6Chemical Bonding

📚
CHEMICAL BOND
▢ Definition: Force holding atoms together in molecule/compound.
▢ Formation:
  • •
    Attraction > repulsion
  • •
    Energy ↓ → exothermic
  • •
    Stability ↑
▢ Valency: No. of e⁻ lost/gained/shared by atom.
▢ Classical Valency: No. of H atoms or 2 × no. of O atoms combining with element.
▢ Exception: H follows duplet rule, not octet rule.
▢ Source:
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ELECTRONEGATIVITY
▢ Definition: Tendency of atom in compound to attract shared e⁻ pair.
▢ Bond Prediction:

Table 1: Electronegativity Difference

\(\Delta EN\)
Bond type
Example
\(> 1.8\)
Ionic/electrovalent
\(NaCl\)
\(0.4-1.8\)
Polar covalent
\(HCl,\ H_2O,\ NH_3\)
\(0-0.4\)
Non-polar covalent
\(H_2,\ Cl_2,\ CH_4\)
▢ Important:
  • •
    Polarity of bond \(\propto \Delta EN\)
  • •
    EN of non-metal generally > 2
  • •
    EN of alkali metal generally < 1
  • •
    \(CsF\) → highest ionic character
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TYPES OF CHEMICAL BOND
▢ Primary Bonds:
  • •
    Ionic/electrovalent
  • •
    Covalent
  • •
    Coordinate/dative
▢ Secondary Bonds:
  • •
    Van der Waals force
  • •
    Hydrogen bond
  • •
    Metallic bond
▢ Note: Primary bonds → valency; secondary bonds → not responsible for valency.
📚
IONIC BOND
▢ Definition: Complete transfer of e⁻ from valence shell of electropositive atom to electronegative atom.
▢ Formation:
  • •
    Metal loses e⁻ → cation
  • •
    Non-metal gains e⁻ → anion
  • •
    Electrostatic attraction → ionic bond
▢ Favoured By:
  • •
    Low ionization energy of cation-forming atom
  • •
    High electron affinity of anion-forming atom
  • •
    High lattice energy
  • •
    Low charge on ions
  • •
    Large cation + small anion
▢ Characteristics:
  • •
    Ionic, not molecular
  • •
    Non-directional electrostatic force
  • •
    No space isomerism
  • •
    High m.p. and b.p.
  • •
    Soluble in high dielectric solvent: water
  • •
    Insoluble in non-polar solvent: benzene, ether
  • •
    Solid → poor conductor
  • •
    Molten/aqueous → good conductor
  • •
    Stability due to lattice energy
  • •
    Solubility due to hydration energy
▢ Exceptions:
  • •
    s-block metals usually ionic except Li, Be
  • •
    p-block and d-block metals usually covalent except hydrated form
📚
FAJAN RULE
▢ Core: No bond is 100% ionic; ionic bond has covalent character.
▢ Covalent Character Increases With:
  • •
    Small cation size → high polarising power
  • •
    Large anion size → high polarisability
  • •
    High charge on cation/anion
  • •
    Cation with 18e⁻ shell > cation with 8e⁻ shell
▢ Orders:
❖ Lithium Halides Covalent Character: \(LiI > LiBr > LiCl > LiF\)
❖ Lithium Halides Melting Point: \(LiI < LiBr < LiCl < LiF\)
❖ Halides Covalent Character: \(NaCl < MgCl_2 < AlCl_3\)
❖ CuCl vs NaCl: \(CuCl\) more covalent than \(NaCl\)
📚
COVALENT BOND
▢ Definition: Mutual sharing of e⁻ between atoms.
▢ Examples: \(H_2,\ Cl_2,\ CH_4,\ H_2O,\ CO_2\)
▢ Characteristics:
  • •
    Molecular nature
  • •
    Low m.p. and b.p. than ionic compounds
  • •
    Directional bond
  • •
    Shows isomerism/stereoisomerism
  • •
    Generally soluble in non-polar solvents
  • •
    Poor conductor
  • •
    Van der Waals force between molecules
▢ Types:
❖ Non-polar Covalent: \(H_2,\ Cl_2,\ O_2,\ CH_4\)
❖ Polar Covalent: \(HCl,\ HBr,\ H_2O,\ NH_3\)
❖ Non-polar Molecule with Polar Bonds: \(CO_2,\ BeF_2,\ BF_3,\ CCl_4\)
📚
COORDINATE COVALENT BOND
▢ Definition: Shared e⁻ pair donated by one atom only.
▢ Donor: Atom donating lone pair.
▢ Acceptor: Atom accepting lone pair.
▢ Examples: \(NH_4^+,\ BF_4^-,\ H_3O^+,\ SO_2,\ H_2SO_4\)
▢ Characteristics:
  • •
    Solid/liquid/gas
  • •
    m.p. and b.p. > covalent but < ionic
  • •
    Soluble in organic solvents
  • •
    Sparingly soluble in water
  • •
    Shows isomerism/stereoisomerism
  • •
    Poor conductor
▢ Compounds With Multiple Bonds:

