrium (forward reaction is faster)

equilibrium constant K; much larger than 1000

only products are present at equilibrium; reaction goes essentially to completion (irreversible reaction)

le chateliers principle (LCP)

when a stress is applied to a system at equilibrium, the equilibrium shifts to relieve the stress to restore an equilibrium

acid

• proton donor

• must contain a H in its formula

• H₂SO₄, HCl, H₂PO₄, HNO₃

base

• proton acceptor

• must contain a lone pair of electrons to bind to the H⁺ ion

• NH₃, CO₃^2-, OH^-

strong acid /base

dissociates completely

weak acid/base

does not dissociate completely

strong acid examples

HCl, HBr, HNO₃, H₃PO₃^-, H₂SO₄

weak acid examples

CH₃COOH, C₆H₅COOH, HNO₂, H₃PO₄

strong base examples

NaOH, KOH, CaO, Mg(OH)₂

weak base examples

NH₃, C₅H₅N

pH

the scale we use to measure how acidic or basic a solution is

• pH = - log₁₀[H₃O⁺]

pOH (potential of hydroxide ion)

a scale used to determine the hydroxide ion concentration in a solution

• pOH = - log₁₀[OH^-]

K_w (ionic product of water)

the equilibrium constant for the self-ionisation reaction of water

• [H₃O⁺] X [OH^-] = 1 X 10^-14 @ 25 degrees Celsius

K_a

acid dissociation constant

• when K_a is large, pK_a is small (strong acid)

• pK_a = -log₁₀ K_a

K_b

base dissociation constant

• when K_b is large = strong base

• pK_b= -log₁₀K_b

equivalence point - titrations

acid and its conjugate base are at completely equal concentrations

buffer

solutions/mixtures that maintain the pH approx. constant despite small addition of an aid or a base

• RESISTS CHANGES IN pH

• usually a mixture of a weak acid and its conjugate base

• When small quantities of H3O+ or OH- are added to the buffer, they cause a small amount of one buffer component to convert into the other

hendersson-hasselbalch equation

blood buffer