Atomic structure

Syllabus points

• Identify the relative charges and relative masses of a proton, a neutron and an electron.
• Deduce the numbers of protons, neutrons and electrons in atoms and ions from proton number and nucleon number.
• Describe the distribution of electrons in shells, sub-shells and orbitals.
• Write electronic configurations, including for ions and for the d-block elements.
• Define and explain first and successive ionisation energies, and the trends across a period and down a group.

The three particles

Everything in this topic follows from three numbers, so they are worth knowing exactly rather than approximately.

The electron's mass is so small that it is usually taken as negligible, which is why the nucleon number (protons + neutrons) is effectively the mass of the atom.

Two definitions that questions rely on:

• Proton number (Z) is the number of protons. It defines the element and never changes in a chemical reaction.
• Nucleon number (A) is protons + neutrons.

Isotopes are atoms of the same element with different numbers of neutrons. They have identical chemical properties, because chemistry is decided by electrons, and different physical properties such as density and rate of diffusion, because those depend on mass.

Working out an ion

For an ion, adjust the electrons and nothing else. In ³⁵Cl⁻, the proton number is 17 and the nucleon number is 35, so:

• protons = 17
• neutrons = 35 - 17 = 18
• electrons = 17 + 1 = 18, because the single negative charge means one electron has been gained

A very common slip is to change the neutron count when forming an ion. Ions form by gaining or losing electrons only.

Shells, sub-shells and orbitals

Electrons occupy shells (numbered 1, 2, 3, ...), each divided into sub-shells (s, p, d, f), each made of orbitals that hold at most two electrons.

An orbital is a region where there is a high probability of finding an electron. An s orbital is spherical; a p orbital is dumbbell-shaped, and the three p orbitals lie along the x, y and z axes at right angles to one another.

Sub-shells fill in order of increasing energy, and the order is not simply numerical:

1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p

The 4s sub-shell is filled before 3d because it is at a lower energy. That single fact is behind several exam favourites.

Writing configurations

Calcium (Z = 20): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²

Iron (Z = 26): 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s²

For ions of d-block elements, electrons are removed from the 4s sub-shell first, even though 4s was filled first. So Fe²⁺ is 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶, not 3d⁴ 4s².

Two elements break the simple pattern because a half-filled or fully filled d sub-shell is more stable:

• Chromium is 3d⁵ 4s¹, not 3d⁴ 4s²
• Copper is 3d¹⁰ 4s¹, not 3d⁹ 4s²

Ionisation energy

First ionisation energy is the energy needed to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions.

Every word of that definition earns marks: one mole, gaseous, and 1+ ions. The equation form is

X(g) → X⁺(g) + e⁻

Three factors decide the size of an ionisation energy:

• Nuclear charge. More protons pull the electrons in more strongly, so the energy is higher.
• Distance from the nucleus. An electron in a shell further out is held less strongly.
• Shielding. Inner shells of electrons repel the outer electron and reduce the pull it feels.

Across a period

First ionisation energy generally increases across a period: the nuclear charge rises while the electrons are added to the same shell, so shielding barely changes and the atomic radius falls.

Two dips break the trend, and both are examined constantly:

• Group 2 to Group 13 (for example Mg to Al). The outer electron in Al is in a 3p orbital, which is at a higher energy than the 3s orbital it would otherwise occupy, so it is easier to remove.
• Group 15 to Group 16 (for example P to S). In S the fourth p electron must pair up in an orbital that already holds one electron, and the repulsion between the pair makes it easier to remove.

Down a group

First ionisation energy decreases down a group. The nuclear charge does rise, but the outer electron is in a shell further from the nucleus and there is far more shielding from the extra inner shells, and those two effects win.

Successive ionisation energies

Removing each further electron takes more energy, because the ion is increasingly positive and holds the remaining electrons more tightly.

A large jump shows that an electron has been taken from a shell closer to the nucleus, and the position of that jump gives the group.

Suppose the first five ionisation energies of an element are 590, 1150, 4940, 6480 and 8120 kJ mol⁻¹. The big jump is between the second and third, so the element has two electrons in its outer shell and is in Group 2.

Common mistakes

• Changing the number of neutrons when an ion forms. Only electrons are gained or lost.
• Writing Fe²⁺ as 3d⁴ 4s². The 4s electrons go first, so it is 3d⁶.
• Leaving "gaseous" or "one mole" out of the ionisation energy definition.
• Saying the atomic radius increases across a period. It decreases, because the nuclear charge grows while shielding stays about the same.
• Explaining the Mg to Al dip with shielding. The reason is the sub-shell: the electron comes from 3p rather than 3s.
• Explaining the P to S dip with nuclear charge. The reason is paired electron repulsion within one p orbital.