1. Introduction to Chemical Bonding
A chemical bond is the attractive force that holds atoms or ions together in a molecule or compound.
Most atoms do not remain completely isolated because they can attain a more stable electronic arrangement by losing, gaining or sharing electrons.
The electrons involved in chemical bonding are mainly the valence electrons, which are present in the outermost shell of an atom.
Why do atoms form bonds?
- To attain a more stable electronic arrangement.
- To reduce the energy of the system.
- To achieve a noble-gas-like configuration in many cases.
- To form stable molecules and compounds.
Main types of chemical bonds
- Ionic bond
- Covalent bond
- Coordinate covalent bond
- Metallic bond
- Hydrogen bond
2. Octet Rule
The octet rule states that atoms tend to gain, lose or share electrons so that they have eight electrons in their valence shell.
The idea is based on the stability of noble gases such as Ne, Ar and Kr.
Example: formation of NaCl
Sodium has electronic configuration:
Sodium can lose one electron:
Chlorine has:
Chlorine gains one electron:
The oppositely charged ions attract each other:
Limitations of the octet rule
- It does not explain all molecules.
- Some molecules have incomplete octets.
- Some elements form expanded octets.
- Odd-electron molecules cannot satisfy the normal octet.
Incomplete octet
Example: BF₃. Boron is surrounded by only six electrons in the simplest Lewis structure.
Expanded octet
Elements from the third period and beyond can sometimes have more than eight electrons around the central atom.
Examples: PCl₅ and SF₆.
Odd-electron molecules
Molecules such as NO contain an odd number of valence electrons, so a complete octet for every atom is not possible.
3. Lewis Structures
A Lewis structure represents valence electrons using dots and shows how atoms share electrons in a molecule.
General steps for drawing a Lewis structure
- Count the total number of valence electrons.
- Choose the central atom.
- Connect the atoms using single bonds.
- Complete the octet of the outer atoms.
- Place remaining electrons on the central atom.
- If necessary, form multiple bonds.
- Check formal charges.
Example: CO₂
Carbon has 4 valence electrons and each oxygen has 6.
The basic arrangement is:
To satisfy the octet of carbon and oxygen, two double bonds are formed:
Thus the Lewis structure of CO₂ contains two C=O double bonds.
Example: NH₃
Nitrogen has 5 valence electrons and three hydrogen atoms contribute 3 electrons.
Three N-H bonds use six electrons. The remaining two electrons form one lone pair on nitrogen.
4. Ionic Bond
An ionic bond is formed by the electrostatic attraction between oppositely charged ions.
Usually one atom loses electron(s) and another atom gains electron(s).
Formation of NaCl
Characteristics of ionic compounds
- They usually have high melting and boiling points.
- They form crystalline solids.
- They are generally hard and brittle.
- Many dissolve in polar solvents such as water.
- They conduct electricity in molten or aqueous state.
- They generally do not conduct electricity in the solid state.
Why does solid NaCl not conduct electricity?
In solid NaCl, Na⁺ and Cl⁻ ions are fixed in the crystal lattice. They cannot move freely. Therefore, electric current cannot pass through the solid easily.
In molten NaCl or aqueous NaCl, ions become mobile and can carry electric charge.
5. Covalent Bond
A covalent bond is formed when two atoms share one or more pairs of electrons.
Single covalent bond
One pair of electrons is shared.
Double covalent bond
Two pairs of electrons are shared.
Triple covalent bond
Three pairs of electrons are shared.
Examples
| Molecule | Bond | Number of shared electron pairs |
|---|---|---|
| H₂ | Single | 1 |
| O₂ | Double | 2 |
| N₂ | Triple | 3 |
Coordinate covalent bond
In a coordinate covalent bond, both electrons in the shared pair are supplied by the same atom.
This is explained in more detail in the next section.
6. Coordinate Covalent Bond
A coordinate covalent bond is a covalent bond in which the two bonding electrons are donated by one atom.
The atom donating the electron pair is called the donor, while the atom accepting the pair is the acceptor.
Example: formation of NH₄⁺
Ammonia contains a lone pair on nitrogen:
The H⁺ ion has no electrons available for forming the bond. Nitrogen donates its lone pair to H⁺.
The newly formed bond is represented as a coordinate bond during formation.
Example: H₃O⁺
Oxygen donates one lone pair to the proton.
