Group 13 Elements: The Boron Family — Explained

On this page

1. 1. Conceptual Foundation and Electronic Configuration:
2. 2. Atomic and Ionic Radii:
3. 3. Ionization Enthalpy:
4. 4. Electronegativity:
5. 5. Oxidation States and the Inert Pair Effect:
6. 6. Chemical Properties:
7. 7. Anomalous Behavior of Boron:
8. 8. Important Compounds of Boron:
9. 9. Important Compounds of Aluminium:
10. 10. NEET-Specific Angle:

Detailed Explanation

The Group 13 elements, often termed the Boron family, represent a fascinating transition in chemical properties within the p-block. Comprising Boron (B), Aluminium (Al), Gallium (Ga), Indium (In), and Thallium (Tl), these elements share a common valence shell electronic configuration of $ns^2np^1$. This configuration dictates their primary chemical behavior, primarily the tendency to exhibit a +3 oxidation state.

1. Conceptual Foundation and Electronic Configuration:

All Group 13 elements have three valence electrons. Their general electronic configuration is $[Noble,Gas]ns^2np^1$. For example, Boron is $[He]2s^22p^1$, Aluminium is $[Ne]3s^23p^1$, and so on. The presence of a single electron in the p-orbital and two electrons in the s-orbital is key to their chemistry.

2. Atomic and Ionic Radii:

Initially, as we move from Boron to Aluminium, the atomic radius increases significantly due to the addition of a new electron shell. However, the trend is not uniform thereafter. From Aluminium to Gallium, the atomic radius actually decreases slightly (Al: 143 pm, Ga: 135 pm).

This unexpected contraction is attributed to the poor shielding effect of the ten d-electrons in Gallium, which are present in its penultimate shell. These d-electrons do not effectively shield the valence electrons from the increased nuclear charge, leading to a stronger attraction and a smaller atomic size.

Subsequently, from Gallium to Indium and then to Thallium, the atomic radius generally increases, but the increase from Indium to Thallium is also smaller than expected due to the presence of 14 f-electrons in Thallium, which also exhibit poor shielding.

3. Ionization Enthalpy:

The ionization enthalpy (IE) trend in Group 13 is highly irregular. Generally, IE is expected to decrease down a group as atomic size increases and valence electrons are further from the nucleus. While Boron has the highest IE, the decrease from Boron to Aluminium is significant.

However, from Aluminium to Gallium, the IE increases (Al: 577 kJ/mol, Ga: 579 kJ/mol). This is a direct consequence of the smaller atomic size of Gallium and the increased effective nuclear charge due to poor d-electron shielding.

Similarly, the IE decreases from Gallium to Indium but then increases again from Indium to Thallium (In: 558 kJ/mol, Tl: 589 kJ/mol) due to the poor shielding of f-electrons in Thallium, leading to a higher effective nuclear charge.

This irregular trend is a hallmark of the p-block elements following the d-block and f-block contractions.

4. Electronegativity:

Electronegativity generally decreases from Boron to Aluminium. However, similar to ionization enthalpy, it shows a slight increase from Aluminium to Thallium. This is again explained by the irregular atomic size and effective nuclear charge trends. Boron is the most electronegative in the group, reflecting its non-metallic character.

5. Oxidation States and the Inert Pair Effect:

The most common oxidation state for Group 13 elements is +3, arising from the loss or sharing of all three valence electrons ($ns^2np^1$). However, as we move down the group, the stability of the +1 oxidation state increases, particularly for Indium and Thallium.

This phenomenon is known as the inert pair effect. It refers to the reluctance of the $ns^2$ electrons to participate in bonding. For heavier elements like Thallium, the $ns^2$ electrons are more tightly held by the nucleus due to the poor shielding of intervening d and f electrons, making it energetically unfavorable to unpair and promote them for bonding.

Consequently, Tl(I) compounds are more stable than Tl(III) compounds, whereas for Boron and Aluminium, +3 is overwhelmingly stable. For Gallium and Indium, both +1 and +3 oxidation states exist, with +3 being more stable.

6. Chemical Properties:

Reactivity towards Air: — Boron is unreactive in crystalline form but reacts with air at high temperatures to form $B_2O_3$ and $BN$. Aluminium forms a thin, protective oxide layer ($Al_2O_3$) on its surface, which prevents further corrosion. Amorphous boron and aluminium powder burn in air to form oxides.

