Strength of Acids and Bases (VCE SSCE Chemistry): Revision Notes

Strength of Acids and Bases

Introduction to acid and base strength

When equal concentrations of different acids react with zinc metal, they don't all react at the same rate. Some acids produce bubbles more vigorously than others, indicating that they are stronger acids. This observation tells us that acids vary in their strength, even when they are present in the same concentration.

According to the Brønsted-Lowry theory, acids are substances that donate protons (H⁺ ions) and bases are substances that accept protons. The strength of an acid refers to how easily it can donate a proton to a base. Similarly, the strength of a base measures how readily it can accept a proton from an acid. When we compare acids and bases in aqueous solutions, we look at how they interact with water molecules to determine their strength.

Different acids with the same concentration don't all produce the same concentration of hydronium ions (H₃O⁺) in solution. This means they will have different pH values, even though the original acid concentration was identical. The key difference lies in the extent to which each acid donates its protons.

Strong acids

Strong acids are excellent proton donors. When a strong acid dissolves in water, it breaks apart completely into ions - virtually no intact acid molecules remain in the solution. This complete breakdown is called complete ionisation.

The three most common strong acids you need to know are:

• Hydrochloric acid (HCl)
• Nitric acid (HNO₃)
• Sulfuric acid (H₂SO₄)

When hydrogen chloride gas bubbles through water, every HCl molecule donates its proton to a water molecule. Similarly, pure nitric acid and sulfuric acid are covalent molecular compounds that completely ionise when they dissolve in water. We can represent these ionisation reactions with chemical equations:

$\text{HCl(g)} + \text{H}_2\text{O(l)} \rightarrow \text{H}_3\text{O}^+\text{(aq)} + \text{Cl}^-\text{(aq)}$

$\text{HNO}_3\text{(l)} + \text{H}_2\text{O(l)} \rightarrow \text{H}_3\text{O}^+\text{(aq)} + \text{NO}_3^-\text{(aq)}$

$\text{H}_2\text{SO}_4\text{(l)} + \text{H}_2\text{O(l)} \rightarrow \text{H}_3\text{O}^+\text{(aq)} + \text{HSO}_4^-\text{(aq)}$

Weak acids

Unlike strong acids, weak acids only partially break apart in water. Most of the acid molecules remain intact, and only a small proportion donate their protons to water molecules at any given time. This incomplete breakdown is called partial ionisation.

Ethanoic acid (CH₃COOH), found in vinegar, is a common weak acid. When pure ethanoic acid dissolves in water, it exists mostly as CH₃COOH molecules. Even in a concentrated 1.0 M solution of ethanoic acid, only about 0.004 M of the acid molecules have donated their protons to form hydronium ions (H₃O⁺) and ethanoate ions (CH₃COO⁻). This means that at 25°C, less than 1% of the ethanoic acid molecules are ionised.

The partial ionisation of ethanoic acid is represented using a double arrow (⇌):

$\text{CH}_3\text{COOH(aq)} + \text{H}_2\text{O(l)} \rightleftharpoons \text{CH}_3\text{COO}^-\text{(aq)} + \text{H}_3\text{O}^+\text{(aq)}$

The double arrow indicates that this is a reversible reaction. As ethanoic acid molecules donate protons to form products, some of the products are simultaneously recombining to reform ethanoic acid molecules. An equilibrium is established where both forward and reverse reactions occur, but most molecules remain in the un-ionised form.

Strong bases

Strong bases are excellent proton acceptors - they readily take protons from acids. The oxide ion (O²⁻) is an example of a strong base. When sodium oxide (Na₂O), an ionic compound, dissolves in water, it releases oxide ions that react completely with water molecules:

$\text{O}^{2-}\text{(aq)} + \text{H}_2\text{O(l)} \rightarrow \text{OH}^-\text{(aq)} + \text{OH}^-\text{(aq)}$

In this reaction, the oxide ion acts as a base by accepting a proton from water (which acts as an acid). The reaction goes to completion, as shown by the single arrow. Each oxide ion produces two hydroxide ions.

