Chemistry · Ch 8 — Elements of Group 1 and 2
Chemical properties of elements of group 1 and group 2
Chemical properties of elements of group 1 and group 2
Both the alkali metals and the alkaline earth metals are highly reactive, which is exactly why they are always found in nature already combined with other elements rather than free. Their reactivity stems from their generally low ionization enthalpies; within each group, reactivity increases going down as atomic radius grows and ionization enthalpy correspondingly falls. Six characteristic reaction classes illustrate this reactivity — reaction with oxygen/air, with water, with hydrogen, with halogens, the metals' reducing behaviour (measured via standard electrode potential), and their distinctive ability to dissolve in liquid ammonia — each detailed as a separate note against this section, since between them they cover most of these two groups' defining chemistry. In every case the alkali metals react somewhat more vigorously, and via somewhat simpler ionic products, than …
Worked out. Group 1: on exposure to air, alkali metals rapidly lose their lustre as an oxide, peroxide or (in some cases) superoxide layer forms. Lithium gives only the simple oxide, 2Li + O2 → 2Li2O (lithium oxide); sodium gives the peroxide, 2Na + O2 → Na2O2 (sodium peroxide); potassium gives the superoxide, K + O2 → KO2 (potassium superoxide). Potassium superoxide can absorb CO2 while releasing O2 at the same time, 4KO2 + 2CO2 → 2K2CO3 + 3O2, which is why KO2 is used in breathing equipment for mountaineers, submarines and space applications. Group 1 oxides are strongly basic and dissolve in water to give strong alkalis, e.g. 2Li2O(s) + H2O(l) → 2LiOH(aq). The reaction of Na and K with oxygen is highly exothermic and both metals catch fire on exposure to air. …
Worked out. Group 1: lithium, sodium and potassium all float on water (buoyed by the hydrogen bubbles released). Lithium reacts only slowly, but sodium and potassium react vigorously and, because the reaction is strongly exothermic, catch fire when placed in water: 2Na + 2H2O → 2NaOH + H2. …
Worked out. Group 1: alkali metals react with hydrogen at high temperature to form ionic (saline) hydrides, 2M + H2 → 2M⁺H⁻ (e.g. 2Na + H2 → 2NaH).
Group 2: every group 2 metal except beryllium forms an MH2-type hydride on heating with hydrogen, M + H2 → MH2. …
Worked out. Worked problem: NaCl is an ionic compound, but LiCl has some covalent character — explain. Solution: the Li⁺ ion is very small, giving it a high charge density and a strong ability to distort (polarize) the electron cloud of the large, easily-deformed chloride ion — this electron-cloud distortion is what gives the Li–Cl bond its partial covalent character. Na⁺ is much larger than Li⁺ and so cannot distort the chloride ion's electron cloud in the same way, leaving NaCl a straightforwardly ionic …
Worked out. Group 1: all alkali metals react vigorously with halogens to give ionic halide salts of general formula MX, 2M + X2 → 2M⁺X⁻ (e.g. 2Na + Cl2 → 2NaCl).
Group 2: all alkaline earth metals combine with halogens at high temperature to form the corresponding halides of general formula MX2, M + X2 → MX2 (e.g. Ca + Cl2 → CaCl2). As with the other reaction classes in this section, the alkali and alkaline earth metals behave as strongly electropositive metals reacting with strongly electronegative non-metal halogens to give simple ionic salts, though — as section 8.2.5 discusses — the smallest members, lithium and beryllium, deviate from purel …
Worked out. The reducing power of a group 1 or 2 metal is measured by its standard electrode potential E° for M⁺(aq) + e⁻ → M(s) (group 1) or the equivalent two-electron process (group 2). All alkali metals have high negative E° values, showing they are strong reducing agents; within the group, lithium is the single most powerful reducing agent and sodium the least powerful (see Table 8.4). All alkaline earth metals likewise have high negative E° values (Table 8.5) and are strong reducing agents, though on the whole slightly weaker reductan …
Worked out. Group 1: alkali metals dissolve in liquid ammonia to give deep-blue, electrically conducting solutions, M + (x+y)NH3 → [M(NH3)x]⁺ + [e(NH3)y]⁻; the free 'ammoniated electron' is what gives the deep blue colour, and the solutions are paramagnetic. On standing, these solutions slowly liberate hydrogen and form the metal amide, turning from blue to bronze and becoming diamagnetic: M⁺(am) + e⁻(am) + NH3(l) → MNH2(am) + H2(g) (here '(am)' denotes a species dissolved in ammonia). …