Definition and Properties of Acids
This topic covers the essential nature of acids, including their definition as proton donors, their characteristic reactions that form salts, and the key distinctions between strength and concentration. Understanding these concepts and the logarithmic pH scale is fundamental for both qualitative and quantitative chemistry problems.
Part of the ESAT Chemistry syllabus — revision for the Engineering and Science Admissions Test (ESAT), the UAT-UK admissions test for Cambridge, Imperial, Oxford and UCL.
Key points
- Acids react with specific substance types to form salts: with reactive metals they produce hydrogen gas, with carbonates they produce carbon dioxide, and with bases (metal oxides/hydroxides) they produce water.
- The terms 'strong' and 'weak' refer to the degree of dissociation of an acid in water, whereas 'concentrated' and 'dilute' refer to the amount of acid dissolved per unit volume.
- The pH scale is logarithmic. A change of 1 pH unit (e.g., from pH 4 to pH 3) represents a 10-fold change in the concentration of H⁺ ions.
- The nature of oxides is predictable: non-metal oxides (e.g., CO₂, SO₂) are typically acidic, while metal oxides are typically basic.
- Not all metals react with acids. A metal must be more reactive than hydrogen in the reactivity series to displace H₂ gas from an acid.
- Acids can be monoprotic (donating one H⁺, like HCl), diprotic (donating two H⁺, like H₂SO₄), or triprotic (donating three H⁺, like H₃PO₄), which is crucial for stoichiometric calculations.
Formulae
Acid + Reactive Metal → Salt + Hydrogen Predicting the products when a metal like magnesium or zinc is added to an acid. Note the production of H₂ gas.
Acid + Metal Carbonate → Salt + Water + Carbon Dioxide Predicting the products when an acid reacts with a carbonate, such as limestone (CaCO₃). Note the production of CO₂ gas.
Acid + Metal Oxide / Metal Hydroxide → Salt + Water For any neutralisation reaction between an acid and a base. This is the fundamental reaction for titrations.
Definitions
- Acid
- A substance that acts as a proton (H⁺) donor, or one that produces hydrogen ions (H⁺) when dissolved in water.
- Strong Acid
- An acid that fully ionises in aqueous solution, meaning almost all its molecules release their H⁺ ions. Example: HCl.
- Weak Acid
- An acid that only partially ionises in aqueous solution, establishing an equilibrium between the intact molecules and their ions. Example: CH₃COOH.
- Salt
- An ionic compound formed when the H⁺ ion of an acid is replaced by a metal ion or another positive ion like ammonium (NH₄⁺).
Worked example
A 100 cm³ solution of nitric acid (HNO₃) has a pH of 1. A separate 100 cm³ solution of sulfuric acid (H₂SO₄) has a pH of 1. Which solution requires a greater volume of 0.1 mol/dm³ NaOH to be completely neutralised, and why?
- 1
Both solutions have a pH of 1.
This means the concentration of H⁺ ions in both is 10⁻¹ mol/dm³, which is 0.1 mol/dm³.
- 2
Nitric acid (HNO₃) is a strong monoprotic acid.
It dissociates as HNO₃ → H⁺ + NO₃⁻.
Therefore, a H⁺ concentration of 0.1 mol/dm³ corresponds to a HNO₃ concentration of 0.1 mol/dm³.
- 3
Sulfuric acid (H₂SO₄) is a strong diprotic acid.
It dissociates as H₂SO₄ → 2H⁺ + SO₄²⁻.
Therefore, a H⁺ concentration of 0.1 mol/dm³ corresponds to a H₂SO₄ concentration of 0.1 / 2 = 0.05 mol/dm³.
- 4
Calculate the moles of acid in each solution.
Moles = concentration x volumeVolume = 100 cm³ = 0.1 dm³ - 5 Moles HNO₃ = 0.1 mol/dm³ × 0.1 dm³ = 0.01 mol
- 6 Moles H₂SO₄ = 0.05 mol/dm³ × 0.1 dm³ = 0.005 mol
- 7
For neutralisation, we consider the total available protons.
Moles H⁺ from HNO₃ = 0.01 mol.
Moles H⁺ from H₂SO₄ = 2 × moles of H₂SO₄ = 2 × 0.005 = 0.01 mol.
The number of moles of H⁺ to be neutralised is identical in both solutions.
- 8
Since the moles of H⁺ are the same in both solutions (0.01 mol) and the concentration of the NaOH base is the same, they will require the exact same volume of NaOH to be neutralised.
Answer: Both solutions require the same volume of NaOH. Although the concentration of the H₂SO₄ acid is half that of the HNO₃, its diprotic nature means that the total number of moles of H⁺ ions is identical in both solutions for the given volume.
Common mistakes
- ×Forgetting reactivity constraints: Do not assume all metals react with acids. Metals below hydrogen in the reactivity series (e.g., Cu, Ag, Au) will not displace H₂ gas from dilute acids.
- ×Incorrect salt formulae: Always balance the charges of the ions when writing the formula for a salt. For example, the salt from Al and H₂SO₄ is Al₂(SO₄)₃, not AlSO₄, because the aluminium ion is Al³⁺ and the sulfate ion is SO₄²⁻.
- ×Confusing strength and concentration: A dilute solution of a strong acid (like 0.001M HCl) can have a higher pH than a concentrated solution of a weak acid (like 1.0M CH₃COOH). The terms are independent.
- ×Ignoring the number of acidic protons: When performing neutralisation calculations, always account for whether an acid is mono-, di-, or triprotic. One mole of H₂SO₄ provides two moles of H⁺ for reaction.
No-calculator tips
- ✓pH to [H⁺] conversion: For integer pH values, the conversion is simple. pH = X means [H⁺] = 10⁻ˣ mol/dm³. For example, pH 4 is [H⁺] = 10⁻⁴ mol/dm³.
- ✓Powers of 10 for dilution: Diluting a strong acid by a factor of 10 increases its pH by 1. Diluting by 100 increases pH by 2. This is a quick way to check dilution calculations.
- ✓Stoichiometry by inspection: For neutralisation, look at the number of H⁺ and OH⁻. To neutralise H₂SO₄ with NaOH, you can see you'll need two NaOH for every one H₂SO₄ just by looking at the formulae (2 H⁺ vs 1 OH⁻). This helps set up molar ratios without writing a full equation.