$4B(s) + 3O_2(g) \xrightarrow{heat} 2B_2O_3(s)$ $2Al(s) + \frac{3}{2}O_2(g) \xrightarrow{heat} Al_2O_3(s)$

Reactivity towards Acids and Bases: — Boron is unreactive with non-oxidizing acids but reacts with strong oxidizing acids like $HNO_3$ and $H_2SO_4$ upon heating. Aluminium is amphoteric; it reacts with both acids and bases.

$2Al(s) + 6HCl(aq) \rightarrow 2AlCl_3(aq) + 3H_2(g)$ $2Al(s) + 2NaOH(aq) + 6H_2O(l) \rightarrow 2Na[Al(OH)_4](aq) + 3H_2(g)$ (Sodium tetrahydroxoaluminate(III)) Gallium also exhibits amphoteric behavior.

Reactivity towards Halogens: — These elements react with halogens to form trihalides ($MX_3$).

$2E(s) + 3X_2(g) \rightarrow 2EX_3(s)$ (where E = Group 13 element, X = halogen) Most trihalides are covalent, especially for Boron and Aluminium, and act as Lewis acids due to their electron-deficient nature.

7. Anomalous Behavior of Boron:

Boron, the first member, exhibits properties significantly different from its heavier congeners, a common trend for the first element in many p-block groups. Key reasons include:

Small Size and High Ionization Enthalpy: — These factors lead to a high charge-to-radius ratio.
High Electronegativity: — Boron is more electronegative than Al, Ga, In, Tl.
Absence of d-orbitals: — Boron cannot expand its octet, limiting its maximum covalency to four. Other elements can utilize vacant d-orbitals.
Non-metallic Character: — Boron is a non-metal, forming covalent compounds, while others are metallic.
Forms Electron-Deficient Compounds: — Boron compounds like $BF_3$ and $B_2H_6$ are electron-deficient and act as Lewis acids.
Forms Stable Hydrides (Boranes): — Boron forms a variety of complex hydrides called boranes, e.g., diborane ($B_2H_6$), which have unique 'banana' or three-center two-electron bonds.
Diagonal Relationship with Silicon: — Boron shows similarities with Silicon (Group 14, Period 3) in properties like forming covalent compounds, acidic oxides, and halides that hydrolyze.

8. Important Compounds of Boron:

Borax ($Na_2B_4O_7 \cdot 10H_2O$): — A white crystalline solid, it's the most important boron mineral. It contains tetranuclear units $[B_4O_5(OH)_4]^{2-}$. When heated, it swells, loses water, and forms a transparent glassy bead of sodium metaborate ($NaBO_2$) and boric anhydride ($B_2O_3$). This is the basis of the borax bead test.

$Na_2B_4O_7 \cdot 10H_2O \xrightarrow{heat} Na_2B_4O_7 \xrightarrow{heat} 2NaBO_2 + B_2O_3$

Boric Acid ($H_3BO_3$ or $B(OH)_3$): — A white crystalline solid with a soapy touch. It is a weak monobasic Lewis acid, not a protic acid. It accepts an electron pair from $OH^-$ ions from water.

$B(OH)_3 + H_2O \rightleftharpoons [B(OH)_4]^- + H^+$ It has a layered structure with planar $BO_3$ units linked by hydrogen bonds.

Diborane ($B_2H_6$): — The simplest boron hydride. It is an electron-deficient compound. Its structure involves two $BH_2$ units linked by two bridging hydrogen atoms. The bridging hydrogen atoms are involved in 'banana bonds' or three-center two-electron bonds (3c-2e bonds). Each boron atom is $sp^3$ hybridized. It reacts with water to form boric acid and hydrogen gas.

$B_2H_6 + 6H_2O \rightarrow 2H_3BO_3 + 6H_2$ It also undergoes cleavage reactions with Lewis bases, e.g., $B_2H_6 + 2NH_3 \rightarrow [BH_2(NH_3)_2]^+[BH_4]^-$.

