AQA GCSE Combined Chemistry Paper 1: Atomic Structure Revision

AQA GCSE Combined Chemistry Paper 1: Atomic Structure Revision

Beyond Science explores Topic 1 from the AQA GCSE Combined Chemistry syllabus, providing demystified atomic structure revision for students. This blog breaks down atomic structure revision, detailing:

  • Atoms
  • Elements
  • Isotopes
  • Relative atomic mass
  • Electronic structure
  • The history of the atom
  • Compounds
  • Chemical equations
  • Mixtures
  • The Perodic Table

Are you ready? Let’s get started on some atomic structure revision…

Atomic Structure Revision


Atoms are very small. They have a radius of about 1 Γ— 10-10m.

Almost all of the mass of an atom is in the centre (nucleus). Contained in the nucleus are subatomic particles called protons and neutrons. Smaller particles called electrons are found orbiting the nucleus in shells or energy levels. Electrons are negatively charged. 

ParticleRelative MassCharge
ElectronVery small-1

Overall, atoms have no charge; the number of protons is equal to the number of electrons. 


Elements are substances made of only one type of atom. Elements can be represented as  symbols: 

  • N = Nitrogen 
  • F = Fluorine 
  • Zn = Zinc 
  • Ca = Calcium 
  • Na = Sodium

The symbols on the periodic table tell you some information about each element. 

Relative atomic mass & atomic number

The atomic number is the number of protons in an atom of the element. The relative atomic mass is the average mass of the element that takes into account the abundance of the isotopes of the element.


Isotopes are atoms of the same element with different numbers of neutrons. They have the same atomic number but a different relative atomic mass. The number of neutrons in an atom of an element can be calculated by taking the atomic number away from the relative atomic mass. 


111 – 1 = 0

112 – 1 = 1

113 – 1 = 2

Relative Atomic Mass

To calculate the relative atomic mass of an element, we use the following equation: 

Relative atomic mass

For example, chlorine has two isotopes. 75% of these are 35Cl and 25% are 37Cl. 

Relative atomic mass

Therefore, the relative atomic mass of chlorine is 35.5. 

Electronic Structure

Electrons are found in energy levels (shells). There is a maximum of two electrons in the innermost shell, up to eight in the second shell and up to eight in the third shell. 

For example, a sodium atom has an electronic structure of 2, 8, 1. This means there are two electrons in the first shell, eight in the second shell and one in the third shell. 

Remember: The lowest energy level (innermost shell) is filled first. Electrons can only begin to occupy the next energy level once the previous shell is full. 

History of the Atom

The model of the atom has developed over time as new experimental evidence has been discovered. 

ScientistDateAtomic Model
John DaltonStart of the 19th centuryAtoms were first described as tiny solid spheres that could not be divided. 
JJ Thompson1897The discovery of the electron led to the Plum Pudding Model. This suggested that the atom is a ball of charge with negative electrons embedded in it. 
Ernest Rutherford1909Alpha scattering experiment disproved the plum pudding model. The new nuclear model suggested that mass is concentrated at the centre (nucleus) and the nucleus is positively charged.
Niels BohrAround 1911Suggested that electrons orbit the nucleus at specific distances. 
James ChadwickAround 1940Experimental evidence showed that there are neutrons in the nucleus.


Compounds are substances made up of two or more different elements chemically bonded together. Examples of compounds are carbon dioxide and magnesium oxide. Compounds can be represented by chemical formulae. Carbon dioxide has the chemical formula CO2, which tells us that it is made up of one atom of carbon and two atoms of oxygen.

Some other examples of chemical formulae are:

  • NaCl
  • HCl
  • H2O
  • Na2O
  • H2SO4

The atoms in a compound are held together by chemical bonds and are difficult to separate.

Chemical Equations

A chemical reaction can be shown by using a word equation:  

magnesium + oxygen β†’  magnesium oxide  

On the left-hand side are the reactants and on the right-hand side are the products.  

This reaction can also be shown by a symbol equation:

2Mg + O2 β†’ 2MgO  

No atoms are lost or made during a chemical reaction so the mass of the products equals the mass of the reactants. Therefore, equations need to be balanced, so there are the same number of atoms on each side of the equation. To do this, numbers are put in front of the formulae. 

e.g. CH4 + O2  β†’  H2O + CO2 

This equation is unbalanced. There are four atoms of hydrogen on the left and only two atoms of hydrogen on the right, so we need to place a 2 in front of the formula H2O on the right-hand side. 

