Acids, Bases and pH (OCR A-Level Chemistry A): Revision Notes

Brønsted–Lowry Acids and Bases

Introduction to acid-base models

Understanding acids and bases has evolved significantly over time. Early scientists recognised that acids taste sour and bases (or alkalis) come from plant ashes. As chemistry developed, several key models emerged to explain acid-base behaviour.

The Arrhenius model (1884) was an important early theory that described:

• Acids dissociate and release $\text{H}^+$ ions in aqueous solution
• Alkalis dissociate and release $\text{OH}^-$ ions in aqueous solution
• Neutralisation occurs when $\text{H}^+$ ions react with $\text{OH}^-$ ions to form water

However, this model was limited to aqueous solutions only. The Brønsted-Lowry model, developed in 1923, extended this understanding by emphasising the central role of proton transfer between chemical species.

Brønsted–Lowry definitions

The Brønsted-Lowry model provides a more comprehensive way to understand acid-base reactions by focusing on what happens to protons (hydrogen ions) during chemical reactions.

What is a Brønsted–Lowry acid?

A Brønsted–Lowry acid is a species that donates a proton ($\text{H}^+$) to another species.

When an acid dissolves in water, it releases protons. For example, hydrochloric acid dissociates as follows:

$\text{HCl(aq)} \rightleftharpoons \text{H}^+\text{(aq)} + \text{Cl}^-\text{(aq)}$

In this reaction, HCl is acting as an acid because it donates a proton.

What is a Brønsted–Lowry base?

A Brønsted–Lowry base is a species that accepts a proton ($\text{H}^+$) from another species.

Bases receive protons during reactions. Common examples include hydroxide ions ($\text{OH}^-$) and ammonia ($\text{NH}_3$). When a base accepts a proton, it forms a new species.

Conjugate acid-base pairs

One of the most important concepts in Brønsted-Lowry theory is the idea of conjugate acid-base pairs. These pairs are intimately connected through proton transfer.

Understanding conjugate pairs

A conjugate acid-base pair consists of two species that differ by one proton ($\text{H}^+$). The two species can be interconverted by transferring a single proton between them.

Let's examine the dissociation of hydrochloric acid:

$\text{HCl(aq)} \rightleftharpoons \text{H}^+\text{(aq)} + \text{Cl}^-\text{(aq)}$

Although this equation shows an equilibrium, HCl is a strong acid, so the position of equilibrium lies far to the right. We can often use a single arrow to show the reaction goes essentially to completion:

$\text{HCl(aq)} \rightarrow \text{H}^+\text{(aq)} + \text{Cl}^-\text{(aq)}$

In this reaction:

• In the forward direction: HCl releases a proton to form its conjugate base, $\text{Cl}^-$
• In the reverse direction: $\text{Cl}^-$ accepts a proton to form its conjugate acid, HCl

HCl and $\text{Cl}^-$ therefore form a conjugate acid-base pair.

Examples of conjugate acid-base pairs

Different acids form different conjugate bases when they donate a proton. Here are some common examples:

Identifying acid-base pairs in equilibria

Many acid-base reactions involve two conjugate acid-base pairs because both the forward and reverse reactions involve proton transfer.

Consider the reaction between hydrochloric acid and hydroxide ions:

$\text{HCl(aq)} + \text{OH}^-\text{(aq)} \rightleftharpoons \text{H}_2\text{O(l)} + \text{Cl}^-\text{(aq)}$

We can label the species as follows:

• Acid 1: HCl (donates a proton)
• Base 2: $\text{OH}^-$ (accepts a proton)
• Acid 2: $\text{H}_2\text{O}$ (formed when base 2 accepts the proton)
• Base 1: $\text{Cl}^-$ (formed when acid 1 donates the proton)

The two conjugate acid-base pairs are:

1. HCl and $\text{Cl}^-$ (acid 1 - base 1)
2. $\text{H}_2\text{O}$ and $\text{OH}^-$ (acid 2 - base 2)

The hydronium ion, $\text{H}_3\text{O}^+$(aq)

In aqueous solutions, acid dissociation requires water to be present because protons must be transferred to a base. Water itself can act as a Brønsted-Lowry base by accepting protons.

When hydrochloric acid dissolves in water, the following equilibrium is established:

$\text{HCl(aq)} + \text{H}_2\text{O(l)} \rightleftharpoons \text{H}_3\text{O}^+\text{(aq)} + \text{Cl}^-\text{(aq)}$

In this reaction:

• HCl acts as an acid (donates a proton)
• $\text{H}_2\text{O}$ acts as a base (accepts a proton)
• $\text{H}_3\text{O}^+$ is formed (the conjugate acid of water)
• $\text{Cl}^-$ is formed (the conjugate base of HCl)

The species $\text{H}_3\text{O}^+$ is called the hydronium ion (or sometimes the oxonium ion). This is the actual acidic species present in all aqueous acid solutions. The hydronium ion is the active acid ingredient that gives acidic solutions their characteristic properties.

Simplification in equations

Although $\text{H}_3\text{O}^+$(aq) is the correct representation of the acidic species in water, it is common practice to use the simplified form $\text{H}^+$(aq) in equations. This simplification makes equations easier to write and understand.

Monobasic, dibasic, and tribasic acids

Acids differ in the number of hydrogen ions they can donate per molecule. This property is described using the terms monobasic, dibasic, and tribasic.

Definitions

The terms refer to the total number of hydrogen ions in the acid molecule that can be replaced per molecule during an acid-base reaction. Typically, these hydrogen ions are replaced by metal ions or ammonium ions ($\text{NH}_4^+$) when the acid reacts to form a salt.