9. Important Compounds of Aluminium:

Aluminium Oxide ($Al_2O_3$, Alumina): — A very stable compound. It is amphoteric, reacting with both acids and bases. Different forms exist, e.g., $\alpha$-alumina (corundum) is very hard.

$Al_2O_3(s) + 6HCl(aq) \rightarrow 2AlCl_3(aq) + 3H_2O(l)$ $Al_2O_3(s) + 2NaOH(aq) + 3H_2O(l) \rightarrow 2Na[Al(OH)_4](aq)$

Aluminium Chloride ($AlCl_3$): — Anhydrous $AlCl_3$ is a covalent compound with a dimeric structure ($Al_2Cl_6$) in the vapor phase and a polymeric structure in the solid state. It is a strong Lewis acid due to the incomplete octet of aluminium and is used as a catalyst in Friedel-Crafts reactions. It fumes in moist air due to hydrolysis.

$AlCl_3 + 3H_2O \rightarrow Al(OH)_3 + 3HCl$ Hydrated $AlCl_3 \cdot 6H_2O$ is ionic.

10. NEET-Specific Angle:

For NEET, the focus should be on:

Trends and Exceptions: — Especially the irregular trends in atomic radii and ionization enthalpy, and the explanation for these (d- and f-block contraction, poor shielding).
Inert Pair Effect: — Its definition, cause, and consequences for the stability of +1 vs +3 oxidation states, particularly for Tl.
Anomalous Behavior of Boron: — Its non-metallic nature, electron deficiency, maximum covalency of 4, and diagonal relationship with Silicon.
Structures and Bonding: — Especially diborane (banana bonds, 3c-2e bonds, $sp^3$ hybridization). Boric acid's layered structure and its nature as a Lewis acid.
Reactions: — Amphoteric nature of Al and Ga, borax bead test, hydrolysis of $AlCl_3$, reaction of diborane with water.
Lewis Acid Character: — $BF_3$, $AlCl_3$, $B(OH)_3$ are important examples. The relative Lewis acid strength of boron trihalides ($BI_3 > BBr_3 > BCl_3 > BF_3$) due to back-bonding. $BF_3$ is the weakest Lewis acid due to effective $ppi-ppi$ back-bonding from F to B.

Understanding these specific points, along with the general properties and compounds, will provide a strong foundation for tackling NEET questions on Group 13 elements.

Often confused with

Side-by-side differences the NEET paper likes to test.

Group 13 Elements: The Boron Family vs Aluminium

Aspect	Group 13 Elements: The Boron Family	Aluminium
Nature	Non-metal	Metal
Physical State	Hard, black solid (amorphous) or crystalline	Soft, silvery-white solid
Oxidation State	Predominantly +3, but often electron-deficient	Exclusively +3 in stable compounds
Covalency	Maximum covalency of 4 (due to absence of d-orbitals)	Can exhibit covalency of 4 or 6 (due to presence of vacant d-orbitals)
Oxides	Acidic ($B_2O_3$)	Amphoteric ($Al_2O_3$)
Hydroxides	Weak Lewis acid ($B(OH)_3$)	Amphoteric ($Al(OH)_3$)
Halides	Monomeric, strong Lewis acids ($BF_3$)	Dimeric ($Al_2Cl_6$) in vapor, polymeric in solid, strong Lewis acids

Boron, the first member of Group 13, stands in stark contrast to Aluminium, its immediate successor. Boron is a non-metal, forming predominantly covalent and electron-deficient compounds, with a maximum covalency of four due to the absence of d-orbitals.

Its oxide is acidic, and boric acid acts as a Lewis acid. Aluminium, on the other hand, is a typical metal, though amphoteric, and can expand its covalency to six by utilizing its vacant d-orbitals. These fundamental differences highlight the anomalous behavior of Boron within its own group, a common trend for the first element of many p-block groups.

Why it is tested: For NEET, understanding the distinct properties of Boron versus Aluminium is crucial. Questions often test the anomalous behavior of Boron, its electron-deficient nature, Lewis acidity, and the differences in bonding and reactivity compared to Aluminium. The amphoteric nature of Aluminium and its compounds is also a frequently tested concept, making this comparison highly relevant for conceptual clarity and problem-solving.

Questions students ask

6 answered on this topic.

What is the inert pair effect and how does it influence the chemistry of Group 13 elements?

