Year 12 Chemistry

Historical Development of Acid/Base Theories

He classified all chemicals into three categories – acids, bases and salts

He believed that all acids contained oxygen and it was this that gave them their sour taste

Flaw: not all acids contain oxygen and metal oxides form bases

Antoine Lavoisier (1743-1794)

Showed that all acids do not contain oxygen

Proposed that acids are hydrogen containing materials following the discovery of HCl

Flaw: not all substances that contain hydrogen are acids

Humphrey Davy (1778-1829)

Acids dissociate in water forming H+ as one product

Bases dissociate in water forming OH- as one product

Neutralisation involves the reaction of H+ and OH- forming a salt in water

Flaws: theories only apply to aqueous

solutionsSome substances such as NH3 are

bases and do not contain OH-Relative strengths not addressedAmphoteric substances not addressed

Svante Arrhenius (1859-1927)

An acid is a proton (H+) donorA base is a proton acceptorExamples:

HCl + H2O <===> H3O+ + Cl¯NH3 + H2O <===> NH4

+ + OH¯

Bronsted-Lowry (~1923)

Identify the acids/bases

Any other acids/bases here?

B-L ExamplesHCl + H2O H3O+ + Cl-

NH3 + H2SO4 NH4+ + HSO4-

HBr + NH2+ NH3+ + Br-

Identify the B-L acids and bases in each of these reactions.

Amphiprotic substances are those that can act as bases and acids. They can donate or accept protons.

Water is an obvious example

Notice in the previous slide that water reacts with both acids and bases.

Amphiprotic Substances

Bisulfate is another amphiprotic substance. Construct chemical equations to show this property.

Exercise

Lewis TheoryAn acid is an electron pair receptorA base is an electron pair donor

Gilbert Lewis (~1920)

Note that BF3 would not be an acid under the B-L Theory as there is no H+ to donate.

Other Lewis examples

A cobalt metal complex has 6 dative bonds formed by the donation of e- pairs from the ammonia molecules (Lewis bases)

Which of these substances are Lewis bases? Lewis acids?

Exercise

Water reacts with carbon dioxide to form carbonic acid. Write out the structures and show how the electrons are transferred, thereby identifying the Lewis acid and Lewis base.

Acid 1 + Base 2 Base 1 + Acid 2

What does this mean?An acid reacts and forms a conjugate base

which can also accept a protonA base reacts and forms a conjugate acid which

can donate a protonConjugate pairs differ only by one H+

Bronsted-Lowry (conjugates)

Conjugate pair 1 Conjugate

pair 2

Example:HNO3 + H2O <===> H3O+ + NO3¯

Identify the conjugate pairs in this reaction

Bronsted-Lowry (conjugates)

Here, nitric acid and the nitrate ion are conjugates and water and the hydronium ion are also conjugates

Strong vs. WeakStrong acidscompletely dissociate in water

HA + H2O H3O+ + A-

In general, these reactions are reversible, but for a strong acid the equilibrium is far right. So,

HA + H2O H3O+ + A-

Examples: HCl, HNO3, H2SO4

Weak acidsPartially dissociate in water

HA + H2O H3O+ + A- There is an equilibrium established with weak acids, which means that there is less H3O+ ions in solution.

CH3COOH + H2O H3O+ + CH3COO-

Examples: CH3COOH, H2CO3

What do you think is the difference in electrical conductivity?

Strong vs. WeakStrong basescompletely dissociate in water

NaOH Na+ + OH-

Again, as with acids, the equilibrium is far right. So,

1 mol NaOH 1 mol OH-

Examples: group I hydroxides, Ba(OH)2

Weak basesPartially dissociate in water

NH3 + H2O NH4+ + OH-

Only about 1% of ammonia dissociates into hydroxide ions.

1 mol NH3 << 1mol OH-

Examples: NH3, other aminesAgain, the electrical conductivity is greater for strong bases

Predicting EquilibriumThe direction of acid-base equilibria is

away from the stronger acid base side and towards the weaker acid base sideThe stronger the acid, the weaker its

conjugate baseThe stronger the base, the weaker its

conjugate acid

Bronsted-Lowry (equilibrium)

Reactions that proceed to a large extent:A strong acid will force the equilibrium in the

opposite direction (in this case, forward or right)

HCl + H2O <===> H3O+ + Cl¯ Reactions that proceed to a small extent:

If the weaker of the two acids and the weaker of the two bases are reactants (appear on the left side of the equation), the reaction is said to proceed to only a small extent:

NH3 + H2O <===> NH4+ + OH¯

Identify the conjugate acid base pairs in each reaction.

Bronsted-Lowry (equilibrium)