• Monobasic acid: Can donate one proton per molecule
• Dibasic acid: Can donate two protons per molecule
• Tribasic acid: Can donate three protons per molecule

The number of hydrogen atoms in the formula gives you a clue about the type of acid, though you must only count hydrogen atoms that can be released as $\text{H}^+$ ions.

Important notes about basicity

Examples of acid reactions based on basicity

The basicity of an acid determines the stoichiometry of its reactions with bases. Consider neutralisation reactions with sodium hydroxide:

Monobasic acid ($\text{HNO}_3$): $\text{HNO}_3\text{(aq)} + \text{NaOH(aq)} \rightarrow \text{NaNO}_3\text{(aq)} + \text{H}_2\text{O(l)}$

One mole of monobasic acid reacts with one mole of NaOH.

Dibasic acid ($\text{H}_2\text{SO}_4$): $\text{H}_2\text{SO}_4\text{(aq)} + 2\text{NaOH(aq)} \rightarrow \text{Na}_2\text{SO}_4\text{(aq)} + 2\text{H}_2\text{O(l)}$

One mole of dibasic acid reacts with two moles of NaOH.

Tribasic acid ($\text{H}_3\text{PO}_4$): $\text{H}_3\text{PO}_4\text{(aq)} + 3\text{NaOH(aq)} \rightarrow \text{Na}_3\text{PO}_4\text{(aq)} + 3\text{H}_2\text{O(l)}$

One mole of tribasic acid reacts with three moles of NaOH.

The role of $\text{H}^+$ in acid reactions

The hydrogen ion (or more accurately, the hydronium ion) is the active species in all acid reactions. Understanding this allows us to write generalised ionic equations that apply to any acid.

Reactions of acids with metals

Dilute acids react with reactive metals to produce salts and hydrogen gas:

$\text{acid} + \text{metal} \rightarrow \text{salt} + \text{hydrogen}$

For example, magnesium reacts with dilute hydrochloric acid:

Full equation: $2\text{HCl(aq)} + \text{Mg(s)} \rightarrow \text{MgCl}_2\text{(aq)} + \text{H}_2\text{(g)}$

Ionic equation: $2\text{H}^+\text{(aq)} + \text{Mg(s)} \rightarrow \text{Mg}^{2+}\text{(aq)} + \text{H}_2\text{(g)}$

Notice that in the ionic equation, we've cancelled the spectator ions (in this case, $\text{Cl}^-$). The ionic equation shows that $\text{H}^+$ ions are the active species from the acid.

The same ionic equation applies for the reaction of magnesium with dilute sulfuric acid:

Full equation: $\text{H}_2\text{SO}_4\text{(aq)} + \text{Mg(s)} \rightarrow \text{MgSO}_4\text{(aq)} + \text{H}_2\text{(g)}$

Ionic equation: $2\text{H}^+\text{(aq)} + \text{Mg(s)} \rightarrow \text{Mg}^{2+}\text{(aq)} + \text{H}_2\text{(g)}$

This demonstrates that any dilute acid will react with magnesium in the same way because $\text{H}^+$ is the reactive species.

Redox reactions between acids and metals

Some metals undergo redox reactions with dilute acids. For example, zinc reacts with any dilute acid:

$2\text{H}^+\text{(aq)} + \text{Zn(s)} \rightarrow \text{Zn}^{2+}\text{(aq)} + \text{H}_2\text{(g)}$

In this reaction:

• Zinc is oxidised (loses electrons): $\text{Zn} \rightarrow \text{Zn}^{2+} + 2e^-$
• Hydrogen ions are reduced (gain electrons): $2\text{H}^+ + 2e^- \rightarrow \text{H}_2$

Reactions of acids with metal oxides

Metal oxides and hydroxides are bases that neutralise acids to form a salt and water only:

$\text{acid} + \text{base} \rightarrow \text{salt} + \text{water}$

For example, magnesium oxide reacts with any dilute acid:

$2\text{H}^+\text{(aq)} + \text{MgO(s)} \rightarrow \text{Mg}^{2+}\text{(aq)} + \text{H}_2\text{O(l)}$

Reactions of acids with carbonates

Carbonates are bases that react with acids to produce a salt, water, and carbon dioxide:

$\text{acid} + \text{carbonate} \rightarrow \text{salt} + \text{water} + \text{carbon dioxide}$

For solid copper(II) carbonate reacting with any acid:

$2\text{H}^+\text{(aq)} + \text{CuCO}_3\text{(s)} \rightarrow \text{Cu}^{2+}\text{(aq)} + \text{H}_2\text{O(l)} + \text{CO}_2\text{(g)}$

For aqueous carbonates, the spectator ions cancel out to give a simpler ionic equation:

$2\text{H}^+\text{(aq)} + \text{CO}_3^{2-}\text{(aq)} \rightarrow \text{H}_2\text{O(l)} + \text{CO}_2\text{(g)}$

Reactions of acids with alkalis

Alkalis are soluble bases. When an acid reacts with an alkali, both species are in solution, so we can write a very simple ionic equation:

$\text{acid} + \text{alkali} \rightarrow \text{salt} + \text{water}$

The ionic equation for neutralisation of $\text{H}^+$(aq) ions by $\text{OH}^-$(aq) ions is:

$\text{H}^+\text{(aq)} + \text{OH}^-\text{(aq)} \rightarrow \text{H}_2\text{O(l)}$

Writing full equations from ionic equations

While ionic equations are useful for showing the essential chemistry, you must also be able to write full balanced equations. To do this:

1. Start with the ionic equation
2. Add back the spectator ions to balance the charges
3. Ensure all atoms are balanced

Volume calculations in neutralisation

Understanding the basicity of acids is essential for stoichiometric calculations.