UNIT-2: Acids and Bases and Inorganic Polymers

1. Brönsted-Lowry Concept of Acid-Base Reactions

The Brönsted-Lowry concept, proposed independently by Johannes Nicolaus Brönsted and Thomas Martin Lowry in 1923, defines acid-base chemistry based entirely on proton transfer.

Brönsted-Lowry Acid: A chemical species (molecule or ion) that acts as a proton (H+) donor.
Brönsted-Lowry Base: A chemical species (molecule or ion) that acts as a proton (H+) acceptor.

Conjugate Acid-Base Pairs

When a Brönsted acid loses a proton, it forms a species capable of accepting a proton, known as its conjugate base. Conversely, when a Brönsted base accepts a proton, it forms its conjugate acid.

General Reaction Equilibrium:

Acid-1 + Base-2 <=> Base-1 + Acid-2

• Example 1: Reaction of Hydrochloric Acid with Water

HCl + H2O <=> Cl- + H3O+

Here, HCl donates a proton to H2O. HCl is Acid-1, H2O is Base-2, Cl- is Conjugate Base-1, and H3O+ is Conjugate Acid-2.

• Example 2: Reaction of Ammonia with Water

NH3 + H2O <=> NH4+ + OH-

Here, H2O donates a proton to NH3. NH3 is Base-1, H2O is Acid-2, NH4+ is Conjugate Acid-1, and OH- is Conjugate Base-2.

Amphiprotic (Amphoteric) Species

Species that can act as both proton donors and proton acceptors depending on the reaction environment are called amphiprotic.

Common examples include H2O, HCO3-, HSO4-, and NH3.

• Acting as an acid: HCO3- + OH- <=> CO3^2- + H2O
• Acting as a base: HCO3- + H3O+ <=> H2CO3 + H2O

2. The Solvated Proton

A free proton (H+) is a bare atomic nucleus with an extremely small radius (approximately 1.5 x 10^-15 meters) and an exceptionally high charge density. Consequently, a free proton cannot exist independently in liquid media or condensed phases.

Structure of Solvated Protons in Aqueous Solution

In aqueous solution, the bare proton immediately attaches to a water molecule via a coordinate covalent bond to form the hydronium ion (H3O+).

H+ + H2O -> H3O+

The hydronium ion adopts a trigonal pyramidal geometry with sp3 hybridization at the central oxygen atom, leaving one lone pair of electrons.

Higher Hydrates of the Proton

Due to strong hydrogen bonding, H3O+ associates further with additional water molecules to form distinct hydrated clusters:

1. Zundel Cation [H5O2]+: A proton symmetric between two water molecules, represented as H2O--H+--OH2.
2. Eigen Cation [H9O4]+: A central H3O+ ion hydrogen-bonded to three surrounding water molecules in a trigonal arrangement.

Proton Solvation in Non-Aqueous Solvents

The nature of the solvated proton depends strongly on the donor power (basicity) of the solvent:

• In Liquid Ammonia: H+ + NH3 -> NH4+ (Ammonium ion)
• In Glacial Acetic Acid: H+ + CH3COOH -> CH3COOH2+ (Acetonium ion)
• In Liquid Sulfur Dioxide: Proton solvation is minimal due to very poor basicity of SO2.

3. Relative Strength of Acids

The relative strength of an acid refers to its quantitative tendency to donate a proton to a reference base (usually water). It is expressed quantitatively by the acid dissociation constant (Ka) or its negative logarithm (pKa).

pKa = -log10(Ka)

A higher Ka value or lower pKa value corresponds to a stronger acid.

Factors Influencing Relative Acid Strength

1. Electronegativity of the Central Atom: Across a period in the periodic table, as the electronegativity of the element attached to hydrogen increases, the polar character of the H-X bond increases, making proton release easier.
   Order of acidity: CH4 < NH3 < H2O < HF
2. Bond Strength and Ionic Size: Down a group in the periodic table, bond dissociation energy dominates over electronegativity. Larger atomic size leads to longer, weaker H-X bonds.
   Order of hydrohalic acid strength: HF < HCl < HBr < HI
3. Oxidation State of Central Atom in Oxoacids: Higher oxidation state of the central atom increases electron withdrawal from the O-H bond, weakening it and stabilizing the resulting conjugate anion via charge dispersion.
   Order of chloric oxoacids strength: HClO < HClO2 < HClO3 < HClO4
4. Inductive Effect: Electron-withdrawing substituents (-I effect) increase acid strength by stabilizing the conjugate base, whereas electron-donating groups (+I effect) decrease acidity.
   Order of acidity: CH3COOH < ClCH2COOH < Cl2CHCOOH < Cl3CCOOH
5. Resonance Stabilization of Conjugate Base: If the negative charge remaining after proton loss can be delocalized across multiple atoms via resonance, the conjugate base becomes more stable, enhancing the acidity of the parent compound.

