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The Reactivity Series Explained (GCSE Chemistry)

What the reactivity series is:

The reactivity series of metals is a chart showing how reactive metals are, with the least reactive metals placed at the bottom of the chart and the most reactive metals placed at the top.

A table outlining the reactivity of various metals with cold water and dilute acids, indicating the reaction types and their reactivity levels from most to least reactive.

Why metals react differently:

The reactivity of a metal is determined by how easily the metal loses electrons; the more easily a metal loses electrons, the higher up it is placed in the Reactivity series table.

Memory tricks for remembering the reactivity series:

There are a variety of ways to remember the reactivity series, one way being the use of mnemonics:

“Please Stop Calling Me A Zebra, I Like Her Calling Me Smart Goat“

“Penguins Swim Like Crazy, Making A Zoo In Cold, Snowy Greenland”

Displacement reactions

Example 1:

For the following reaction, we need to decide if the magnesium and copper atoms swap places when magnesium metal strips are added to a blue copper sulphate solution:

Mg (s) + CuSO₄ (aq) →

magnesium + copper sulfate →

Chart illustrating the reactivity of various metals, with potassium at the top as the most reactive and gold at the bottom as the least reactive.

Since the magnesium is above the copper on the metal reactivity series table, magnesium is more reactive than copper so this means the magnesium atoms (Mg0) release 2 electrons to become completely dissolved in solution and become Mg2+ ions that bond with the water (H2O) and SO₄2- ions while the copper ions (Cu2+) accept the 2 electrons and become copper metal (Cu0) which drops at of solution and forms at the bottom (see the equation below):

Mg (s) + CuSO₄ (aq) → Cu (s) + MgSO₄ (aq)

magnesium + copper sulfate → copper + magnesium sulfate

Additionally, you will start to notice the blue copper sulfate solution starts getting paler when magnesium strips are added to the solution until you get a transparent magnesium sulfate solution.

Example 2:

The following is an example of the highly exothermic Thermite reaction where iron (III) oxide reacts with fine aluminium powder to produce iron and aluminium oxide

Fe2O3 + 2Al → 2Fe + Al2O3

iron(III) oxide + aluminium → iron + aluminium oxide

WARNING: don’t attempt the Thermite experiment unless you have proper scientific lab training, understand the health and safety aspects of this experiment, understand proper chemical disposal procedures, undertaken risk assessments etc.

A vertical chart displaying metals arranged by reactivity, with the most reactive metals listed at the top, including potassium and sodium, and the least reactive, such as lead and copper, at the bottom. The image indicates which metals can be extracted using carbon and which cannot.

This experiment proceeds because iron is below the aluminium in the reactivity series so the aluminium atoms (Al0) lose 3 electrons to become Al3+ and the Fe3+ ions in Fe2O3 accepts the 3 electrons to form iron metal (Fe0). The Fe3+ ions in Fe2O3 are swapped by the Al3+ ions to form Al2O3.

Exam questions:

An examination paper titled 'Reactivity of Metals (F)' with questions related to the reaction between zinc and copper sulfate solution. The first question asks for the type of reaction and includes answer options: combustion, decomposition, and displacement. The second part requires the calculation of the percentage by mass of copper in copper sulfate, providing relative atomic and formula mass values.
A GCSE exam paper question focusing on electrolysis and the extraction of metals, including a table for products of electrolysis and a chemical equation to balance.
A worksheet page showing a chemistry question about calculating the relative formula mass of aluminium oxide (Al2O3) and a reactivity series of metals including potassium, lithium, carbon, zinc, tin, and gold, along with extraction methods.
A worksheet question on displacement reactions, specifically focusing on the extraction of iron from iron oxide using carbon. It includes parts for balancing a chemical equation, explaining the reduction of iron oxide, and calculating the relative formula mass of Fe2O3.
A chemistry exam paper displaying a question about the reactivity of metals, including a chemical reaction involving copper oxide and hydrogen, calculations for percentage atom economy, and an investigation of four different metals by a student.
A table displaying the reactivity results of four metals (A, B, C, D) with various metal sulfate solutions, indicating whether displacement reactions occurred.
An educational worksheet on the reactivity of metals with hydrochloric acid, featuring a diagram (Figure 1) illustrating the rate of bubbling in test tubes labelled A, B, C, and D.

