H 1.1 Periodicity

Periodic table

A chart listing elements in order of increasing atomic number, arranged into groups and periods with recurring chemical properties.

Group

A vertical column in the periodic table; elements share similar outer-shell electron configurations and properties.

Period

A horizontal row in the periodic table; elements have the same number of electron shells and show gradual changes in properties.

Covalent radius

Half the distance between the centers of two covalently bonded atoms; a measure of atom size (units: pm or Å).

Shielding effect

Inner electron shells shield outer electrons from the full positive charge of the nucleus, reducing the effective nuclear attraction.

Nuclear charge

The total positive charge of the nucleus (number of protons); increases with atomic number.

First ionisation energy (IE1)

Energy required to remove one mole of electrons from one mole of free gaseous atoms; always endothermic.

Ionisation energy trend down a group

Decreases down a group due to

1. increasing shielding and larger atomic size.

2. Outer e⁻ is less strongly attracted to nucleus

Ionisation energy trend across a period

Increases across a period as

1. nuclear charge increases and

2. Outer e⁻ is more strongly attracted to nucleus

Electronegativity

A measure of the attraction an atom has for electrons in a bond; quantified on the Pauling scale.

Pauling scale

A scale for electronegativity values, with fluorine assigned the highest value (4.0) as a reference.

Electron affinity

Energy released when a neutral atom in the gaseous state gains an electron to form an anion.

Covalent molecular

Molecules held together by covalent bonds with discrete molecular units (e.g., H2, O2, N2).

Covalent network

A solid where atoms are linked by an extended network of covalent bonds (e.g., diamond, graphite, Si).

Metallic bonding

Bonding in metals where a lattice of positive ions is surrounded by a 'sea' of delocalised electrons.

Allotrope

Different structural forms of the same element. (eg. diamond, graphite…)

Monoatomic elements

Noble gases; exist as single atoms and are generally unreactive due to full outer electron shells.

Ionic radius

Radius of an ion; often differs from covalent radius; cations (positive ions) are typically smaller, anions (negative ions) slightly larger than their neutral atoms.

Density

Mass per unit volume; in a period, density of metallic elements tends to rise toward the centre due to mass increase and radius changes.

Ionisation equation (example)

General form for the first ionisation energy: X(g) → X+(g) + e−.

Covalent radius trends across periods

Decreases across a period due to increasing nuclear charge pulling outer electrons closer, with shielding constant.

Covalent radius trends down groups

Increases down a group as additional electron shells increase shielding → reduce the nuclear pull on outer electrons.

Noble gases and covalent radii

Noble gases do not form covalent bonds with themselves, so they have no covalent radii values.

Electronegativity trend down a group

Decreases down a group due to

1. increasing shielding and larger atomic size

2. Shared e⁻ are less strongly attracted to the nucleus.

Electronegativity trend across a period

Increases across a period as

1. nuclear charge increases and

2. Shared e⁻ are more strongly attracted to the nucleus

Which of the first 20 elements form molecules?

H₂, N₂, O₂, F₂, Cl₂, P₄, S₈, C₆₀ (carbon as fullerene)

Why does graphite conduct electricity?

Each carbon atom is bonded covalently to 3 other carbon atoms, meaning that there are some free electrons that can move freely and conduct electricity.

Ionic bond

An electrostatic attraction between positively and negatively charged ions.

Why the third ionisation energy of Magnesium is so much higher than the second ionisation energy?

1. 3rd electron is removed from the shell closer/full/more stable to the nucleus. AND

2. 3rd  is less shielded from the nuclear charge and more energy is required to remove it.