AQA Chemistry Paper 1: Chemical Changes GCSE Revision

AQA Chemistry Paper 1: Chemical Changes GCSE Revision

Beyond Science tackles chemical changes GCSE revision, as part of AQA Chemistry Paper 1. Here, we’ll explore every facet of chemical changes.

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The Reactivity Series

The reactivity series shows the reactivity of metals with other substances. This is determined by how easily they form positive ions by losing electrons. The higher the metal is in the reactivity series the more reactive it is. We can use a rhyme to remember the order of the metals.

  • Purple (Potassium)
  • Slime (Sodium)
  • Can (Calcium)
  • Make (Magnesium)
  • A (Aluminium)
  • Careless (Carbon)
  • Zebra (Zinc)
  • Insane (Iron)
  • Try (Tin)
  • Learning (Lead)
  • How (Hydrogen)
  • Camels (Copper)
  • Surprise (Silver)
  • Gorillas (Gold)

Reactions of Metals with Oxygen

Metals react with oxygen to make a metal oxide. As the metal gains oxygen this is known as an oxidation reaction. If we look at the following example of burning magnesium:

Magnesium + oxygen β†’ magnesium oxide

In terms of oxygen gain or loss, if a metal gains oxygen we call it an oxidation reaction, if a compound loses oxygen we call it a reduction reaction.

Reactions of Metals with Water

Metals, when reacted with water, produce a metal hydroxide and hydrogen:

If we take the example of lithium we can show this as:

lithium + water β†’  lithium hydroxide + hydrogen

To give this as a balanced equation gives us:

2Li + 2H2O β†’  2LiOH + H2

The more reactive the metal is, the faster the reaction. However, less reactive metals such as copper or iron do not react with water.

Reactions of Acids

Reactions of Metals with Dilute Acid:

Metals, when reacted with acids, produce a salt and hydrogen:

If we take the example of sodium and hydrochloric acid:

sodium + hydrochloric acid β†’ sodium chloride + hydrogen

We can write this as a balanced equation:

2Na + 2HCl β†’ 2NaCl + H2

Metals that are below hydrogen in the reactivity series do not react with dilute acids

The general formula for the reaction between an acid and a metal is:

Acid + Metal β†’  Salt + Hydrogen

For example:

hydrochloric acid + sodium β†’  sodium chloride + hydrogen

We can write this as a balanced equation:

2HCl + 2Na β†’  2NaCl + H2

Acids and Alkalis or Neutralisation

When an acid reacts with an alkali, a neutralisation reaction takes place and salt and water is produced.

The general formula for this kind of reaction is:

Acid + Alkali β†’ Salt + Water

For example:

hydrochloric acid + sodium hydroxide β†’ sodium chloride + water

We can write this as a balanced equation:

HCl + NaOH β†’ NaCl + H2O

Naming Salts

The first part of a salt’s name comes from the metal in the metal carbonate, oxide or hydroxide.

The second part comes from the acid that was used to make it.

The table below shows the salt produced by three different acids:

Acid usedSalt produced
hydrochloricchloride
nitricnitrate
sulfuricsulfate

So for example if we reacted magnesium with each of these acids we would produce:

  • Magnesium chloride
  • Magnesium nitrate
  • Magnesium sulfate

Redox Reactions

When metals react with acids, they undergo a redox reaction. A redox reaction occurs when both oxidation and reduction take place at the same time. A redox reaction occurs when electrons are transferred. We can remember the process of redox reactions through the following rhyme:

Oxidation

Is

Loss

Reduction

Is

Gain

The following reaction between calcium and dilute acid shows a redox reaction

Calcium atoms are oxidised to Ca2+ ions when they react with dilute acid

2H+ + Ca β†’ Ca2+ + H2 

The ionic equation can be further split into two half equations.

