9.3.D - Acid Behaviour

Our ideas about acids and bases have changed over time as chemists have improved their understanding of these substances. The ancients classified acids as those substances that produced a sour taste on the tongue while nowadays we consider acids as those substances that donate protons and bases as those that accept them in chemical reactions. This theory is known as the Bronsted-Lowry theory and allows us to conceptualise the behaviour of acids and bases at increasingly sophisticated levels of understanding. Seeing acids in this way accounts for the behaviour of buffer solutions which resist changes in pH as well as the varying pH of salt solutions.

In this unit you will learn:

  • About the historical development of ideas about acids and bases,
  • How acids and bases behave in chemical reactions in terms of the Bronsted-Lowry theory,
  • Why different salt solutions can have varying pH values, and
  • How buffer solutions work.

Acid-Base Theories

Ancient people noticed that certain substances produced a sharp taste when placed on the tongue. Citrus fruits and vinegar tasted sour. The latin word for sour is acere.

In 1776, Lavoiser discovered that many compounds that contain oxygen had acidic properties when dissolved in water. He defined an acid as a non-metal compound containing oxygen.

In 1810, Davy showed that hydrogen chloride had acidic properties and contained no oxygen. He redefined an acid as a substance that contained replaceable hydrogen atoms as hydrogen was evolved when acids reacted with active metals such as magnesium.

In 1894, Arrhenius extended this theory to explain that acids contain hydrogen that is lost from a molecule as a hydrogen ion. An acid is, therefore, a substance that provides hydrogen ions is aqueous solution and a base was one that provided hydroxide ions. His theory was important in the development of the pH concept, which is defined in terms of hydrogen ion concentration. Arrhenius realised that strong and weak acids differed according to the degree of ionisation and that neutralisation reactions could be understood in terms of the reaction between hydrogen and hydroxide ions. His theory, however, could not explain why carbonates and metal oxides had basic properties and why some salt solutions could be acidic and basic.

In 1923, Bronsted in Denmark and Lowry in England independently proposed a new theory. An acid was a proton donor and a base was a proton acceptor. In neutralisation reactions, a proton is transferred from the acid to the base. As a result, a new base, the conjugate base, and a new acid, the conjugate acid, was formed.


Bronsted-Lowry Theory

The Bronsted-Lowry definition of acids and bases is named after Johannes Bronsted and Thomas Lowry, who independently proposed it in 1923. A Bronsted-Lowry acid is defined as any substance that can donate a hydrogen ion (proton) and a Bronsted-Lowry base is any substance that can accept a proton. Thus, according to the Bronsted-Lowry definition, acids and bases must come in what are called conjugate acid-base pairs. After the acid has lost its proton it becomes a conjugate base. After the base has accepted the proton, it becomes an conjugate acid.

Hydrochloric acid acts as a Bronsted-Lowry acid in the following equation and water is acting as the base. A proton is transferred from the acid to the base and the hydronium ion is formed as the conjugate acid and the chloride ion is formed as the conjugate base.

Let's next consider the ionisation of the weak acid ethanoic (acetic) acid in water.

Here, ethanoic/acetic acid (CH3COOH) is an acid because it donates a proton and the ethanoate/acetate ion (CH3COO-) is its conjugate base because it can now accept a proton. Note that water and the hydronium ion also form an acid-base conjugate pair. In this case, water is the base and the hydronium ion is its corresponding conjugate acid. Because ethanoic acid is a weak acid, its conjugate base, the ethanoate ion, will be strong.

The hydronium ion rather than the hydrogen ion has been used to denote the nature of hydrogen ions in water in the above reaction. Without the use of the hydronium ion in this equation, it would be difficult to see conjugate acid-base pairs.

Similarly when the weak base ammonia (NH3) is dissolved in water the following equilibrium reaction occurs.

Here, ammonia is the base and the ammonium ion (NH4+) is its conjugate acid. Similarly, water acts as an acid and the hydroxide ion acts as its conjugate base. Because ammonia is a weak base, its corresponding conjugate acid, the ammonium ion, will be strong.

Another important advantage of the Bronsted-Lowry definition is that we are not limited to water as the solvent. Consider the reaction that occurs when hydrochloric acid is dissolved in ammonia.

HCl (g) + NH3 (g) ⇒ NH4Cl (s)

Here, hydrogen chloride acts as the acid with the chloride ion as its conjugate base. Ammonia acts as the base with the ammonium ion as its conjugate acid.

An interesting ambiguity comes up within the Bronsted-Lowry definition, namely that some species can act either as an acid or a base. Such species are called amphiprotic substances. An example is the hydrogen carbonate ion (HCO3-). When dissolved in water, two possible reactions can occur.

HCO3- (aq) + H2O (aq) ⇔ CO32- (aq) + H3O+ (aq)


HCO3- (aq) + H2O (aq) ⇔ H2CO3 (aq) + OH- (aq)

In the first of these the hydrogen carbonate ion acts as the acid with the carbonate ion (CO32-) as its conjugate base. In the second the hydrogen carbonate ion acts as a base with carbonic acid (H2CO3) as its conjugate acid. Other examples of amphiprotic substances include the hydrogen sulfate ion, the dihydrogen phosphate ion and the water molecule.

Conjugate Acids and Bases: Introduction to conjugate acids and bases

Kahn Academy introduction to conjugate acids and bases.


pH of Salt Solutions

A salt is formed during the neutralisation reaction between an acid and a base. Salts are actually defined as acids with their hydrogen ion replaced by a metal ion. One would expect salts to be neutral in aqueous solution as they generally don't contain hydrogen or hydroxide ions. However, the Bronsted-Lowry definition of acids and bases does permit some salts to be acidic or basic. The ions of some salts react with water in solution to varying degrees and this is what contributes to the acidic or basic pH values for some salts.