Reference: AQA GCSE Chemistry Topic 4: Chemical Changes Revision – PMT

Exam answers:

A document displaying a mark scheme for a science question, detailing calculations involving displacement, percentage, volume of copper sulfate, mass, temperature change, concentration, line of best fit, and characterizing a reaction as endothermic.
An educational chemistry exam question on the reactivity of metals, including sections on molten compounds, electrolytic products, and calculations related to aluminium oxide.
An examination question page detailing chemical reactions and calculations, including equations, explanations of concepts like oxygen loss, atom economy, and reactivity of metals.
A diagram and text layout for an exam question on reactivity of metals, including pH levels and corresponding universal indicator colours. The layout includes sections labelled (a) to (f), asking for specific chemical reactions and items such as 'neutralisation' and 'burette'.
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Science Revision

Electrolysis Explained for GCSE Students

Definition of electrolysis

Electrolysis is the breaking down of an electrolyte using an applied electric current.

An electrolyte is an ionic substance that has been dissolved in a solvent or has been melted.

Electrolysis of molten compounds

The ionic compound is heated to a very high temperature until it reaches liquid state. The metal cations are positively charged metal ions. They will move towards the negatively charged electrode, called a cathode. These ions become metal atoms at the cathode.

The non-metal anions are negatively charged ions. They will move towards the positively charged electrode, called an anode. These ions become non-metal atoms at the anode.

Fig.1 – electrolysis of molten lead (II) bromide

Diagram showing the electrolysis of molten lead(II) bromide, including a DC power supply, positive lead ions attracted to the negative electrode, and negative bromide ions attracted to the positive electrode.

Electrolysis of aqueous solutions

The ionic compound is dissolved in water. In water, there are H+ and OH– ions existing and they will compete with the metal cations and non-metal anions.

Fig.2 – electrolysis of aqueous copper sulfate

Diagram illustrating a copper electrolysis setup showing an anode, cathode, and copper(II) sulphate solution, with reactions for oxygen gas production and copper metal deposition.

Half-equations explained

To figure out what is produced at the cathode (negative electrode) and at the anode (positive electrode), follow the two rules below:

Rule 1: At the cathode, check the reactivity of the metal ion. If it is more reactive than hydrogen, the hydrogen ions are reduced to form a hydrogen molecule.

Chart showing metals arranged by reactivity, from most reactive potassium at the top to least reactive gold at the bottom.

Rule 2: At the anode, the anions are oxidised in the following order:

Halide ion > hydroxide ion > all other negatively charged ions

Table showing negative ions from electrolytes and the corresponding elements released at the anode.

If there are no halide ions present to be oxidised at the anion, then the hydroxide ions are next to be oxidised and the following half-equation describes what happens to the hydroxide ion that is oxidised at the anode:

4OH– → 2H2O + O2 + 4e–

Practice questions

A worksheet on electrolysis and metal extraction, featuring multiple choice and table completion questions about molten substances, electrode products, and balancing a chemical equation.
A chemistry exam question focusing on the calculation of the relative formula mass of aluminium oxide and a diagram illustrating the reactivity series of metals, including potassium, lithium, carbon, zinc, tin, and gold, with instructions for predicting extraction methods.

Source for Questions papers and Answers:

AQA GCSE Chemistry Topic 4: Chemical Changes Revision – PMT

Answers to practice questions:

A mark scheme for an exam question on electrolysis, detailing the movement of ions, products at electrodes, and calculations related to electrochemical processes.
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Science Revision

Photosynthesis Explained Step-by-Step (GCSE Biology)

What is photosynthesis?

Photosynthesis is a chemical process by which plants convert carbon dioxide and water into glucose and oxygen using light energy from the Sun. Photosynthesis occurs in the chloroplasts found in plant cells.

The photosynthesis equation

Here is the chemical equation for photosynthesis:

Limiting factors

A limiting factor is a factor that affects the rate of photosynthesis of a plant.

Limiting factors in photosynthesis are:

  • light intensity
  • the concentration of carbon dioxide
  • the temperature of its surroundings.

Example GCSE exam question

Source:

4.1 Photosynthesis (F) QP.pdf

Source:

4.1 Photosynthesis (F) MS.pdf

Link to more papers using GCSE AQA Biology:

AQA GCSE Biology Topic 4: Bioenergetics Revision – PMT

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Science Revision

Understanding Chemical Equations

What chemical equations represent

Scientists use symbols in chemical equations to show what reactant(s) and product(s) are involved in a chemical reaction, what direction the reaction proceeds in, what physical states the reactant(s) and product(s) are in and what reaction conditions are involved (e.g. temperature, presence of a catalyst, time etc). This helps scientists to understand chemical equations and how to apply the equation practically if they want to carry out the reaction to obtain a product (or group of products) for commercial and research purposes.