7. Formal Charge
Formal charge is the hypothetical charge assigned to an atom in a Lewis structure by assuming that bonding electrons are shared equally.
Where:
- V = valence electrons of the free atom
- N = non-bonding electrons
- B = bonding electrons
Example: nitrogen in NH₃
Nitrogen has 5 valence electrons.
It has two non-bonding electrons and six bonding electrons.
Therefore nitrogen has formal charge zero in NH₃.
8. Fajan's Rule
Fajan's rule helps explain the covalent character present in compounds that are mainly considered ionic.
The basic idea is that a cation can distort or polarize the electron cloud of an anion.
Covalent character increases when:
- The cation is small.
- The cation has a high positive charge.
- The anion is large.
- The anion has a high negative charge.
A small highly charged cation has high polarizing power. A large anion is more easily polarized.
Example
AlCl₃ has greater covalent character than NaCl because Al³⁺ has much greater polarizing power than Na⁺.
9. Lattice Energy
Lattice energy is associated with the formation or separation of an ionic crystal lattice.
Depending on the convention used, lattice energy may be defined as the energy released when one mole of an ionic solid forms from gaseous ions, or the energy required to separate the solid into gaseous ions.
Therefore, always check which convention is being used in the question.
Factors affecting lattice energy
- Charge on ions
- Size of ions
- Distance between ions
Higher ionic charge generally produces stronger electrostatic attraction and greater lattice energy magnitude.
Smaller ions can approach each other more closely, increasing electrostatic attraction.
10. Theories of Covalent Bonding
Valence Bond Theory
Valence Bond Theory explains a covalent bond as the result of overlap between atomic orbitals containing electrons with suitable spin.
Greater overlap generally means a stronger bond.
Types of orbital overlap
Overlap can occur along the internuclear axis or sideways.
Sigma bond
A sigma bond is formed by head-on overlap of orbitals along the internuclear axis.
Pi bond
A pi bond is formed by sideways overlap of parallel p orbitals.
11. VSEPR Theory
VSEPR stands for Valence Shell Electron Pair Repulsion.
According to this theory, electron pairs around the central atom repel each other and arrange themselves as far apart as possible.
Order of repulsion
Lone pairs occupy more space around the central atom than bonding pairs. Therefore, lone pairs produce greater repulsion.
Example: CH₄
Carbon has four bonding pairs and no lone pair. The four pairs arrange themselves tetrahedrally.
Example: NH₃
Nitrogen has three bonding pairs and one lone pair.
Example: H₂O
Oxygen has two bonding pairs and two lone pairs.
12. Molecular Shapes
| Species | Electron pairs / domains | Shape | Approx. angle |
|---|---|---|---|
| BeCl₂ | 2 bonding pairs | Linear | 180° |
| BF₃ | 3 bonding pairs | Trigonal planar | 120° |
| CH₄ | 4 bonding pairs | Tetrahedral | 109.5° |
| NH₃ | 3 bonding + 1 lone pair | Trigonal pyramidal | 107° |
| H₂O | 2 bonding + 2 lone pairs | Bent | 104.5° |
| PCl₅ | 5 bonding pairs | Trigonal bipyramidal | 90°, 120° |
| SF₆ | 6 bonding pairs | Octahedral | 90° |
13. Hybridization
Hybridization is the mixing of atomic orbitals of comparable energy to form new equivalent hybrid orbitals suitable for bonding.
sp hybridization
One s orbital mixes with one p orbital.
Geometry: linear
Bond angle: 180°
Example: BeCl₂ and carbon atoms in ethyne.
sp² hybridization
Geometry: trigonal planar
Bond angle: approximately 120°
Example: BF₃ and carbon atoms in ethene.
sp³ hybridization
Basic geometry: tetrahedral
Example: CH₄.
sp³d hybridization
Geometry: trigonal bipyramidal.
Example: PCl₅.
sp³d² hybridization
Geometry: octahedral.
Example: SF₆.
sp² → trigonal planar → 120°
sp³ → tetrahedral → 109.5°
sp³d → trigonal bipyramidal
sp³d² → octahedral
Hybridization and carbon compounds
| Compound | Carbon hybridization | Basic geometry |
|---|---|---|
| CH₄ | sp³ | Tetrahedral |
| C₂H₄ | sp² | Trigonal planar around each C |
| C₂H₂ | sp | Linear |
14. Polarity of Chemical Bonds and Molecules
A covalent bond becomes polar when the bonded atoms have different electronegativities and therefore attract the shared electrons unequally.
