Ionisation and dissociation

Drop a pinch of salt into water and a bulb lights up when wired across the solution; drop in the same mass of sugar and the bulb stays dark. Both dissolve completely, so why does only one solution carry a current? The answer sits at the particle level: some substances break apart into charged ions when they meet water, and only free-moving ions can carry an electric charge through a liquid. This splitting-into-ions is called ionisation (or dissociation), the mechanism behind the Arrhenius picture of acids and bases that anchors Unit 3 of ICSE Class 10 Chemistry. This article stays qualitative, as the syllabus demands — no $K_a$ , $K_b$ , or Ostwald's dilution law — but goes one step further than a bare definition, showing how the degree of ionisation lets you compare two electrolytes numerically and predict which one lights the bulb brighter.

The core insight

Electrolytes split into free-moving cations and anions when dissolved in water — completely for strong acids and bases, only partially for weak ones — and it is the resulting concentration of free ions, not the electrolyte's identity or concentration alone, that fixes a solution's conductivity, reactivity, and pH.

The core idea

Electrolytes, ions, and the two words for the same idea. Acids, bases, and salts are collectively called electrolytes because their solutions conduct electricity. When an electrolyte dissolves in water, it breaks up into charged particles called ions: positively charged cations and negatively charged anions. For a substance that is already an ionic lattice in the solid state (such as NaCl or NaOH), this break-up is usually called dissociation — the ions were always there, and water simply frees them from the lattice. For a substance that exists as neutral molecules before dissolving (such as HCl gas), and only forms ions on contact with water, the process is usually called ionisation — new ions are created that did not exist before. At Class 10 level the two words are used interchangeably for acids and bases dissolving in water.

\text{HCl (molecule)} \xrightarrow{\text{water}} \text{H}^+(aq) + \text{Cl}^-(aq)

\text{NaOH (ionic solid)} \xrightarrow{\text{water}} \text{Na}^+(aq) + \text{OH}^-(aq)

Degree of ionisation: how much, not just whether. Not every dissolved electrolyte splits to the same extent. The degree of ionisation is the fraction of dissolved molecules that actually break into free ions:

A strong acid or base ionises almost completely — nearly every dissolved molecule exists as separate ions in solution.
A weak acid or base ionises only to a small extent — most of the dissolved substance remains as intact, un-ionised molecules, with only a small proportion present as free ions at any instant, in a dynamic equilibrium between molecules and ions.

This one idea — how much has actually split into ions, not how much was dissolved — is the thread that ties together every property below.

Figure

Two beakers of equal concentration compared, a strong acid beaker showing nearly all particles as separated ion pairs and a weak acid beaker showing mostly intact molecules with only a few separated ions, with a brighter bulb over the strong acid beaker

What the free-ion concentration then explains:

Electrical conductivity: a solution conducts only because free-moving ions carry charge through it. More free ions (strong electrolyte) means better conductivity than fewer free ions (weak electrolyte) at the same dissolved concentration; a non-electrolyte like sugar solution, which does not ionise at all, does not conduct.
Reactivity: reactions of acids depend on the concentration of free $\text{H}^+$ ions available, so a strong acid generally reacts faster with a reactive metal (releasing $\text{H}_2$ faster) than an equally concentrated weak acid.
pH placement: a strong acid is more acidic (lower pH) than a weak acid of the same concentration, because it supplies more free $\text{H}^+$ ions.

Water is not a passive bystander in any of this: ionisation happens specifically in water (or another ionising solvent). Dry, water-free HCl gas does not ionise and, notably, does not turn dry litmus paper red — the process needs water to proceed at all.

Derivation ADegree of ionisation and the concentration of free ions

The goal is to turn "strong acids ionise more than weak acids" into a number you can compute, by defining the degree of ionisation as a fraction and using it to find the concentration of free $\text{H}^+$ ions produced by a given dissolved concentration of a monoprotic acid — the calculation that explains why a strong acid's bulb glows brighter than a weak acid's at the same molar concentration.

Step 1. Define the degree of ionisation, $\alpha$ , as the fraction of dissolved molecules that have split into ions at equilibrium: $\alpha = \dfrac{\text{number of molecules ionised}}{\text{total number of molecules dissolved}}$ . As a fraction, $\alpha$ is dimensionless and lies between $0$ (no ionisation) and $1$ (complete ionisation).

Step 2. Connect $\alpha$ to concentration. If a monoprotic acid $\text{HA}$ is dissolved to total concentration $C$ (mol/dm³) and a fraction $\alpha$ ionises as $\text{HA} \rightarrow \text{H}^+ + \text{A}^-$ , the free $\text{H}^+$ concentration is that fraction of $C$ :

[\text{H}^+] = \alpha \, C.

where: $[\text{H}^+]$ is the free hydrogen-ion concentration (mol/dm³), $\alpha$ is the degree of ionisation (no units), and $C$ is the total dissolved concentration (mol/dm³). Each ionised molecule releases one $\text{H}^+$ and one $\text{A}^-$ , so $[\text{A}^-] = [\text{H}^+] = \alpha C$ too.

