Chemical changes 

Reactivity of Metals

Metals can react with oxygen to form metal oxides. To form a metal oxide, an oxidation reaction occurs as oxygen is added to the metal. 

2Cu + O₂ ⟶ 2CuO

Metal oxides can also lose their oxygen which is known as a reduction reaction. 

ZnO + C ⟶ Zn + CO

The Reactivity Series 

The reactivity series places metals in order of reactivity based on their reactions with acids and water. When a metal reacts, it loses electrons to form positive metal ions. The ease at which it can do this determines its reactivity- the easier a metal loses electrons, the more reactive it is. 

Hydrogen is placed in the reactivity series because metals only react with acid if the metal is more reactive than hydrogen. Carbon is placed in the reactivity series because it allows us to see whether a metal oxide can be reduced by carbon or not for metal extraction.

Metal Extraction

Many metals exist as ores in the Earth’s crust. Ores are rocks containing enough metal that it is economically worthwhile extracting them. Usually, the ore is an oxide of the metal, for example, aluminium oxide. 

The method used to extract these metals from their ores depends on the reactivity of the metal:

• Metals above carbon in the reactivity series are extracted using electrolysis 
• Metals below carbon in the reactivity series are extracted using carbon as a reducing agent.

metal oxide + carbon   →   metal   +   carbon dioxide

Some metals, such as gold, are unreactive and found as an uncombined element so do not need to be chemically extracted. These are known as native metals. 

What is displacement?

Displacement involves a more reactive element taking the place of a less reactive element in its compound.

zinc + copper oxide → zinc oxide + copper

The zinc, which is higher up the reactivity series than copper, displaces the copper from its oxide. For more on displacement visit:

• 4.1.2 The Reactivity Series

Oxidation & Reduction in Terms of Electrons

During this process, electrons are lost and gained. The loss of electrons is called oxidation, and the gain of electrons is called reduction. We can write ionic equations and half equations for displacement reactions. 

The word equation for the displacement of copper in copper sulfate by magnesium is:

Magnesium + copper sulfate  →  magnesium sulfate + copper 

The ions involved in the equation are:

Mg (s) + Cu²⁺ (aq) + SO₄^2- (aq) → Mg²⁺ (aq) + SO₄^2- (aq) + Cu (s)

To write the ionic equation, remove the ions that appear on both sides of the equation (known as spectator ions). In this case, it is the sulfate ions:

Mg (s) + Cu²⁺ (aq)  → Mg²⁺ (aq)  + Cu (s)

To write the half equations, identify the species that has been oxidised and reduced:

Mg → Mg²⁺ + 2e^– 

Cu²⁺ + 2e^– → Cu

• Magnesium atoms lose electrons so are oxidised 
• Copper ions gain electrons so are reduced

Reactions of Acids

Metals above hydrogen in the reactivity series will react with dilute acids to produce a salt and hydrogen:

metal + acid ⟶ salt + hydrogen

The reaction between a metal and hydrogen is a redox reaction meaning oxidation (loss of electrons) and reduction (gain of electrons) occur at the same time. 

The reaction between zinc and hydrochloric acid produces zinc chloride and hydrogen 

Zn + 2HCl ⟶ ZnCl₂ + H₂

The ionic equation is:

Zn + 2H⁺ ⟶ Zn₂₊ + H₂

The half equations are:

Zn → Zn²⁺ + 2e^–

2H⁺ + 2e^– → H₂

• Zinc atoms lose electrons so they are oxidised 
• Hydrogen ions gain electrons so are reduced

Acids can also react with bases, including metal oxides, metal hydroxides and metal carbonates in neutralisation reactions. The products formed depends on the reactants used:

• metal oxide + acid →  salt + water
• metal hydroxide + acid → salt + water
• metal carbonate + acid → salt + water + carbon dioxide

Examples of salts include sodium chloride and copper(II) sulfate. To name a salt, the first part comes from the metal in the reactant, and the second part comes from the acid used: 

• Hydrochloric acid form chlorides
• Nitric acid forms nitrates
• Sulfuric acid forms sulfates

For example, to make sodium chloride and copper(II) sulfate the following reactions could be done:

copper oxide + sulfuric acid  →   copper(II) sulfate + water

sodium hydroxide + hydrochloric acid →  sodium chloride + water 

The Preparation of a Soluble Salt 

A soluble salt is made by the reaction between an acid and an insoluble base. 

A method for the preparation of copper(II) sulfate is: 

• Add sulfuric acid to a beaker and warm gently with a bunsen burner 
• Add copper(II) oxide in excess to the acid to ensure all of the acid has reacted 
• Filter the solution into a conical flask to remove the excess copper oxide 
• Heat the solution gently to evaporate some of the water 
• Leave the remaining solution in a warm place to dry and crystallise
   
• More information about how to prepare copper(II) sulfate: "4.2.5 Required Practical: Preparation of a Soluble Salt"

The pH Scale & Hydrogen Ions 

Acids form hydrogen ions, H⁺, when they are added to water, and alkalis form hydroxide ions, OH^–. It is the presence of H⁺ ions which makes a solution acidic, and OH^– ions which makes a solution alkaline. During neutralisation, the H⁺ ions in an acid react with the OH^– ions in an alkali to form a neutral solution. 