Polar covalent bond
In HCl, chlorine is more electronegative than hydrogen.
The bond has a dipole.
Non-polar covalent bond
In H₂, both atoms have the same electronegativity.
The electrons are shared equally, so the bond is non-polar.
Polar bond does not always mean polar molecule
CO₂ contains polar C=O bonds, but the molecule is linear and the bond dipoles cancel each other.
Therefore CO₂ is overall non-polar.
H₂O contains polar O-H bonds and has a bent shape. The dipoles do not cancel, so water is polar.
15. Resonance
Resonance occurs when one Lewis structure cannot adequately represent the actual electron distribution of a molecule or ion.
More than one valid contributing structure can then be drawn.
Example: carbonate ion
The carbonate ion, CO₃²⁻, can be represented using several equivalent resonance structures in which the position of the double bond changes.
The real structure is not rapidly changing between these forms. Instead, the actual molecule is represented by a resonance hybrid.
Effects of resonance
- It can increase stability.
- It can make equivalent bonds have similar properties.
- It can distribute charge over several atoms.
- It can affect bond length and bond strength.
16. Important Bond Parameters
1. Bond length
Bond length is the average distance between the nuclei of two bonded atoms.
Generally:
So a triple bond is generally shorter than a double bond, and a double bond is shorter than a single bond between the same atoms.
2. Bond angle
Bond angle is the angle between two bonds around a central atom.
Examples:
3. Bond energy
Bond energy is the energy required to break a bond under specified conditions.
Stronger bonds generally require more energy to break.
4. Bond order
Bond order represents the number of bonding electron pairs shared between two atoms in simple Lewis structures.
17. Molecular Orbital Theory
Molecular Orbital Theory explains bonding by considering the combination of atomic orbitals to form molecular orbitals that belong to the whole molecule.
Formation of molecular orbitals
When two atomic orbitals combine, molecular orbitals are formed. They can be bonding or antibonding.
Bonding molecular orbital
A bonding molecular orbital has lower energy than the original atomic orbitals and generally increases stability.
Antibonding molecular orbital
An antibonding molecular orbital has higher energy and tends to oppose bonding.
Antibonding orbitals are represented by an asterisk:
Bond order in MOT
Where:
- Nb = number of electrons in bonding molecular orbitals
- Na = number of electrons in antibonding molecular orbitals
Example: H₂
H₂ has two electrons, both occupying the bonding σ1s orbital.
Therefore H₂ is stable.
Magnetic behaviour
A species with unpaired electrons is paramagnetic.
A species with all electrons paired is diamagnetic.
18. Hydrogen Bonding
Hydrogen bonding is an attractive interaction involving hydrogen attached to a highly electronegative atom such as N, O or F and another electronegative atom having a lone pair.
Types of hydrogen bonding
Intermolecular hydrogen bonding
It occurs between different molecules.
Example: water molecules interact with each other through hydrogen bonds.
Intramolecular hydrogen bonding
It occurs within the same molecule.
Certain compounds can form an internal hydrogen bond when the appropriate donor and acceptor groups are close enough.
Effects of hydrogen bonding
- Raises boiling point in many substances.
- Influences solubility.
- Affects viscosity.
- Helps determine molecular association.
- Is important in biological molecules.
Why does water have a relatively high boiling point?
Water molecules form extensive hydrogen bonding. Extra energy is required to separate the molecules sufficiently during boiling.
19. Metallic Bonding
Metallic bonding is the attraction between positively charged metal centres and delocalized electrons in a metal.
The outer electrons in a metal are not restricted to one particular atom. They can move through the metallic structure.
Properties explained by metallic bonding
- Electrical conductivity
- Thermal conductivity
- Malleability
- Ductility
- Metallic lustre
The mobility of electrons is particularly important in explaining the electrical conductivity of metals.
20. Important Comparisons
Ionic bond vs Covalent bond
| Point | Ionic bond | Covalent bond |
|---|---|---|
| Basic process | Electron transfer followed by electrostatic attraction | Electron sharing |
| Particles | Ions | Molecules or covalent networks |
| Typical melting point | Generally high | Often lower for molecular substances |
| Electrical conduction | Molten/aqueous forms usually conduct | Usually poor for molecular covalent substances |
Bond order and bond length
Bond pair and lone pair
| Bond pair | Lone pair |
|---|---|
| Shared between atoms | Located mainly on one atom |
| Forms a covalent bond | Does not directly form a bond |
| Usually occupies less repulsive space | Produces stronger repulsion |
21. Worked Problems
Calculate the formal charge on nitrogen in NH₃.
