Main-group chemistry I: hydrogen, the s-block and the boron and carbon groups

The principles of the first nine lessons now pay off. In this lesson and the next we survey the main-group elements group by group, and at each step you should be able to explain what you see with ideas you already have: atomic size and effective nuclear charge (lesson 2), lattice energy (lesson 3), VSEPR and multicentre bonding (lessons 4 and 7), acid–base character (lesson 8) and Frost diagrams (lesson 9). We start with hydrogen, move through the s-block metals of groups 1 and 2, and finish with groups 13 and 14, where the change from metal to nonmetal down and across the table is most dramatic.

Hydrogen

Hydrogen is the most abundant element in the universe and the third most abundant atom at the Earth's surface, mostly in water. It does not fit neatly into any group. Like the alkali metals it has one valence electron and forms H⁺; like the halogens it is one electron short of a noble-gas configuration and forms H⁻. Its ionization energy (1312 kJ/mol) is much higher than those of the alkali metals, and its electron affinity (−73 kJ/mol) much smaller in magnitude than those of the halogens, so its chemistry is mainly covalent.

Hydrogen has three isotopes: protium ¹H, deuterium ²H (D, discovered by Harold Urey in 1931) and radioactive tritium ³H (T, half-life 12.3 years). Because D has twice the mass of H, its compounds show the largest isotope effects of any element: D₂O boils at 101.4 °C, and C–D bonds break more slowly than C–H bonds, a kinetic isotope effect used to study reaction mechanisms. Heavy water is used as the moderator in CANDU nuclear power reactors because deuterium absorbs few neutrons.

Hydrides

The binary compounds of hydrogen fall into three broad classes that follow the electronegativity of the partner. Saline (ionic) hydrides form with the most electropositive metals, groups 1 and 2 except beryllium: NaH, CaH₂. They contain H⁻ ions in a rock-salt or similar lattice, react with water to release hydrogen (NaH + H₂O → NaOH + H₂) and are used as strong bases and drying agents. Molecular (covalent) hydrides form with the p-block elements: B₂H₆, CH₄, NH₃, H₂O, HF and their heavier analogues. Metallic (interstitial) hydrides form with many d-block metals: palladium absorbs hundreds of times its own volume of hydrogen to form PdHₓ, with H atoms in octahedral holes of the metal lattice, a property used to purify hydrogen.

Among the molecular hydrides, the boiling points of NH₃, H₂O and HF are anomalously high compared with the heavier hydrides of the same groups, because these molecules form hydrogen bonds. Water's boiling point, 100 °C, is more than 150 °C above what the trend from H₂Te, H₂Se and H₂S would predict. Hydrogen bonding is responsible for the open structure of ice (less dense than liquid water), for the double helix of DNA and for much of the behaviour of water as a solvent.

Most hydrogen today is made by steam reforming of natural gas (CH₄ + H₂O → CO + 3H₂, followed by the water–gas shift CO + H₂O → CO₂ + H₂), which releases carbon dioxide. Hydrogen made by electrolysis of water using renewable electricity produces no direct emissions and is being developed for steelmaking, ammonia synthesis and fuel cells, although it is still considerably more expensive. Most hydrogen is used to make ammonia (lesson 11) and to refine petroleum.

The alkali metals (group 1)

Lithium, sodium, potassium, rubidium and cesium are soft, silvery metals with a single ns¹ valence electron. They have the lowest ionization energies in their periods, form M⁺ ions almost exclusively, and are powerful reducing agents that react with water, faster as one goes down the group: lithium fizzes, sodium melts and darts across the surface, potassium ignites the hydrogen released, and cesium explodes. Their melting points fall down the group, from 180.5 °C for lithium to 28.4 °C for cesium, because the metallic bond weakens as the atoms grow and the single valence electron is spread over a larger volume.

One property bucks the trend. The standard potential of Li⁺/Li (−3.04 V) is more negative than those of Na⁺/Na (−2.71 V) and K⁺/K (−2.93 V), even though lithium has the highest ionization energy. The explanation is hydration energy: the small Li⁺ ion is very strongly hydrated (−520 kJ/mol, against −405 for Na⁺), which more than makes up for its higher ionization energy. This is why lithium is the anode material of choice for high-voltage batteries. The development of the lithium-ion battery, in which Li⁺ shuttles between a graphite anode and a metal oxide cathode such as LiCoO₂, earned John Goodenough, Stanley Whittingham and Akira Yoshino the Nobel Prize in Chemistry in 2019.