4. Types of Acid-Base Reactions

Brönsted-Lowry acid-base interactions encompass several distinct reaction classes based on the nature of reactants and solvent media:

1. Protolysis (Proton Transfer) Reactions

Reactions involving the direct transfer of a proton from an acid to a base without radical or redox changes.

HCl + NH3 -> NH4+ + Cl-

2. Neutralization Reactions

Reaction between equivalent amounts of an acid and a base to yield a salt and solvent molecules.

H3O+ + OH- -> 2 H2O

3. Displacement Reactions

A stronger acid displaces a weaker acid from its salt, or a stronger base displaces a weaker base from its salt.

CH3COONa + HCl -> CH3COOH + NaCl (Acid displacement)

NH4Cl + NaOH -> NH3 + H2O + NaCl (Base displacement)

4. Gas-Phase Acid-Base Reactions

Acid-base proton transfers taking place completely in the absence of any solvent phase.

NH3(g) + HCl(g) -> NH4Cl(s)

5. Solvolytic / Non-Aqueous Acid-Base Reactions

Reactions taking place in solvents other than water (such as liquid NH3, anhydrous H2SO4, or liquid SO2).

NH4+ + NH2- -> 2 NH3 (Neutralization in liquid ammonia)

5. Levelling Solvents and the Levelling Effect

The intrinsic strengths of strong acids or strong bases can be obscured depending on the solvent used. This phenomenon is known as the levelling effect.

Levelling Effect: The phenomenon wherein all acids stronger than the conjugate acid of the solvent are completely dissociated and appear to have identical strength in that solvent.

Levelling Solvents

A solvent is called a levelling solvent for a group of acids if it possesses sufficient basicity to accept protons completely from all those acids, reducing their effective strengths to that of the solvent's solvated proton (e.g., H3O+ in water).

• Water as a Levelling Solvent: Mineral acids such as HCl, HNO3, HClO4, and H2SO4 all appear equally strong in aqueous solution because all are completely converted to H3O+.
• Liquid Ammonia as a Strong Levelling Solvent: Because NH3 is more basic than H2O, even weak acids like CH3COOH behave as strong acids in liquid NH3.

Differentiating Solvents

A solvent is called a differentiating solvent if it is weakly basic (or strongly acidic), allowing different strong acids to dissociate to different extents, thereby enabling the measurement of their relative intrinsic strengths.

Example: When dissolved in glacial acetic acid (a weakly basic solvent), the relative acid strengths can be differentiated as:

HClO4 > HBr > H2SO4 > HCl > HNO3

6. Lewis Acid-Base Concept

In 1923, G. N. Lewis proposed a broader electronic theory of acids and bases independent of protons or solvent systems.

Lewis Acid: Any chemical species (neutral or charged) that can accept an electron pair from a donor species (electrophile).
Lewis Base: Any chemical species (neutral or charged) that can donate an electron pair to an acceptor species (nucleophile).

Mechanism of Lewis Acid-Base Reaction

A Lewis acid-base reaction involves the formation of a coordinate covalent (dative) bond between the donor (base) and acceptor (acid) to yield an acid-base adduct or complex.

A + :B -> A<--B (Adduct)

Example: BF3 + :NH3 -> F3B<--NH3

7. Classification of Lewis Acids

Lewis acids are categorized into five major classes based on their electronic structure and electron-accepting mechanism:

1. Molecules with Incomplete Octet Central Atoms: Electron-deficient compounds whose central atom lacks a complete valence octet.
   Examples: BF3, BCl3, AlCl3, GaCl3, BH3.
2. Simple Metallic Cations: Positively charged metal ions that accept electron pairs into empty valence orbitals.
   Examples: H+, Ag+, Cu2+, Fe3+, Al3+, Mg2+.
   Observation: Cations with higher charge and smaller radii act as stronger Lewis acids.
3. Molecules with Central Atoms Having Vacant d-Orbitals: Central atoms that already possess an octet but can expand their coordination sphere using empty outer d-orbitals.
   Examples: SiF4, SnCl4, SF4, PF5, TiCl4.
   Reaction: SiF4 + 2 F- -> [SiF6]^2-
4. Molecules with Polar Multiple Bonds Between Atoms of Different Electronegativity: Multiple bonds where electron density can shift away upon attack by a nucleophile.
   Examples: CO2, SO2, SO3.
   Reaction: O=C=O + OH- -> HCO3-
5. Elements with Electron Deficiencies / Open Valence Shells: Neutral elemental species that accept electron pairs during reactions.
   Examples: O, S (e.g., in oxidation/sulfidation pathways).

8. Hard and Soft Acids and Bases (HSAB)

Ralph Pearson in 1963 introduced the Hard and Soft Acids and Bases (HSAB) principle to rationalize the stability of coordination compounds and the direction of chemical reactions.

HSAB Principle: Hard acids prefer to bind to hard bases, and soft acids prefer to bind to soft bases. Hard-hard interactions are predominantly ionic, whereas soft-soft interactions are predominantly covalent.

Characteristics of Hard and Soft Species

Classification Table of Acids and Bases

9. Applications of the HSAB Principle

The HSAB principle provides Qualitative explanations for various chemical phenomena:

1. Relative Stability of Complexes:
   [AlF6]^3- is extremely stable because Al3+ (hard acid) interacts with F- (hard base). Conversely, [AlI6]^3- does not readily form because I- is a soft base.
2. Predicting Direction of Chemical Reactions:
   Reactions shift toward combinations that pair hard with hard and soft with soft.
   HgF2 + BeI2 -> BeF2 + HgI2
   Explanation: Be2+ (hard acid) prefers F- (hard base), forming BeF2. Hg2+ (soft acid) prefers I- (soft base), forming HgI2.
3. Solubility of Salts in Water:
   AgI is virtually insoluble in water (Ksp ~ 8.5 x 10^-17) because Ag+ (soft acid) and I- (soft base) form a stable, strongly covalent interaction. AgF is highly soluble because Ag+ (soft) and F- (hard) form a less stable pairing easily hydrated by water.
4. Occurrence of Ores and Minerals in Nature:
   Hard metal cations (Mg2+, Ca2+, Al3+, Ti4+) occur naturally as oxides, carbonates, or fluorides (hard bases). Soft metal cations (Cu+, Ag+, Hg2+, Pb2+) occur as sulfide ores (S^2- is a soft base).
5. Symbiosis Concept:
   Hard ligands tend to cluster on a central atom already bonded to hard ligands; soft ligands cluster on centers bonded to soft ligands. For example, BF3 accepts F- to form [BF4]-, while BH3 accepts H- to form [BH4]-.
6. Heavy Metal Toxicity:
   Soft heavy metal ions like Hg2+, Pb2+, and Cd2+ are toxic because they bind strongly to soft sulfur donor atoms in cellular enzymes (sulfhydryl groups -SH), deactivating vital metabolic catalysts.

10. Types of Inorganic Polymers

Inorganic polymers are high-molecular-weight macromolecules whose backbone chain consists of repeating covalent linkages between non-carbon atoms, although organic side-chains may be present.

Classification Based on Backbone Composition

1. Wholly Inorganic Polymers: Both the main chain and side groups are entirely devoid of carbon atoms.
   Examples: Polymeric sulfur, Polyphosphazenes [N=PCl2]n, Silicates.
2. Organometallic / Partially Inorganic Polymers: The structural backbone consists of inorganic elements (e.g., Si-O, Si-N, P-N), but organic functional groups (e.g., methyl, phenyl) are attached as side chains.
   Examples: Silicones (Polysiloxanes), Polysilanes.

Classification Based on Structural Dimensionality

• One-Dimensional (Linear / Branched): Single chain molecules, flexible and often soluble or meltable (e.g., linear silicones, polyphosphazenes).
• Two-Dimensional (Sheet / Layered): Cross-linked in two dimensions to form continuous planar networks (e.g., mica, talc, layered silicates, graphite).
• Three-Dimensional (Network): Highly cross-linked, rigid 3D framework polymers (e.g., quartz, zeolites, cross-linked silicone resins).