Table 1: Bond Types in Compounds

Compound
Bond present
\(NaCl,\ KNO_3\)
Ionic + covalent
\(NH_4Cl,\ CuSO_4,\ K_4[Fe(CN)_6]\)
Ionic + covalent + coordinate
\(CuSO_4\cdot5H_2O\)
Ionic + covalent + coordinate + hydrogen
📚
DIPOLE MOMENT
▢ Formula: \(\mu = q \times d\)
▢ Meaning: Degree of polarity in polar compound.
▢ Direction: Negative pole → positive pole.
▢ Nature: Vector quantity.
▢ Unit: Debye \((D)\)
▢ Rules:
  • •
    Greater \(\mu\) → greater polarity
  • •
    Symmetrical molecule → net \(\mu = 0\)
  • •
    cis-isomer \(\mu\) > trans-isomer \(\mu\)
  • •
    o-isomer \(\mu\) > m-isomer \(\mu\) > p-isomer \(\mu\)
▢ Examples:
❖ CO₂: \(\mu = 0\), linear, polar bonds cancel.
❖ CCl₄: \(\mu = 0\), tetrahedral symmetry.
❖ SO₂: Polar; lone pair present.
❖ SO₃: Non-polar; no lone pair, symmetrical.
❖ BF₃: \(\mu = 0\); trigonal planar.
❖ NH₃ vs NF₃: \(NH_3\) more polar than \(NF_3\).
▢ CH₃Cl Series: \(CH_3Cl > CH_2Cl_2 > CHCl_3 > CCl_4\) generally for dipole moment.
📚
HYBRIDIZATION
▢ Definition: Mixing of atomic orbitals of comparable energy → equal number of identical hybrid orbitals.
▢ Formula:
❖ Neutral Molecule: \(x = \frac{1}{2}(a+b)\)
❖ Cation: \(x = \frac{1}{2}(a+b-c)\)
❖ Anion: \(x = \frac{1}{2}(a+b+d)\)
▢ Symbols:
  • •
    \(a\) = valence e⁻ of central atom
  • •
    \(b\) = monovalent peripheral atoms
  • •
    \(c\) = positive charge
  • •
    \(d\) = negative charge
  • •
    \(x\) = no. of hybrid orbitals
▢ Hybridization Table:

Table 1: Hybridization, Geometry, Bond Angle

x
Hybridization
Geometry/Shape
Bond angle
Examples
2
\(sp\)
Linear
\(180^\circ\)
\(BeCl_2,\ CO_2,\ C_2H_2\)
3
\(sp^2\)
Trigonal planar / angular
\(120^\circ\)
\(BF_3,\ AlCl_3,\ SO_2,\ NO_2^-\)
4
\(sp^3\)
Tetrahedral / pyramidal / angular
\(109^\circ28'\)
\(CH_4,\ NH_3,\ H_2O\)
5
\(sp^3d\)
Trigonal bipyramidal
\(90^\circ,\ 120^\circ\)
\(PCl_5,\ SF_4,\ ClF_3,\ XeF_2\)
6
\(sp^3d^2\)
Octahedral
\(90^\circ\)
\(SF_6,\ IF_5,\ XeF_4\)
7
\(sp^3d^3\)
Pentagonal bipyramidal
\(72^\circ,\ 90^\circ\)
\(IF_7\)
▢ Important Points:
  • •
    Benzene → all C \(sp^2\)
  • •
    1,3-butadiene → all C \(sp^2\)
  • •
    Diamond → C \(sp^3\)
  • •
    Graphite → C \(sp^2\)
  • •
    \(CO_2\) → linear
  • •
    \(CO_3^{2-}\) → trigonal planar
  • •
    \(PCl_5\) → trigonal bipyramidal
  • •
    \(XeF_2\) → linear due to 3 lone pairs
  • •
    \(XeF_4\) → square planar due to 2 lone pairs
  • •
    In aniline/phenol, N/O lone pair involved in resonance → \(sp^2\)
▢ Multiple Central Atoms Rule:
  • •
    1 \(\pi\)-bond by atom → \(sp^2\)
  • •
    2 \(\pi\)-bonds by atom → \(sp\)
  • •
    Resonance lone pair not counted in hybridization
▢ Carbocation/Carbanion/Free Radical:

