Revision notes for OCR GCSE Chemistry Atomic structure. Open the guide for explanations and worked examples. Written against the OCR GCSE Chemistry (J248) specification, so the content matches what's examinable rather than general Chemistry background.
Atomic structure
What you'll learn
- How scientists’ ideas about atoms changed from Dalton to Bohr.
- What protons, neutrons and electrons are, including their charges and masses.
- How tiny atoms are, and why most of an atom is empty space.
- How to calculate numbers of protons, neutrons and electrons in atoms, isotopes and ions.
Starting point: what is an atom?
An atom is the smallest part of an element that still has that element’s chemical properties. An element is a substance made from only one type of atom.
A molecule is a group of atoms joined together. For example, oxygen gas, O₂, is made from oxygen molecules, and each molecule contains two oxygen atoms.
Common Mistake
Atoms and molecules are not the same thing
An atom is one particle of an element. A molecule contains two or more atoms joined together. Some elements exist as molecules, such as O₂ and N₂, but they are still made from atoms.
Models of the atom change over time
A scientific model is a simplified idea or picture used to explain observations. Models are useful, but they can change when new evidence is found.
Definition
Atomic model
An atomic model is a scientific explanation of what atoms are like inside. It has changed over time because experiments have provided new evidence.
The story of atomic structure is a great example of science improving by testing ideas. Each model was useful for its time, but later evidence showed it needed changing.
Dalton: solid sphere model
John Dalton described atoms as tiny, solid spheres. In this model, atoms were indivisible, meaning they could not be split into smaller parts.
This was a good starting model, but later scientists discovered atoms contain smaller particles.
Thomson: plum pudding model
J. J. Thomson discovered the electron, a tiny negatively charged particle. This showed atoms were not indivisible.
Thomson suggested the plum pudding model: a positively charged sphere with negative electrons embedded in it, like raisins in a pudding.
Geiger and Marsden: alpha particle scattering
Hans Geiger and Ernest Marsden fired positively charged alpha particles at very thin gold foil. Most alpha particles passed straight through, some were deflected, and a very small number bounced back.
That result was surprising. If the plum pudding model were correct, the positive charge would be spread out, so alpha particles should not be strongly deflected.
Example
Explaining alpha particle scattering
- Most alpha particles passed straight through the gold foil, so most of the atom must be empty space.
- Some alpha particles were deflected, so there must be a concentrated positive region that repels positive alpha particles.
- A tiny number bounced back, so that positive region must be very small, very dense, and contain most of the atom’s mass.
Rutherford and Bohr
Ernest Rutherford used the alpha scattering evidence to suggest the nuclear model: atoms have a tiny, dense, positively charged nucleus at the centre, with electrons around it.
Niels Bohr improved the model by suggesting electrons are arranged in fixed shells or energy levels around the nucleus. This is the GCSE model you use for many chemistry explanations.
Key Idea
Why the atomic model changed
The atomic model changed because new experimental evidence could not be explained by older models. Better evidence led to better models.
The structure of the atom today
The GCSE model of an atom has a central nucleus surrounded by electrons. The nucleus contains protons and neutrons.
Definition
Subatomic particles
Subatomic particles are particles smaller than an atom. The three you need here are protons, neutrons and electrons.
Protons, neutrons and electrons
A proton is a positively charged particle found in the nucleus.
A neutron is a neutral particle found in the nucleus. Neutral means it has no charge.
An electron is a negatively charged particle found outside the nucleus, arranged in shells.
For GCSE, learn these relative charges and approximate relative masses:
- Proton: relative charge +1; relative mass about 1.
- Neutron: relative charge 0; relative mass about 1.
- Electron: relative charge -1; relative mass about 11836\frac{1}{1836}18361, so it is often treated as 0 when finding mass number.
Because protons and neutrons have much more mass than electrons, almost all the mass of an atom is in the nucleus.
Key Idea
Charge and mass in an atom
The nucleus is positive because it contains protons. The atom’s mass is mostly in the nucleus because protons and neutrons are much heavier than electrons.
How small is an atom?
Atoms are extremely small. Typical atomic radii and bond lengths are about 1×10−101 \times 10^{-10}1×10−10 metres (m). This is 0.1 nanometres (nm), because 1 nm is 1×10−91 \times 10^{-9}1×10−9 m.
A small molecule is also on this tiny scale, though molecules can be larger depending on how many atoms they contain.
The nucleus is much smaller than the whole atom. Diagrams often draw the nucleus far too large so you can see it.