Step 3. Apply it to a strong acid. For HCl, $\alpha \approx 1$ , so $[\text{H}^+] \approx C$ — free-ion concentration nearly equals dissolved concentration.

Step 4. Apply it to a weak acid. For acetic acid, $\alpha$ is small, so $[\text{H}^+] = \alpha C \ll C$ — most of the acid stays un-ionised. Since conductivity rises with free-ion concentration, $[\text{H}^+] = \alpha C$ is exactly why a strong acid outperforms a weak acid of the same $C$ on the conductivity bulb.

Worked example

Question: $0.1\ \text{mol/dm}^3$ acetic acid ( $\text{CH}_3\text{COOH}$ ) has a degree of ionisation $\alpha = 1.3\% = 0.013$ at this concentration. An equal volume of $0.1\ \text{mol/dm}^3$ hydrochloric acid, for comparison, ionises essentially completely ( $\alpha \approx 1$ ). Find the concentration of free $\text{H}^+$ ions in each solution, and state how many times more free ions the HCl solution contains.

Step 1. Write down what is known for acetic acid: $C = 0.1\ \text{mol/dm}^3$ , $\alpha = 0.013$ .

Step 2. Apply $[\text{H}^+] = \alpha C$ for acetic acid: $[\text{H}^+] = 0.013 \times 0.1 = 1.3\times10^{-3}\ \text{mol/dm}^3$ . Only about $1.3$ molecules in every $100$ dissolved have actually split into ions at any moment.

Step 3. Apply the same relation for HCl, with $\alpha \approx 1$ : $[\text{H}^+] = 1 \times 0.1 = 0.1\ \text{mol/dm}^3$ — essentially the full dissolved concentration appears as free ions.

Step 4. Compare the two: $\dfrac{[\text{H}^+]_{\text{HCl}}}{[\text{H}^+]_{\text{acetic}}} = \dfrac{0.1}{1.3\times10^{-3}} \approx 77$ . The HCl solution carries roughly 77 times more free $\text{H}^+$ ions than the acetic acid solution at the same $0.1\ \text{mol/dm}^3$ concentration — which is exactly why HCl lights the conductivity bulb far more brightly, reacts with metals far faster, and gives a much lower pH, even though both solutions started out at the identical dissolved concentration.

Common traps

Believing "strong acid" means "highly concentrated acid" — strength is about $\alpha$ , the fraction that ionises, and is completely independent of $C$ , the amount dissolved; a dilute strong acid can still have $\alpha \approx 1$ .
Assuming all acids conduct electricity equally well simply because they are all "acids" — conductivity tracks $[\text{H}^+] = \alpha C$ , which differs sharply between strong and weak electrolytes at the same $C$ .
Thinking dry (water-free) HCl gas is acidic on its own — it does not turn dry litmus paper red, because ionisation, and therefore the release of free $\text{H}^+$ , requires water.
Treating "ionisation" and "dissociation" as always distinct processes — at this level both simply mean formation of free ions in solution and are used interchangeably for acids and bases.

Ask the tutor: want to work through the same alpha-times-C calculation for a strong or weak base, or see how degree of ionisation is actually measured in a lab experiment?

If degree of ionisation is a fraction, does a more concentrated weak acid ionise more?

Not quite as you might expect. Increasing the total $C$ of a weak acid usually makes $\alpha$ itself slightly smaller (more molecules compete for the same water). What actually rises is the free-ion concentration $[\text{H}^+] = \alpha C$ , since the drop in $\alpha$ is much smaller than the rise in $C$ . So a concentrated weak acid can conduct better than a dilute one of the same acid — but it will still conduct worse than an equally concentrated strong acid, whose $\alpha$ stays near $1$ .

Is ionisation only for acids and bases, or does it apply to salts too?

Salts undergo the same process. A salt like NaCl is already built from $\text{Na}^+$ and $\text{Cl}^-$ ions in its solid lattice, so on dissolving, water simply separates the existing ions apart — usually the case labelled dissociation. Most common salts dissociate almost completely, giving high $\alpha$ and good conductivity, like a strong acid or base. This article focuses on acids and bases because that is where ICSE expects the Arrhenius definitions, but the underlying idea is identical for salts.

Why does the equation use H+ specifically, and what about OH- for bases?

$[\text{H}^+] = \alpha C$ is written for an acid because acids release $\text{H}^+$ on ionisation. A base behaves identically with the opposite ion: for $\text{BOH} \rightarrow \text{B}^+ + \text{OH}^-$ , the same logic gives $[\text{OH}^-] = \alpha C$ . A strong base such as NaOH has $\alpha \approx 1$ , so nearly all its dissolved concentration appears as free $\text{OH}^-$ ; a weak base such as ammonia solution has a small $\alpha$ , so only a small slice is free at any instant — exactly mirroring the acid case.

Ionisation and dissociation

The core idea

Worked example

Common traps

Before this

Where this leads