4.2 Reactions and Compounds of Alkali and Alkaline Earth Metals

Properties and Trends

Characteristics of Alkali and Alkaline Earth Metals

Alkali metals (Group 1: Li, Na, K, Rb, Cs, Fr) and alkaline earth metals (Group 2: Be, Mg, Ca, Sr, Ba, Ra) are among the most reactive elements on the periodic table. Their reactivity stems from their electron configurations: alkali metals have a single valence electron ($ns^1$) and form $M^+$ ions, while alkaline earth metals have two valence electrons ($ns^2$) and form $M^{2+}$ ions. Losing these electrons to achieve a noble gas configuration requires relatively little energy.

Physically, these metals share several traits:

• Soft, silvery-white metals with low melting and boiling points (Na melts at just 97.8°C; compare that to a transition metal like Fe at 1538°C)
• Low densities, especially for the lighter members (Li, Na, and K are less dense than water)
• Excellent conductors of heat and electricity due to metallic bonding with delocalized electrons

Reactivity and Periodic Trends

Reactivity increases going down both groups. Two factors drive this: atomic radius increases and ionization energy decreases as you add electron shells. The valence electron sits farther from the nucleus and is easier to remove, so heavier metals react more readily.

• Alkali metals react vigorously with water, producing hydrogen gas and the metal hydroxide: $2M(s) + 2H_2O(l) \rightarrow 2MOH(aq) + H_2(g)$. Li reacts gently, Na reacts energetically, and K ignites the evolved hydrogen.
• Alkaline earth metals also react with water but less vigorously. Mg reacts very slowly with cold water (though readily with steam), while Ca reacts steadily with cold water.
• Electronegativity decreases down each group, making the heavier elements even more electropositive and their bonding more purely ionic.

Flame tests are a classic way to identify these metals. Each element emits a characteristic color when its electrons are excited: Na gives a persistent yellow flame, K gives lilac, Ca gives orange-red, Sr gives crimson, and Ba gives green.

Diagonal Relationships and Analytical Techniques

A diagonal relationship exists between certain elements in adjacent groups and periods that share surprisingly similar chemistry. The most important example is Li and Mg. Both have high charge densities (Li is small with +1; Mg is slightly larger but with +2), which leads to comparable behavior: both form nitrides directly with $N_2$, both form covalent organometallic compounds, and their carbonates decompose on heating (unlike the other alkali metal carbonates). A similar diagonal relationship connects Be with Al.

For quantitative identification of these metals beyond flame tests:

• Atomic absorption spectroscopy (AAS) measures the absorption of light by free metal atoms in a flame or graphite furnace, giving precise concentration data
• X-ray fluorescence (XRF) spectroscopy provides non-destructive elemental analysis by measuring characteristic X-rays emitted from a sample

Oxides and Hydroxides

Oxides and Their Reactivity

The type of oxide an alkali metal forms depends on the metal and the oxygen supply. This is a point that trips people up, so pay attention to the pattern:

• Li forms the normal oxide $Li_2O$ (even with excess $O_2$)
• Na preferentially forms the peroxide $Na_2O_2$
• K, Rb, Cs form superoxides ($KO_2$, $RbO_2$, $CsO_2$)

The trend reflects the increasing size of the cation stabilizing progressively larger anions ($O^{2-} < O_2^{2-} < O_2^{-}$). Normal oxides of all these metals can be made with a limited oxygen supply, but the thermodynamically preferred product shifts as the cation grows.

Alkaline earth metal oxides are typically produced by:

1. Direct combination of the metal with $O_2$
2. Thermal decomposition of carbonates: $CaCO_3(s) \xrightarrow{\Delta} CaO(s) + CO_2(g)$

Both groups of oxides are basic and react with water to form hydroxides: $CaO(s) + H_2O(l) \rightarrow Ca(OH)_2(aq)$. The exception is BeO, which is amphoteric, dissolving in both acids and strong bases. This ties back to beryllium's high charge density and diagonal relationship with aluminum (whose oxide is also amphoteric).

Applications of these oxides include ceramics, glass production ($Na_2O$ as a flux), and catalysis.

Peroxides and Superoxides

• Peroxides contain the $O_2^{2-}$ ion (O-O single bond). $Na_2O_2$ is used in "oxygen candles" for submarines and spacecraft, reacting with $CO_2$ and moisture to regenerate $O_2$.
• Superoxides contain the $O_2^{-}$ ion (one-electron reduction of $O_2$). $KO_2$ is used in self-contained breathing apparatus: $4KO_2 + 2CO_2 \rightarrow 2K_2CO_3 + 3O_2$. This reaction both absorbs exhaled $CO_2$ and releases $O_2$.
• Hydrogen peroxide ($H_2O_2$) is related but distinct. It decomposes to water and oxygen ($2H_2O_2 \rightarrow 2H_2O + O_2$) and is widely used as a bleaching agent and disinfectant.