Salt from a Weak Acid and Strong Base

Let's consider the salt formed in the neutralisation reaction between ethanoic acid and sodium hydroxide. The equation for the neuralisation reaction is:

CH3COOH (aq) + NaOH (aq) ⇒ NaCH3COO (aq) + H2O (l)

Spectator ions are those that don't actually participate in the reaction. In the above reaction, sodium and ethanoate ions are unchanged as the start as ions and remain as the same ions in the solution after the neutralisation is complete. The hydrogen and hydroxide ions react to form water and create a covalent bond in the process. It is the specatator ions though in this reaction that make up the salt and they react to varying degrees with the water in the solution to give it a pH of around 9.

To account for this pH we must determine to what extent each ion will react with water. Considering the ethanoate ion first, the equation for its reaction with water is shown below.

CH3COO- (aq) + H2O (l) ⇔ CH3COOH (aq) + OH- (aq)

This is an equilibrium system which favours the products over the reactants. You will remember that the conjugate base of a weak acid is a strong conjugate base. The ethanoate ion is a strong conjugate base which accounts for the equilibrium favouring the products and the presence of a significant concentration of hydroxide ions present in the solution.

Considering now the sodium ion and its reaction with water, it would not react with the water.

Na+ (aq) + H2O (l) ⇔ NO REACTION

Since sodium hydroxide is a very strong base, the sodium ion is a very weak conjugate acid. For this reason the reaction does not occur.

For sodium ethanoate, the concentration of hydroxide ions in the salt solution is higher than the concentration of hydrogen ions so the salt will have a pH of around 9. In a titration between a weak acid and a strong base, this is also why the equivalence point occurs at around the same pH.

Salt from a Weak Base and a Strong Acid

Let's now consider the salt ammonium chloride (NH4Cl) formed through the neutralisation reaction of hydrochloric acid and ammonia (NH3). The reactions for each ion with water would be:

Cl- (aq) + H2O (aq) ⇔ NO REACTION

NH4+ (aq) + H2O (l) ⇔ NH3 (aq) + H3O+ (aq)

The chloride ion is an extremely weak conjugate base so the first hydrolysis reaction does not occur at all. Since ammonia is a weak base, the ammonium ion is a strong conjugate acid, so the second hydrolysis reaction has an equilibrium that favours the products and produces a significant concentration of hydronium ion in the solution. This gives the salt an acidic pH of around 5 as is also why the equivalence point for a titration between a weak base and strong acid occurs at around the same pH.

Salts from Strong Acids and Strong Bases

Salts made from strong acids and strong bases will have neither hydronium or hydroxide ions in solution because for the reasons outlined above, the hydrolysis reactions do not occur. Sodium chloride (NaCl) has sodium and chloride ions that do not react with water at all as they are both weak conjugates. Such salts generally have a neutral pH.

Salts from Weak Acids and Weak Bases

For salts of both weak acids and weak bases, the situation is more complex. However, if we assume that the equilibrium for both hydrolysis reactions lies to the same extent to the right, then the presence of equal amounts of hydroxide and hydronium ions will cancel each other out. This will result in a salt solution with a pH of seven. Should one conjugate be stronger than the other though, the pH of the solution will be slightly acidic or basic depending on which conjugate is stronger.


A salt solution will have a:

  • pH > 7 if formed from a weak acid and a strong base.

  • pH = 7 if formed from a strong acid and a strong base or an equally weak acid and base.

  • pH < 7 if formed from a strong acid and a weak base.



Buffer solutions usually consist of stoichiometrically equivalent amounts of a weak acid and its conjugate base in equilibrium. Buffer solutions are solutions which resist changes in pH upon addition of small amounts of acid or base. The resistive action is the result of the equilibrium which is set up between the weak acid and its conjugate base. If an alkali is added to the solution, hydronium ions mop it up. These ions are regenerated as the equilibrium moves to the right and some of the acid is broken down in to hydrogen ions and anions. If an acid is added, the anions from the conjugate base simply combine with the hydrogen ions and once again pH is restored.

A simple buffer can be made by mixing equal volumes 0.1 mol L-1 solutions of ethanoic acid (CH3COOH) and sodium ethanoate (NaCH3COO). This mixture will keep the solution at a pH of about 5. A titration curve for the reaction between a weak acid and a strong base shows a buffering region before the equivalence point. This is due to the presence of the weak acid and its conjugatebase in the conical flask.

Buffers are found in many natural systems like blood and cell cytoplasm and they are also used in man-made systems like aquaria.

Human Blood

By far the most important buffer for maintaining acid-base balance in the blood is the H2CO3/HCO3- buffer used to maintain the pH of human blood at a value close to 7.4. There are several weak acids at play In this buffer system, but one of them is the equilibrium between carbonic acid and its conjugate base the hydrogen carbonate ion. The simultaneous equilibrium reaction of interest is:

H2CO3 (aq) + H2O (aq) ⇔ HCO3- (aq) + H3O+ (aq)

When excess hydronium ions enter the system, the concentration of these ions increases and the equilibrium shifts to the left to partially counteract the change. This reduces the concentration of hydronium ions and returns the pH to its original level.

When excess hydroxide ions enter the system, the hydroxide ions react with the hydronium ions effectively reducing the concentration of hydronium ions. The equilibrium shifts to the right to partially counteract the change in hydronium ion concentration and this restores the pH to its original level.

Buffers are sensitive and only work within certain limits. If the influx of hydronium or hydroxide ions into a system is too large, the buffer will not be able to restore the pH.