The calcium atoms lose electrons, They are oxidised by the hydrogen ions:

Ca -2e β†’ Ca2+ 

The hydrogen ions gain electrons. They are reduced by the calcium ions:

2H+ + 2e→ H2

Reactions with Bases

The general formula for the reaction between an acid and a metal oxide is:

Acid + Metal Oxide β†’ Salt + Water

The salt that is produced depends upon the acid and the metal ion, one example is the reaction between sulfuric acid and copper oxide shown below:

sulfuric acid + copper oxide β†’ copper sulfate + water

This can be shown as:

H2SO4 + CuO β†’ CuSO4 + H2O

Reactions with Carbonates

The general formula for the reaction between an acid and a metal carbonate is:

Acid + Metal Carbonate ——-> Salt + Water + Carbon Dioxide

Again the salt what is produced depends on the acid and the metal ion, one example is the reaction between hydrochloric acid and calcium carbonate shown below:

hydrochloric acid + calcium carbonate β†’ calcium chloride + water + carbon dioxide

This can be shown as:

2HCl+CaCO3–>CaCl+H2O+CO2

pH Scale

In aqueous solutions, acids produce H+ ions and alkalis produce OH ions.

Neutral solutions are pH7 and are neither acids nor alkalis.

Chemical Changes GCSE Revision: pH scale

Strong and Weak Acids

A strong acid completely ionise in a solution, for example:

HCl β†’ H+ + Cl

Hydrochloric acid is able to completely dissociate in solution to form hydrogen ions and chloride ions.

Other strong acids include:

  • Nitric acid (HNO3
  • Sulfuric acid (H2SO4)

Weak acids only partially dissociate, for example:

Ethanoic acid partially dissociates to form a hydrogen ion and an ethanoate ion. This is also reversible.

CH3COOH ⇄ CH3COO + H+

Electrolysis

Electrolysis is the splitting up of an ionic substance using electricity.

On setting up an electric circuit for electrolysis, two electrodes are required to be placed in the electrolyte (a molten or dissolved ionic compound).

The electrodes are conducting rods. One of the rods is connected to the positive terminal and the other to the negative terminal of a power supply.

The electrodes are inert (this means they do not react with the electrolyte) and are often made from graphite or platinum.)

During the process of electrolysis, opposites attract. The positively charged ions will be attracted towards the negatively charged electrode.

At the positive electrode (called the anode) negative ions lose electrons to become elements. This is an oxidation reaction.

At the negative electrode (called the cathode) positive ions gain electrons to become elements. This is a reduction reaction.

Electrolysis of Aqueous solutions

Chemical Changes GCSE Revision: Electrolysis

Gases may be given off or metals deposited at the electrodes. This is dependent on the reactivity of the elements involved.

If the metal is more reactive than hydrogen in the reactivity series, then hydrogen will be produced at the negative cathode.

At the positive anode, negatively charged ions lose electrons. This is called oxidation and we say that the ions have been oxidised.

At the negative cathode, positively charged ions gain electrons. This is called reduction and we say that the ions have been reduced.

Using Electrolysis to Extract Metals

Metals are extracted by electrolysis if the metal in question reacts with carbon or if it is too reactive to be extracted by reduction with carbon. During the extraction process, large quantities of energy are used to melt the compounds.

Aluminium is obtained by the process of electrolysis. Aluminium oxide has a high melting point and melting it would use a large amount of energy. This would increase the cost of the process, therefore molten cryolite is added to aluminium oxide to lower the melting point and therefore the cost.

Electrolysis of Molten Ionic Compounds – Lead Bromide

Lead bromide is an ionic substance. Ionic substances, when solid, are not able to conduct electricity. When molten or in a solution, the ions are free to move and carry a charge.

The positive lead ions are attracted towards the negative cathode at the same time as the negative bromide ions are attracted towards the positive anode.

Oxidation is the loss of electrons and reduction is the gaining of electrons. We represent what is happening at the electrode by using half equations.

The lead ions are attracted towards the negative electrode. When the lead ions (Pb2+) reach the cathode, each ion gains two electrons and becomes a neutral atom. We say that the lead ions have been reduced.

Pb2+–> 2e + Pb

The bromide ions are attracted towards the positive electrode. When the bromide ions (Br ) reach the anode, each ion loses one electron to become a neutral atom. Two bromine atoms are then able to bond together to form the covalent molecule Br2-.

2Br β†’ Br2 + 2e

Chemical Changes GCSE Revision Material from Beyond

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Bicarbonate of Soda and Vinegar Experiment Pack

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