Definitions of acids and bases

According to the Arrhenius theory, an acid is a substance that produces hydrogen ions when dissolved in water (in reality these attach to a water molecule as the hydronium ion, $\text{H}_3\text{O}^+$), and a base is a substance that produces hydroxide ions in water. A base that also dissolves in water is given the special name alkali — so every alkali is a base, but a base is only an alkali if it is soluble.

where: the (aq) tag marks an ion surrounded by water molecules, and the arrow shows the compound splitting apart, not merely dissolving as whole molecules. Two counting terms describe how many ions a formula unit can give. Basicity of an acid is the number of replaceable hydrogen atoms one molecule can furnish: HCl is monobasic, $\text{H}_2\text{SO}_4$ is dibasic, $\text{H}_3\text{PO}_4$ is tribasic. Acidity of a base is the number of hydroxide ions (or H⁺-accepting sites) one formula unit can furnish: NaOH is monoacidic, $\text{Ca(OH)}_2$ is diacidic. The names look swapped at first glance but are the standard ICSE terms — basicity always refers to acids, acidity always refers to bases.

Strong and weak describe the degree of ionisation in water, a separate idea from how much acid or base is dissolved. A strong acid or base breaks up almost completely into ions (HCl, $\text{H}_2\text{SO}_4$, $\text{HNO}_3$; NaOH, KOH). A weak acid or base ionises only partly, so most molecules stay un-ionised (acetic acid $\text{CH}_3\text{COOH}$, carbonic acid $\text{H}_2\text{CO}_3$; ammonium hydroxide $\text{NH}_4\text{OH}$). No numerical ionisation constant is needed at this level — only the qualitative "almost all" versus "only a little" picture.

Step 1. Recognise that "acid" and "base" in this syllabus are not defined by taste or feel, but by a single behavioural test: which ion does the substance hand over to water once dissolved. Every property that follows — sour taste and red litmus for acids, bitter taste and blue litmus for alkalis, the fizzing of an acid with a reactive metal, the neutralisation of one by the other — is a consequence of this ion-producing definition, not the definition itself.

Step 2. Separate "produces $\text{OH}^-$" (makes something a base) from "dissolves in water" (makes a base an alkali). These are two independent questions. Copper(II) hydroxide, $\text{Cu(OH)}_2$, would still react with an acid like a base does, but because it barely dissolves in water it never delivers free $\text{OH}^-$ ions into solution, so it is correctly called an insoluble base rather than an alkali.

Step 3. Separate "how completely it ionises" (strong versus weak) from "how much is dissolved" (concentration). A beaker of dilute HCl still has almost every one of its few molecules broken into ions, while a beaker of concentrated acetic acid still has most of its many molecules sitting un-ionised. Strength is a property of the substance's chemical behaviour in water; concentration is a property of how the solution was prepared.

Step 4. Use basicity and acidity purely as counting tools, tied to the formula, not to strength. A weak acid can still be polybasic (carbonic acid $\text{H}_2\text{CO}_3$ is dibasic despite being weak), and a strong base can be monoacidic (NaOH). Keeping "how many ions per formula unit" separate from "what fraction of molecules ionise" avoids the most common ICSE mix-up in this chapter.

Step 1. Classify hydrochloric acid and acetic acid as strong or weak and explain the difference using ionisation. Both are acids because both release $\text{H}^+$ in water, so the definition test is satisfied by each; the question is how completely each one does so.

Step 2. HCl is a strong acid — essentially all dissolved HCl molecules break apart: $\text{HCl} \rightarrow \text{H}^+ + \text{Cl}^-$. A solution of HCl therefore contains a very high concentration of free $\text{H}^+$ ions relative to the amount of acid that was dissolved.

Step 3. Acetic acid is a weak acid — only a small fraction of its molecules ionise: $\text{CH}_3\text{COOH} \rightleftharpoons \text{CH}_3\text{COO}^- + \text{H}^+$. Most acetic acid stays as un-ionised $\text{CH}_3\text{COOH}$ molecules, so the free $\text{H}^+$ concentration is much lower even for the same starting amount of acid.

Step 4. Consequence: equal molar amounts of HCl and acetic acid do not behave equally "acidic." The HCl solution conducts electricity better, reacts faster with a reactive metal to release $\text{H}_2$ gas, and turns litmus red more sharply, because it has a much higher free $\text{H}^+$ concentration — even though both acids are monobasic by formula.

Competitive stretch

Is ammonia solution ($\text{NH}_4\text{OH}$) an alkali, and what is the acidity of $\text{Ca(OH)}_2$? $\text{NH}_4\text{OH} \rightleftharpoons \text{NH}_4^+ + \text{OH}^-$ produces $\text{OH}^-$, so it is a base, and since it is soluble it is also an alkali — though a weak one, since only a small fraction ionises. $\text{Ca(OH)}_2 \rightarrow \text{Ca}^{2+} + 2\text{OH}^-$ furnishes two hydroxide ions per formula unit, so its acidity is 2 — it is a diacidic base.

• Confusing "base" with "alkali" — every alkali is a base, but a base is only an alkali if it dissolves in water; insoluble bases like $\text{Fe(OH)}_3$ or $\text{Cu(OH)}_2$ are never alkalis.
• Thinking "strong acid" means "concentrated acid" — strength is the degree of ionisation, concentration is how much is dissolved per unit volume; a dilute strong acid still ionises almost completely.
• Swapping basicity (a property of acids, counting replaceable H) with acidity (a property of bases, counting available $\text{OH}^-$) — the names sound reversed but are standard terminology.
• Assuming all acids contain oxygen or all bases are metal hydroxides — ammonia-based and organic acids do not fit these shortcuts.

Ask the tutor: Want to practise classifying five more everyday substances as strong/weak acid or base and stating their basicity or acidity?