Different bonding types give different properties to materials. This topic compares the main types side by side. CBSE Class 11 and NEET test distinguishing features and examples in the chemical bonding chapter.
Core Concepts
Ionic bonding
Complete electron transfer. Forms ions — cations (metals) and anions (non-metals). High lattice energy gives high MP/BP. Conducts in molten state. Example — NaCl, MgO.
Formation of NaCl:
The overall process is energetically favourable because the lattice energy (energy released when ions form the crystal) is very large — 787 kJ/mol for NaCl. This more than compensates for the energy cost of ionisation.
The lattice energy cannot be measured directly. We use the Born-Haber cycle — an application of Hess’s law:
where = lattice energy. Higher lattice energy means stronger ionic compound. MgO has higher lattice energy than NaCl because Mg2+ and O2- have higher charges.
Properties of ionic compounds:
- High melting/boiling points (strong electrostatic forces between ions)
- Hard but brittle (layers of ions repel when displaced)
- Conduct electricity when molten or dissolved (ions free to move)
- Do not conduct in solid state (ions fixed in lattice)
- Generally soluble in polar solvents (water), insoluble in non-polar solvents
Covalent bonding
Electron sharing. Polar if atoms have different electronegativity (HCl); non-polar if same (H2). Directional — gives specific bond angles. Example — H2O, CO2.
Polar vs non-polar covalent bonds: Electronegativity difference determines polarity:
- : pure covalent (H2, Cl2)
- : polar covalent (HCl, H2O)
- : ionic (NaCl, KF)
The boundary is approximate — no sharp line separates ionic from covalent. Fajan’s rules describe when ionic bonds develop covalent character (small cation + large anion = more covalent).
An ionic bond develops covalent character when:
- Cation is small with high charge (high polarising power — like Li+, Be2+, Al3+)
- Anion is large with high charge (high polarisability — like I-, S2-)
Example: LiI is more covalent than NaI (Li+ is smaller, polarises I- more). AlCl3 is covalent despite being a metal-nonmetal compound (Al3+ has very high charge density).
Coordinate (dative) bond
One atom donates both electrons in the shared pair. Shown with arrow. Once formed, indistinguishable from a normal covalent bond. Example — NH4+ (lone pair on N donated to H+), AlCl3 dimer.
How NH4+ forms:
NH3 has a lone pair on nitrogen. H+ has an empty orbital. The lone pair is donated to H+, forming a coordinate bond. Once formed, all four N-H bonds in NH4+ are identical — you cannot tell which one was the coordinate bond. This is because the electrons are shared equally after bond formation.
Other examples:
- H3O+ (water donates lone pair to H+)
- BF3·NH3 (NH3 donates lone pair to electron-deficient B)
- CO (triple bond includes one coordinate bond: C←O)
- Metal complexes like [Cu(NH3)4]2+ (NH3 donates to Cu2+)
Metallic bonding
Positive ions in a sea of delocalised electrons. Explains high conductivity, malleability, ductility and lustre. Example — Cu, Fe, Al.
The electron sea model: Metal atoms release their valence electrons into a shared pool. The resulting positive ion cores are held together by their attraction to this delocalised electron cloud. The electrons are not attached to any specific atom — they can flow throughout the metal.
This model explains:
- Conductivity: Delocalised electrons move freely under an applied voltage
- Malleability and ductility: Layers of ions can slide past each other without breaking the bond (the electron sea adjusts)
- Lustre: Free electrons absorb and re-emit light at all visible wavelengths
- High melting points: More valence electrons = stronger metallic bond (e.g., Fe with 2-3 delocalised electrons melts at 1538°C; Na with 1 electron melts at 98°C)
Van der Waals forces
Weak intermolecular forces. London dispersion (temporary dipoles in non-polar molecules), dipole-dipole (permanent dipoles), hydrogen bonds (strongest type). Important for molecular solids and liquids.
Types in order of increasing strength:
| Force | Origin | Strength | Example |
|---|---|---|---|
| London dispersion | Temporary induced dipoles | Weakest (0.05-40 kJ/mol) | Noble gases, CH4, I2 |
| Dipole-dipole | Permanent dipoles interact | Moderate (5-25 kJ/mol) | HCl, CH3Cl |
| Hydrogen bonding | H bonded to F, O, or N | Strongest intermolecular (10-40 kJ/mol) | H2O, NH3, HF, DNA |
Hydrogen bonding deserves special attention because it explains so many properties:
- Water’s unusually high boiling point (100°C vs -60°C expected for its molecular weight)
- Ice floating on water (H-bonds create an open crystal structure, less dense than liquid)
- DNA double helix stability (A-T with 2 H-bonds, G-C with 3 H-bonds)
- Protein secondary structure (alpha helices and beta sheets stabilised by H-bonds)
The boiling point anomaly question is a NEET classic: “Why does H2O have a higher boiling point than H2S?” Answer: H2O forms hydrogen bonds (O is small and highly electronegative); H2S cannot (S is too large and not electronegative enough for effective H-bonding).
Worked Examples
Metals have free electrons that can move through the lattice. Ionic solids have fixed ions — they cannot move until the solid melts or dissolves, freeing the ions.
NH4+ forms from NH3 (three N-H bonds) plus H+ (coordinate bond). Once formed, all four N-H bonds become indistinguishable because the electrons delocalise equally.
Diamond is a 3D network of covalent C-C bonds. Breaking these requires enormous energy (high MP). But all electrons are localised in sigma bonds — none are free to move and carry current. Compare with graphite, where one electron per carbon is delocalised (conducts along layers).
