Topic 12.1 Strong and weak acids

Proposed the theory into how acids and bases react with each other (in 1923)

Acid + conjugate base Base + conjugate acid (strong reactactant forms weak opposite product)

Donates one proton (HCl is monoprotic/monobasic- as 1 mol of acid can neutralise 1 mol OH- ions)

Donates two protons (Sulfuric acid is diuretic/dibasic)

A substance that can act as both an acid and a base (water is anphoteric)

-Weak acid NH3 + H2O ⇌ NH4+ + OH- Base Acid Acid Base (NH3 and NH4 conjugate pair) (H2O and OH- conjugate pair)

Strong acids are almost completely dissociated in solution. –> Reverse reaction doesn’t take place so single direction arrow is used (HCl –> H+) –> pH increases by a factor of one for each 10x decrease of conc

Weak acids only partly dissociate <10% (a significant amount of undissociated acid is present). –> Represented using a reversible reaction arrow (HA ⇌ H+ + A-)–>(CH3COOH ⇌ CH3COO- + H+) –> pH increases 0.5 for each 10x decrease in concentration.

• Ka is the acid dissociation constant. • Used to quantify the degree to which an acid dissociates.

pH is related to the hydrogen ion concentration (moldm^-3) of a solution. pH = -log[H+] (eg: pH = -log0.100 pH = 1)

[H+] = 10^-pH

-Need to know the degree of dissociation of the acid -Use Ka the acid dissociation constant to show this ____________

pKa = -logKa -Larger Ka = stronger acid -Smaller Ka = weaker acid

Water ionises to form some H+ and OH- ions H2O ⇌ H+ + OH- Kc = [H+][OH-] / [H2O]

• Kw = [H+][OH-] • Kw is the ionic product of water. • It has a constant value at a given temperature (298K).

• Sodium hydroxide is a strong base. • It is completely dissociated in water.

• They resist changes in pH • A buffer solutions pH is almost unchanged on the addition of small amounts of acid or base • A buffer solution can be made from a weak acid and a salt of the acid.

• A weak acid with its conjugate base • A weak base with its conjugate acid

pH = pka + log [acid]/[base] (?) –> pH = (-logKa) + log [salt]/ [acid] [H+] = Ka x [acid]/[salt] pH = -log[H+]

Acid dissociation constant. HA(aq) ⇌ H+(aq) + A-(aq)

When an acid is dissolved in water, it produces so many hydrogen ions that the small contribution of the water is insignificant, unless the acid concentration is very small. However the fact that water ionises is the reason why even the most alkene solutions contain some hydrogen ions. Sodium hydroxide is a strong base, so in dilute aqueous solutions we can consider it ions to be completely dissociated.

• Proton donor • Positive charge/becomes more negative • Higher pH- weaker the acid • carboxylic- weak acid • Others strong

• Proton acceptor • Lone pair/becomes more positive • Higher pH- stronger base • OH- strong base • NH3- weak base

The equivalence point is where the moles of acid and the moles of base would neutralize each other according to the chemical reaction. (when an acid and base have reacted together in the exact proportions as dictatef by the stoichiometric equation)