M6-S1: Acid and Base Theories
Explore the changes in definitions and models of an acid and a base over time to explain the limitations of each model, including but not limited to:
- Arrhenius’ theory
- Brønsted–Lowry theory (ACSCH064, ACSCH067)
Definitions of Acids and Bases
Oxygen Theory of Acids (1776)
- French chemist Antoine Lavoisier showed that many non-metal compounds containing oxygen produced acidic solutions when dissolved in water. For example, oxides of carbon, sulfur and phosphorus.
$$CO_2(aq)+H_2O(l) \rightleftharpoons HCO_3(aq)$$
- Lavoisier hypothesised that the presence of oxygen atoms in these compounds gave them their acidic properties.
- Limitation: many experiments showed acidic properties in substances which do not contain oxygen e.g. HCl
Hydrogen Theory of Acids (1810)
- Davy electrolysed samples of hydrochloric acid and showed that it produced hydrogen gas and chlorine gas, but not oxygen gas disproves Lavoisier’s oxygen theory of acids.
- Later experiments by other chemists showed that other acids e.g. hydrocyanic acid (HCN), also contained no oxygen but did contain hydrogen. As a result, Davy proposed that the presence of hydrogen in acids gave them their acidic properties.
- Limitation: Davy’s theory did not explain why many compounds with hydrogen atoms were not acidic e.g. methane (CH4).
Extension of Davy’s Hydrogen Theory of Acids (1838)
- German chemist Justus von Liebig stated that acids were substances that contained replaceable hydrogens. He reasoned this by saying that when acids react with metals, the metals replace the hydrogen atoms in the acid to form a salt.
$$2Na(s)+HCl(g) \rightarrow 2NaCl(s)+H_2(s)$$
- Limitation: Von Liebig’s ideas failed to account for the production of gases such as nitrogen dioxide (rather than hydrogen gas) when nitric acid reacts with metals.
Arrhenius’ Hydrogen & Hydroxide Ions Theory (1887)
Swedish scientist Svante Arrhenius describes:
- Acids as substances that dissociate to produce hydrogen ion (more correctly hydronium ion) in water.
$$HCl(aq) \rightarrow H^+(aq) + Cl^-(aq)$$
- Bases as substances that dissociate to produce hydroxide ion in water.
$$NaOH(aq) \rightarrow Na^+(aq) + OH^-(aq)$$
- Arrhenius’ theory explains neutralisation between acids and bases. He proposed that when neutralisation occurs, the hydrogen ions (produced by acids), reacts with hydroxide ions (produced by bases), to form neutral water
$$HCl(aq) + NaOH(aq) \rightarrow NaCl(aq) + H_2O(l)$$
- Arrhenius could not explain why certain compounds do not contain hydroxide ions, despite displaying basic properties. For example, oxides of metals and carbonates:
$$2HCl(aq) + CaCO_3(s) \rightarrow CaCl_2(aq) + H_2O(l) + CO_2(g)$$
Calcium carbonate (CaCO3) is basic in nature but it does not contain OH–.
- Arrhenius could not explain why some neutrally-charged salts e.g. ZnCl2 are acidic in solution whilst others are basic e.g. Na2
- Arrhenius could not explain acid-base reactions that do not occur in aqueous solutions. For example, the reaction between gaseous ammonia and hydrogen chloride is an example of acid-base reaction.
Brønsted-Lowry Theory (1923)
- The Brønsted-Lowry definition describes acids and bases by their For example, the theory describes:
- Acids as protons donors, any compound that can donate protons regardless of its state of matter or the nature of the solvent.
- Bases as protons acceptors, any compound that can accept protons regardless of its state of matter or nature of the solvent.
- Brønsted-Lowry assigned a role to the solvent. The theory outlines the important role of water as an ionising solvent. Water can act has both an acid (proton donors) and a base (proton acceptors).
- Proton transfer between acids and/or bases can occur in non-aqueous solutions.
$$HCl(aq) + H_2O(l) \rightarrow H_3O^+(aq) + Cl^-(aq)$$
(Here HCl is a proton donor since it deprotonates, and water is a proton acceptor since it protonates)
$$NH_3(aq) + H_2O(l) \leftrightharpoons NH_4^+(aq) + OH^-(aq)$$
$$(Here NH_ is the proton acceptor since it protonates and water is a donor since it deprotonates)$$
- Acid-base pairs, conjugate acid and bases
Consequently, the Brønsted-Lowry theory introduces the concept of acid-base pairs. When a Brønsted-Lowry acid donates its proton to the base, the resultant anion of the acid is deprotonated. This species can act as a Brønsted-Lowry base (conjugate base) by accepting a proton to reform the original acid molecule.
$$HF(aq) + H_2O(l) \leftrightharpoons H_3O^+(aq) + F^-(aq)$$
The acid-conjugate-base concept only applies to weak acids because the deprotonated anion of strong acids e.g. Cl– has very little to no tendency of accepting a proton. In other words, if the Ka of an acid is too high or Kb of its conjugate base is too low, the equilibrium has little reversibility.
A weak base, after accepting a proton, forms a conjugate acid. For example, protonation of ammonia by water:
$$NH_3(aq) + H_2O(l) \leftrightharpoons NH_4^+(aq) + OH^-(aq)$$
Arrhenius versus Brønsted-Lowry Theory
Overall, the Brønsted-Lowry theory is better than the Arrhenius theory because it can explain the acid and base nature of more species. For example:
- Bases that do not contain hydroxide ions e.g. ammonia, NH3. In Arrhenius’ theory, ammonia is not considered a base because it does not contain nor dissociate into hydroxide ions (OH–).
$$ NH_3(aq) + H_2O(l) \leftrightharpoons NH_4^+(aq) + OH^-(aq)$$
- Acid and base reactions that occur in non-aqueous solution, namely the reaction between ammonia and hydroxide chloride gas. In this reaction, ammonia accepts a proton from hydrogen chloride (base) while hydrochloride donates a proton to ammonia (acid).
$$NH_3(g) + HCl(g) \leftrightharpoons NH_4Cl(s)
The Arrhenius theory does not explain the acid base nature of the two reactants above because they are in gaseous states (non-aqueous).
Major Advantages of the Brønsted-Lowry Theory
The Brønsted-Lowry definition of acids and bases offers many advantages over the Arrhenius and operational definitions.
- It expands the list of potential acids to include positive and negative ions, as well as neutral molecules. Therefore, it can explain a wider range of acids.
- It explains the role of water in acid-base reactions: Water accepts H+ ions from acids to form the H3O+
- It can be expanded to include solvents other than water and reactions that occur in the gas or solid phases.
- It links acids and bases into conjugate acid-base pairs. This concept
- explains the acidity and basicity of acidic and basic salts respectively
- explains the relative strengths of acids and bases according to the strength of their conjugate partner
- Explains the acidic and basic properties of amphiprotic substances
Limitations of the Brønsted-Lowry Theory
- Does not explain the acidity of acidic oxides such as SO2 and SO3
- Does not explain the basicity of basic oxides such as MgO and CaO
- Does not explain reactions between acidic and basic oxides as they do not involve proton transfer. For example, $$SO_3(g) + CaO(s) \rightarrow CaSO_4(s)$$
- The above limitations are explained in the Lewis Theory of acid and bases which states:
- Acids are electron pair acceptors
- Bases are electron pair donors