Theories of acids & bases

8.1 Theories of acids and bases (1.5 hours)

Pause for thought

Have some sympathy for Svante Arrhenius, the Swedish chemist who won the Nobel Prize for Chemistry in 1903. Sub-topic 8.1 is titled "Theories of acids and bases" and yet the only theory covered in the subsequent 'Understandings' is the BrØnsted-Lowry theory. Arrhenius's ionic theory is implied throughout all of the sub-topics 8.1 - 8.5 and in the AHL sub-topics in Topic 18 but it is not specifically mentioned. The Lewis theory of acids and bases is only covered at Higher Level. There are actually many theories of acids and bases. Early theories such as acids having a sour taste and Lavoisier's theory of an acid being the oxide of a non-metal and water are only of passing historical interest but Arrhenius's ionic theory is clearly still very important. This theory states that acids are substances that dissociate to give hydrogen ions in aqueous solution. Most of the other sub-topics in Topics 8 and 18 are in fact much more based on this definition than either the BrØnsted-Lowry or Lewis definitions so it seems strange that it has been left out of 8.1. For example, the pH scale, K_a and pK_a calculations etc. are all concerned with the concentration of hydrogen ions in aqueous solution.

Although Johannes BrØnsted and Thomas Lowry are usually lumped together they worked independently. BrØnsted (see right), who was a Dane, and Lowry from the UK, first proposed their theory in 1923. It considers the deprotonation of acids to form conjugate bases, or the protonation of bases to form conjugate acids. One of the biggest differences between the ionic theory and the BrØnsted-Lowry theory is when it is applied to water. Under the ionic theory pure water is neutral but under the BrØnsted-Lowry theory it can either be acidic or basic depending upon whether it is donating or accepting a proton. This property of water acting as either an acid or a base by donating or accepting a proton is described as amphiprotic. All amphiprotic substances can also be described as amphoteric (can act as either an acid or a base) but the reverse statement is not true as amphiprotic only applies to the BrØnsted-Lowry definition. Metals oxides, such as aluminium oxide, Al₂O₃, can react with acids or bases to forms salts and are therefore amphoteric but they cannot be described as amphiprotic as they are not able to donate a proton.

Nature of science

The fact that both hydrochloric acid, HCl and prussic acid, HCN do not contain oxygen provides a good example of how theories such as Lavoisier's theory (that acids are oxides of non-metals with water) can be falsified and allow other theories to develop. This also provides an example of how theories are superseded.
Reference can also be made to the public understanding of science as terms such as 'the acid test' or 'litmus test' are often used in a non-scientific context.

Learning outcomes

After studying this topic students should be able to:

Understand:

Brønsted–Lowry acids are proton/H+ donors and Brønsted–Lowry bases are proton/H+ acceptors.
The term amphiprotic refers to species that can act as both Brønsted–Lowry acids and bases.
A conjugate acid-base pair differ by a single proton.

Apply their knowledge to:

Deduce the Brønsted–Lowry acids and bases in a chemical reaction.
Deduce the conjugate acids or conjugate bases in a chemical reaction.

Clarification notes

Lewis theory is not required in this sub-topic.
The precise location of the proton transferred should be shown. For example, CH₃COOH/CH₃COO^– rather than C₂H₄O₂/C₂H₃O₂^–.
Students should know that a proton in aqueous solution can be represented as either H⁺(aq) or H₃O⁺(aq).
Students should know the difference between the terms amphoteric and amphiprotic.

International-mindedness

Acid–base theory has been developed by scientists from many different countries and its vocabulary has been influenced by their languages. For example, acidus means sour in Latin, alkali is derived from the Arabic word for calcined ashes and oxygene (thought by Lavoisier to be responsible for the acidic properties of a compound) means acid-forming in Greek.

Teaching tips

From earlier work most students already have the idea that an acid gives up hydrogen ions in aqueous solution and gives a solution with a pH less than 7. You can build on this to try to show exactly what is happening in terms of proton transfer (BrØnsted-Lowry).

I use water firstly with hydrogen chloride then with ammonia to show that it can act either as an acid or a base and explain why. This leads to the idea that in a reversible acid-base reaction there are actually two acid-base conjugate pairs. Stress that it is important to show where the proton that is transferred originates from. Thus in ethanoic acid the methyl hydrogen atoms are bonded too strongly to break so the proton comes from the -OH group. Stress too that a conjugate base will always have one proton less than the acid it originates from and a conjugate acid will always have one proton more than the base it originates from.

Stress the difference between amphiprotic which only applies to BrØnsted-Lowry acids and bases and amphoteric which is a more general terms and covers all substances (whether or not they contain a proton) which can act both as an acid or a base.

Study guide

Page 59

Questions

For ten 'quiz' multiple choice questions with the answers explained see MC test: Theories of acids & bases.

For short-answer questions which can be set as an assignment for a test, homework or given for self study together with model answers see Theories of acids & bases questions.