Atomic Structure
Fundamental Particles, Isotopes and Relative Atomic Mass
OxfordAQA International AS and A-level Chemistry
The model of the atom
An atom is a nucleus containing protons and neutrons, surrounded by electrons. That picture has not always been the accepted one: understanding of atomic structure has evolved over time.
The fundamental particles
| Particle | Relative charge | Relative mass |
|---|---|---|
| Proton | +1 | 1 |
| Neutron | 0 | 1 |
| Electron | −1 | 1/1836 |
Write the relative mass of the electron as 1/1836 or as negligible. Never write 0 for the relative mass of an electron.
A proton is 1836 times heavier than an electron, so on a scale where the electron is given a relative mass of 1 instead, the proton is 1836 and the neutron is 1838.
Counting protons, neutrons and electrons
The mass number, A, is the number of protons plus neutrons in the nucleus of an atom. It is a number, not a mass and not an average. The atomic number, Z, is the number of protons.
- Protons = Z, and neutrons = A minus Z, in an atom and in an ion alike.
- Electrons = Z minus the charge, so a positive ion has fewer electrons than protons and a negative ion has more.
| Species | Protons | Neutrons | Electrons |
|---|---|---|---|
| 7634Se | 34 | 42 | 34 |
| 7734Se2− | 34 | 43 | 36 |
| 6529Cu+ | 29 | 36 | 28 |
Where extra protons and neutrons are added to a known atom, the new proton count gives the element and its symbol, and the protons and neutrons together give the mass number. Write the symbol and the mass number, and never an atomic number that does not belong to the symbol beside it.
The Periodic Table is in order of atomic number, not relative atomic mass: each element has one more proton than the one before it.
Isotopes
- Isotopes are atoms of one element with the same number of protons and a different number of neutrons.
- They react in the same way because they have the same electron configuration.
- They have the same atomic radius because they have the same number of protons and the same number of electrons.
How an atom reacts depends on its electrons alone, so for two isotopes reacting in the same way write same electron configuration and nothing else: protons and neutrons are not part of that answer. Size depends on the pull of the nucleus on those electrons, so for the same atomic radius the answer is both halves, same number of protons and same number of electrons. Never write that the isotope with more neutrons has the larger atomic radius: a neutron is uncharged, so it makes no difference to the size of the atom.
Relative atomic mass
Relative atomic mass is the average (mean) mass of one atom of the element divided by one twelfth of the mass of one atom of carbon-12. Relative isotopic mass is the same comparison made for one atom of a single isotope, rather than for the average across all of them.
The top line and the bottom line must describe the same kind of thing. Never write the mass of one atom over one twelfth of the mass of one mole of carbon-12: one line is then a single atom and the other a mole, and the two cannot be compared. The twelve belongs on one line only, so multiplying the top line by 12 goes with a bottom line of the mass of one atom of carbon-12, and never with a top line multiplied by 12 with one twelfth underneath.
Calculating relative atomic mass from abundances
- The denominator is the total abundance, which is 100 only where the abundances are percentages adding up to 100. Add the abundance column up as it stands, whether it holds percentages, relative intensities or the terms of a ratio, and divide by that total.
- Give Ar to one decimal place. Never write a relative atomic mass rounded to a whole number.
Working backwards to a missing abundance or mass number
- Find any abundance that is missing, by subtracting the abundances given from the total.
- Call the unknown abundance x, and write the other one as the total minus x.
- Write the Ar expression in x, set it equal to the Ar given, and solve for x.
- Where the unknown is a mass number, round the answer to the nearest integer.
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