Step-by-step balancing method with chemical equation example 1

Step 1) Here is an example of the combustion of methane, CH4 , which is an unbalanced equation:

CH4​(g) + …O2​(g) → CO2​(g)+ …H2​O(l)

(g) means the physical state of the chemical is in a gaseous state while (l) means the physical state of the chemical is in a liquid state. So methane, oxygen and carbon dioxide are in a gaseous state while water is in a liquid state.

The small number 4, which is called subscript 4, from CH4 , means there are four Hydrogen atoms covalently bonded to a Carbon atom.

Subscript 2 from O2 means there are two Oxygen atoms covalently bonded to each other.

Subscript 2 from CO2 means there are two Oxygen atoms each covalently bonded to a Carbon atom.

Subscript 2 from H2O means there are two Hydrogen atoms each covalently bonded to an Oxygen atom.

Step 2)

Next step I would do to balance the chemical equation is I would list the type of atoms and their number on the left hand side (LHS) of the equation and on the right hand side (RHS) of the equation.

CH4​(g) + …O2​(g) → CO2​(g)+ …H2​O(l)

Step 3)

I can see on the LHS there are 4 hydrogen atoms while on the RHS there are 2 hydrogen atoms. We can’t change subscript 2 on H2​O to subscript 4 due to the way 2 hydrogen atoms are each bonded to an oxygen atom, we can instead double the number of water molecules to get 4 hydrogen atoms on the RHS. When we double the number of water molecules on the RHS, we also increase the number oxygen atoms on the RHS as well to get 4 oxygen atoms.

Step 4)

The last step is to now double the number of oxygen atoms on the LHS to get a total of 4 oxygen atoms.

Step 5)

So now the the complete balanced equation is:

CH4​(g) + 2O2​(g) → CO2​(g)+ 2H2​O(l)

What this equation tells the scientist is that 1 mole of methane molecule will react with 2 moles of oxygen molecules to produce 1 mole of carbon dioxide and 2 moles of water molecules. If you don’t know what a mole is, click on the following link below:

What is a mole in chemistry explained simply

Congratulations! If you have read this far and followed the steps without any difficulty, you have fully understood balancing the chemical equation for combustion of methane.

References:

Combustion of hydrocarbon fuels – Polluting the atmosphere – AQA – GCSE Chemistry (Single Science) Revision – AQA – BBC Bitesize

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Science Revision

What is a mole in chemistry explained simply

What is a mole?

Ever wondered what a mole is? A mole seems complicated but

1 mole = 6.02214076 × 1023 atoms, molecules, compounds or ions

Mole explained in simple terms

Think of 1 mole as a way to group a very large amount of atoms/ ions/ molecules/ compounds similar to how we group people e.g. 1 class = 30 students , 1 football team (on the pitch) = 11 players.

Mole and its relationship to atomic mass unit of an element

If you look at the Periodic Table for an element, for example Iron (symbol Fe), you will notice the number 55.845 which is its atomic mass number; atomic mass number is g/mol units and that means that if you want 1 mole of Iron, you will need 55.845 g of it.

Fig.1

Fig. 2

Fig. 3

Let’s take a look at another example which is Carbon-12. Carbon-12 (which has an atomic mass of 12 g/mol) helps to provide a concrete example of 1 mole because 12 grams of carbon-12 is equal to 1 mole of carbon atoms (think of coal as that is made mostly of carbon). Other substances will have different masses to make up 1 mole e.g. because a sodium atom has an atomic mass unit of 22.99 g/mol , you would need 22.99 g of sodium metal to make 1 mole of sodium atoms. Another example is calcium which has an atomic mass unit of 40.08 g/mol so you would need 40 g of calcium metal to make 1 mole of calcium atoms.

Fig. 4 – Picture showing 12g of carbon measured which is equivalent to 1 mole of carbon-12 atoms

Example of a past paper question involving moles

Let’s have a look at a past paper question example from PhysicsAndMathsTutor.com based on a GCSE AQA Chemistry past paper:

Below is an example Mark Scheme of how to solve it:

So in Step 1 of the mark scheme, we want to find out the number of molecules from 1 mole of carbon atoms, not from 70 carbon molecules so first step is to calculate 1÷70 = 0.0142857 moles.

Step 2:

if 1 mole = 6.02214076 × 1023 atoms

then

0.0142857 x 1 moles = 0.0142857 x 6.02214076 × 1023 atoms

0.0142857 moles = 8.6 x 1021 atoms