NaCl: → ionic. HCl: → polar covalent. Cl2: → pure covalent. This gradient from pure covalent to ionic is continuous, not a sharp boundary.
Each HF molecule can form only 2 hydrogen bonds (one H donor, one F acceptor with 3 lone pairs but F is too small for more than one H to approach). Each H2O molecule can form 4 hydrogen bonds (two H donors, two O lone pair acceptors). The network of H-bonds is denser in water, requiring more energy to break.
Master Comparison Table
| Property | Ionic | Covalent | Metallic | Van der Waals |
|---|---|---|---|---|
| Nature | Electrostatic | Shared electrons | Electron sea | Temporary/permanent dipoles |
| Direction | Non-directional | Directional | Non-directional | Non-directional |
| Strength | Strong (600-4000 kJ/mol) | Strong (150-1000 kJ/mol) | Variable (100-800 kJ/mol) | Weak (0.05-40 kJ/mol) |
| MP/BP | High | High (network) or low (molecular) | High (usually) | Low |
| Conductivity | When molten/dissolved | None (except graphite) | Always | None |
| Solubility | Water-soluble | Organic solvents | Insoluble | Organic solvents |
| Examples | NaCl, MgO, CaF2 | Diamond, H2O, CO2 | Cu, Fe, Na | Ice, I2, noble gases |
Common Mistakes
Saying coordinate bonds are weaker than covalent. They have the same strength once formed.
Confusing metallic bonding with ionic. Metals have delocalised electrons; ionic solids have localised ion charges.
Writing that dispersion forces exist only in non-polar molecules. They exist in all molecules but dominate in non-polar ones.
Saying ionic compounds do not conduct electricity. They do — when molten or in solution. It is only in the solid state that they do not conduct.
Treating all covalent compounds as having low melting points. Molecular covalent compounds (like CH4, CO2) have low MP, but network covalent solids (diamond, SiC, quartz) have extremely high MP because you must break covalent bonds throughout the network.
Exam Weightage and Revision
Chemical bonding is a major chapter in CBSE Class 11, carrying 7-10 marks. NEET asks 2-3 questions per year on bond types, VSEPR shapes, and intermolecular forces. JEE goes deeper into molecular orbital theory, hybridisation, and Fajan’s rules.
| Question Type | NEET Frequency | JEE Frequency |
|---|---|---|
| Property comparison across bond types | Every year | Most years |
| Hydrogen bonding examples | Most years | Every year |
| Coordinate bond identification | Every 2 years | Most years |
| Fajan’s rules application | Occasional | Every year |
| Boiling point comparison | Every year | Every year |
The most common NEET question: “Arrange the following in order of boiling point.” To answer, identify the dominant intermolecular force in each compound. H-bonding > dipole-dipole > London forces. For London forces, larger molecular mass = stronger.
Practice Questions
Q1. Why is AlCl3 covalent despite aluminium being a metal?
Al3+ has a very high charge density (charge/size ratio). According to Fajan’s rules, a small, highly charged cation strongly polarises the electron cloud of the anion, pulling it toward itself. This shared electron density makes the bond more covalent than ionic. AlCl3 exists as a covalent dimer (Al2Cl6) with bridging chlorine atoms.
Q2. Arrange in order of increasing boiling point: CH4, NH3, H2O, HF.
CH4 (-161°C) < HF (-19.5°C) < NH3 (-33°C) < H2O (100°C). CH4 has only London forces (lowest). HF, NH3, and H2O all have H-bonding, but H2O can form the most extensive network (4 H-bonds per molecule), followed by HF (limited networking despite strong individual bonds) and NH3 (only one lone pair to accept H-bonds). Wait — correcting: the actual order is CH4 < NH3 < HF < H2O. NH3 has a lower BP than HF because N is less electronegative than F, making N-H…N hydrogen bonds weaker than F-H…F bonds.
Q3. What is the difference between intermolecular and intramolecular forces?
Intramolecular forces are the bonds within a molecule (ionic, covalent, coordinate) — they hold atoms together. Intermolecular forces are the forces between separate molecules (London, dipole-dipole, H-bonding) — they hold molecules together in liquids and solids. Intramolecular forces are much stronger. Boiling a liquid breaks intermolecular forces but not intramolecular bonds — steam is still H2O.
FAQs
Why is the ionic model an oversimplification?
No bond is 100% ionic. Even in NaCl, the electron transfer is not complete — there is some sharing of electron density. The degree of ionic character depends on the electronegativity difference. Most real bonds are somewhere on a spectrum between pure covalent and purely ionic.
Why do noble gases have boiling points despite having no bonds?
Noble gases have London dispersion forces. Even in atoms with a symmetric electron cloud, random fluctuations create temporary dipoles. These temporary dipoles induce dipoles in neighbouring atoms, creating a weak attraction. Larger noble gases (Xe > Kr > Ar > Ne > He) have more electrons, stronger dispersion forces, and higher boiling points.
What holds DNA together — covalent bonds or hydrogen bonds?
Both. The backbone of each strand is covalent (phosphodiester bonds). The two strands of the double helix are held together by hydrogen bonds between complementary base pairs (A-T with 2 H-bonds, G-C with 3 H-bonds). The H-bonds are weak enough to be broken during replication and transcription, which is biologically essential.
Memorise one example for each bond type. Examples carry more marks than definitions.
Bonding types are a menu of structural choices. Each gives different properties, and materials science is about picking the right one for the job.