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Constituents of the atom
Three particles, two numbers and one notation carry the whole of atomic bookkeeping. Add the idea of specific charge and you can tell any nucleus, ion or particle apart with a single division.
Builds on SI units and prefixes.
IN THIS TOPIC
- Quote the charge and mass of the proton, neutron and electron in SI and relative units.
- Calculate the specific charge of particles, nuclei and ions.
- Use nuclide notation with Z and A, and explain what isotopes are and why isotopic data is useful.
WHAT YOU PROBABLY THINK
Atoms of the same element are identical.
Three particles, two units each
The simple model of the atom needs three players. The proton: charge +1.60 × 10−19 C, mass 1.67 × 10−27 kg. The neutron: no charge, mass 1.67 × 10−27 kg, a shade heavier than the proton. The electron: charge −1.60 × 10−19 C, mass 9.11 × 10−31 kg, nearly two thousand times lighter.
Alongside the SI values sit relative units: charges of +1, 0 and −1 (in units of e), masses of 1, 1 and about 1/1800. Relative units keep nuclear equations tidy; the SI values do the calculating. Quote whichever the question speaks in, and be ready to translate.
Specific charge
The specific charge of a particle is its charge divided by its mass:
in coulombs per kilogram. It works for anything: a lone particle, a bare nucleus, an ion. For a nucleus, Q is the proton count times e and m is (very nearly) the nucleon count times the nucleon mass; for an ion, count the electrons that remain. The electron holds the record, about 1.76 × 1011 C kg−1, because nothing else packs a whole unit of charge into so little mass. Exam questions love specific charge precisely because it audits both of your counts at once.
The notation
A nuclide is written with the nucleon number A (protons plus neutrons) as a leading superscript and the proton number Z as a leading subscript on the element symbol. Z is the element's identity: change Z and you have changed element. The neutron count is never written, because it is always A − Z. One housekeeping note: the atomic mass unit belongs to the Year 13 Nuclear unit, so masses here stay in kilograms.
Isotopes
Isotopes are atoms with the same Z but different A: the same element, the same chemistry, a different neutron count and so a different nuclear character. That last difference is what makes isotopic data useful: the ratio of carbon-14 to carbon-12 in dead material falls with age and dates archaeological finds, and the isotopic signature of a sample can identify where it came from. Chemistry cannot tell isotopes apart; their nuclei can.
THE EXAM BIT
- Relative charge and mass answer one kind of question, SI values another. Read which is asked for; a charge of “+1 C” for a proton is wrong by nineteen orders of magnitude.
- Specific charge is charge over mass, with both counted honestly: for an ion, subtract or add the electrons; for a nucleus, there are none.
- The neutron's specific charge is zero. It has mass but no charge, and stating that in one line is a full-mark answer.
- In nuclide notation, top is A, bottom is Z, and neutrons are A − Z. The examiners' favourite trap is asking for neutrons and marking the subtraction.
- Isotopes: same proton number, different nucleon number. Both halves of the sentence are required, every time.
CHECK YOURSELF
Calculate the specific charge of an iron-56 nucleus (Z = 26). Take e = 1.60 × 10−19 C and the nucleon mass as 1.67 × 10−27 kg.
Show a hint
Charge counts protons only; mass counts every nucleon.
Show the answer
Charge: Q = 26 × 1.60 × 10−19 = 4.16 × 10−18 C.
Mass: m = 56 × 1.67 × 10−27 = 9.35 × 10−26 kg.
Specific charge: Q/m = 4.4 × 107 C kg−1. Smaller than a proton's alone, because thirty neutrons added mass without adding any charge.
Z is the identity, A is the headcount.
Specific charge audits both at once.
No animated video for this topic yet; these notes stand alone. InkPhysics on YouTube.