Welcome to Acids, Bases, and Salts!
Welcome to one of the most exciting and foundational topics in your CCEA GCSE Chemistry course! From the sour taste of lemons to the sting of a bee, and from cleaning products under your sink to life-saving medicines, acids and bases are everywhere in our everyday lives.
In this chapter, you will learn what makes a chemical an acid or an alkali, how they behave, how to measure their strength using indicators and the pH scale, and how they react together to create useful compounds called salts.
Don't worry if chemical equations seem tricky at first — we will break down every single reaction step-by-step so you feel totally confident for your exams!
1. What are Acids, Bases, and Alkalis?
To understand this topic, let's start by defining our three main characters: acids, bases, and alkalis.
Acids
An acid is a substance that dissolves in water to produce hydrogen ions, written as \(H^+\).
Common laboratory acids include:
• Hydrochloric acid: \(HCl\)
• Sulfuric acid: \(H_2SO_4\)
• Nitric acid: \(HNO_3\)
• Ethanoic acid: \(CH_3COOH\) (found in vinegar)
Memory Tip: Whenever you see a formula starting with \(H\) (like \(HCl\) or \(HNO_3\)), it is almost certainly an acid producing \(H^+\) ions in water!
Bases and Alkalis
A base is a substance that reacts with an acid to form a salt and water only. Bases are typically metal oxides (such as copper(II) oxide, \(CuO\)) or metal hydroxides (such as iron(III) hydroxide, \(Fe(OH)_3\)).
An alkali is simply a soluble base — a base that dissolves in water. When an alkali dissolves in water, it produces hydroxide ions, written as \(OH^-\).
Common alkalis include:
• Sodium hydroxide: \(NaOH\)
• Potassium hydroxide: \(KOH\)
• Calcium hydroxide: \(Ca(OH)_2\)
• Ammonia solution: \(NH_3\text{ (aq)}\) or \(NH_4OH\)
The Analogy to Remember: Think of bases as "shoes" and alkalis as "trainers". All trainers are shoes, but not all shoes are trainers! In the same way, all alkalis are bases, but not all bases are alkalis (because not all bases dissolve in water).
Key Takeaway
Acids produce \(H^+\) ions in aqueous solution.
Alkalis are soluble bases that produce \(OH^-\) ions in aqueous solution.
2. The pH Scale and Indicators
The pH Scale
The pH scale runs from \(0\) to \(14\) and measures how acidic or alkaline an aqueous solution is:
• pH \(0\) to \(2\): Strongly acidic (e.g. \(HCl\), stomach acid)
• pH \(3\) to \(6\): Weakly acidic (e.g. ethanoic acid, orange juice, milk)
• pH \(7\): Neutral (e.g. pure water, sodium chloride solution)
• pH \(8\) to \(11\): Weakly alkaline (e.g. baking soda, ammonia solution)
• pH \(12\) to \(14\): Strongly alkaline (e.g. concentrated \(NaOH\), bleach)
Indicators and Their Colour Changes
An indicator is a dye that changes colour depending on whether it is in an acidic, neutral, or alkaline solution. For CCEA GCSE, you must know the exact colour changes for four key indicators:
1. Universal Indicator:
• Red in strong acid (\(\text{pH } 0-2\))
• Orange/Yellow in weak acid (\(\text{pH } 3-6\))
• Green in neutral (\(\text{pH } 7\))
• Blue in weak alkali (\(\text{pH } 8-11\))
• Purple/Violet in strong alkali (\(\text{pH } 12-14\))
2. Litmus (paper or solution):
• Red in acid
• Blue in alkali
• (Purple in neutral solution)
3. Phenolphthalein:
• Colourless in acid and neutral solutions
• Pink (or magenta) in alkali
4. Methyl Orange:
• Red in acid
• Yellow in alkali and neutral solutions (often appears orange at exact neutral end-point)
Common Mistake to Avoid: Do not confuse methyl orange and phenolphthalein colours! Remember: Phenolphthalein turns Pink in alkali.