Weak bases

Weak bases only partially accept protons from acids. At any given moment, only a small proportion of weak base molecules have accepted a proton.

Ammonia (NH₃) is the most common weak base. When ammonia dissolves in water, it accepts a proton from water molecules to form ammonium ions and hydroxide ions:

$\text{NH}_3\text{(aq)} + \text{H}_2\text{O(l)} \rightleftharpoons \text{NH}_4^+\text{(aq)} + \text{OH}^-\text{(aq)}$

In this reaction, ammonia acts as a base (it gains a proton), while water acts as an acid (it donates a proton). The double arrow shows that this is a reversible reaction. Even in a concentrated 1.0 M ammonia solution, most of the dissolved substance exists as NH₃ molecules. Only a small number of ammonium ions (NH₄⁺) and hydroxide ions (OH⁻) are present because the ionisation is incomplete.

Relative strength of conjugate acid-base pairs

Conjugate acids and bases are pairs of species that differ by exactly one proton (H⁺). When an acid donates a proton, what remains is its conjugate base. When a base accepts a proton, it becomes its conjugate acid.

There is an important inverse relationship between the strength of an acid and its conjugate base: the stronger an acid is, the weaker its conjugate base will be. Similarly, the stronger a base is, the weaker its conjugate acid will be.

We can arrange common acids and bases in order of their relative strength. At the top of the list are the strongest acids, which have the weakest conjugate bases. At the bottom are the weakest acids (strongest bases). Here are some key points from the strength comparison:

• Strong acids (HCl, H₂SO₄, HNO₃) have very weak conjugate bases
• H₃O⁺ is in the middle - moderately strong acid
• Weak acids (HF, CH₃COOH, H₂CO₃) have moderately strong conjugate bases
• Very weak acids (H₂O, OH⁻) have very strong conjugate bases

This relationship helps predict the direction of acid-base reactions. Reactions tend to proceed from the stronger acid and base towards the weaker acid and base.

Strength versus concentration

It is crucial not to confuse the terms 'strong' and 'weak' with 'concentrated' and 'dilute'. These terms describe different properties of acid and base solutions.

This means you can have four different types of acid solutions:

1. Strong and concentrated: A large amount of strong acid dissolved in water, all of it ionised (e.g., concentrated HCl)
2. Strong and dilute: A small amount of strong acid dissolved in water, all of it ionised (e.g., dilute HCl)
3. Weak and concentrated: A large amount of weak acid dissolved in water, only partially ionised (e.g., concentrated CH₃COOH)
4. Weak and dilute: A small amount of weak acid dissolved in water, only partially ionised (e.g., dilute CH₃COOH)

To give solutions a precise, quantitative description, we state their concentration in units such as mol L⁻¹ or g L⁻¹, rather than using qualitative terms like 'strong', 'weak', 'concentrated', or 'dilute'.

Super acids

Fluorosulfuric acid (HSO₃F) is one of the strongest acids known to chemistry. Its molecular structure is similar to sulfuric acid, but with one key difference - a highly electronegative fluorine atom replaces one of the hydroxyl groups.

This fluorine atom makes the oxygen-hydrogen bond much more polarised than in sulfuric acid. As a result, the acidic proton transfers to a base far more easily, making fluorosulfuric acid an extremely strong acid.

Common strong and weak acids and bases

Here is a summary table of common acids and bases you should know:

Strong acids	Weak acids	Strong bases	Weak bases
Hydrochloric acid, HCl	Ethanoic acid, CH₃COOH	Sodium hydroxide, NaOH	Ammonia, NH₃
Sulfuric acid, H₂SO₄	Carbonic acid, H₂CO₃	Potassium hydroxide, KOH
Nitric acid, HNO₃	Phosphoric acid, H₃PO₄	Calcium hydroxide, Ca(OH)₂