Tip
Diagram scale
Atomic diagrams are not usually drawn to scale. If the atom were drawn at a realistic scale, the nucleus would be a tiny dot with lots of empty space around it.
As a separate Chemistry J248 student, it is worth being comfortable with rough scale estimates using powers of ten.
Example
Comparing atom and nucleus size
Suppose an atom has a radius of 1×10−101 \times 10^{-10}1×10−10 m and its nucleus has a radius of about 1×10−141 \times 10^{-14}1×10−14 m.
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Compare the two radii by dividing the atom radius by the nucleus radius.
1×10−10 m1×10−14 m\frac{1 \times 10^{-10}\ \text{m}}{1 \times 10^{-14}\ \text{m}}1×10−14 m1×10−10 m -
Divide the powers of ten by subtracting the indices: −10−(−14)=4-10 - (-14) = 4−10−(−14)=4.
1×10−101×10−14=1×104\frac{1 \times 10^{-10}}{1 \times 10^{-14}} = 1 \times 10^{4}1×10−141×10−10=1×104 -
So the atom’s radius is about 10,000 times larger than the nucleus radius.
Atomic number and mass number
Every element is defined by its number of protons. If the number of protons changes, it becomes a different element.
Definition
Atomic number
The atomic number is the number of protons in the nucleus of an atom. It is often shown using the symbol ZZZ.
In a neutral atom, the number of electrons equals the number of protons. This is because the positive charges from protons cancel out the negative charges from electrons.
Definition
Mass number
The mass number is the total number of protons and neutrons in the nucleus. It is often shown using the symbol AAA.
So:
number of neutrons=mass number−atomic number\text{number of neutrons} = \text{mass number} - \text{atomic number}number of neutrons=mass number−atomic numberStandard nuclear notation shows an atom like this:
ZAX {}^{A}_{Z}\text{X} ZAXHere, X is the chemical symbol, AAA is the mass number, and ZZZ is the atomic number.
For example, 612C{}^{12}_{6}\text{C}612C is carbon with mass number 12 and atomic number 6.
Isotopes
Definition
Isotope
Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons.
For example, carbon-12 and carbon-14 are isotopes. Both are carbon, so both have 6 protons. Carbon-12 has 6 neutrons, while carbon-14 has 8 neutrons.
Common Mistake
Neutral does not mean protons equal neutrons
A neutral atom has equal numbers of protons and electrons. It does not have to have equal numbers of protons and neutrons.
Ions
An ion is a charged particle formed when an atom or group of atoms gains or loses electrons.
Definition
Ion
An ion is an atom or group of atoms with an overall electric charge because it has gained or lost electrons.
If an atom loses electrons, it becomes a positive ion because it has more protons than electrons.
If an atom gains electrons, it becomes a negative ion because it has more electrons than protons.
Common Mistake
Ions form by electron changes
Atoms become ions by gaining or losing electrons, not by changing the number of protons. Changing the number of protons would change the element.
Calculating particles in atoms and ions
To calculate subatomic particles:
- Protons = atomic number.
- Neutrons = mass number − atomic number.
- Electrons = protons for a neutral atom.
- For positive ions, subtract electrons.
- For negative ions, add electrons.
Example
Counting particles in an ion
Find the number of protons, neutrons and electrons in 1735Cl−{}^{35}_{17}\text{Cl}^{-}1735Cl−.
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The atomic number is 17, so the atom has 17 protons.
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The mass number is 35, so calculate neutrons:
35−17=1835 - 17 = 1835−17=18It has 18 neutrons.
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The ion has a -1 charge, so it has gained 1 electron compared with a neutral chlorine atom:
17+1=1817 + 1 = 1817+1=18It has 18 electrons.
Tip
Quick ion check
Positive ion: fewer electrons than protons. Negative ion: more electrons than protons.
Exam technique
In the exam
- For atomic structure calculations, identify the atomic number first: this gives the number of protons.
- Use mass number − atomic number to find neutrons; do not assume protons and neutrons are equal.
- For ions, change only the number of electrons: subtract electrons for positive ions and add electrons for negative ions.
Self review
Check yourself
- Why did Rutherford’s nuclear model replace Thomson’s plum pudding model?
- What are the relative charges and approximate relative masses of protons, neutrons and electrons?
- How many protons, neutrons and electrons are in 1224Mg2+{}^{24}_{12}\text{Mg}^{2+}1224Mg2+?
Recap questions
Test yourself with 5 quick questions on this guide. Answer them all correctly to complete it.
Particles
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