Learn Unit 11-Some p -Block Elements with Free Lessons & Tips

The variation in oxidation states across the elements of a group or period often reflects underlying trends in electronic configurations and chemical reactivity. Let's explore the pattern of oxidation states for the elements:

(i) Boron (B) to Thallium (Tl):

• Boron (B): Boron typically exhibits an oxidation state of +3 in its compounds, as it prefers to lose its three valence electrons to achieve a stable electronic configuration.
• Aluminum (Al): Aluminum predominantly shows an oxidation state of +3 in its compounds, similar to boron, due to the loss of its three valence electrons.
• Gallium (Ga): Gallium mainly displays an oxidation state of +3, following the trend set by boron and aluminum.
• Indium (In): Indium is more versatile in its oxidation states compared to the preceding elements. It can exhibit both +1 and +3 oxidation states, with the +3 oxidation state being more common.
• Thallium (Tl): Thallium, like indium, shows a wider range of oxidation states. It can exhibit +1, +3, and even +5 oxidation states. However, the +1 oxidation state is more stable and common than the +3 or +5 states.

The trend across this group shows a general progression from +3 oxidation state (B, Al, Ga) to a more varied set of oxidation states as we move down the group, with elements like indium and thallium displaying a wider range of oxidation states due to the increasing ease of losing electrons as we move down the group.

(ii) Carbon (C) to Lead (Pb):

• Carbon (C): Carbon typically exhibits oxidation states of -4 (in compounds like methane) to +4 (in compounds like carbon dioxide). However, its most common oxidation states are +4 in carbon dioxide and -4 in methane.
• Silicon (Si): Silicon primarily shows oxidation states of +4 (in compounds like silicon dioxide) and sometimes +2 (in compounds like silane, although less common).
• Germanium (Ge): Germanium mainly exhibits an oxidation state of +4, following the trend set by carbon and silicon.
• Tin (Sn): Tin is more versatile in its oxidation states compared to the preceding elements. It can exhibit oxidation states of +2 and +4, with the +2 oxidation state being more common.
• Lead (Pb): Lead, like tin, can exhibit multiple oxidation states. It commonly shows oxidation states of +2 and +4, with the +2 oxidation state being more stable and common.

The trend across this group shows a general progression from a wider range of oxidation states (-4 to +4) for carbon to a narrower range of oxidation states (+2 to +4) for lead. This narrowing occurs due to the increasing size and decreasing electronegativity of the elements as we move down the group, making it less favorable for higher oxidation states to be stabilized.

The stability of BCl3 compared to TlCl3 can be attributed to several factors:

1. Electronegativity: Boron (B) is more electronegative than Thallium (Tl), which means that in BCl3, the bonding electrons are pulled closer to the boron atom, resulting in stronger B-Cl bonds. In contrast, Tl has relatively low electronegativity, leading to weaker Tl-Cl bonds in TlCl3.

2. Size of the central atom: Boron is smaller in size compared to thallium. In BCl3, the smaller size of boron allows for stronger bonds because the bonding electrons are held closer to the nucleus, leading to more effective overlap between the atomic orbitals, resulting in stronger bonding. In TlCl3, the larger size of thallium results in weaker bonds due to decreased effective overlap of atomic orbitals.

3. Steric effects: Thallium's larger size leads to greater steric hindrance compared to boron. This steric hindrance can destabilize the Tl-Cl bonds in TlCl3, making them more prone to breaking compared to the bonds in BCl3.

4. Polarity: BCl3 is a planar molecule with trigonal planar geometry, while TlCl3 adopts a distorted trigonal pyramidal geometry due to the lone pair on the thallium atom. This lone pair contributes to the polarity of TlCl3, making it more prone to hydrolysis and other reactions compared to the non-polar BCl3.

In summary, the combination of higher electronegativity, smaller size, and less steric hindrance in BCl3 compared to TlCl3 leads to stronger and more stable bonds in BCl3.

Boron trifluoride, BF3BF3, behaves as a Lewis acid because it can accept a pair of electrons from a Lewis base to form a coordinate covalent bond. In BF3BF3, boron has an incomplete octet, meaning it has only six electrons in its valence shell instead of the stable eight. This makes it electron deficient and eager to accept electron pairs to complete its octet.

