CC3 · Atomic structure · Topic 1 — Key concepts in chemistry | GCSE Edexcel Combined Science

Revise the key ideas

Dalton modelled atoms as tiny spheres and proposed that elements contain characteristic atoms. Later evidence showed atoms have smaller particles and isotopes of an element can have different masses.
J. J. Thomson discovered the negatively charged electron. His plum-pudding model placed electrons within a spread-out region of positive charge.
In the alpha-scattering experiment, most alpha particles passed straight through thin gold foil. Some changed direction (were deflected), and very few bounced back. This did not fit a model with positive charge spread throughout the atom.
Most particles passing through suggested the atom is mostly empty space. Strong deflections suggested a small, dense, positively charged nucleus containing most of the mass.
Evidence supported a small positive nucleus, replacing plum pudding.
Later developments established protons, neutrons and electrons occupying energy levels around the nucleus. New evidence can change a scientific model; a model is not a literal photograph.
Atoms are roughly 10⁻¹⁰ m across; nuclei are much smaller, roughly 10⁻¹⁵ to 10⁻¹⁴ m. Simple diagrams hugely exaggerate nuclear size and do not represent electrons as real orbiting planets.

Protons have relative charge +1 and relative mass 1. Neutrons have charge 0 and relative mass 1. Both are in the nucleus.
Electrons have relative charge −1 and relative mass about 1/1836, usually treated as negligible in mass-number calculations. They occupy shells or energy levels around the nucleus.
Protons and neutrons are nuclear; electrons occupy surrounding energy levels.
A neutral atom has equal numbers of protons and electrons. Charges cancel; neutrons do not supply negative charge.
Atomic number, Z, is the number of protons and identifies the element. Every carbon atom has six protons, whatever its isotope.
Mass number, A, is the total number of protons and neutrons in one nucleus. It is always a whole number. Relative atomic mass is different: it is an average that takes account of the element’s isotopes.
In notation ²³₁₁Na, the upper number is mass number 23 and lower number atomic number 11: 11 protons, 12 neutrons, and 11 electrons for the neutral atom.
Mass number belongs to one isotope, not a weighted isotope average.

Number of neutrons = mass number − atomic number. For carbon-14: 14 − 6 = 8 neutrons.
For a neutral atom, number of electrons = atomic number. Always check whether the question instead describes an ion.
An ion forms when electrons are lost or gained; its proton count remains the same. Changing protons would change the element.
A sodium ion, Na⁺, has 11 protons but 10 electrons. A chloride ion, Cl⁻, has 17 protons but 18 electrons.
For a positive ion, subtract its charge number from the neutral atom’s electron count: a 2+ ion has lost two electrons. For a negative ion, add the charge number: a 2− ion has gained two. Forming these ions does not change the neutron count.
If an atom has 13 protons and 14 neutrons, its mass number is 27 and the element is aluminium. Its neutral electron count is 13.

Isotopes are atoms of the same element with the same proton number but different neutron numbers, so their mass numbers differ.
Carbon-12, carbon-13 and carbon-14 each have six protons but six, seven and eight neutrons respectively. They are all carbon.
Same element, different neutron numbers.
Neutral isotopes of an element have the same number of electrons and electronic arrangement, giving very similar chemical behaviour.
Isotopes can differ in physical properties because their masses differ; some isotopes are radioactive, but not all are.
Isotopes are not ions: isotope differences concern neutrons, while ordinary ion differences concern electrons.
An element can occur as a naturally occurring mixture of isotopes. Their proportions affect the relative atomic mass in the periodic table.

Relative atomic mass, Ar, is the weighted mean mass of atoms of an element compared with one twelfth of the mass of a carbon-12 atom. It has no unit.
For isotope percentages, Ar = sum of (isotope mass × percentage abundance) ÷ 100. At GCSE, use mass numbers as isotope masses when supplied that way.
For chlorine with 75% chlorine-35 and 25% chlorine-37: Ar = (35 × 75 + 37 × 25) ÷ 100 = 35.5.
The result is nearer 35 because chlorine-35 is more abundant.
For relative abundances rather than percentages, divide the weighted total by the total abundance. A 3:1 ratio gives (35 × 3 + 37 × 1) ÷ 4 = 35.5.
An Ar that is not a whole number does not mean an atom has part of a neutron. It is an average of the masses of different isotopes, taking account of how common each isotope is.
Check that your answer lies between the isotope masses and nearer the more abundant isotope. An ordinary unweighted average works only when abundances are equal.