Table 1: Carbon Species Hybridization

Species
σ bonds + lone/unpaired e⁻
Hybridization
Carbocation \((C^+)\)
3σ + 0 lp
\(sp^2\)
Carbanion \((C^-)\)
3σ + 1 lp
\(sp^3\)
Free radical \((C\cdot)\)
3σ + 1 unpaired e⁻
\(sp^2\)
📚
VSEPR THEORY
▢ Given By: Gillespie and Nyholm.
▢ Principle: Electron pairs around central atom repel and arrange maximum distance apart.
▢ Repulsion Order: \(lp-lp > lp-bp > bp-bp\)
▢ Bond Angle:
  • •
    More lone pairs → bond angle ↓
  • •
    Same hybridization + shape: more electronegative central atom → bond angle ↑
  • •
    Bulky groups may increase bond angle due to steric repulsion
▢ Orders:
❖ Hydrides of Group 16: \(H_2O > H_2S > H_2Se > H_2Te\)
❖ Group 15 Hydrides: \(NH_3 > PH_3 > AsH_3 > BiH_3\)
❖ Nitrogen Species: \(NH_4^+ > NH_3 > NH_2^-\)
▢ Limitations:
  • •
    Cannot explain highly polar bond species properly
  • •
    Cannot explain extensive delocalized \(\pi\)-system
  • •
    Cannot explain inert pair effect species
📚
VALENCE BOND THEORY
▢ Given By: Heitler and London.
▢ Postulates:
  • •
    Bond formation requires unpaired e⁻
  • •
    Bond formed by pairing of e⁻
  • •
    Maximum e⁻ density lies between bonded atoms
  • •
    Atoms tend to complete shell
▢ Limitations:
  • •
    Cannot explain odd e⁻ molecules: \(O_2,\ NO,\ H_2^+\)
  • •
    Cannot explain coordinate bond and \(\pi\)-bond properly
  • •
    Cannot explain stereochemistry
  • •
    Predicts \(B_2\) diamagnetic, but actually paramagnetic
  • •
    Cannot explain paramagnetism of \(O_2\)
  • •
    Cannot explain existence of \(H_2^+\), \(He_2^+\)
📚
MOLECULAR ORBITAL THEORY
▢ Given By: Hund and Mulliken.
▢ Principle: Atomic orbitals combine → molecular orbitals over whole molecule.
▢ Number Rule: No. of MOs formed = No. of AOs mixed.
▢ Types:
  • •
    Bonding MO → lower energy
  • •
    Antibonding MO → higher energy
▢ Bond Order: \(B.O. = \frac{1}{2}(n_b - n_a)\)
▢ Interpretation:
  • •
    \(B.O. = 0\) → molecule does not exist
  • •
    \(B.O. > 0\) → stable molecule
  • •
    \(B.O.\) ↑ → stability ↑, bond energy ↑, ionization energy ↑
  • •
    \(B.O.\) ↑ → bond length ↓
  • •
    Unpaired e⁻ → paramagnetic
  • •
    No unpaired e⁻ → diamagnetic
▢ Bond Order Table:

Table 1: Common Molecules/Ions

Species
Bond order
Magnetic nature
\(O_2\)
2
Paramagnetic
\(O_2^+\)
2.5
Paramagnetic
\(O_2^-\)
1.5
Paramagnetic
\(O_2^{2-}\)
1
Diamagnetic
\(N_2\)
3
Diamagnetic
\(N_2^+\)
2.5
Paramagnetic
\(N_2^-\)
2.5
Paramagnetic
\(NO\)
2.5
Paramagnetic
\(NO^+\)
3
Diamagnetic
\(NO^-\)
2
Paramagnetic
\(CO\)
3
Diamagnetic
▢ Oxygen Species:
❖ Bond Energy Increasing: \(O_2^{2-} < O_2^- < O_2 < O_2^+\)
❖ Bond Length Decreasing: \(O_2^{2-} > O_2^- > O_2 > O_2^+\)
▢ Bond Order Values:
  • •
    \(B.O.=0\): \(He_2,\ Be_2,\ Ne_2\) → not exist
  • •
    \(B.O.=1\): \(H_2,\ Li_2,\ B_2,\ F_2\)
  • •
    \(B.O.=2\): \(C_2,\ O_2\)
  • •
    \(B.O.=2.5\): \(NO,\ O_2^+,\ N_2^+\)
  • •
    \(B.O.=3\): \(N_2,\ CN^-,\ CO,\ NO^+\)
📚
SIGMA AND PI BONDS
▢ Sigma Bond:
  • •
    Head-on overlap
  • •
    Electron density along internuclear axis
  • •
    Stronger than \(\pi\)-bond
  • •
    Shape of molecule decided by σ-bonds
  • •
    H atom always forms σ-bond
▢ Pi Bond:
  • •
    Sidewise overlap
  • •
    More reactive than σ-bond
  • •
    Present in multiple bonds
▢ Multiple Bond Rule: In multiple bond between 2 atoms: one σ-bond + remaining \(\pi\)-bonds.
▢ Examples:
❖ Acetylene: 3σ + 2π
❖ CO: 1σ + 1π + 1 coordinate bond
❖ Benzene C-C Bond Order: 1.5
❖ 1-butene-3-yne: 7σ + 3π
📚
BOND LENGTH AND BOND ENERGY
▢ Bond Length: Internuclear distance between bonded atoms.
▢ Bond Energy: Energy required to break 1 mole of bonds.
▢ General Relations:
  • •
    Bond order ↑ → bond length ↓
  • •
    Bond order ↑ → bond energy ↑
  • •
    Bond length ↑ → bond energy ↓
  • •
    s-character ↑ → bond energy ↑
  • •
    Lone pair on bonded atoms ↑ → bond energy ↓
▢ Orders:
❖ C-N Bonds Energy: \(C\equiv N > C=N > C-N\)
❖ C-H Strength: C-H strongest than C-N, C-O, C-C
❖ Halogen Bond Energy: \(Cl-Cl > Br-Br > F-F > I-I\)
❖ Hydrogen Halides Bond Energy: \(H-F > H-Cl > H-Br > H-I\)
❖ Carbon Hybrid Bond Energy: \(sp > sp^2 > sp^3\)
📚
HYDROGEN BOND
▢ Definition: Electrostatic attraction between covalently bonded H and electronegative atom \((F,O,N)\) of same/different molecule.
▢ Requirement:
  • •
    H attached to F/O/N
  • •
    \(\Delta EN > 0.7\)
  • •
    Small electronegative atom
  • •
    Usually F, O, N only
▢ Nature:
  • •
    Special case of dipole-dipole/Debye force
  • •
    Stronger than Van der Waals
  • •
    Weaker than covalent bond
  • •
    Strength: about 10–40 kJ mol⁻¹
▢ Types:
❖ Intermolecular:
◉ Definition: Between molecules.
◉ Examples: \(H_2O,\ HF,\ NH_3,\ alcohols,\ acids\)
◉ Effect: m.p., b.p., viscosity, surface tension, solubility ↑
❖ Intramolecular:
◉ Definition: Within same molecule.
◉ Also Called: Chelation.
◉ Effect: Water solubility ↓, m.p./b.p. ↓
◉ Examples: o-nitrophenol, o-hydroxybenzaldehyde, chloral hydrate
▢ Effects:
  • •
    Water liquid despite low molecular weight
  • •
    Alcohol b.p. > ether/thiol
  • •
    Polyhydric alcohol b.p. high
  • •
    1° and 2° amines show H-bonding; 3° amines do not
  • •
    Monosaccharides, disaccharides, urea highly water-soluble
  • •
    RCOOH forms dimer in benzene
▢ Boiling Point Orders:
❖ Hydrogen Halides: \(HF > HI > HBr > HCl\)