Alkali metals dissolve in liquid ammonia to give deep blue solutions containing solvated electrons, e⁻(NH₃)ₙ. At high concentration the solutions turn bronze and conduct like metals. The solvated electron is a powerful reducing agent used in the Birch reduction of aromatic rings. Large organic ligands can capture alkali metal ions: crown ethers, discovered by Charles Pedersen in 1967, and the three-dimensional cryptands of Jean-Marie Lehn bind Na⁺ or K⁺ selectively according to cavity size. Pedersen, Lehn and Donald Cram shared the Nobel Prize in Chemistry in 1987. With a cryptand to hold Na⁺, James Dye made salts containing the sodide anion Na⁻ in 1974, alkalides in which an alkali metal is the anion.

Lithium differs from its heavier congeners in ways that resemble magnesium, the element diagonally below and to the right. Both form normal oxides (Li₂O, MgO) on burning rather than peroxides or superoxides, both form nitrides directly from N₂ (Li₃N, Mg₃N₂), and both have carbonates that decompose easily to the oxide. This diagonal relationship, also seen between beryllium and aluminium and between boron and silicon, arises because moving one step right increases charge density and electronegativity while moving one step down decreases them, and the two changes roughly cancel.

The alkaline earth metals (group 2)

Beryllium, magnesium, calcium, strontium and barium have two ns² valence electrons and form M²⁺ ions. They are harder, denser and higher melting than the alkali metals, because two electrons per atom contribute to metallic bonding. Beryllium is an exception to the ionic chemistry of the group: Be²⁺ is so small and highly charged that its compounds are largely covalent. BeCl₂ is a chain polymer in the solid, with each beryllium bridged to its neighbours by chlorides, and a monomeric linear molecule in the vapour, an electron-deficient species with only four electrons around beryllium. Beryllium and its compounds are highly toxic when inhaled as dust.

Magnesium and calcium are essential to life. Magnesium sits at the centre of the chlorophyll molecule, where it is coordinated by four nitrogen atoms of a porphyrin-like ring, and Mg²⁺ binds ATP in every cell. Calcium phosphate in the form of hydroxyapatite, Ca₅(PO₄)₃OH, is the mineral of bones and teeth, and calcium carbonate forms limestone, marble, chalk and the shells of marine organisms. Organomagnesium halides, the Grignard reagents RMgX discovered by Victor Grignard (Nobel Prize 1912), are among the most widely used reagents in organic synthesis.

Trends in group 2 compounds follow the lattice and hydration arguments of lesson 3. The thermal stability of carbonates and nitrates increases down the group, the solubility of hydroxides increases (Mg(OH)₂ is nearly insoluble and used as an antacid, while Ba(OH)₂ is quite soluble), and the solubility of sulfates decreases. Flame tests distinguish the ions by the light emitted by excited atoms: lithium crimson, sodium yellow-orange, potassium lilac, calcium orange-red, strontium crimson and barium pale green, the colours of fireworks.

Group 13: boron, aluminium, gallium, indium and thallium

Group 13 shows the full change from nonmetal to metal. Boron is a hard, black, semiconducting metalloid whose crystal structures are built from B₁₂ icosahedra. Its compounds are covalent and electron deficient: BF₃ and BCl₃ are trigonal planar Lewis acids (lesson 8), and the boron hydrides are held together by 3c–2e bonds (lesson 7). Borax, Na₂B₄O₅(OH)₄·8H₂O, is the main ore; boric acid is a mild antiseptic and flame retardant; and borosilicate glass, which contains about 13 % B₂O₃, expands little on heating and is used for laboratory glassware. Boron nitride, BN, is isoelectronic with carbon: hexagonal BN has a layered structure like graphite and is a white lubricant and electrical insulator, and cubic BN, made at high pressure, is second only to diamond in hardness.

Aluminium is the most abundant metal in the Earth's crust. It is a reactive metal, but a thin, tough film of Al₂O₃ protects it from corrosion, which is why it is used in aircraft, cans and window frames. Its oxide and hydroxide are amphoteric, dissolving in acid to give [Al(H₂O)₆]³⁺ and in base to give [Al(OH)₄]⁻, a property exploited in the Bayer process, which purifies bauxite by dissolving alumina in hot sodium hydroxide and leaving iron oxides behind. The metal is then made by electrolysis (lesson 9). Aluminium chloride exists as Al₂Cl₆ dimers in the vapour and in non-polar solvents, with two bridging chlorides providing the missing electron pairs; its strong Lewis acidity makes it a catalyst in organic chemistry.

