Trends in the Periodic Table

Atomic Radius

Atomic radius can be defined as the distance between the nucleus and the valence shell of an atom.

Atomic radius decreases across periods. The nuclear charge increases but the number of core electrons remains the same. Therefore the core charge increases, increasing the force of attraction of the valence electrons to the nucleus. Since the valence electrons are now closer to the nucleus, the atomic radius decreases.

Atomic radius increases down groups. Down groups, the number of shells in an atom increases, so the distance between the valence shell and the nucleus increases.

Electronegativity

Electronegativity can be thought of as the ability of an atom to attract other electrons. Atoms with higher electronegativities attract more electrons.

Electronegativity increases across periods. Since the atomic radius decreases across periods, and smaller atoms will attract electrons more easily, the electronegativity will increase across periods.

Electronegativity decreases down groups. Since the atomic radius increases down groups, and smaller atoms will attract electrons more easily, the electronegativity will decrease down groups.

Metallic character

Metallic character is defined as the chemical properties of metals. Normally this can be defined as a readiness to lose electrons.

Metallic character decreases across periods. This is because the number of electrons in the valence shell increases, meaning that atoms will more readily accept electrons rather than lose them.

Metallic character increases down groups. This is because the atomic radius increases down groups, meaning the valence electrons are easier to lose, increasing the readiness of the atom to lose valence electrons.

Ionisation energy

The first ionisation energy is defined as the amount of energy required to remove a single electron from the valence shell of an atom. The reaction is as follows:

$\text{X}+\text{energy}⟶{\text{X}}^{+}+{\text{e}}^{-}$

The n-th ionisation energy is the amount of energy required to remove an electron from the atom with a charge of (n-1). For example, the 2nd ionisation energy is as follows:

${\text{X}}^{+}+\text{energy}⟶{\text{X}}^{2+}+{\text{e}}^{-}$

Essentially the n-th ionisation energy can be defined as the amount of energy used in the equation:

${\text{X}}^{\left(\text{n}-1\right)+}+\text{energy}⟶{\text{X}}^{\text{n}+}+{\text{e}}^{-}$

The ionisation energy increases across periods. Since there is a higher nuclear charge attracting the valence electrons, the energy required to remove an electron increases.

The ionisation energy decreases down groups. The valence electrons are now further away from the nucleus since there are more shells, so there is a weaker force of attraction, meaning that the amount of energy required to remove one decreases.

Melting point across Period 3

Period 3 consists of the following elements:

$\text{Na},\text{Mg},\text{Al},\text{Si},\text{P},\text{S},\text{Cl},\text{Ar}$.

The first three are all metals, and the core charge increases from left to right, so the melting points increase from sodium through aluminium.

Silicon is covalent network, so it has the highest melting point.

Phosphorus, sulfur, chlorine and argon are all covalent molecular. They have chemical formulae ${\text{P}}_4^{},{\text{S}}_8^{},{\text{Cl}}_2^{},\text{Ar}$ respectively. Since the strength of a VDW bond is proportional to the number of electrons, argon has a lower melting point that chlorine, with a lower melting point than phosphorus, which is lower than sulfur.