Given:
Formula:
Substitution:
Therefore, the formal charge on nitrogen is zero.
A molecule has 8 electrons in bonding molecular orbitals and 2 electrons in antibonding molecular orbitals. Find the bond order.
Formula:
Substitution:
Therefore, the molecule has bond order 3 and contains a triple-bond-like bond.
Determine the number of sigma and pi bonds in ethene, C₂H₄.
Structure:
The C=C double bond contains:
There are four C-H bonds:
Therefore:
Determine the hybridization of carbon in C₂H₂.
Ethyne contains a C≡C triple bond.
Each carbon forms two sigma-bonding directions: one toward hydrogen and one toward the other carbon.
Therefore each carbon is sp hybridized.
Predict the shape of NH₃.
Nitrogen has five valence electrons. In NH₃, it forms three N-H bonds and retains one lone pair.
Four electron domains give tetrahedral electron-pair geometry. Because one position contains a lone pair, the molecular shape becomes trigonal pyramidal.
22. Common Mistakes
Ionic bonding involves electron transfer followed by electrostatic attraction. Covalent bonding involves sharing of electrons.
BF₃, PCl₅, SF₆ and odd-electron species show that there are important exceptions.
Lone pairs strongly affect molecular shape and bond angle.
Molecular geometry must also be considered. CO₂ has polar bonds but is overall non-polar because the dipoles cancel.
Never count a double bond as two sigma bonds.
Resonance structures are contributing representations of one actual electronic structure.
NH₃ has tetrahedral electron-pair geometry but trigonal pyramidal molecular shape.
23. TU Examination Preparation
For Chemical Bonding, prepare definitions first and then practise structures, comparisons and explanations. Numerical questions are usually easier when the formula and units are written clearly.
Important 2-Mark Questions
- Define chemical bond.
- State the octet rule.
- What is a coordinate bond?
- Define formal charge.
- What is resonance?
- Define hybridization.
- What is a sigma bond?
- What is a pi bond?
- Define hydrogen bonding.
- What is lattice energy?
Important 5-Mark Questions
- Explain the octet rule and its limitations.
- Explain VSEPR theory with suitable examples.
- Explain sp, sp² and sp³ hybridization.
- Explain Fajan's rule.
- Differentiate ionic and covalent bonds.
- Explain hydrogen bonding and its effects.
- Explain Lewis structures with examples.
Important 10-Mark Questions
- Explain Valence Bond Theory and discuss sigma and pi bond formation.
- Explain VSEPR theory and predict the shapes of CH₄, NH₃, H₂O, PCl₅ and SF₆.
- Explain Molecular Orbital Theory and calculate bond order for suitable molecules.
- Explain different types of chemical bonds with examples and compare their properties.
- Explain hybridization and its relation to molecular geometry.
24. Quick Revision
- Valence electrons are mainly involved in chemical bonding.
- Ionic bonding involves electron transfer and attraction between oppositely charged ions.
- Covalent bonding involves sharing of electron pairs.
- Coordinate bonding occurs when both bonding electrons are supplied by one atom.
- Octet rule is useful but has exceptions.
- Formal charge helps in evaluating Lewis structures.
- Small highly charged cations increase polarizing power.
- VSEPR explains molecular shapes using electron-pair repulsion.
- Lone pair repulsion is stronger than bond-pair repulsion.
- sp → linear → 180°.
- sp² → trigonal planar → 120°.
- sp³ → tetrahedral → 109.5°.
- sp³d → trigonal bipyramidal.
- sp³d² → octahedral.
- Single bond = 1σ.
- Double bond = 1σ + 1π.
- Triple bond = 1σ + 2π.
- Polar bonds do not necessarily make the whole molecule polar.
- Resonance represents delocalization of electrons.
- Higher bond order generally means shorter and stronger bonds.
- Bond order in MOT = 1/2(Nb − Na).
- Unpaired electrons → paramagnetic.
- All electrons paired → diamagnetic.
- Hydrogen bonding strongly affects boiling point and other physical properties.