Gallium melts at 29.8 °C, low enough to melt in the hand, and remains liquid to about 2400 °C, the longest liquid range of any element. Gallium arsenide, GaAs, and gallium nitride, GaN, are semiconductors used in light-emitting diodes, lasers and high-frequency electronics; the development of efficient blue LEDs based on GaN won Isamu Akasaki, Hiroshi Amano and Shuji Nakamura the Nobel Prize in Physics in 2014. Indium tin oxide, a transparent conductor, coats almost every touchscreen and flat-panel display.

The inert-pair effect

Thallium, at the bottom of group 13, is most stable as Tl⁺ rather than Tl³⁺. Tl³⁺ is a strong oxidizing agent (E°(Tl³⁺/Tl⁺) = +1.25 V), while the corresponding Al³⁺ and Ga³⁺ show no tendency to be reduced to the +1 state. The same pattern, a growing preference for an oxidation state two below the group maximum, appears in lead (Pb²⁺ more stable than Pb⁴⁺) and bismuth (Bi³⁺ more stable than Bi⁵⁺). This inert-pair effect does not mean the 6s² electrons are chemically inert. Its causes are a combination of factors: the 6s electrons are held more tightly than expected, because of poor shielding by the filled 4f and 5d shells and the relativistic contraction of the 6s orbital (lesson 2); and the bonds formed by these large atoms are weak, so the energy released by forming two extra bonds does not repay the cost of ionizing the s² pair. Lead dioxide, PbO₂, which contains Pb(IV), is therefore a strong oxidizing agent, and its reduction to PbSO₄ supplies the energy of the lead–acid car battery.

Group 14: carbon, silicon, germanium, tin and lead

Carbon is unique in its ability to form long chains and rings with itself (catenation), because C–C bonds are strong (346 kJ/mol) and comparable in strength to C–O and C–H bonds; this is the basis of organic chemistry. Its allotropes have very different structures and properties. Diamond is a three-dimensional network of sp³ carbons, the hardest natural material and the best thermal conductor known at room temperature. Graphite consists of sheets of sp² carbons with delocalized π electrons, which makes it a soft, slippery conductor. Fullerenes, closed cages such as C₆₀ with the shape of a soccer ball, were discovered in 1985 by Harold Kroto, Robert Curl and Richard Smalley (Nobel Prize 1996). Carbon nanotubes, rolled-up graphite sheets, were described by Sumio Iijima in 1991, and graphene, a single sheet of graphite, was isolated by Andre Geim and Konstantin Novoselov in 2004 using adhesive tape (Nobel Prize in Physics 2010).

Silicon, the second most abundant element in the crust, never occurs free; it is found as silica, SiO₂, and as silicates. Unlike CO₂, a gas of small molecules with C=O double bonds, SiO₂ is a network solid of SiO₄ tetrahedra sharing corners, because Si–O single bonds (452 kJ/mol) are much stronger than half a Si=O double bond would be: silicon's larger 3p orbitals overlap poorly in π bonds. The silicate minerals are built from the same SiO₄ tetrahedra joined in different ways: isolated tetrahedra in olivine, single chains in pyroxenes, double chains in amphiboles (including asbestos), sheets in micas and clays, and three-dimensional frameworks in quartz and the feldspars. Replacing some Si⁴⁺ by Al³⁺ gives the aluminosilicate frameworks of zeolites, with cages and channels of molecular size that act as molecular sieves, ion exchangers in detergents and acid catalysts in oil refining. Silicones, polymers with Si–O–Si backbones and organic side groups, are used as sealants, lubricants and medical implants.

Elemental silicon, purified to better than one impurity atom in a billion, is the basis of the semiconductor industry (lesson 12). Germanium was the material of the first transistor in 1947. Tin has two common allotropes: metallic white tin is stable above 13.2 °C, and brittle, non-metallic grey tin, with the diamond structure, below it. The slow transformation in cold weather, called tin pest, can crumble tin objects. Lead is a dense, soft, toxic metal whose compounds were once widely used in paints and gasoline; it is now restricted in most uses because of its neurological effects, especially on children.