11. Structural Aspects and Applications of Silicones and Siloxanes

Silicones (or organopolysiloxanes) are synthetic organosilicon polymers containing alternating silicon and oxygen atoms in the backbone (-Si-O-Si-) with organic side groups attached directly to the silicon atoms.

General Formula

[R2SiO]n, where R represents alkyl or aryl groups such as methyl (-CH3) or phenyl (-C6H5).

Synthesis of Silicones

Synthesis involves three major stages:

1. Preparation of Chlorosilanes (Direct Process):
   Reaction of methyl chloride with elemental silicon at ~300 °C in the presence of a copper catalyst:
   2 CH3Cl + Si -(Cu)-> (CH3)2SiCl2 (Dimethyldichlorosilane)
2. Hydrolysis of Alkylchlorosilanes:
   (CH3)2SiCl2 + 2 H2O -> (CH3)2Si(OH)2 + 2 HCl
   The resulting silanols are unstable intermediates that undergo rapid condensation.
3. Polymerization via Condensation:
   n (CH3)2Si(OH)2 -> [-(CH3)2Si-O-]n + n H2O

Chain length and branching are controlled by mixing chlorosilanes of different functionalities:

• R3SiCl (Monochlorosilane): Acts as a chain end-blocker / terminator.
• R2SiCl2 (Dichlorosilane): Produces linear chain polymers.
• RSiCl3 (Trichlorosilane): Produces 3D cross-linked resins.

Structural Aspects

• Siloxane Backbone (-Si-O-Si-): Silicon is sp3 hybridized with tetrahedral orientation. The Si-O bond is partially ionic (high thermal stability) with a flexible Si-O-Si bond angle (130°-150°).
• Hydrophobic Outer Shell: Non-polar organic alkyl groups orient outward, imparting exceptional water repellency and low surface tension.

Types and Applications of Silicones

12. Borazines (Inorganic Benzene)

Borazine, B3N3H6, is a cyclic inorganic compound consisting of alternating boron and nitrogen atoms in a six-membered ring. Because it is isoelectronic and isostructural with benzene (C6H6), it is commonly known as "Inorganic Benzene".

Synthesis of Borazine

1. Reaction of Diborane and Ammonia:
   3 B2H6 + 6 NH3 -(180-200 °C)-> 2 B3N3H6 + 12 H2
2. Laboratory Preparation from Boron Trichloride:
   3 BCl3 + 3 NH4Cl -> Cl3B3N3H3 + 9 HCl
   Cl3B3N3H3 + 3 NaBH4 -> B3N3H6 + 3 NaCl + 3 BH3

Structural Comparison: Borazine vs. Benzene

Both borazine and benzene possess planar hexagonal six-membered rings with sp2 hybridized ring atoms and delocalized pi-electrons.

Chemical Reactivity of Borazine

Due to partial ionic character of the B-N bonds (B^+ delta - N^- delta), borazine is significantly more reactive than benzene:

• Addition Reaction with HCl: B3N3H6 + 3 HCl -> B3N3H9Cl3 (H adds to N, Cl adds to B)
• Hydrolysis: B3N3H6 + 6 H2O -> 3 H3BO3 + 3 NH3 + 3 H2

Applications of Borazine

• Precursor for synthesizing specialized Boron Nitride (BN) ceramics and thin films.
• Preparation of high-temperature resistant ceramic coatings and nanocomposites.

13. Silicates: Structure, Classification, and Applications

Silicates are inorganic compounds composed of silicon and oxygen anions, constituting over 90% of the Earth's crust.

Basic Structural Unit

The fundamental building block of all silicates is the orthosilicate tetrahedron [SiO4]^4-.

• Silicon is at the center, bonded to four oxygen atoms at the corners of a regular tetrahedron.
• Silicon undergoes sp3 hybridization.
• Silicate structures are classified based on how many oxygen vertices are shared between adjacent [SiO4] tetrahedra.

Classification of Silicates

Applications of Silicates

1. Construction and Building Materials: Portland cement, concrete, bricks, and structural glass are silicate-based materials.
2. Zeolites as Catalysts and Molecular Sieves: Porous 3D aluminosilicates with uniform pore dimensions are used as molecular sieves, ion-exchange water softeners, and cracking catalysts in petroleum refining.
3. Ceramics and Glass Industry: Kaolin clay and feldspar are primary raw materials for porcelain, pottery, and optical fibers.
4. Thermal and Electrical Insulation: Mica and asbestos are widely used as heat-resistant and electrical insulating barriers.