JEE Chemistry Chemical Bonding and Molecular Structure Notes Set 10

Ionic Bond and Lattice Enthalpy

An ionic compound forms readily when (1) the electronegativity difference is large, (2) the ionisation enthalpy of the metal is low, (3) the electron gain enthalpy of the non-metal is highly negative, and (4) the lattice enthalpy is high. Lattice enthalpy is the energy needed to separate one mole of a solid ionic compound into gaseous ions; it rises with ionic charge and falls with ionic size.
Ionic compounds are solids (not gases) at room temperature, are almost non-conducting as solids but conduct when molten or dissolved, and dissolve in polar solvents of high dielectric constant such as water. Small, highly charged ions are strongly hydrated: the hydrated radius increases from \( \ce{Sr^2+} \) to \( \ce{Ca^2+} \), \( \ce{Mg^2+} \) and \( \ce{Be^2+} \), so the ionic mobility in water falls in the same order, \( \ce{Sr^2+} \) > \( \ce{Ca^2+} \) > \( \ce{Mg^2+} \) > \( \ce{Be^2+} \).

Isomorphism

Crystals of different compounds with the same crystal form are isomorphous (for example \( \ce{FeSO4.7H2O} \) and \( \ce{MgSO4.7H2O} \), and all alums). The conditions are: the same formula type, ions of similar (not too different) size, similar shapes of the cations and of the anions, and similar polarisation properties. \( \ce{SO4^2-} \) and \( \ce{MnO4^-} \) are both tetrahedral, so \( \ce{BaSO4} \) and \( \ce{KMnO4} \) are isomorphous. \( \ce{NaNO3} \) and \( \ce{NaClO3} \) have the same formula type but are not isomorphous because \( \ce{NO3^-} \) is planar and \( \ce{ClO3^-} \) is pyramidal.

Fajans' Rules

A cation polarises the neighbouring anion by distorting its electron cloud; the more the distortion, the more covalent character the compound has. Covalent character increases with: (1) smaller cation size (\( \ce{BeCl2} \) > \( \ce{MgCl2} \) > \( \ce{CaCl2} \) > \( \ce{SrCl2} \) > \( \ce{BaCl2} \)); (2) larger anion size (LiF < LiCl < LiBr < LiI); (3) higher charge on the cation (\( \ce{NaCl} \) < \( \ce{MgCl2} \) < \( \ce{AlCl3} \)); (4) higher charge on the anion (\( \ce{AlF3} \) < \( \ce{Al2O3} \) < AlN); (5) a pseudo-inert-gas (18-electron) cation, which polarises more than an inert-gas (8-electron) cation of similar size, so \( \ce{CuCl} \) is more covalent than \( \ce{NaCl} \).

Applications of Fajans' Rules

Solubility: \( \ce{Ag2S} \) is less soluble than \( \ce{Ag2O} \) (larger, more polarisable \( \ce{S^2-} \)), and \( \ce{Fe(OH)3} \) is less soluble than \( \ce{Fe(OH)2} \) (\( \ce{Fe^3+} \) is smaller and more highly charged). AgI is less soluble than AgF, which is ionic and soluble.
Colour: larger anions are polarised more, so their electrons can absorb visible light. AgCl is white while AgBr and AgI are yellow; \( \ce{SnCl2} \) is white and \( \ce{SnI2} \) is red-orange; \( \ce{PbCl2} \) is white and \( \ce{PbI2} \) is yellow.
Melting point: more covalent means lower melting point. Melting points rise from \( \ce{BeCl2} \) to \( \ce{BaCl2} \) (more ionic) and fall from \( \ce{CaF2} \) to \( \ce{CaI2} \) (anion polarisation grows). \( \ce{AlCl3} \) melts far lower than \( \ce{NaCl} \). Among \( \ce{NaCl} \), \( \ce{MgCl2} \), \( \ce{AlCl3} \) and LiCl the ionic \( \ce{NaCl} \) melts highest.

