Trend across Group 1: The first ionisation energy decreases across Group 1. This is because the atomic radius increases across a period. As the atomic radius increases, the valence electrons are further from the nucleus and shielded by more inner electrons. This means the electrostatic attraction between the nucleus and the valence electrons is weaker, making it easier to remove an electron and therefore lowering the first ionisation energy.

Trend down Group 2: The first ionisation energy increases down Group 2. This is because the atomic radius increases down a group. As the atomic radius increases, the valence electrons are further from the nucleus, but the effective nuclear charge experienced by the valence electrons remains the same. However, the increased distance between the nucleus and the valence electrons means the electrostatic attraction is weaker, so it requires more energy to remove an electron. Therefore, the first ionisation energy increases down the group.

Elements within the same group (vertical column) of the periodic table have the same number of valence electrons – electrons in the outermost shell. The chemical properties of an element are primarily determined by the number and arrangement of its valence electrons, which dictate how it interacts with other atoms.

For example, Group 1 elements (alkali metals) all have one valence electron in an s orbital (e.g., Group 1: [Noble Gas] ns¹). This single electron is easily lost, leading to a strong tendency to form +1 ions and react readily with non-metals. Group 17 elements (halogens) have seven valence electrons (e.g., Group 17: [Noble Gas] ns²np⁵). They readily gain one electron to achieve a stable noble gas configuration, forming -1 ions and reacting readily with metals.

The similar electronic configurations within a group result in similar reactivity patterns, ionization energies, and electronegativities. The shielding effect of inner electrons is relatively constant across a group, meaning the valence electrons experience a similar effective nuclear charge. This leads to similar bond-forming tendencies.

The periodic table is extremely useful for predicting the properties of elements because it shows a clear relationship between an element's position and its chemical and physical properties. This relationship arises from the recurring patterns in electron configurations.

Properties that can be predicted include:

• Metallic Character: Elements on the left side of the periodic table (metals) tend to be more metallic, meaning they readily lose electrons. As you move across a period, metallic character decreases. Elements on the right side of the table (non-metals) are generally non-metallic.
• Reactivity: Metals in Group 1 (alkali metals) are highly reactive, while metals in Group 2 (alkaline earth metals) are less reactive. Non-metals are generally unreactive.
• Ionization Energy: The ionization energy (the energy required to remove an electron) generally increases across a period and decreases down a group. This is because the effective nuclear charge increases across a period, making it harder to remove an electron.
• Atomic Radius: Atomic radius generally decreases across a period and increases down a group. This is due to increasing effective nuclear charge across a period and increasing shielding by inner electrons down a group.
• Electronegativity: Electronegativity generally increases across a period and decreases down a group. This is related to the ability of an atom to attract electrons in a chemical bond.

The periodic table helps with this prediction by showing the group an element belongs to. For example, knowing an element is in Group 1 tells you it's an alkali metal and therefore likely to be highly reactive and readily lose an electron. Knowing an element is in Period 3 tells you it has 3 electron shells, which influences its chemical properties.