The reactivity of these metals increases as you move down a group. Therefore, reactivity increases from Aluminium to Magnesium to Sodium. This is because reactivity increases as the outermost electron is further from the nucleus and therefore less tightly held. Sodium (Na) has an electronic configuration of 2, 8, 1. It readily loses this single electron to form a positive ion (Na+), making it highly reactive. Magnesium (Mg) has an electronic configuration of 2, 8, 2. It has two valence electrons, which are held more tightly than a single electron, so it is less reactive than Sodium. Aluminium (Al) has an electronic configuration of 2, 8, 3. It has three valence electrons, which are held even more tightly than two, making it the least reactive of the three. The larger atomic radius of the metals as you go down the group also contributes to lower reactivity, as the valence electrons are further from the nucleus and less strongly attracted.

The charge of an ion formed from an element in Group 1 is always +1 because these elements have only one valence electron in their outer shell. To achieve a stable noble gas configuration, they readily lose this single electron. This results in the formation of a positive ion with a +1 charge. The process can be illustrated as follows:

Diagram:

Explanation of Diagram: The diagram shows a sodium atom (Na) with its electron configuration (1s² 2s² 2p⁶ 3s¹). It depicts the loss of the 3s¹ electron to form a sodium ion (Na⁺) with a 2s² 2p⁶ electron configuration. The loss of one electron results in a +1 charge. The element readily loses this electron due to its low ionization energy.

The metallic character of elements generally increases as you move down a group. This is because the atomic size increases. Here's the explanation:

• Atomic Size: As you move down a group, the number of electron shells increases. This means the outermost electrons are further from the nucleus, resulting in a larger atomic radius.
• Ionization Energy: With a larger atomic radius, the valence electrons are further from the nucleus and are less tightly held. Therefore, it requires less energy to remove an electron, meaning the ionization energy decreases.
• Metallic Bonding: A larger atomic radius means the delocalised electrons are further from the nucleus and are less strongly attracted to the nucleus. This results in a weaker metallic bond. However, the increased number of valence electrons also contributes to the delocalisation of electrons, leading to a greater overall metallic character.

In essence, the increased atomic size and decreased ionization energy down a group facilitate the delocalisation of electrons, leading to a stronger metallic bond and increased metallic character.