The inert pair effect describes the reluctance of the outermost $ns^2$ electrons to participate in chemical bonding. As we move down Group 13, particularly for heavier elements like Indium and Thallium, the effective nuclear charge increases due to the poor shielding provided by the intervening d and f electrons.

This causes the $ns^2$ electrons to be held more tightly by the nucleus, making them less available for bonding. Consequently, the stability of the +1 oxidation state (where only the $np^1$ electron is involved) increases down the group, while the stability of the +3 oxidation state (involving all three valence electrons) decreases.

Thallium(I) compounds are significantly more stable than Thallium(III) compounds.

Why does Boron show anomalous behavior compared to other Group 13 elements?

Boron exhibits anomalous behavior primarily due to its exceptionally small atomic size, very high ionization enthalpy, and high electronegativity. Unlike its heavier congeners, Boron is a non-metal and forms predominantly covalent compounds.

It lacks d-orbitals, limiting its maximum covalency to four (e.g., in $[BF_4]^-$). Its compounds are often electron-deficient, acting as strong Lewis acids. These characteristics lead to significant differences in its physical and chemical properties compared to the metallic and larger elements like Aluminium, Gallium, Indium, and Thallium.

Explain the structure and bonding in Diborane ($B_2H_6$).

Diborane ($B_2H_6$) has a unique electron-deficient structure. It consists of two $BH_2$ units connected by two bridging hydrogen atoms. Each boron atom is $sp^3$ hybridized. There are four terminal B-H bonds, which are conventional two-center two-electron (2c-2e) bonds.

The two bridging B-H-B bonds are special; they are three-center two-electron (3c-2e) bonds, often called 'banana bonds'. In these bonds, two electrons are shared among three atoms (two boron atoms and one hydrogen atom).

This arrangement allows each boron atom to achieve a stable octet-like configuration despite having only six valence electrons in total for the molecule.

Why is Boric acid considered a Lewis acid and not a protic acid?

Boric acid ($H_3BO_3$ or $B(OH)_3$) is a weak monobasic acid, but it does not donate a proton directly. Instead, it acts as a Lewis acid by accepting a lone pair of electrons from a hydroxyl ion ($OH^-$) from water.

The boron atom in $B(OH)_3$ has an incomplete octet (only six electrons in its valence shell), making it electron-deficient and capable of accepting an electron pair. The reaction is $B(OH)_3 + H_2O \rightleftharpoons [B(OH)_4]^- + H^+$.

This mechanism distinguishes it from typical Brønsted-Lowry acids that directly donate protons.

What is the amphoteric nature of Aluminium and its compounds?

Aluminium and its compounds like aluminium oxide ($Al_2O_3$) and aluminium hydroxide ($Al(OH)_3$) are amphoteric, meaning they can react with both acids and bases. When reacting with acids, they behave as bases, forming salts and water.

For example, $Al_2O_3 + 6HCl \rightarrow 2AlCl_3 + 3H_2O$. When reacting with strong bases, they behave as acids, forming complex salts. For instance, $Al_2O_3 + 2NaOH + 3H_2O \rightarrow 2Na[Al(OH)_4]$ (sodium tetrahydroxoaluminate(III)).

This dual reactivity is characteristic of elements whose oxides and hydroxides lie on the borderline between acidic and basic behavior.

How does the Lewis acid strength of boron trihalides vary?

The Lewis acid strength of boron trihalides ($BX_3$) follows the order $BI_3 > BBr_3 > BCl_3 > BF_3$. This trend is contrary to what might be expected based on electronegativity (F is most electronegative, so $BF_3$ should be the strongest Lewis acid).

The explanation lies in the concept of back-bonding or $ppi-ppi$ overlap. The empty p-orbital on boron can accept electron density from the lone pairs on the halogen atoms. Fluorine, being small and highly electronegative, has its $2p$ orbitals of appropriate size and energy to effectively overlap with the $2p$ orbital of boron, leading to significant back-bonding.

This back-bonding reduces the electron deficiency of boron in $BF_3$, making it a weaker Lewis acid. As the size of the halogen increases (Cl, Br, I), the $ppi-ppi$ overlap becomes less effective, reducing back-bonding and thus increasing the electron deficiency of boron, making the Lewis acid stronger.