When a Lewis base approaches BF3BF3, such as a molecule with a lone pair of electrons, like NH3NH3 or H2OH2O, the boron atom can accept the lone pair of electrons from the Lewis base, forming a coordinate covalent bond. This results in the formation of a complex, with the boron atom surrounded by more than the usual number of electron pairs, thus satisfying the octet rule.

So, in summary, BF3BF3 behaves as a Lewis acid because it can accept electron pairs from Lewis bases due to its electron deficiency.

BCl3 (boron trichloride) and CCl4 (carbon tetrachloride) have different behaviors when they come into contact with water due to their molecular structures and properties.

1. BCl3 (Boron Trichloride):

   • BCl3 is a Lewis acid, meaning it can accept a pair of electrons to form a covalent bond.
   • When BCl3 reacts with water, it acts as a Lewis acid and accepts a lone pair of electrons from water molecules to form hydrochloric acid (HCl) and boric acid (H3BO3). The reaction can be represented as: BCl3+3H2O→B(OH)3+3HClBCl3+3H2O→B(OH)3+3HCl
   • The reaction is highly exothermic and produces acidic solutions due to the formation of HCl. Boric acid formed is weakly acidic and can further react with water to form boric acid solutions.
   • Thus, BCl3 hydrolyzes vigorously in water to produce acidic solutions.

2. CCl4 (Carbon Tetrachloride):

   • CCl4 is a non-polar molecule composed of carbon and four chlorine atoms arranged symmetrically around the carbon atom.
   • It lacks any polar bonds or a permanent dipole moment, making it highly insoluble in water.
   • Due to the lack of polarity, CCl4 molecules do not interact strongly with water molecules through hydrogen bonding or dipole-dipole interactions.
   • Consequently, when CCl4 is mixed with water, it forms a separate layer on top of the water due to its lower density, and they do not readily mix or react. This property makes it useful as a non-polar solvent.

In summary, BCl3 reacts vigorously with water, undergoing hydrolysis to form acidic solutions due to its Lewis acidic nature, while CCl4 remains largely unreactive and immiscible with water due to its non-polar nature.

Boric acid is technically a weak Lewis acid rather than a protonic acid. Let me break it down:

1. Protonic Acid: Protonic acids, also known as Bronsted acids, are substances that can donate a proton (H⁺ ion) to another substance. In simpler terms, they are acids that readily release hydrogen ions in solution. Examples include hydrochloric acid (HCl) and sulfuric acid (H₂SO₄).

2. Lewis Acid: In contrast, Lewis acids are substances that can accept a pair of electrons. They don't necessarily need to donate protons; instead, they can form a coordinate covalent bond by accepting an electron pair from another substance. Boric acid falls into this category.

Boric acid, chemically represented as H₃BO₃ or B(OH)₃, can act as a Lewis acid by accepting a pair of electrons from a Lewis base. Its behavior as an acid is due to the ability of the boron atom to accept an electron pair from a base, forming a coordinate covalent bond. This property allows it to react with substances like alcohols or water to form borate ions.

While boric acid can behave as an acid in certain reactions, it's not as strong as typical protonic acids like hydrochloric acid. Its acidic properties are more subtle and primarily manifest in reactions where it acts as a Lewis acid.

When boric acid (H3BO3) is heated, it undergoes several chemical changes. Initially, boric acid dehydrates upon heating, losing water molecules to form metaboric acid (HBO2):

2H3BO3 (boric acid) -> H2B4O7 (metaboric acid) + H2O

Further heating leads to the conversion of metaboric acid into various polymeric forms of boric anhydride or tetraboric acid (H2B4O7):

4H2B4O7 (metaboric acid) -> 2B2O3 (boric anhydride) + 5H2O

The boron oxide formed can further polymerize to form complex boron oxide networks at higher temperatures.

The exact products and reactions may vary depending on the specific conditions of heating, such as temperature and presence of other substances. Boric acid's thermal decomposition is utilized in various industrial processes, including the production of boron-containing compounds and ceramics.