❖ Group 16 Hydrides: \(H_2O > H_2Te > H_2Se > H_2S\)
❖ HF, H₂O, NH₃: \(H_2O > HF > NH_3\)
▢ DNA and Protein:
  • •
    DNA base pairs stabilized by H-bond
  • •
    Secondary structure of protein stabilized by H-bond
📚
VAN DER WAALS FORCE
▢ Definition: Weakest intermolecular force.
▢ Types:
❖ Ion-Dipole: Present in electrolytic solution; strongest Van der Waals force.
❖ Dipole-Dipole: Debye force; between polar molecules.
❖ Instantaneous/Induced Dipole: London/dispersion force; weakest; between non-polar species.
▢ Depends On:
  • •
    Size ↑ → force ↑
  • •
    Surface area ↑ → force ↑
  • •
    Polarity ↑ → force ↑
  • •
    Pressure ↑ → force ↑
  • •
    Temperature ↑ → force ↓
▢ Examples:
❖ Noble Gases Force Order: \(He < Ne < Kr < Xe\)
❖ Pentane Isomers b.p.: n-pentane > isopentane > neopentane
❖ Ideal Gas Condition: High temperature + low pressure
📚
METALLIC BOND
▢ Concept: Positive metal kernels + delocalized mobile valence e⁻ sea.
▢ Definition: Net attraction between free electrons and metal cations.
▢ Strength Depends On:
  • •
    No. of free e⁻ ↑ → metallic bond ↑
  • •
    Atomic size ↓ → metallic bond ↑
  • •
    d-block metals hard due to additional covalent character
▢ Orders:
❖ Hardness: \(Na < Mg < Al\)
❖ Alkali Metals Hardness: \(Li > Na > K > Rb > Cs\)
📚
OCTET RULE DEVIATIONS
▢ Incomplete Octet:
❖ Definition: Electron-deficient molecules.
❖ Examples: \(BeCl_2,\ BF_3,\ B_2H_6,\ CO,\ NO\)
▢ Expanded Octet:
❖ Definition: Central atom has > 8 valence e⁻.
❖ Examples:
  • •
    \(PCl_5\): P has 10 e⁻
  • •
    \(ClF_3\): Cl has 10 e⁻
  • •
    \(SF_6\): S has 12 e⁻
  • •
    \(OsO_4\): Os has 16 e⁻
▢ Odd Electron Bonds:
❖ Definition: Molecules/ions with odd e⁻.
❖ Examples:
◉ \(NO\): 1 unpaired e⁻
◉ \(NO_2\): 1 unpaired e⁻
◉ \(O_2^-\): 1 unpaired e⁻
◉ \(OH\): 1 unpaired e⁻
📚
SPECIAL HIGH-YIELD POINTS
▢ Read and Digest:
    ◉ **type: bullet
  1. \(KH_2\) Fremy's salt exists due to H-bond.
  2. \(HF_2^-\) has strongest H-bond.
  3. \(HCl_2^-\) cannot exist due to absence of H-bond.
  4. Banana bond in diborane: 3 atoms + 2 e⁻.
  5. \(NH_4Cl\) has ionic + covalent + coordinate bonds.
  6. \([Cu(NH_3)*4]SO_4\) has ionic + covalent + coordinate bonds.
  7. \(NaCN\) and \(NH_4Cl\) contain ionic, covalent, coordinate bonds.
  8. \(Cl_2\) non-polar; insoluble in water due to no H-bonding.
  9. \(H_2O\) liquid; \(H_2S\) gas due to H-bonding in water.
  10. Oxygen forms diads.
  11. Hyperconjugation = delocalization of σ and π bonds.
  12. Acetylene has 3σ and 2π bonds.
  13. Solid NaCl poor conductor; molten NaCl good conductor.
  14. Born-Haber cycle gives stability of ionic compound.
  15. Graphite sheets held by Van der Waals force.
  16. Diamond carbon atoms joined by covalent bonds.
  17. \(H_2O\) angular and dipolar; \(BeF_2\) linear and non-polar.
  18. Maximum H-bonds by water molecule in ice = 4.
  19. \(sp^3d\) uses \(d*{z^2}\).
  20. \(d^2sp^3\) octahedral uses \(d*{x^2-y^2}\), \(d*{z^2}\).
  21. Blue vitriol: 4 water molecules coordinate to \(Cu^{2+}\), 5th water by H-bond.
  22. \(CCl_4\) covalent; \(CaH_2\) electrovalent.
  23. Bond order of C-C in benzene = 1.5.
