Category: Periodic Table

Going down Group VII, reactivity decreases while melting and boiling point increases.

Going down Group I, reactivity increases, while melting point decreases.

Elements in the same Period have the same number of electronic shells, but the proton number of the elements increase from left to right.

As number of protons increase, the attractive force of the positive protons on the negative electrons also increases, pulling the electronic shells closer together.

Hence atomic radius decreases from left to right in the same Period.

Group VII elements exist as diatomic molecules with simple molecular structure.

Going down Group VII, the molecular mass of the element increases.

An increasing molecular mass results in increasing intermolecular force of attraction.

The higher the intermolecular force of attraction, the higher the energy needed to break these bonds, hence higher melting point and boiling point.

Going down Group I, the number of electronic shells increases, and atomic radius increases. As the valence electrons get further from the nucleus, the attractive force between the positive protons in the nucleus and the negative valence electrons decreases.

Hence, going down Group I, the metallic bonding formed between the positive protons in the nucleus and the delocalized valence electrons becomes weaker.

The weaker the metallic bond, the lesser the energy needed to break these bonds, and the lower the melting point and boiling point of these elements.

– Metal piece darts about quickly around the surface of the water

– Effervescence observed (due to hydrogen gas produced in water)

– Hissing sound heard (due to the vigorous release of hydrogen gas)

– Metal piece becomes smaller in size

– Flame observed (for some metals)

– Lithium (no flame)

– Sodium (yellow flame)

– Potassium (lilac flame)

– pH of solution increases from pH7 to pH14 (due to the formation of a strong alkali)

From Na to Si, the melting points increase, with a sharp increase from Na to Mg, and Al to Si, reaching a maximum for Si.

From Si to P, there is a sharp decrease in melting point, followed by a small increase in melting point from P to S, and a decrease in melting point from S to Ar.

Na, Mg, Al all have metallic structure. Moving from Na to Mg to Al, the number of protons and electrons increase, so there is an increasing electrostatic force of attraction between the positive ions and delocalised electrons, pulling the electrons closer to the nucleus, making the metallic bonding stronger. An increasing amount of energy is needed to break the increasing strength of the metallic bonds during melting. This results in the increasing trend of melting points from Na to Al.

Silicon exists as a giant molecular structure. A lot more energy is needed to break the large number of covalent bonds in the giant molecular structure, resulting in Silicon’s very high melting point.

Phosphorus, sulfur, chlorine and argon are non-metals which have simple covalent structures, with weak intermolecular forces of attraction. These forces require less energy to break during melting, resulting in their relatively lower melting points.

Phosphorus exists as P₄ molecules, sulfur exists as S₈ molecules, chlorine exists as Cl₂ molecules and argon exists as individual atoms. The strength of the intermolecular forces of attraction decreases as the size of the molecule decreases, so melting points decrease from S to P to Cl to Ar. This explains the slight increase in melting point from phosphorus to sulfur, as sulfur molecules are bigger in mass than phosphorus molecules.

Going down Group VII, the number of electronic shells increases, and atomic radius increases. As the valence electrons get further from the nucleus, the attractive force between the positive protons in the nucleus and the negative valence electrons decreases.

Hence, going down Group VII, it becomes more difficult for the atom to gain a valence electron to form an ion, causing reactivity to decrease.

The atoms of Group O elements have a stable electronic configuration, hence these atoms have no tendency to gain or lose electrons to form ions, making them chemically inert.