Predicting Reactivity

• Lithium: Fizzes gently and skates slowly across the water's surface until it disappears.
• Sodium: Melts into a shiny molten ball, fizzes vigorously, and moves rapidly across the surface.
• Potassium: Reacts explosively, igniting the hydrogen gas with a distinctive lilac flame.

They also tarnish rapidly in the air to form metal oxides and react vigorously when heated in chlorine gas to form white chloride salts. The ease of ionization determines their reactivity, which relies entirely on their atomic structure.

Group 7: The Halogens

Why is chlorine used to kill bacteria in swimming pools, but iodine is safe enough to put directly on minor cuts? This difference in severity comes down to their position in Group 7. These elements are non-metals known as the halogens, and they exist as diatomic molecules (such as $\text{Cl}_2$ or $\text{Br}_2$). They all have seven outer shell electrons and react by gaining one electron to form a $1-$ ion, known as a halide.

For Group 7, the reactivity trend is the exact opposite of Group 1: reactivity decreases as you move down the group. Their oxidising power gets progressively weaker from fluorine down to astatine.

As you descend the group, the atoms get larger and their intermolecular forces increase, causing clear physical trends:

• Melting and boiling points increase.
• Room temperature states change from gas (fluorine, chlorine) to liquid (bromine) to solid (iodine, astatine).
• Colours darken: pale yellow gas (fluorine) to pale green gas (chlorine) to red-brown liquid (bromine) to a grey-black solid (iodine).

A more reactive halogen will always displace a less reactive halogen from an aqueous halide salt solution. For example, if chlorine gas is bubbled through a solution containing bromide ions, the chlorine displaces the bromine:

Explaining Trends: The Atomic Structure "Golden Chain"

Understanding atomic structure explains why the reactivity trends in Group 1 and Group 7 go in completely opposite directions. As you move down any group in the Periodic Table, two things happen: the atomic radius increases because each successive element has an extra electron shell, and nuclear shielding increases because there are more inner shells sitting between the nucleus and the outer electrons.

Even though the nuclear charge (the total number of protons) also increases down a group, the combined effect of the larger radius and extra shielding outweighs this higher charge. This "outweighing principle" results in a weaker net electrostatic attraction between the positive nucleus and the outer electrons.

• In Group 1: The weaker attraction means the single outer electron is lost more easily. Therefore, reactivity increases down the group.
• In Group 7: The weaker attraction means it is harder for the nucleus to attract and gain an incoming electron. Therefore, reactivity decreases down the group.

Predicting Unknown Elements

Chemists were able to predict the exact properties of missing elements years before they were actually discovered, simply by looking at the Periodic Table. In your exam, you will be expected to use known trends to forecast the chemical and physical behaviour of elements you may not have studied in class and justify those predictions using atomic structure.

Elements in the same group share similar chemical properties because they have the same number of outer shell electrons. Group 0 elements (the Noble Gases) are inert (unreactive) because they already possess a stable, full outer shell of electrons.

Worked Example

Based on the trends in Group 1 and Group 7, predict the physical state of Astatine at room temperature, and predict the observations if Caesium is dropped into a trough of water. Justify your prediction for Caesium using atomic structure.

Step 1: Use Group 7 trends for Astatine.

• Trend: Melting/boiling points increase down the group. Iodine (above Astatine) is a solid.
• Prediction: Astatine will be a solid at room temperature (likely black/dark grey).

Step 2: Predict and justify the reactivity for Caesium.

• Prediction: Caesium will react instantly and explosively upon contact with water, reacting far more violently than potassium.
• Justification: Caesium is further down Group 1 than potassium. It has more electron shells, meaning it has a larger atomic radius and more nuclear shielding. These factors outweigh the increase in nuclear charge, leading to a weaker electrostatic attraction between the positive nucleus and the outer electron. Therefore, the outer electron is lost more easily, resulting in higher reactivity.

Step 3: Construct the balanced symbol equation for Caesium.

• Group 1 reactions with water follow the 2:2:2:1 stoichiometry:
• $2\text{Cs(s)} + 2\text{H}_2\text{O(l)} \rightarrow 2\text{CsOH(aq)} + \text{H}_2\text{(g)}$