Key Takeaway
pH below \(7\) is acidic (\(H^+\) ions present), pH \(7\) is neutral, and pH above \(7\) is alkaline (\(OH^-\) ions present). Universal indicator gives a range of colours, while single indicators like phenolphthalein and methyl orange give distinct, sharp colour changes.
3. Reactions of Acids
Acids take part in three classic chemical reactions. Learning these general word equations is the secret to mastering this chapter!
Reaction 1: Acid + Metal \(\rightarrow\) Salt + Hydrogen
Mnemonic: MASH (Metal + Acid \(\rightarrow\) Salt + Hydrogen)
When a reactive metal reacts with an acid, it bubbles and releases hydrogen gas.
• Word equation: \(\text{Magnesium} + \text{Hydrochloric acid} \rightarrow \text{Magnesium chloride} + \text{Hydrogen}\)
• Balanced symbol equation: \(Mg\text{ (s)} + 2HCl\text{ (aq)} \rightarrow MgCl_2\text{ (aq)} + H_2\text{ (g)}\)
Observations to note:
• Grey metal solid disappears / dissolves
• Effervescence / bubbles of colourless gas produced
• The reaction mixture gets warm (exothermic reaction)
Test for Hydrogen Gas (\(H_2\)): Apply a lit wooden splint to the gas. A positive result is a squeaky pop sound.
Reaction 2: Acid + Base/Alkali \(\rightarrow\) Salt + Water (Neutralisation)
When an acid reacts with a metal oxide or metal hydroxide, a neutralisation reaction occurs to produce a salt and water.
• Word equation (with insoluble base): \(\text{Copper(II) oxide} + \text{Sulfuric acid} \rightarrow \text{Copper(II) sulfate} + \text{Water}\)
• Balanced symbol equation: \(CuO\text{ (s)} + H_2SO_4\text{ (aq)} \rightarrow CuSO_4\text{ (aq)} + H_2O\text{ (l)}\)
• Word equation (with alkali): \(\text{Sodium hydroxide} + \text{Hydrochloric acid} \rightarrow \text{Sodium chloride} + \text{Water}\)
• Balanced symbol equation: \(NaOH\text{ (aq)} + HCl\text{ (aq)} \rightarrow NaCl\text{ (aq)} + H_2O\text{ (l)}\)
The Ionic Equation for Neutralisation:
In all aqueous acid-alkali neutralisation reactions, hydrogen ions from the acid react with hydroxide ions from the alkali to form water molecules:
\(H^+\text{ (aq)} + OH^-\text{ (aq)} \rightarrow H_2O\text{ (l)}\)
Reaction 3: Acid + Metal Carbonate \(\rightarrow\) Salt + Water + Carbon Dioxide
When an acid reacts with a metal carbonate or metal hydrogencarbonate, it fizzes and releases carbon dioxide gas.
• Word equation: \(\text{Calcium carbonate} + \text{Hydrochloric acid} \rightarrow \text{Calcium chloride} + \text{Water} + \text{Carbon dioxide}\)
• Balanced symbol equation: \(CaCO_3\text{ (s)} + 2HCl\text{ (aq)} \rightarrow CaCl_2\text{ (aq)} + H_2O\text{ (l)} + CO_2\text{ (g)}\)
Observations to note:
• White solid disappears / dissolves
• Vigorous effervescence / bubbles of gas produced
• Colourless solution formed
Test for Carbon Dioxide Gas (\(CO_2\)): Bubble the gas through limewater (calcium hydroxide solution). The limewater turns from colourless to milky/cloudy white.
