Group 13 — Boron Family (General)

Members and Electronic Configuration

Group 13 of the periodic table is headed by boron and comprises five elements studied here: boron ($\ce{B}$), aluminium ($\ce{Al}$), gallium ($\ce{Ga}$), indium ($\ce{In}$) and thallium ($\ce{Tl}$). The synthetically prepared radioactive element nihonium ($\ce{Nh}$, $Z=113$) also belongs to the group, but its chemistry is not established and it is left out of property discussions. The family spans the full metallic gradient: boron is a typical non-metal, aluminium is a metal with several chemical similarities to boron, and gallium, indium and thallium are almost exclusively metallic in character.

The defining feature of the group is the outer electronic configuration $ns^2np^1$ — three valence electrons available for bonding. What changes down the group is the inner core beneath those valence electrons. Boron and aluminium sit on a noble-gas core; gallium and indium carry an additional ten $d$-electrons; thallium carries ten $d$- plus fourteen $f$-electrons. This difference in core makes the electronic structure of the heavier members far more complex than that of the group 1 and group 2 elements, and it is the root cause of the irregularities in their atomic and chemical properties.

Atomic and Physical Data

The atomic and physical constants of the group, taken from the supplement's Table 11.2, are the quantitative backbone for every trend discussed below. Read the table not as numbers to memorise but as evidence: the deviations in radius and ionisation enthalpy are exactly where the inner-core $d$- and $f$-electrons leave their fingerprints.

Physical character. Boron is non-metallic — an extremely hard, black solid existing in several allotropic forms; its very strong crystalline lattice gives it an unusually high melting point. The remaining members are soft metals of low melting point and high electrical conductivity. Gallium is a striking exception with a melting point of only 303 K, so it can melt in the hand and exist as a liquid in summer; its very high boiling point (2676 K) makes it useful for measuring high temperatures. Density rises steadily from boron to thallium.

Trend in Atomic Radii

On descending a group one extra electron shell is added at each step, so atomic radius is expected to increase smoothly. The boron family obeys this only partly. The genuine anomaly is that the atomic radius of gallium (135 pm) is smaller than that of aluminium (143 pm) — a reversal of the expected trend.

The cause lies in the inner core. Between aluminium and gallium, ten $3d$-electrons are inserted. These $d$-electrons screen the nuclear charge poorly, so the outermost electrons of gallium feel a higher effective nuclear charge and are drawn inward, shrinking the atom. The same poor screening by intervening $d$- and $f$-electrons recurs lower in the group and underlies the irregular ionisation pattern discussed next.

Ionisation Enthalpy and Electronegativity

Ionisation enthalpy does not fall smoothly down the group as the simple "size increases, electrons leave easily" rule would predict. The decrease from $\ce{B}$ to $\ce{Al}$ is associated with the expected increase in size. But there is a clear discontinuity between $\ce{Al}$ and $\ce{Ga}$, and again between $\ce{In}$ and $\ce{Tl}$: at these steps the newly added $d$- (and later $f$-) electrons screen poorly and fail to offset the larger nuclear charge, so ionisation enthalpy does not drop as expected. For each element the order is, predictably, $\Delta_iH_1 < \Delta_iH_2 < \Delta_iH_3$, and the sum of the first three is very high — a fact that decides their chemistry.

Electronegativity follows a matching zig-zag: it first decreases from $\ce{B}$ ($2.0$) to $\ce{Al}$ ($1.5$), then increases marginally down to thallium ($1.8$). This non-monotonic behaviour is again a consequence of the discrepancies in atomic size produced by the inner $d$- and $f$-cores.

Oxidation States and the Inert Pair Effect

The group oxidation state is $+3$, matching the three valence electrons. For the lighter members this is the dominant state. As we go down, however, a $+1$ state — two units below the group state — appears and grows in importance. This is the celebrated inert pair effect.