  24. \(SiF_4\) has zero dipole moment.
  25. Stability/m.p./b.p. of ionic solid \(\propto\) lattice energy.
Q1.
Chemical bond implies
Q2.
Chemical bond formation takes place when
Q3.
Formation of chemical bond is accompanied by
Q4.
Stability of ionic compound is due to
Q5.
Fused ionic compounds
Q6.
Most favourable condition for electrovalency is
Q7.
Sodium chloride is ionic because
Q8.
Covalent compounds have low m.p. because
Q9.
Covalent compounds are soluble in
Q10.
Which is least ionic?
Q11.
Double bond connecting two atoms shows sharing of
Q12.
Compound containing both ionic and covalent bonds is
Q13.
Coordinate covalent compounds are formed by
Q14.
Type of bonds present in CuSO₄·5H₂O are
Q15.
Bonds present in ammonium chloride are
Q16.
Which compound contains coordinate covalent bond?
Q17.
Hybridisation involves
Q18.
Structure of ammonia is
Q19.
Structure of SO₂ is
Q20.
Carbon atoms in diamond are bonded in configuration
Q21.
Hybridisation in SO₃ is
Q22.
In NO₃⁻, central atom uses hybrid orbitals of type
Q23.
BF₃ shows hybridisation
Q24.
Structure of ICl₃ is
Q25.
Pair of species having identical shape is
Q26.
Correct decreasing sequence of bond angle of hydrides is
Q27.
Strongest bond among given is
Q28.
Longest bond is
Q29.
Covalent bond formed by axial overlap is
Q30.
Hydrogen bonding is maximum in
Q31.
Hydrogen bond is strongest in
Q32.
Compound containing species KHF₂ has
Q33.
Strength of hydrogen bond is intermediate between
Q34.
Coupling between base units of DNA is through
Q35.
Water is liquid and H₂S is gas because water has
Q36.
NH₃ has much higher boiling point than PH₃ because
Q37.
Correct order of decreasing polarizability is
Q38.
Which molecule has zero dipole moment?
Q39.
Which molecule has maximum dipole moment?
Q40.
The ion isoelectronic with CO is
Q41.
Bond order increases when
Q42.
VSEPR theory cannot explain
Q43.
Which molecule has planar structure?
Q44.
Which is non-polar molecule?
Q45.
Which has square planar shape?
Q46.
Strongest bond is
Q1.
Which of the following is most volatile?
📅MOE 2062
Q2.
Which of the following compounds has two lone pairs of electron and 2 bond pairs?
📅MOE 2060
Q3.
Which does not obey octet rule?
📅IOM 2008•MOE 2003
Q4.
An element forms XCl3, X2O5, XO2 but can't form XCl5. The element may be
📅B.E. 2009
Q5.
Which of the following elements shows sp, sp2 and sp3 hybridization?
📅MOE 2002
Q6.
Which of the following is incorrect?
📅MOE 2002
Q7.
Which of the following is non-planar?
📅MOE 2002
Q8.
The valency of C in CO is 2. The valency of C in CO2 is 4. The valency of C in acetylene is
📅MOE 2002
Q9.
When s and p orbitals hybridise, then
📅IOM 1997
Q10.
The electronic configuration of an element is 2, 8, 1. Which of the following statement is true?
📅IOM 1996
Q11.
The no. of 'σ' and 'π' bonds in 'C' atoms of CaC2 is
📅BPKIHS 2006
Q12.
Which of the following is a pseudo halide?
📅BPKIHS 1994
Q13.
The shape of XeF2 molecule is
📅BPKIHS
Q14.
Which of the following hybridization is seen in CO2 molecule?
📅BPKIHS 2006
Q15.
Which of the following is not true about CCl4?
📅BPKIHS 2005
Q16.
The hybridization of H2O is
Q17.
Diamond has
📅BPKIHS
Q18.
Shape of Ammonia is
📅BPKIHS