Naming the Salt Formed
A salt gets its name from two parts:
1. The first part comes from the metal in the base, metal, or carbonate (e.g. Sodium, Copper, Calcium).
2. The second part comes from the acid used:
• Hydrochloric acid (\(HCl\)) makes chlorides (\(Cl^-\))
• Sulfuric acid (\(H_2SO_4\)) makes sulfates (\(SO_4^{2-}\))
• Nitric acid (\(HNO_3\)) makes nitrates (\(NO_3^-\))
• Ethanoic acid (\(CH_3COOH\)) makes ethanoates (\(CH_3COO^-\))
Example: Copper oxide + Nitric acid \(\rightarrow\) Copper nitrate + Water
Key Takeaway
1. Metal + Acid \(\rightarrow\) Salt + Hydrogen (\(H_2\) pops with a lit splint)
2. Base + Acid \(\rightarrow\) Salt + Water (Ionic equation: \(H^+ + OH^- \rightarrow H_2O\))
3. Carbonate + Acid \(\rightarrow\) Salt + Water + Carbon Dioxide (\(CO_2\) turns limewater milky)
4. Preparing Soluble Salts in the Laboratory
A core practical requirement in CCEA GCSE Chemistry is preparing a pure, dry sample of a soluble salt using an insoluble base or carbonate and an acid (for example, making copper(II) sulfate crystals, \(CuSO_4 \cdot 5H_2O\), from copper(II) oxide and sulfuric acid).
Step-by-Step Method: The Excess Solid Method
Step 1: Reaction and Heating
Measure a set volume of acid (e.g. \(H_2SO_4\)) into a beaker and warm it gently using a Bunsen burner. Add the insoluble solid (e.g. black \(CuO\) powder) one spatula at a time while stirring with a glass rod.
• Why warm the acid? To speed up the rate of reaction.
• Why add excess solid? To make sure all the acid is completely reacted and used up so that no unreacted acid contaminates our final salt crystals!
Step 2: Filtration
Filter the warm mixture through filter paper in a filter funnel into a conical flask.
• Why filter? To remove the unreacted excess insoluble solid (residue). The pure salt solution collects in the flask as the filtrate (a clear blue solution of \(CuSO_4\)).
Step 3: Evaporation / Concentration
Pour the clear filtrate into an evaporating basin. Heat the solution over a water bath or gentle Bunsen flame to evaporate water until the crystallisation point is reached.
• How do we test for the crystallisation point? Dip a cool glass rod into the solution; if small crystals form on the rod when removed, the solution is ready.
Step 4: Cooling and Crystallisation
Turn off the heat and leave the concentrated solution to cool slowly at room temperature.
• Why cool slowly? Slower cooling allows larger, well-shaped crystals to grow.
Step 5: Drying the Crystals
Filter off the crystals and gently pat them dry between two sheets of clean filter paper, or leave them in a warm desiccator / drying oven.
• Common Mistake: Do not heat strongly with a direct flame to dry them, or hydrated salts will lose their water of crystallisation and turn into an anhydrous powder!
Key Takeaway
To make a pure soluble salt from an insoluble reactant: React in excess \(\rightarrow\) Filter \(\rightarrow\) Evaporate to crystallisation point \(\rightarrow\) Cool to crystallise \(\rightarrow\) Filter and dry with filter paper.
5. Quick Summary and Exam Revision Checklist
Make sure you have memorised these essential facts before your exam:
• Acids contain \(H^+\) ions; alkalis contain \(OH^-\) ions.
• Strong acids have pH \(0-2\); strong alkalis have pH \(12-14\); neutral solutions have pH \(7\).
• Litmus is red in acid, blue in alkali.
• Phenolphthalein is colourless in acid, pink in alkali.
• Methyl orange is red in acid, yellow in alkali.
• Neutralisation ionic equation: \(H^+\text{ (aq)} + OH^-\text{ (aq)} \rightarrow H_2O\text{ (l)}\).
• Metal + Acid \(\rightarrow\) Salt + \(H_2\)
• Base + Acid \(\rightarrow\) Salt + \(H_2O\)
• Carbonate + Acid \(\rightarrow\) Salt + \(H_2O\) + \(CO_2\)
• Always add excess insoluble solid when preparing crystals to ensure all acid is neutralised!