Common Acids and Bases in the Home

Common Acids and Bases in the Home

The Properties of Acids and Bases

The Properties of Acids and Bases

The Arrhenius Theory of Acids and Bases

The first person to recognize the essential nature of acids and bases was the Swedish chemist Svante Arrhenius (1859–1927).

Arrhenius theory - a theory stating that, in an aqueous solution:

  • an acid is a substance that produces hydrogen ions,
  • a base is a substance that produces hydroxide ions.

Arrhenius Theory

Arrhenius Theory

The Bronsted–Lowry Theory of Acids and Bases

A more general explanation of acids and bases was suggested by the Danish chemist Johannes Bronsted (1879–1947) and the English chemist Thomas Lowry (1874–1936). 

Bronsted–Lowry theory - a theory stating that:

  • an acid is a hydrogen ion (proton) donor,
  • a base is a hydrogen ion (proton) acceptor.

The Bronsted–Lowry Theory and Acidic Solutions

HCl + H2O ⇌ Cl- + H3O+

The Arrhenius theory predicts that from hydrogen chloride, in an aqueous solution, hydrogen ions and chloride ions will be produced.

Experimental evidence suggests that the hydrogen ion is not stable on its own.

The hydrogen ion appears to react with a water molecule to form a hydronium ion, H3O+ (aq).

The water molecule acts as a Bronsted–Lowry base in this reaction because its oxygen atom has lone electron pairs.

One of these electron pairs forms a coordinate covalent bond with the “donated” hydrogen ion, forming a hydronium ion.

The reaction of hydrogen chloride and water is reversible and results in a dynamic equilibrium in an aqueous solution.

The dissolved hydronium ions are responsible for the solution’s acidic properties

Bronsted–Lowry bases contain at least one atom with one or more lone electron pairs (most often, O, N, or P).

The Brønsted–Lowry Theory and Acidic Solutions

The Bronsted–Lowry Theory and Basic Solutions

NH3 (g) + H2O (l) ⇌ NH4+ (aq) + OH- (aq)

In this reaction, a water molecule donates a hydrogen ion to the ammonia molecule, forming an ammonium ion, NH4+ (aq), and a hydroxide ion, OH- (aq).

The water molecule acts as a Bronsted–Lowry acid (a proton donor) and the ammonia acts as a base (a proton acceptor) according to the Bronsted–Lowry theory.

The hydroxide ions are responsible for the solution’s basic properties.

The Brønsted–Lowry Theory and Basic Solutions

Conjugate Acid–Base Pairs

Acid–base reactions are reversible, a hydrogen ion (proton) transfer may occur in the forward reaction and also in the reverse reaction.

Thus, there is a Bronsted–Lowry acid (hydrogen ion donor) and a Bronsted–Lowry base (hydrogen ion acceptor) on each side of the reaction equation.

Conjugate acid–base pairs

The conjugate acid–base pairs in the equations representing the reactions between (a) a hypothetical acid, HA(aq), and water and (b) a hypothetical base, B (aq), and water.

Conjugate base - the substance that forms when an acid loses a hydrogen ion (proton).

Conjugate acid - the substance that forms when a base, according to the Bronsted–Lowry theory, accepts a hydrogen ion (proton).

Conjugate acid–base pair – two substances related to each other by the donating and accepting of a single hydrogen ion.

For any acid–base reaction, there will always be one conjugate acid–base pair made up of an acid and its conjugate base and another conjugate acid–base pair made up of a base and its conjugate acid.

The Bronsted–Lowry Theory and Non-aqueous Reactions

The Arrhenius theory assumes that acid–base reactions occur in aqueous solutions. 

The Bronsted–Lowry theory is not limited to aqueous solutions because it can be extended to reactions in other states.

Amphiprotic (Amphoteric) Substances

A substance can be classified as an acid or base, according to the Bronsted–Lowry theory, only for a specific reaction.

A substance may act as an acid in one reaction and a base in another reaction.

Amphiprotic (or amphoteric) substance is a substance that may act as a Bronsted–Lowry acid in some reactions and as a Bronsted–Lowry base in others.