CH4 + O2  β†’  2H2O + CO2 

Now, there is an equal number of hydrogen atoms on both sides of the equation. However, there are now four atoms of oxygen on the right and only two atoms of oxygen on the left. If we place a 2 in front of the formula O2 on the left-hand side, the equation is now balanced. 

CH4 + 2O2 β†’ 2H2O + CO2 



  • A mixture consists of two or more substances not chemically combined together. In a mixture there are no chemical bonds between the different substances, so the substances are easy to separate. Examples of mixtures are air and salt water. Different methods can be used to separate different mixtures:


  • Used to separate an insoluble solid from a liquid, e.g. sand and water. 


  • Used to obtain a sample of pure salt from a solution.

Simple Distillation

  • Used to separate and collect a solvent from a solution, e.g. salt water. 

Fractional Distillation

  • Used to separate miscible liquids with different boiling points, for example separating crude oil into fractions.


  • Used to separate out mixtures of coloured dyes.

The Periodic Table

In the early 1800s, scientists attempted to arrange elements in order of their atomic weights. The early periodic tables were not complete because some of the elements had not been found, and some elements were placed in inappropriate groups. 

Dimitri Mendeleev (1869) left gaps in the periodic table and changed the order of some of the elements. 

The gaps show that he believed there were undiscovered elements. He was right! 

The discovery of isotopes explained why the order based on atomic weights was not always correct. 

The Modern Periodic Table

  • Elements are arranged in order of atomic number.
  • Elements are arranged in columns, known as groups.
  • Elements in the same group have the same number of electrons in their outer shell. This gives them similar chemical properties. 
  • The rows are called periods: the period number shows the number of electron shells.
Atomic Structure Revision: The periodic table of elements

Metals and Non-Metals

The majority of elements are metals. Metals are found on the left of the periodic table and non-metals are on the right.  


Metals are elements that react to form positive ions. They are typically strong, malleable and good conductors of electricity and heat. They are usually solids at room temperature, with the exception of mercury, which is a liquid. 


Non-metals usually form negative ions. They are generally dull, brittle and poor conductors of electricity. Most non-metals are gases at room temperature, but some are solid and one (bromine) is a liquid. 

Group 1: Alkali Metals

The alkali metals (Group 1 elements) are soft, very reactive metals. Atoms of these elements all have one electron in their outer shell, making them very reactive. Group 1 metals have a low density and low melting point compared to other metals. They include:

  • Lithium
  • Sodium
  • Potassium

They form ionic compounds with non-metals by losing the electron from their outer shell. As you go down the group, the elements become more reactive. This is because the number of shells increases and it becomes easier to lose the electron in the outer shell as it is further from the positive nucleus. 

Alkali metals react with water to produce metal hydroxides and hydrogen

potassium + water β†’  potassium hydroxide + hydrogen

2K + 2H2O β†’ 2KOH + H2

Alkali metals react with chlorine to produce metal chlorides. 

sodium + chlorine β†’  sodium chloride

2Na + Cl2 β†’ 2NaCl

Alkali metals react with oxygen to form metal oxides.

lithium + oxygen β†’ lithium oxide

4Li + O2 β†’ 2Li2O

Group 7: Halogens 

The halogens (Group 7 elements) are non-metals. Atoms of these elements all have seven electrons in their outer shell. They include:

  • Fluorine
  • Chlorine
  • Bromine
  • Iodine

Group 7 elements exist as molecules made of pairs of atoms. The atoms in a molecule are covalently bonded due to a shared pair of electrons. 

The halogens can form ionic compounds with metals by gaining an electron to complete their outer shell. As you go down the group, the elements become less reactive. It becomes harder for them to gain an extra electron because the outer shell is further away from the positive nucleus. The melting and boiling points of the halogens increase as you go down the group.  

An example of a reaction between a Group 1 element and a Group 7 element is shown below.

potassium + bromine β†’ potassium bromide

2K + Br2 β†’ 2KBr

A more reactive halogen can displace a less reactive halogen from an aqueous solution of its salt. 

For example, chlorine is more reactive than bromine, so chlorine can displace bromine from potassium bromide solution.

potassium bromide + chlorine β†’  potassium chloride + bromine

2KBr + Cl2 β†’ 2KCl + Br2

Group 0: Noble Gases

The noble gases (Group 0 elements) include:

  • Helium
  • Neon
  • Argon 

They are unreactive as they have full outer shells, which makes them very stable. They are all colourless gases at room temperature. 

The boiling points all increase as you go down the group – they have greater intermolecular forces because of the increase in the number of electrons.

Atomic Structure Revision from Beyond

Atomic Structure Revision from Beyond

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