Chemistry 112

Lewis Acids and Bases

Acid: a substance that can accept an electron pair.

Base: a substance that can donate an electron pair.

Acid/base reaction: donation of an electron pair to create a new covalent bond.

This definition has no solvent or phase restrictions and is not dependent on any single chemical species. This is the most general acid/base definition we've looked at.

Identifying Lewis acid/base reactions usually means working with Lewis dot structures (not surprisingly, it's the same Lewis) so we can find electron pairs and new covalent bonds.

We will use Lewis acid/base concepts later in the course when we study complex ions:



4 :NH3(aq)


Lewis Acid

Lewis Base

Complex Ion

More Brønsted–Lowry acids and bases.

Consider the two examples:

When HCl is dissolved in water, essentially 100% of the H+ ions are transferred to the water to form hydronium ions:

HCl(aq) + H2O(l) H3O+(aq) + Cl(aq)

In contrast, acetic acid, HC2H3O2, only partially transfers its hydrogen ion to water, so the reaction must be considered an equilibrium:

HC2H3O2(aq) + H2O(l) H3O+(aq) + C2H3O2(aq)

We distinguish between these two kinds of acids as strong or weak:

Strong acids completely transfer their hydrogen ion to water in aqueous solution. All reactions involving strong acids in aqueous solution go to completion.

Weak acids only partially transfer their hydrogen ion to water in aqueous solution. These reactions are equilbria.

Similar observations and statements can be made for bases.

Because bases are often ionic salts, writing their Bronsted–Lowry acid base reactions can be a little tricky:

NaOH(aq)+HOH(l) HOH(l) + Na+(aq) + OH(aq)

The color-coding shows the hydrogen ion transfer, but since water shows up on both sides, it cancels giving the net equation:

NaOH(aq) Na+(aq) + OH(aq)

Thus, when we write Brønsted–Lowry for ionic salts, even though the H+ transfer may not be explicitly written, we understand that it does take place.

For bases that are not ionic salts, this problem does not exist:

C6H5NH2(aq) + H2O(l) C6H5NH3+(aq) + OH(aq)


Fortunately, there are a very limited number of strong acids and bases; we must memorize these:

Strong Acids

HCl    Hydrochloric acid

HBr    Hydrobromic acid

HI    Hydroiodic acid

HNO3    Nitric acid

HClO4    Perchloric acid

H2SO4    Sulfuric acid (1st hydrogen ion only)

Strong Bases

Group 1 Hydroxides

LiOH    Lithium hydroxide

NaOH    Sodium hydroxide

KOH    Potassium hydroxide

RbOH    Rubidium hydroxide

CsOH    Cesium hydroxide

Group 2 Hydroxides (except Be(OH)2)

Mg(OH)2    Magnesium hydroxide

Ca(OH)2    Calcium hydroxide

Sr(OH)2    Strontium hydroxide

Ba(OH)2    Barium hydroxide

The group 2 hydroxides are not very water soluble so are not often used as bases.

In any acid–base reaction we write, we must correctly distinguish if the reaction goes to completion or if it is an equilibrium.

Any reaction that uses either a strong acid or a strong base goes to completion.

All other acid–base reactions are equilibria.


Complete and balance:

HBr(aq) + NH3(aq) ?

H3PO4(aq) + F(aq) ?

A problem:

For the reverse reaction, HF is the H+ donor and H2PO4 is the H+ acceptor.

So how do we label the acid and base in an equilibrium situation?

We define two new terms:

Conjugate acid: the H+ donor written on the products side of the reaction.

Conjugate base: the H+ acceptor written on the products side of the reaction.

These terms are used even if the reaction is not an equilibrium.

Thus, in the above example:

H3PO4 is the acid.

F is the base.

HF is the conjugate acid.

H2PO4 is the conjugate base.

HF/F are called a conjugate acid/base pair.

H3PO4/ H2PO4 are also a conjugate acid/base pair.

Conjugate acid/base pairs differ only by an H+.