Formation and Properties of Hydroxides

Alkali metal hydroxides (NaOH, KOH) are strong bases that are highly soluble in water. They dissociate completely, making them useful wherever a concentrated source of $OH^-$ is needed.

Alkaline earth metal hydroxides show a solubility trend: $Be(OH)_2$ is essentially insoluble and amphoteric, $Mg(OH)_2$ is sparingly soluble (the basis of "milk of magnesia"), and solubility increases down the group through $Ca(OH)_2$, $Sr(OH)_2$, to $Ba(OH)_2$, which is moderately soluble and a strong base.

Key reactions and applications:

• Neutralization with acids: $NaOH(aq) + HCl(aq) \rightarrow NaCl(aq) + H_2O(l)$
• NaOH is used in soap production (saponification), paper manufacturing (Kraft process), and water treatment
• Industrial NaOH is produced primarily by the chloralkali process, the electrolysis of brine (concentrated NaCl solution)

Salts and Hydrides

Halides and Their Applications

Alkali and alkaline earth metals form ionic halides with the halogens. Most adopt high-symmetry crystal structures (NaCl adopts a face-centered cubic lattice). The ionic character of these halides generally increases with larger cation size and larger anion size, though small, high-charge-density cations like $Be^{2+}$ and $Li^+$ can introduce significant covalent character (Fajans' rules).

Notable halides and their uses:

• NaCl (table salt): essential for biological function, food preservation, and as feedstock for the chloralkali process
• $CaCl_2$: de-icing agent (exothermic dissolution) and drying agent in organic synthesis
• LiBr: employed in absorption refrigeration systems due to its highly hygroscopic nature
• KI: added to table salt as a dietary iodine supplement to prevent goiter

Hydrides and Their Reactivity

Metal hydrides contain the hydride ion ($H^-$), where hydrogen has gained an electron. The s-block metals form two types:

• Ionic (saline) hydrides: formed by the heavier alkali metals (Na, K, Rb, Cs) and the heavier alkaline earth metals (Ca, Sr, Ba). These are high-melting crystalline solids with NaCl-type structures.
• Covalent/polymeric hydrides: $BeH_2$ and $MgH_2$ have significant covalent character due to the high charge density of $Be^{2+}$ and $Mg^{2+}$. $LiH$ is sometimes grouped here as well, though it is borderline ionic.

All s-block hydrides react with water, releasing $H_2$:

$NaH(s) + H_2O(l) \rightarrow NaOH(aq) + H_2(g)$

This makes them powerful bases and reducing agents. NaH is widely used in organic synthesis as a strong, non-nucleophilic base. $CaH_2$ ("calcium hydride") is a convenient drying agent for organic solvents because it reacts with trace water but is easier to handle than Na metal.

Carbonates, Nitrates, and Sulfates

Carbonates form by reaction of metal oxides or hydroxides with $CO_2$. Their thermal stability is a high-yield topic:

• Alkali metal carbonates (except $Li_2CO_3$) are thermally stable and do not decompose at Bunsen burner temperatures. $Li_2CO_3$ decomposes because $Li^+$ is small enough to polarize the large $CO_3^{2-}$ ion.
• Alkaline earth metal carbonates all decompose on heating, with decomposition temperature increasing down the group: $MgCO_3$ decomposes around 350°C, while $BaCO_3$ requires about 1360°C. Again, smaller cations polarize the carbonate ion more effectively, destabilizing it.
• $Na_2CO_3$ (washing soda) is used in glass production and water softening.

Nitrates are all highly soluble. On heating, alkali metal nitrates decompose to nitrites and $O_2$ ($2KNO_3 \rightarrow 2KNO_2 + O_2$), except $LiNO_3$, which decomposes further to the oxide. Alkaline earth metal nitrates decompose directly to the oxide, $NO_2$, and $O_2$. Nitrates are used in fertilizers and pyrotechnics.

Sulfates show a solubility trend opposite to the hydroxides for Group 2: solubility decreases down the group. $MgSO_4$ (Epsom salt) is very soluble, $CaSO_4$ (gypsum) is sparingly soluble, and $BaSO_4$ is essentially insoluble. This insolubility is why $BaSO_4$ is safe to use as a contrast agent in X-ray imaging despite barium's toxicity: it passes through the GI tract without dissolving.