BF3, or boron trifluoride, has a trigonal planar shape. Each fluorine atom forms a single bond with the central boron atom, resulting in three bonding pairs of electrons around boron. Since there are no lone pairs on boron, the shape is trigonal planar.

BH4−, or tetrahydroborate ion, has a tetrahedral shape. Boron in BH4− has four hydrogen atoms bonded to it, resulting in four bonding pairs of electrons around boron. There are no lone pairs on boron, so the shape is tetrahedral.

In both cases, the hybridization of boron can be determined by counting the number of regions of electron density (bonding pairs and lone pairs) around the boron atom.

For BF3:

• There are 3 regions of electron density (3 bonding pairs), indicating sp2 hybridization.

For BH4−:

• There are 4 regions of electron density (4 bonding pairs), indicating sp3 hybridization.

1. "Aluminum's amphoteric nature is evident when it reacts with both acids and bases. When it encounters an acidic environment, such as hydrochloric acid, it forms aluminum chloride and hydrogen gas. On the other hand, in a basic solution like sodium hydroxide, aluminum reacts to form sodium aluminate and hydrogen gas. This ability to react with both acidic and basic substances showcases its amphoteric behavior."

2. "One clear demonstration of aluminum's amphoteric nature is its reaction with both strong acids and strong bases. When aluminum reacts with an acid like sulfuric acid, it produces aluminum sulfate and hydrogen gas. Conversely, when it interacts with a strong base like potassium hydroxide, it yields potassium aluminate and hydrogen gas. This dual reactivity illustrates its characteristic amphoteric behavior."

3. "Aluminum exhibits its amphoteric properties through its reactions with both acidic and basic solutions. For instance, when it reacts with an acid such as nitric acid, it forms aluminum nitrate and releases hydrogen gas. Similarly, when it comes into contact with a base like calcium hydroxide, it produces calcium aluminate and liberates hydrogen gas. These reactions clearly demonstrate aluminum's ability to behave as both an acid and a base."

4. "The amphoteric nature of aluminum becomes apparent in its reactions with acids and bases. When treated with an acid like hydrochloric acid, aluminum undergoes a reaction to form aluminum chloride and hydrogen gas. In contrast, when exposed to a base such as sodium hydroxide, aluminum reacts to produce sodium aluminate and hydrogen gas. These reactions underscore aluminum's dual propensity to interact with both acidic and basic environments, thus exhibiting its amphoteric character."

Electron deficient compounds are molecules that possess fewer electrons than what is typically expected based on the octet rule or the duet rule in the case of hydrogen. These compounds often exhibit incomplete valence electron shells and tend to be highly reactive, seeking to gain electrons to achieve a more stable electronic configuration.

BCl3 (boron trichloride) and SiCl4 (silicon tetrachloride) are indeed examples of electron deficient species. Let's examine each:

1. BCl3 (boron trichloride):

   • Boron has three valence electrons, and in BCl3, it forms three covalent bonds with chlorine atoms.
   • However, boron itself only has six electrons around it, leaving it short of the octet rule. Thus, BCl3 is electron deficient.
   • Due to its electron deficiency, BCl3 is highly reactive and acts as a Lewis acid, readily accepting a pair of electrons from a Lewis base.

2. SiCl4 (silicon tetrachloride):

   • Silicon has four valence electrons, and in SiCl4, it forms four covalent bonds with chlorine atoms.
   • Similar to boron, silicon ends up with only eight electrons around it, which is still short of a complete octet.
   • Therefore, SiCl4 is also electron deficient.
   • Silicon tetrachloride is also reactive due to its electron deficiency, though it's not as reactive as boron trichloride.

In both cases, the central atom (boron or silicon) lacks a full complement of valence electrons, making these molecules electron deficient. This electron deficiency makes them susceptible to reacting with species that can donate electron pairs, making them important reagents in various chemical processes.

This completes the resonance structure for the bicarbonate ion (HCO₃⁻).

Keep in mind that in both cases, the actual structure of the ion is a hybrid of these resonance structures, with the true structure being an average of the contributing resonance forms.