3.1.1 Atomic Structure

Protons

A positively charged sub-atomic particle

Charge = +1

Relative mass = +1

Neutron

A neutral charged sub-atomic particle

Charge = 0

Relative mass = +1

Electron

A negatively charged sub-atomic particle

Charge = -1

Relative mass = 1/1836

Determines an element's chemical properties

Nucleons

Sub-atomic particles found in the nucleus (protons and neutrons)

Mass number

The total number of protons and neutrons in the nucleus of an atom

Atomic number

The number of protons in the nucleus of the atom

It identifies the element

Ion

A charged particle formed when one or more electrons are lost or gained by an atom

Isotope

Atoms of the same element that have different numbers of neutrons but the same number of protons and electrons

Relative atomic mass (Ar)

The weighted average of the masses of its isotopes relative to 1/12 of the mass of a carbon-12 atom

Relative isotopic mass

The mass of the isotope relative to 1/12 of the mass of a carbon-12 atom

Relative formula mass (Mr)

The sum of all relative atomic masses of the atoms making up a compound

For both ionic and covalent substances, unlike relative molecular mass which is only referred to when specifying covalent compounds

Time of flight mass spectrometer

Orbital

A region of a sub-level that contains a maximum of 2 electrons, as long as they have an opposite spin

S orbitals

Sphere-shaped with 1 orbital

Hence, it can hold 2 electrons

P orbitals

Dumbbell-shaped with 3 orbitals

Hence, it can hold 6 electrons

D orbitals

Clover-shaped with 5 orbitals

Hence, it can hold 10 electrons

Sub-shell notation

Ne = 1s² 2s² 2p⁶

Coefficient = energy level/shell

Letter = sub-shell

Squared number = number of electrons

Magnesium atom vs ion notation

Mg = 1s² 2s² 2p⁶ 3s²

Mg²⁺ = 1s² 2s² 2p⁶

Chlorine atom vs ion notation

Cl = 1s² 2s² 2p⁶ 3s² 3p⁵

Cl⁻ = 1s² 2s² 2p⁶ 3s² 3p⁶

First ionisation energy

The energy needed to remove the outermost electron from each atom in one mole of gaseous atoms to form one mole of gaseous positive ions

First ionisation of oxygen equation

O₍g₎ → O⁺₍g₎ + e⁻

Ionisation energy = +1314 kJmol⁻¹

Rules for ionisation energy equations

You must always use the gas symbol, as ionisation energies are measured for gaseous atoms

Always refer to 1 mole of atoms

The lower the ionisation energy, the easier it is to form a positive ion

Factors affecting ionisation energies - nuclear charge

Nuclear charge; the more protons there are in the nucleus, the more positively charged the nucleus is, hence a stronger attraction for the electrons, equalling a higher ionisation energy

Factors affecting ionisation energies - distance from the nucleus

Distance from the nucleus; an electron close to the nucleus will be much more strongly attracted to the nucleus than the one further away

Factors affecting ionisation energies - shielding

Shielding; as the number of electrons between the outer electron and the nucleus increases, the outer electrons feel less attraction to the nucleus

Successive ionisation energy equations

A⁽ⁿ⁻¹⁾⁺₍g₎ → A⁽ⁿ⁾⁺₍g₎ + e⁻

(e.g. O⁴⁺₍g₎ → O⁵⁺₍g₎ + e⁻)

Ionisation trends down Group 2 (Be → Ba)

As each element going down Group 2 has one more electron shell than the previous, the shielding from the extra shell would reduce the attraction of the outer electrons to the nucleus, resulting in a lower ionisation energy ass you go down the group

Ionisation trends across periods (Na → Ar)

As you move across a period, the ionisation energy increases, given that the number of protons is increasing, hence a stronger nuclear attraction

Reason for a drop in ionisation energy across periods (Na → Ar)

Mg = 1s² 2s² 2p⁶ 3s² 1ˢᵗ ionisation energy = +738kJmol⁻¹

Al = 1s² 2s² 2p⁶ 3s² 3p¹ 1ˢᵗ ionisation energy = +578kJmol⁻¹

Aluminium's outer electron is in a 3p orbital, which is further from the nucleus, hence the energy needed to remove it is lower

Moreover, the 3p orbital has more shielding provided by the 3s electrons, resulting in less attraction between the nucleus and the outer electron, making it easier to lose

Evidence for electron orbitals

The trend across periods, specifically Na → Ar, provides evidence for electron sub-shells