Models and conjugate pairs
Vinegar on your chips at the chipper is an acid you can eat. The acid in a car battery will burn your skin. Both are "acids", so calling something acidic is not enough: we need a model that explains what happens at particle level, and that model must separate how strong an acid is from how concentrated it is.
From Arrhenius to Brønsted–Lowry
Hold an open bottle of concentrated ammonia near an open bottle of concentrated hydrochloric acid and a dense white smoke of solid ammonium chloride forms in the air. There is no liquid water and no solution to release ions. How can we call this a neutralisation?
The first modern theory was proposed by Svante Arrhenius in 1884. It defines acids and bases by what they release when dissolved in water.
This works well for "classic" acids and bases, but it leaves out obvious cases. Ammonia, $\mathrm{NH_3}$, has no $\mathrm{OH^-}$ in its formula, yet it is a well-known base. In 1923 Johannes Brønsted and Thomas Lowry independently proposed a wider model based on the transfer of a proton.
This is why Brønsted–Lowry is the more extensive model: it covers everything Arrhenius does, and reactions in other solvents or in the gas phase as well.
Everyday acids and bases
Acids and bases are all around you. Vinegar (ethanoic acid) and citrus juices are acids. Antacid tablets and toothpaste contain bases. Indicators such as litmus change colour with the type of solution. Neutralisation is used in practice: farmers spread agricultural lime on acidic fields, and an antacid neutralises excess stomach acid in acid indigestion.
Acids also have characteristic reactions. Each of these gives a salt:
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Acid + base → salt + water: $\mathrm{HCl + NaOH \rightarrow NaCl + H_2O}$
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Acid + metal → salt + hydrogen: $\mathrm{Mg + 2HCl \rightarrow MgCl_2 + H_2}$
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Acid + carbonate → salt + water + carbon dioxide: $\mathrm{CaCO_3 + 2HCl \rightarrow CaCl_2 + H_2O + CO_2}$
The last one is why soda bread rises: the acid in buttermilk reacts with the bicarbonate of soda and releases $\mathrm{CO_2}$.
Understanding of the behaviour of acids and bases developed decisively in the late 19th and early 20th centuries. Consider the substances $\mathrm{HCl}$ and…
Understanding of the behaviour of acids and bases developed decisively in the late 19th and early 20th centuries. Consider the substances $\mathrm{HCl}$ and $\mathrm{NH}_3$ at $25\,^{\circ}\mathrm{C}$.
- a) Justify why hydrogen chloride behaves as an acid according to both Arrhenius's theory and the Brønsted-Lowry theory, writing the corresponding chemical equation in aqueous solution.
- b) Explain, with reasoning, why a strict historical interpretation based only on the dissociation of hydroxides does not directly describe ammonia. Write its reaction with water and explain how proton transfer produces hydroxide ions, consistent with the extended aqueous definition of a base.
The concept of neutralisation has complementary interpretations depending on the theoretical framework used. a) Describe the neutralisation reaction in aqueous solution…
The concept of neutralisation has complementary interpretations depending on the theoretical framework used.
- a) Describe the neutralisation reaction in aqueous solution between solutions of sodium hydroxide, $\mathrm{NaOH}$, and nitric acid, $\mathrm{HNO}_3$, writing the net ionic equation and justifying the process according to Arrhenius's definition.
- b) Now consider the reaction between the gases ammonia, $\mathrm{NH}_3(\mathrm{g})$, and hydrogen chloride, $\mathrm{HCl}(\mathrm{g})$, to form the white solid ammonium chloride, $\mathrm{NH}_4\mathrm{Cl}(\mathrm{s})$, in the complete absence of solvent. Explain whether this process is a neutralisation according to Arrhenius and according to Brønsted-Lowry, indicating the acid, the base and the proton transferred.
Conjugate acid–base pairs
Amphiprotic species and the self-ionisation of water
Add bicarbonate of soda to vinegar and it fizzes, because it neutralises the acid. Mix it with a strong base and it reacts again, this time acting as an acid. How can one substance behave as a base towards some reagents and as an acid towards others?
Water itself is amphiprotic, and this has a major consequence: two water molecules can react with each other.
In the self-ionisation, one water molecule acts as an acid (donating $\mathrm{H^+}$) and the other as a base (accepting it). The two conjugate pairs are $\mathrm{H_3O^+/H_2O}$ and $\mathrm{H_2O/OH^-}$.
Acid strength and concentration: two different things
For a weak acid, $K_a = \dfrac{[\mathrm{H_3O^+}][\mathrm{A^-}]}{[\mathrm{HA}]}$. The larger $K_a$ is, the greater the degree of dissociation. For a weak base, $K_b = \dfrac{[\mathrm{BH^+}][\mathrm{OH^-}]}{[\mathrm{B}]}$.
Two independent aqueous solutions of two monoprotic acids, $\mathrm{HA}$ and $\mathrm{HB}$, are prepared at the same initial analytical concentration…
Two independent aqueous solutions of two monoprotic acids, $\mathrm{HA}$ and $\mathrm{HB}$, are prepared at the same initial analytical concentration $c_0 = 0.10\,\mathrm{mol\times L^{-1}}$ and at $25\,^{\circ}\mathrm{C}$. The diagram shows a microscopic sample of identical volume for both containers, in which a total equivalent of $6$ acid molecules was initially introduced into each solution. Observe the distribution of particles at equilibrium illustrated below:
- a) From the count of species present at equilibrium (undissociated molecules and ions produced), determine the percentage or degree of dissociation of each acid.
- b) Deduce, with reasoning, which of the two acids is stronger, justifying whether two solutions with exactly the same initial molar concentration of solute necessarily have the same acidity. Use the ideal model and neglect the autoprotolysis of water where that approximation is valid.
Back to the start: Vinegar is safe to eat because it contains a weak acid ($\mathrm{CH_3COOH}$) that is only slightly dissociated and at low concentration. Battery acid burns the skin because it combines a strong acid ($\mathrm{H_2SO_4}$ dissociates fully in its first step) with a very high concentration.