The mechanism is straightforward. Down the group the intervening $d$- and $f$-orbitals shield poorly, so the increased effective nuclear charge holds the $ns^2$ pair tightly and restricts its participation in bonding; only the lone $np^1$ electron is then used. NIOS §18.9 attributes this reluctance to two energetic factors: the high promotion energy needed to reach the bonding-ready valence state, and the poorer orbital overlap (hence weaker bonds) for the large heavy atoms. The relative stability of the $+1$ state therefore rises in the sequence $\ce{Al} < \ce{Ga} < \ce{In} < \ce{Tl}$. In thallium the $+1$ state is predominant, while $\ce{Tl^3+}$ is a powerful oxidising agent.

The standard electrode potentials make the crossover quantitative. For aluminium $E^\circ(\ce{Al^3+}/\ce{Al}) = -1.66\ \text{V}$, showing a strong drive to form $\ce{Al^3+}(aq)$; aluminium is a highly electropositive metal. For thallium $E^\circ(\ce{Tl^3+}/\ce{Tl}) = +1.26\ \text{V}$, so $\ce{Tl^3+}$ is unstable in solution and a powerful oxidiser — hence $\ce{Tl+}$ is more stable than $\ce{Tl^3+}$. Energetically, compounds in the $+1$ state are more ionic than those in the $+3$ state.

Chemical Reactivity

The high sum of the first three ionisation enthalpies dictates the bonding mode. In boron this sum is so large that $\ce{B^3+}$ ions never form; boron is forced into purely covalent compounds. From $\ce{B}$ to $\ce{Al}$ the sum falls considerably, so aluminium readily forms $\ce{Al^3+}$ and acts as a highly electropositive metal. In trivalent compounds the central atom carries only six electrons (as in $\ce{BF3}$); such electron-deficient molecules accept a lone pair to complete the octet and so behave as Lewis acids, the strength of which falls as size grows down the group.

Reactivity towards air

Crystalline boron is unreactive. Aluminium develops a thin protective oxide layer that shields the metal from further attack. On heating in air, amorphous boron and aluminium form their oxides, and with dinitrogen at high temperature they form nitrides:

$$\ce{4Al + 3O2 ->[\Delta] 2Al2O3} \qquad \ce{2Al + N2 ->[\Delta] 2AlN}$$

The acid–base character of the oxides shifts down the group: $\ce{B2O3}$ is acidic and reacts with basic oxides to give borates; $\ce{Al2O3}$ and $\ce{Ga2O3}$ are amphoteric; the oxides of indium and thallium are basic.

Reactivity towards acids and alkalies

Boron does not react with acids or alkalies even at moderate temperature. Aluminium, by contrast, is amphoteric — it dissolves in both mineral acids and aqueous alkalies, liberating dihydrogen:

$$\ce{2Al + 6HCl -> 2AlCl3 + 3H2}$$

$$\ce{2Al + 2NaOH + 6H2O -> 2Na[Al(OH)4] + 3H2}$$

Concentrated nitric acid, however, renders aluminium passive by building a protective oxide film, which is why $\ce{conc. HNO3}$ can be transported in aluminium containers.

Reactivity towards halogens

All these elements react with halogens to form trihalides (the lone exception being $\ce{TlI3}$, which does not exist as a true thallium(III) iodide):

$$\ce{2E + 3X2 -> 2EX3} \quad (\text{E = element; X = F, Cl, Br, I})$$

The monomeric trihalides are electron deficient and act as strong Lewis acids. Boron's maximum covalence is fixed at four ($\ce{BF4-}$) for want of $d$-orbitals, but heavier members complete their octet by halogen bridging — aluminium chloride dimerises to $\ce{Al2Cl6}$, each aluminium accepting electrons from a bridging chlorine.

Anomalous Behaviour of Boron

As the first member of the group, boron differs markedly from its heavier congeners. The reasons echo the lithium–beryllium diagonal anomaly of the s-block: boron's small size, high ionisation enthalpy, high electronegativity, and crucially the absence of $d$-orbitals in its valence shell.

The lack of $d$-orbitals caps boron's covalence at four — it can hold only four electron pairs, so it forms $\ce{BF4-}$ but never $\ce{BF6^3-}$. Aluminium and the heavier elements, having vacant $d$-orbitals, expand their covalence to six and form ions such as $\ce{AlF6^3-}$. This single structural fact explains a whole cluster of boron's peculiarities and is a favourite NEET discriminator.