Module 3: Chapter 8 (Reactivity Trends)

Alkaline Earth Metals

• reactive metals • never found in their native state • compounds are common throughout nature, e.g. calcium carbonate

• Called a reducing agent as it reduces another species, gains the two outer electrons from the group 2 element, forms a 2+ group 2 ion

• forms a metal oxide, MO • For example, 2Mg + O2 -> 2MgO (solid) • Magnesium is oxidised and Oxygen is reduced

• forms an alkaline hydroxide, M(OH)2, and hydrogen gas • reaction becomes more vigorous down group (more reactive) • for example: Sr + 2H2O -> Sr(OH)2

metal + acid -> salt + hydrogen (MASH) • reactivity increases down the group • e.g. Mg + 2HCl -> MgCl2 + H2

• Ionisation energies decrease down the group because of increasing atomic radius and increasing shielding meaning less nuclear attraction for outer electron and lower ionisation energy • Elements in group 2 react by losing 2 electrons (forming +2 ions), that is two ionisation energies in terms of energy input

• react with water to release OH- ions forming an alkaline solution of a metal hydroxide • CaO (solid) + H2O -> Ca2+ + 2OH- • the group 2 hydroxides are sparingly soluble in water, if the solution is saturated, a solid ppt will form: Ca2+ + 2OH- -> Ca(OH)2

• increases down the group • more soluble = more OH- ions in solution = more alkaline • for example Magnesium hydroxide forms a solution of a low concentration of pH 10 and barium forms a more concentrated solution, pH 13

Agriculture: Ca(OH)2 added to fields to increase the pH of acidic soils: Ca(OH)2 + 2H+ -> Ca2+ + 2H2O Medicine: • Acid in stomach is mainly HCl, can be neutralised • often used as antacids (treating acid indigestion) • Mainly Magnesium and Calcium carbonates used

Halogens

• most reactive non-metal • only occur as halide ions in nature (never in elemental form) • in elemental form, exist as diatomic molecules, X2 • structure is simple molecular

Fluorine = pale yellow gas Chlorine = pale green gas Bromine = an orange-red volatile liquid Iodine = grey-black crystalline solid

• increases down the group • more electrons • stronger London forces • more energy required to break intermolecular forces

• outer shell configuration: s2p5 • gaining one electron achieves noble gas electronic configurations, so redox reactions are the most common reactions for halogens • e.g Cl2 +2e- -> 2Cl- (half equation), oxidising agents as the chlorine is reduced

• reactivity decreases down the group • if a more reactive halogen is added to less reactive halide a displacement reaction will take place (halogen atom displaces ion in solution)

Solution in water: Cl = pale green Br = orange I = brown top layer colour with cyclohexane: Cl = pale green Br = orange I = violet

• increasing atomic radius • increased shielding from more inner shells means less nuclear attraction to capture an electron from another species • reactivity decreases down the group/weaker oxidising agent

A redox reaction in which an element is oxidised and reduced

Cl2 (g) -> H2O (l) -> HClO (aq) + HCl HClO = chloric (1) acid, kills bacteria and acts a weak bleach - Chlorine 0 to +1 and -1 (disproportionation) - Chlorine is used to treat drinking water by killing harmful bacteria, the equation above is the reaction of Cl being added to water

• dilute NaOH will allow more chlorine to dissolve (solubility is lower in water) Cl2 (g) + 2NaOH (aq) -> NaClO (aq) + NaCl (aq) + H20 (l) NaClO (sodium chlorate) is a much stronger bleach, large concentration of ClO- ions, chlorate ions - disproportionation of Cl2, 0 to +1 (in NaClO) and -1

• a toxic gas • will react with methane(/organic hydrocarbons) in water, methane is formed by decaying vegetation, -> chlorinated hydrocarbons are suspected carcinogens • without chlorine in our water we would be at risk of disease such as typhoid and cholera

• white crystalline solids • very soluble in water forming colourless solutions • insoluble in cyclohexane

• acids react with carbonates to form a salt, water and carbon dioxide gas (effervescence) • CO2 produced can be detected by bubbling the gas through lime water, a white precipitate produced (turned limewater cloudy white), calcium carbonate