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Until now we have treated particles as plain featureless circles. In this lesson we look inside the particle. The smallest particle of an element is an atom, and an atom is itself built from even smaller subatomic particles: protons, neutrons and electrons. Learning how these are arranged — a tiny, dense nucleus surrounded by electrons in shells — is the key to the whole of chemistry, because it explains why atoms are neutral, where almost all of an atom's mass sits, and just how empty an atom really is. This lesson is part of Topic C1 of your OCR Gateway Combined Science course.
By the end of this lesson you should be able to describe the structure of an atom, state the relative masses and charges of protons, neutrons and electrons, explain why atoms have no overall charge, and use standard form to compare the size of an atom with the size of its nucleus.
This lesson develops AO1 (recalling the sub-atomic particles with their relative masses and charges) and AO2 (applying standard-form arithmetic to compare the size of an atom with its nucleus).
An atom has two parts:
The nucleus is extremely small compared with the whole atom, yet it holds almost all of the atom's mass, because protons and neutrons are far heavier than electrons. The electrons take up the relatively vast space around the nucleus, which means most of an atom is empty space.
Exam Tip: A labelled atom needs the nucleus (containing protons and neutrons) at the centre and electrons in shells around it. State that the mass is concentrated in the nucleus and that most of the atom is empty space.
Protons, neutrons and electrons differ in two key ways: their mass and their charge. Because the real masses and charges are so tiny, we use relative values (comparing everything with a proton).
| Particle | Where it is | Relative mass | Relative charge |
|---|---|---|---|
| Proton | In the nucleus | 1 | +1 |
| Neutron | In the nucleus | 1 | 0 (neutral) |
| Electron | In shells around the nucleus | 18361 (≈ negligible) | −1 |
The facts to learn are:
Because protons and neutrons each have a relative mass of 1, while electrons have almost none, the mass of an atom is concentrated in its nucleus.
We use relative masses and charges rather than the true values because the real figures are far too small and awkward to handle — a single proton has a mass of roughly 1.7×10−27kg, a number that would be cumbersome in every calculation. By comparing each particle with a proton and calling the proton's mass 1, we get a tidy set of numbers (1, 1 and negligible) that captures everything we need. The relative charges work the same way: instead of quoting a charge in coulombs, we say the proton is +1 and the electron −1, because what matters chemically is simply that they are equal and opposite.
Exam Tip: Learn the table cold. The two facts tested most often are that the electron's mass is negligible (about 1/1836, not zero — but treated as zero for mass calculations) and that the neutron is neutral (charge 0). Mixing up neutron and proton charges is a common error.
A complete atom has no overall electric charge — it is neutral. This is because it contains equal numbers of protons and electrons:
For example, a carbon atom has 6 protons (6×(+1)=+6) and 6 electrons (6×(−1)=−6), giving a total charge of +6+(−6)=0. Neutrons make no difference to the charge because they are neutral.
That is why, in a neutral atom, the number of electrons always equals the number of protons. (If an atom gains or loses electrons it becomes a charged ion, which you meet in the next lesson and again in C2.)
It is worth pausing on why the charges are exactly equal and opposite. The charge on a single proton is the same size as the charge on a single electron — equal in magnitude but opposite in sign — so one proton is precisely balanced by one electron. As long as there are the same number of each, every positive charge is matched by a negative one, and the atom as a whole shows no charge to the outside world. This balance is the normal, stable state of an atom, and it is the starting point for understanding bonding: when atoms lose or gain electrons in C2, it is exactly this balance that is upset, creating charged ions.
A fluorine atom has 9 protons, 10 neutrons and 9 electrons. Show that it is neutral.
Step 1 — total positive charge from protons: 9×(+1)=+9.
Step 2 — total negative charge from electrons: 9×(−1)=−9.
Step 3 — neutrons contribute 0. Add up: +9+(−9)+0=0.
Answer: the charges cancel, so the atom is neutral.
The relative masses in the table also tell you where an atom's mass is stored. Because a proton and a neutron each have a relative mass of 1, the total mass of an atom is, to a very good approximation, simply the number of protons plus the number of neutrons. The electrons, at about 1/1836 each, contribute so little that they can be left out of the count entirely.
It is worth seeing this with a real example. A carbon atom has 6 protons, 6 neutrons and 6 electrons. Adding up the relative masses gives 6+6=12 from the nucleus, and only 6×18361≈0.003 from the electrons — utterly negligible next to 12. So the mass of a carbon atom is taken to be 12, all of it effectively in the nucleus. This is why, in the next lesson, the "mass number" of an atom is defined as the number of protons plus neutrons: the electrons simply do not weigh enough to matter.
The same reasoning explains a common source of confusion. Two atoms can have exactly the same mass even though one has more protons and the other more neutrons, because it is the total of protons and neutrons that fixes the mass. And a positive ion — an atom that has lost one or more electrons — has almost exactly the same mass as the neutral atom it came from, because losing an electron removes hardly any mass at all. Keeping the roles clear (protons and neutrons carry the mass; electrons carry almost none) makes both of these facts obvious rather than surprising.
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