Polarity and Dipole Moment

A bond between unlike atoms is polar. Dipole moment \( \mu = q \times d \) is a vector, measured in debye: 1 D = \( 3.336 \times 10^{-30} \) C m = \( 10^{-18} \) esu cm. The dipole moment of a molecule is the vector sum of its bond moments (and lone-pair moments), so for two bond moments p and q at angle \( \theta \), the resultant is \( R = \sqrt{p^2 + q^2 + 2pq\cos\theta} \).
Polar bonds can still give a non-polar molecule when the shape is symmetrical: zero dipole for \( \ce{BF3} \), \( \ce{CO2} \), \( \ce{SO3} \), \( \ce{CF4} \), \( \ce{CCl4} \), \( \ce{PCl5} \) (gas), \( \ce{SF6} \), \( \ce{XeF2} \), \( \ce{XeF4} \), \( \ce{CS2} \) and \( \ce{BeCl2} \). Angular or unsymmetrical molecules such as \( \ce{SO2} \), \( \ce{H2O} \), \( \ce{NH3} \), \( \ce{NF3} \), \( \ce{SF4} \) and \( \ce{H2S} \) have a non-zero dipole moment. A centre of symmetry means zero dipole moment (for example trans-\( \ce{N2F2} \)).

Dipole Moment Trends

HF (1.92 D) > HCl (1.08) > HBr (0.78) > HI (0.38). \( \ce{CH3Cl} \) (1.86) > \( \ce{CH2Cl2} \) (1.6) > \( \ce{CHCl3} \) (1.0) > \( \ce{CCl4} \) (0). \( \ce{NH3} \) has a larger dipole moment than \( \ce{NF3} \) because in \( \ce{NH3} \) the N-H bond moments point toward N and add to the lone-pair moment, while in \( \ce{NF3} \) the N-F bond moments point toward F and oppose the lone-pair moment. For disubstituted benzenes usually ortho > meta > para, and cis isomers usually exceed trans, with exceptions.
Percentage ionic character = (observed moment / moment for 100% ionic) x 100, where the 100% ionic moment is \( e \times d \) (4.8 D per angstrom of bond length).

Acidic, Basic and Amphoteric Oxides

Non-metal oxides are generally acidic (\( \ce{SO2} \), \( \ce{SO3} \), \( \ce{CO2} \)) and metal oxides generally basic (\( \ce{Na2O} \), \( \ce{CaO} \)). Amphoteric oxides react with both acids and bases: \( \ce{BeO} \), \( \ce{Al2O3} \), \( \ce{ZnO} \), \( \ce{SnO} \), \( \ce{SnO2} \), \( \ce{PbO} \). Neutral oxides such as CO, \( \ce{N2O} \) and NO react with neither. Water is amphiprotic.
Across a period the oxides go from strongly basic to strongly acidic: \( \ce{Na2O} \) (strongly basic), MgO (basic), \( \ce{Al2O3} \) (amphoteric), \( \ce{SiO2} \) (weakly acidic), \( \ce{P4O10} \), \( \ce{SO3} \) and \( \ce{Cl2O7} \) (strongly acidic). Down a group the basic character rises (the alkali metal oxides). For one element, higher oxidation state means more acidic: MnO (basic) < \( \ce{Mn2O3} \) < \( \ce{MnO2} \) (amphoteric) < \( \ce{MnO3} \) < \( \ce{Mn2O7} \) (acidic).

Acid Strength of Oxyacids

Across a period the acid strength rises: \( \ce{H3BO3} \) < \( \ce{H2CO3} \) < \( \ce{HNO3} \), and \( \ce{H2SiO3} \) < \( \ce{H3PO4} \) < \( \ce{H2SO4} \) < \( \ce{HClO4} \). Down a group it falls: \( \ce{HNO3} \) > \( \ce{H3PO4} \) > \( \ce{H3AsO4} \) and \( \ce{HClO4} \) > \( \ce{HBrO4} \) > \( \ce{HIO4} \). For one element, more oxygen atoms or higher oxidation state means a stronger acid: \( \ce{HClO4} \) > \( \ce{HClO3} \) > \( \ce{HClO2} \) > HClO and \( \ce{HNO3} \) > \( \ce{HNO2} \).

Worked Examples

1. Why is AgI yellow although AgCl is white? The large \( \ce{I^-} \) ion is polarised much more by \( \ce{Ag^+} \) than \( \ce{Cl^-} \) is, so the electrons absorb part of the visible light.
2. In water, each O-H bond moment is x and the angle is 104 degrees. If the resultant is 1.85 D (cos 104 = -0.25): \( 1.85^2 = 2x^2 + 2x^2(-0.25) = 1.5x^2 \), so x = 1.51 D.
3. HCl has \( \mu \) = 1.03 D and bond length 1.275 angstrom. For 100% ionic character \( \mu = 4.8 \times 1.275 = 6.12 \) D, so the ionic character is 1.03/6.12 = 17%.