Acids and bases: three definitions

Acids and bases around us

Lemon juice and vinegar taste sour. Soap feels slippery and tastes bitter. Chemists noticed long ago that acids turn blue litmus red, and bases turn red litmus blue. An acid and a base can cancel each other in a neutralisation, giving a salt and water.

Why do they behave like this? Three ideas grew, one after another. Each one includes the one before it.

The Arrhenius idea

In 1884, Svante Arrhenius said:

• An acid gives hydrogen ions, $\text{H}^+$, in water.
• A base gives hydroxide ions, $\text{OH}^-$, in water.

A bare proton cannot exist alone in water. It sticks to a water molecule at once, making the hydronium ion, $\text{H}_3\text{O}^+$. So $\text{H}^+\text{(aq)}$ and $\text{H}_3\text{O}^+$ mean the same thing in this course.

The Arrhenius idea has limits. It works only in water. It cannot explain why ammonia, $\text{NH}_3$, is a base, since ammonia has no $\text{OH}$ group to give.

The Bronsted-Lowry idea

In 1923, Johannes Bronsted and Thomas Lowry looked at what moves between the particles: a proton.

• An acid is a proton donor.
• A base is a proton acceptor.

Ammonia in water fits well:

Water gives a proton, so here water is the acid. Ammonia takes the proton, so it is the base.

Conjugate pairs

Look at the reverse reaction. $\text{NH}_4^+$ can give its proton back, so it is an acid. $\text{OH}^-$ can take it, so it is a base.

An acid that has lost a proton becomes its conjugate base. A base that has gained a proton becomes its conjugate acid. The two members of a pair differ by just one proton.

• $\text{HCl}$ and $\text{Cl}^-$ form a pair. So do $\text{H}_3\text{O}^+$ and $\text{H}_2\text{O}$.
• $\text{NH}_4^+$ and $\text{NH}_3$ form a pair. So do $\text{H}_2\text{O}$ and $\text{OH}^-$.

Water appears as a base with HCl and as an acid with ammonia. A substance that can act both ways is called amphiprotic. The hydrogen carbonate ion, $\text{HCO}_3^-$, is another.

Strong and weak

Strong electrolytes ionise completely in water. Strong acids include $\text{HCl}$, $\text{HBr}$, $\text{HI}$, $\text{HNO}_3$ and $\text{HClO}_4$. Strong bases include $\text{NaOH}$, $\text{KOH}$ and $\text{Ba(OH)}_2$.

Weak electrolytes ionise only a little. In a solution of acetic acid, most molecules stay whole. Only about one in a hundred gives up its proton. Ammonia is a weak base in the same way.

There is a neat rule. A strong acid has a very weak conjugate base. HCl holds its proton so loosely that $\text{Cl}^-$ hardly takes one back. A weak acid holds its proton tightly, so its conjugate base is fairly strong.

Practise with new numbers

The Lewis idea

In 1923, Gilbert Lewis looked at electrons instead of protons.

• A Lewis acid accepts a pair of electrons.
• A Lewis base donates a pair of electrons.

Boron trifluoride has only six electrons round boron. It accepts the lone pair of ammonia:

So $\text{BF}_3$ is a Lewis acid, though it has no proton to give. $\text{AlCl}_3$ and metal ions such as $\text{Co}^{3+}$ are Lewis acids too. Every Bronsted base has a lone pair to take a proton, so it is also a Lewis base.

What makes an acid strong

For the acids HX, two things matter.

• Down a group, bond strength wins. The H-X bond gets longer and weaker from HF to HI. So the acid strength rises: $\text{HF} < \text{HCl} < \text{HBr} < \text{HI}$.
• Across a period, polarity wins. From carbon to fluorine, the atom pulls harder on the shared electrons. The H-A bond grows more polar and gives up $\text{H}^+$ more easily. So $\text{CH}_4 < \text{NH}_3 < \text{H}_2\text{O} < \text{HF}$.

Try it in Python

The program compares a strong and a weak acid of the same concentration. It finds what is left in each beaker.

Common mistakes

• Mixing up the pair. A conjugate pair differs by exactly one proton. $\text{H}_2\text{SO}_4$ and $\text{SO}_4^{2-}$ are not a pair.
• Thinking water is always neutral in its role. With an acid it is a base. With ammonia it is an acid.
• Calling a weak acid dilute. Weak means little ionisation. Dilute means little acid in the water. A concentrated weak acid is possible.
• Forgetting Lewis acids without protons. $\text{BF}_3$ and $\text{AlCl}_3$ have no proton, yet they are acids by Lewis.

Remember these

• Arrhenius: acids give $\text{H}^+$, bases give $\text{OH}^-$, in water only.
• Bronsted-Lowry: acids give protons, bases take them.
• A conjugate pair differs by one proton. Strong acid, very weak conjugate base.
• Lewis: acids accept an electron pair, bases donate one.
• HX strength rises down a group. Across a period it rises with bond polarity.

Check yourself

Try each one on your own first. If you are stuck, open a hint. Then check with Show the answer.

1. What is the conjugate base of $\text{H}_2\text{O}$, and its conjugate acid?

2. Why is ammonia a base by Bronsted and Lowry, but not by Arrhenius?

3. Is $\text{BF}_3$ a Lewis acid or a Lewis base?

4. Which is the stronger acid, HCl or HI? Why?

5. Which species is amphiprotic: $\text{HCO}_3^-$ or $\text{Cl}^-$?