S3.1.4 Trends in chemical reactivity Notes

Reactivity Trends in Group 1 (Alkali Metals) and Group 17 (Halogens)

Why Does Reactivity Increase Down Group 1?

1. The alkali metals are renowned for their high reactivity, which increases as you move down the group.
2. This trend can be explained by examining ionization energy, the energy required to remove the outermost electron from an atom.
3. Ionization energy decreases down the group.
   1. As you move down Group 1, the atomic radius increases.
   2. Each successive element has an additional electron shell, increasing the distance between the nucleus and the valence electron.
   3. This greater distance weakens the electrostatic attraction between the nucleus and the outermost electron.
   4. Additionally, inner electron shells shield the valence electron from the nucleus’s pull, further reducing ionization energy.
4. Lower ionization energy means easier electron loss.
   1. Alkali metals react by losing their single valence electron to form a positive ion (cation).
   2. As it becomes easier to lose this electron, the reactivity of the metal increases.

Applications and Implications of Alkali Metal Reactivity

The high reactivity of alkali metals has significant practical applications but also demands careful handling:

1. Industrial Use: Sodium and potassium are used in heat transfer systems and the production of compounds like fertilizers.
2. Safety Concerns: Alkali metals must be stored under oil or in inert atmospheres to prevent reactions with moisture or oxygen in the air.

Reactivity of Group 17 Elements (Halogens)

Why Does Reactivity Decrease Down Group 17?

1. Halogens are highly reactive non-metals, but their reactivity decreases as you move down the group.
2. This trend is tied to their ability to gain an electron to form a negative ion (anion).
3. Electron affinity decreases down the group.
   1. Halogens react by gaining one electron to achieve a stable octet.
   2. As the atomic radius increases down the group, the added electron experiences weaker attraction to the nucleus due to increased shielding by inner electrons.
   3. This makes it less energetically favorable for halogens lower in the group (like iodine) to gain an electron compared to those higher up (like fluorine).
4. Fluorine is the most reactive halogen.
   1. Fluorine’s small atomic radius and high nuclear charge make it extremely effective at attracting electrons.
   2. In contrast, iodine’s larger size and weaker nuclear attraction result in lower reactivity.

Applications and Implications of Halogen Reactivity

Halogens play critical roles in various industries and biological systems:

1. Disinfection: Chlorine is widely used to disinfect water supplies due to its strong oxidizing ability.
2. Organic Chemistry: Halogens are key reactants in the synthesis of pharmaceuticals and polymers.
3. Biological Importance: Iodine is essential for thyroid function in humans.

Comparing Reactivity Trends in Groups 1 and 17

The contrasting reactivity trends in Groups 1 and 17 provide a fascinating look at how atomic structure influences chemical behavior:

1. Group 1: Reactivity increases down the group as ionization energy decreases.
2. Group 17: Reactivity decreases down the group as electron affinity decreases.

Changes in Metallic Character Between Group 1 and Group 17

1. This property decreases as you move from Group 1 (alkali metals) to Group 17 (halogens) across the periodic table.
2. Group 1 (Highly Metallic):
   1. Alkali metals have low ionization energies and large atomic radii, making it easy for them to lose their single valence electron and exhibit strong metallic behavior.
3. Group 17 (Non-Metallic):
   1. Halogens have high ionization energies, smaller atomic radii, and a stronger effective nuclear charge, making them more likely to gain electrons and behave as non-metals.
4. The shift from highly reactive metals in Group 1 to highly reactive non-metals in Group 17 highlights the decreasing tendency to lose electrons across a period, driven by increasing nuclear attraction and reduced atomic size.