H⁺ (aq) + OH^– (aq) ⟶ H₂O (l)

The pH scale is a measure of hydrogen ion concentration. It indicates whether a substance is an acid, alkali or neutral:

• Acids have a pH below 7
• Alkalis have a pH above 7 
• Neutral substances have a pH of 7. 

The higher the pH the more alkaline it is, the lower the pH the more acidic it is. An indicator or digital pH meter can measure the pH of a substance. A common indicator is universal indicator which changes colour to give an approximate pH as shown below:

The pH scale is logarithmic meaning that each change of 1 on the scale represents a change in concentration by a factor of 10:

• An acid with a pH of 3 has ten times the concentration of H+ ions than an acid of pH 4
• An acid with a pH of 2 has 10 x 10 = 100 times the concentration of H+ ions than an acid with a pH of 4

From this, we can summarise that for two acids of equal concentration, where one is strong and the other is weak, then the strong acid will have a lower pH due to its capacity to dissociate more and hence put more H+ ions into solution than the weak acid

Strong & Weak Acids (Higher Tier Only)

Strong acids will fully ionise in water, so there is a high percentage of H+ ions present in the solution. The higher the percentage of H+ ions, the lower the pH and the stronger the acid (approximately pH 1-3).

HCl ⟶ H⁺ + Cl^–

Weak acids will only partially ionise in water, so there is a lower percentage of H+ ions in the solution. This results in the pH being higher (approximately pH 4-6). Equations to show the ionisation of weak acids use a  ⇌ symbol to indicate the reaction is reversible.

CH₃CH₂COOH ⇌ H⁺ + CH₃CH₂COO^–

What is a titration?

Titrations are used to calculate the concentration of a solution by determining how much alkali is needed to neutralise a quantity of acid and vice versa. 

Method for a strong acid-alkali titration:

• Add hydrochloric acid to the burette, making a note of the initial volume 
• Pipette 25cm³ of sodium hydroxide into a conical flask and place it on a white tile under the burette
• Add a few drops of indicator to the conical flask
• Run the acid from the burette into the conical flask, swirling the flask until the solution changes colour (the end-point)
• Record the volume of acid that was needed to achieve the colour change
• Repeat the titration, adding the solution from the burette one drop at a time towards the end-point until the solution changes colour 
• Repeat until two concordant results are achieved (readings with 0.1 cm³ of each other)

The average volume of hydrochloric acid with the volume and concentration of the sodium hydroxide can be used to determine the concentration of the acid.  For examples of these calculations please visit here: "Titration Calculations"

There are many indicators that can be used, each with a different colour change at the end-point. 

Electrolysis 

What is electrolysis?

Electrolysis is the splitting up of an ionic compound using electricity.  An electrolytic cell is the name given to the set-up used in electrolysis and which consists of the following:

• Electrode: a rod of metal or graphite through which an electric current flows into or out of an electrolyte
• Electrolyte: ionic compound in molten or dissolved solution that conducts the electricity
• Anode: the positive electrode of an electrolysis cell
• Anion: negatively charged ion which is attracted to the anode
• Cathode: the negative electrode of an electrolysis cell
• Cation: positively charged ion which is attracted to the cathode

What happens during the electrolysis of molten ionic compounds?

Molten ionic compounds only contain two types of ions: positively charged metal ions and negatively charged non-metal ions. The positive metal ions will be attracted to the cathode where they gain electrons (reduction) and form a metal. The negative non-metal ions will be attracted to the anode where they lose electrons (oxidation) to form a non-metal. 

What happens during the electrolysis of aqueous ionic compounds?

Aqueous solutions of ionic compounds have been dissolved in water so in addition to the metal and non-metal ions there are also hydrogen ions (H⁺) and hydroxide ions (OH^–). The less reactive ions will be discharged at the electrodes. 

At the cathode:

• The metal ions and H⁺ ions will be attracted to the cathode
• If the metal is above hydrogen in the reactivity series, then hydrogen ions will be discharged at the electrode where they will gain electrons and form hydrogen gas
• If the metal is below hydrogen in the reactivity series, the metal ions will be discharged at the electrode, gain electrons and will form the metal 

At the anode:

• Non-metal ions and OH^– ions will be attracted to the anode
• If there are halide ions present (Group 7 ions) then these ions are discharged at the electrode where they will lose electrons to form a halogen
• If no halide ions are present, OH^– ions will be discharged at the anode, lose electrons and form oxygen gas 

How do I write half equations?

These show the reactions that occur at the anode and cathode. At the anode, ions lose electrons and are oxidised. At the cathode, ions gain electrons so are reduced. During the electrolysis of lead bromide, the following reactions occur at the electrodes:

• Anode: 2Br^– ⟶ Br₂ + 2e^–
• Cathode: Pb²⁺ + 2e^– ⟶ Pb

It is important that the number of atoms/ions are balanced as well as the charges on the ions and electrons.

What keyword definitions do I need to know for chemical changes? 

Some keyword definitions you need to know are:

• Oxidation- the gain of oxygen and loss of electrons
• Reduction- the loss of oxygen and gain of electrons
• Redox- oxidation and reduction occurring simultaneously 
• Displacement- a more reactive element takes the place of a less reactive element 
• Neutralisation- the reaction between an acid and alkali to form a salt and water
• Electrolysis- the splitting up of an ionic compound using electricity 

This is a quick summary of some key concepts on chemical changes - remember to go through the full set of revision notes, which are tailored to your specification, to make sure you know everything you need for your exams!