Investigate the properties of acids and bases and the historical development of the Arrhenius model of acids and bases

What this dot point is asking

NESA wants you to describe the physical and chemical properties shared by acids and bases, recall the Arrhenius model and the reasoning behind it, and explain why later models (Bronsted-Lowry, Lewis) had to extend it. This is the entry point to Module 6 and the foundation for every later calculation, including reactions of acids, strong vs weak ionisation, and Bronsted-Lowry conjugate pairs.

The answer

Observed properties of acids

• Taste sour (citric acid in lemons, ethanoic acid in vinegar). Never taste laboratory chemicals.
• Turn blue litmus red.
• React with active metals (Mg, Zn, Fe) to produce hydrogen gas.
• React with metal carbonates and hydrogencarbonates to produce $CO_2$.
• React with bases to form a salt and water (neutralisation).
• Aqueous solutions conduct electricity (they are electrolytes).
• Have pH less than 7 at 25 degrees C.

Observed properties of bases

• Taste bitter and feel soapy or slippery (do not test by taste or touch).
• Turn red litmus blue.
• React with acids to form a salt and water.
• React with ammonium salts to release ammonia gas.
• Aqueous solutions conduct electricity.
• Have pH greater than 7 at 25 degrees C.

Indicators

An indicator is a weak acid or weak base whose protonated and deprotonated forms have different colours. The colour change occurs across a narrow pH range, usually about 2 pH units wide.

Universal indicator is a mixture of indicators that gives a continuous colour scale from red (very acidic) to violet (very basic).

The Arrhenius model

Svante Arrhenius (1887) proposed that acids and bases are substances that ionise in water.

• Arrhenius acid: a substance that releases $H^+$ in aqueous solution.
• Arrhenius base: a substance that releases $OH^-$ in aqueous solution.
• Neutralisation: $H^+_{(aq)} + OH^-_{(aq)} \rightarrow H_2O_{(l)}$.

The model elegantly explained why all aqueous acids share the same chemistry (because they all produce the same ion, $H^+$) and why all aqueous bases share the same chemistry (because they all produce $OH^-$).

Historical development

The Arrhenius model built on earlier ideas:

• Lavoisier (1780s). Believed all acids contained oxygen (the name "oxygen" means "acid former"). Disproved when Humphry Davy showed that hydrochloric acid contains no oxygen.
• Davy (1810). Proposed that hydrogen is the essential element in acids.
• Liebig (1838). Refined Davy's idea: an acid is a hydrogen-containing compound whose hydrogen can be replaced by a metal.
• Arrhenius (1887). Provided the ionic explanation: acids dissociate in water to give $H^+$.

This progression is a classic example of how a scientific model is refined as new evidence (electrolysis, conductivity, ionic theory) becomes available.

Limitations of the Arrhenius model

Arrhenius works well for simple aqueous acid-base reactions, but it cannot explain:

1. Basic species without hydroxide. Ammonia ($NH_3$) turns litmus blue and reacts with acids, yet contains no $OH^-$. Arrhenius cannot account for its basicity.
2. Non-aqueous acid-base chemistry. $HCl$ reacts with $NH_3$ in the gas phase to form $NH_4Cl$ with no water involved.
3. The hydrated proton. The bare $H^+$ ion never exists in water; it always attaches to a water molecule to form $H_3O^+$. Arrhenius treats $H^+$ as a free species.
4. Acid-base behaviour of salts. Solutions of $NH_4Cl$ are slightly acidic and solutions of $CH_3COONa$ are slightly basic, yet Arrhenius offers no mechanism for these effects.

These limitations motivated the Bronsted-Lowry model (1923), which defines acids as proton donors and bases as proton acceptors.

Worked example

Identify each substance as an Arrhenius acid, an Arrhenius base, neither, or both, and write the relevant ionisation equation if applicable.

(a) $HNO_3$. Acid. HNO_{3(aq)} \rightarrow H^+_{(aq)} + NO_3^-_{(aq)}.

(b) $KOH$. Base. $KOH_{(aq)} \rightarrow K^+_{(aq)} + OH^-_{(aq)}$.

(c) $NH_3$. Not an Arrhenius base (no $OH^-$ in the molecule), but a Bronsted-Lowry base. This is exactly the case that motivated the move away from Arrhenius.

(d) $NaCl$. Neither. It dissociates in water but produces neither $H^+$ nor $OH^-$.

(e) $H_2SO_4$. Acid, diprotic: H_2SO_{4(aq)} \rightarrow 2H^+_{(aq)} + SO_4^{2-}_{(aq)}.

Common traps

Calling ammonia an Arrhenius base. It is not. Ammonia is a Bronsted-Lowry base because it accepts a proton from water: $NH_3 + H_2O \rightleftharpoons NH_4^+ + OH^-$. The $OH^-$ comes from the water, not from $NH_3$ itself.

Forgetting state symbols. Acid-base equations are expected to carry $(aq)$ or $(l)$ on every species.

Treating the model as wrong, rather than limited. Arrhenius is correct for aqueous, ionic acids and bases. It is a special case of the broader Bronsted-Lowry model. Markers want "limited" not "incorrect".

Confusing strength and concentration. A property like "turns blue litmus red" depends on $[H^+]$, which is set by both the strength (degree of ionisation) and the concentration. Cover this carefully on the strong vs weak page.

Listing taste or feel as a test. Never test laboratory chemicals by taste or touch. State the property but not the procedure.

In one sentence

Acids turn blue litmus red, react with metals, carbonates, and bases, and (per Arrhenius) ionise in water to give $H^+$, while bases turn red litmus blue and ionise to give $OH^-$; the Arrhenius model works for aqueous solutions but cannot explain ammonia, non-aqueous chemistry, or the hydrated proton, which is why Bronsted-Lowry replaced it.