Periodic trends in numbers

The code below collects data for group 1 and group 14 and examines the trends discussed in this lesson: falling ionization energy and melting point down group 1, the anomalous potential of lithium, the hydration energy that explains it, and the falling strength of element–element bonds down group 14 that limits catenation and favours the inert-pair state.

# Trends in group 1 and group 14 (approximate literature values).
g1 = {  # symbol: (metallic radius pm, I1 kJ/mol, E0 V, mp C, hydration of M+ kJ/mol)
 'Li': (152, 520, -3.04, 180.5, -520), 'Na': (186, 496, -2.71, 97.8, -405),
 'K': (227, 419, -2.93, 63.5, -321), 'Rb': (248, 403, -2.98, 39.3, -300), 'Cs': (265, 376, -3.03, 28.4, -277)}
print('group 1: radius  I1   E0     mp    hydration')
for s, (r, I, E, mp, h) in g1.items():
    print(f'   {s:2} {r:6} {I:5} {E:6.2f} {mp:6.1f} {h:6}')

# least-squares line of melting point against 1/radius
xs = [1 / v[0] for v in g1.values()]; ys = [v[3] for v in g1.values()]
n = len(xs); mx = sum(xs) / n; my = sum(ys) / n
m = sum((x - mx) * (y - my) for x, y in zip(xs, ys)) / sum((x - mx) ** 2 for x in xs)
b = my - m * mx
r = sum((x - mx) * (y - my) for x, y in zip(xs, ys)) / (sum((x - mx) ** 2 for x in xs) * sum((y - my) ** 2 for y in ys)) ** 0.5
print(f'melting point vs 1/radius: slope {m:.0f} C pm, intercept {b:.0f} C, correlation r = {r:.3f}')

# Why Li has the most negative E0: compare I1 with the hydration energy
for s in ('Li', 'Na', 'K'):
    r_, I, E, mp, h = g1[s]
    print(f'{s}: I1 + hydration = {I + h:+d} kJ/mol (the lower the value, the easier M+(aq) forms)')

# Group 14: element-element single bond enthalpies and catenation
g14 = {'C': 346, 'Si': 226, 'Ge': 188, 'Sn': 151, 'Pb': 100}
EO = {'C': 358, 'Si': 452}
print('\ngroup 14 E-E bond enthalpies (kJ/mol):', g14)
print(f'C-C vs C-O: {g14["C"]} vs {EO["C"]}; Si-Si vs Si-O: {g14["Si"]} vs {EO["Si"]}')
print('ratio E-O/E-E: C %.2f, Si %.2f -> silicon prefers Si-O-Si networks' % (EO['C'] / g14['C'], EO['Si'] / g14['Si']))

group 1: radius  I1   E0     mp    hydration
   Li    152   520  -3.04  180.5   -520
   Na    186   496  -2.71   97.8   -405
   K     227   419  -2.93   63.5   -321
   Rb    248   403  -2.98   39.3   -300
   Cs    265   376  -3.03   28.4   -277
melting point vs 1/radius: slope 52826 C pm, intercept -173 C, correlation r = 0.992
Li: I1 + hydration = +0 kJ/mol (the lower the value, the easier M+(aq) forms)
Na: I1 + hydration = +91 kJ/mol (the lower the value, the easier M+(aq) forms)
K: I1 + hydration = +98 kJ/mol (the lower the value, the easier M+(aq) forms)

group 14 E-E bond enthalpies (kJ/mol): {'C': 346, 'Si': 226, 'Ge': 188, 'Sn': 151, 'Pb': 100}
C-C vs C-O: 346 vs 358; Si-Si vs Si-O: 226 vs 452
ratio E-O/E-E: C 1.03, Si 2.00 -> silicon prefers Si-O-Si networks

The fits and comparisons quantify the qualitative story. Melting point correlates closely with the inverse of the metallic radius. For lithium, the sum of ionization energy and hydration energy is the most favourable of the three metals, which is why its potential is the most negative even though its ionization energy is the highest. Down group 14 the E–E bond enthalpy falls by more than a factor of three, and for silicon the ratio of Si–O to Si–Si bond strength is twice that of carbon, which is why silicon chemistry on Earth is the chemistry of silicates, not of long silicon chains.