Inorganic chemistry

Inorganic chemistry is the study of the properties and behaviour of inorganic compounds, including metals, minerals, and organometallic compounds. At its core is the Periodic Table, which organises elements based on their atomic structure and allows us to predict their chemical and physical properties through a concept known as periodicity.

Periodicity

Periodicity refers to the repeating patterns of properties observed in elements when they are arranged in order of increasing atomic number. Understanding these trends is crucial for predicting the behaviour of elements and their compounds.

Key periodic trends include:

1. Atomic Radius: This is half the distance between the nuclei of two bonded atoms. The atomic radius decreases across a period because the increasing number of protons (increasing nuclear charge) pulls the electrons in the same principal energy level closer. It increases down a group as each new period adds an additional electron shell, increasing the distance of the valence electrons from the nucleus and causing greater electron shielding.
2. First Ionisation Energy: This is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions (X(g) → X⁺(g) + e⁻). It generally increases across a period due to the increasing nuclear charge and decreasing atomic radius, which hold electrons more tightly. It decreases down a group because the outermost electron is further from the nucleus and more shielded, making it easier to remove.
3. Electronegativity: This is a measure of the ability of an atom to attract a bonding pair of electrons in a covalent bond. It increases across a period and decreases down a group, following similar reasoning to ionisation energy.

Group 2: The Alkaline Earth Metals

Group 2 elements (Be, Mg, Ca, Sr, Ba) are reactive metals that form +2 ions. Their reactivity increases down the group.

Reactions:

• With Water: The reaction produces a metal hydroxide and hydrogen gas. The vigour of the reaction increases down the group as the first two ionisation energies decrease. Magnesium reacts very slowly with cold water but readily with steam to form magnesium oxide and hydrogen: Mg(s) + H₂O(g) → MgO(s) + H₂(g). Calcium reacts steadily with cold water: Ca(s) + 2H₂O(l) → Ca(OH)₂(aq) + H₂(g).

Solubility Trends:

• Hydroxides: The solubility of Group 2 hydroxides, like Mg(OH)₂, increases down the group. Magnesium hydroxide is sparingly soluble, whereas barium hydroxide is much more soluble.
• Sulfates: The solubility of Group 2 sulfates, like BaSO₄, decreases down the group. Barium sulfate is famously insoluble and is used in medicine as a 'barium meal' to x-ray the digestive system, as it is opaque to x-rays and non-toxic due to its insolubility.

Group 17: The Halogens

The Halogens (F, Cl, Br, I) exist as diatomic molecules (e.g., Cl₂). They are powerful oxidising agents because they readily accept an electron to form a halide ion (X⁻).

Reactivity Trends:

• The reactivity of halogens decreases down the group. This is because as the atoms get larger, the outermost shell is further from the nucleus and more shielded, making it harder to attract an electron. Fluorine is the most reactive halogen, and iodine is the least.

Displacement Reactions: A more reactive halogen will displace a less reactive halide ion from its aqueous solution. This is a key demonstration of their reactivity trend. For example, chlorine is more reactive than bromine: Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq) The colourless solution turns orange/brown due to the formation of aqueous bromine. However, bromine cannot displace chloride ions from a solution of sodium chloride because it is less reactive.

Reactions of Halide Ions: While halogens are oxidising agents, their ions (halides) can act as reducing agents. The reducing power of halide ions increases down the group (I⁻ > Br⁻ > Cl⁻). This is because the outermost electron in an iodide ion is furthest from the nucleus and most easily lost. This trend can be observed in their reactions with a strong oxidising agent like concentrated sulfuric acid (H₂SO₄).

• NaCl shows no redox reaction, only an acid-base reaction producing HCl gas.
• NaBr reduces concentrated H₂SO₄ to sulfur dioxide (SO₂).
• NaI is a strong enough reducing agent to reduce H₂SO₄ further to sulfur (S) and even hydrogen sulfide (H₂S).