Both water and the hydrogen carbonate ion, HCO3- (aq), are amphiprotic:

Amphiprotic (Amphoteric) Substances - water and the hydrogen carbonate ion

The Acid Ionization Constant, Ka

Acid ionization constant (Ka) - the equilibrium constant for the ionization of an acid; also called the acid dissociation constant.

The reaction of a weak acid, HA, with water forms a dynamic equilibrium involving H3O+ and a conjugate base, A-:

HA (aq) + H2O (l) ⇌ H3O+ (aq) + A- (aq).

An equilibrium law equation for this reaction:

Equilibrium law equation

Simplification of this equation:

1. For convenience, it is possible to write the reaction equation as though the acid, HA(aq), simply separates into H+ (aq) and A- (aq) ions:

HA (aq) ⇌ H+ (aq) + A- (aq)

2. If the concentration of a chemical substance remains constant during a reaction, that substance is omitted from the equilibrium law equation. In a dilute solution, we can assume that the concentration of liquid water remains essentially constant when an acid is dissolved. Thus, H2O (l) is omitted from the equilibrium law equation.

The equilibrium law equation, with these modifications included, is called the acid ionization constant equation:

equilibrium law equation

Some Acid Ionization Constants

Acid Acid ionization constant, Ka
hydrocyanic, HCN(aq) 6.2 x 10-10
benzoic, HC6H5CO2 (aq) 6.3 x 10-5
propanoic, HC3H5O2 (aq) 1.3 x 10-5
ethanoic (acetic), HC2H3O2 (aq) 1.8 x 10-5
hydrofluoric, HF(aq) 6.6 x 10-4
nitrous, HNO2 (aq) 4.6 x 10-4
methanoic (formic), HCHO2 (aq) 1.8 x 10-4


A Competition for Protons

In the general equation for the reaction of an acid with water, there are always two bases:

HA (aq) + H2O (l) ⇌ H3O+ (aq) + A- (aq)

Two bases, H2O (l) and A- (aq), compete for the hydrogen ion.

If H2O has a much greater affinity for H+ than does A- (if it is a stronger base), the equilibrium position will be far to the right. Most of the dissolved acid will be in the ionized form, A- (aq).

Conversely, if A- is a much stronger base than H2O, then the equilibrium position will lie far to the left. As a result, most of the dissolved acid will be present at equilibrium as HA(aq) molecules.

Base ionization constant, Kb

Base ionization constant (Kb) – the equilibrium constant for the ionization of a base; also called the base dissociation constant.

The general equation for the reaction of a base, B, with water:

B (aq) + H2O (l) ⇌ BH+ (aq) + OH- (aq)

Bronsted-Lowry bases react with water to produce OH- (aq) ions and a conjugate acid, which together determine the acid–base properties of the aqueous solution. 

Weak bases, like weak acids, form dynamic equilibria in aqueous solutions.

For the reaction of a generic base with water, the equilibrium law equation, K, is written as follows:

The equilibrium law equation of base

Since the concentration (density) of water is a constant, it can be incorporated into the value of K (just as it was in the equilibrium law equation for Ka).

This yields a new constant, Kb, called the base ionization constant:

Base ionization constant

Weak bases are the conjugate bases of weak acids.

For example, the ethanoate ion, C2H3O2- (aq), is the conjugate base of ethanoic acid, HC2H3O2 (aq). Similarly, the hypochlorite ion, ClO- (aq), is the conjugate base for hypochlorous acid, HClO(aq).

Kb Values of Selected Weak Bases at 25 °C

Name of base Formula Kb
dimethylamine (CH3)2NH (aq) 9.6 x 10-4
butylamine C4H9NH2 (aq) 5.9 x 10-4
methylamine CH3NH2 (aq) 4.4 x 10-4
aniline C6H5NH2 (aq) 4.1 x 10-10
ammonia NH3 (aq) 1.8 x 10-5
hydrazine N2H4(aq) 1.7 x 10-6
morphine C17H19NO3 (aq) 7.5 x 10-7
hypochlorite ion CIO2 (aq) 3.45 x 10-7
pyridine C5H5N(aq) 1.7 x 10-9
ethanoate ion C2H3O2- (aq) 5.6 x 10-10
urea NH2CONH2 (aq) 1.5 x 10-14