How do you tell whether a bond is ionic or covalent?

Compare the electronegativities. A metal bonded to a non-metal with a large difference (about 1.7 or more, a little less for very reactive metals) transfers electrons and forms an ionic bond. Two non-metals share electrons and form a covalent bond, which is polar if the difference is about 0.4 or more.

What does VSEPR theory say?

Electron pairs around a central atom repel one another and arrange themselves as far apart as possible. The number of bonding pairs and lone pairs sets the shape: linear, trigonal planar, tetrahedral, bent, pyramidal and so on.

Why can CO2 be non-polar when its C=O bonds are polar?

CO2 is linear, so the two equal bond dipoles point in opposite directions and cancel. The molecule has no net dipole.

Why is O2 paramagnetic?

Molecular orbital theory puts the last two electrons of O2 singly into the two degenerate pi-star orbitals, so O2 has two unpaired electrons and is attracted to a magnetic field.

How do you find the bond order of a diatomic molecule?

Bond order is half of the difference between the number of electrons in bonding orbitals and in antibonding orbitals.

Home / Chemical bonding lab

Chemical bonding lab

Q: What is electronegativity?

Electronegativity is how strongly an atom in a bond pulls the shared electrons towards itself. The Pauling scale runs from about 0.7 (caesium) to 4.0 (fluorine).

Atoms bond to reach a more stable arrangement, usually a full outer shell. A metal bonded to a non-metal usually transfers electrons (an ionic bond); two non-metals share them (a covalent bond, polar if one atom pulls harder); metals pool theirs in a sea of electrons (a metallic bond). Pick any two elements and watch which one forms, then explore shapes, molecular orbitals and organic molecules. Pitched at class 9 to JEE / NEET.

First atom

Second atom

Or pick from the periodic table

Click two elements: the first goes in the highlighted box, then the second. Hover for the name and electronegativity.

How bonds form: the ideas, with the main points

Class 9–11 — Why atoms bond, and the three types

Main points

Atoms bond because the bonded arrangement has lower energy than the separate atoms.
Kössel–Lewis idea: atoms tend to reach the octet (8 outer electrons; 2 for H, Li, Be: the duet) by gaining, losing or sharing electrons.
Ionic = transfer (metal + non-metal). Covalent = sharing (non-metal + non-metal). Metallic = delocalised electrons shared by a lattice of metal ions.
Electronegativity (Pauling scale, 0.7 Cs to 4.0 F) decides how unequal the sharing is: it increases across a period and decreases down a group.
Rule of thumb: ΔEN < 0.4 non-polar covalent · 0.4 to about 1.7 polar covalent · above about 1.7 ionic. It is a guide: many borderline compounds mix both characters.

Class 9–12 — The ionic bond, lattice enthalpy and Fajans' rules

Main points

Formation: Na(g) → Na⁺(g) + e⁻ costs the ionisation enthalpy; Cl(g) + e⁻ → Cl⁻(g) releases the electron gain enthalpy; the real driving force is the lattice enthalpy released when the gaseous ions pack into a crystal.
Favourable when the cation has a low ionisation enthalpy, the anion a high negative electron gain enthalpy, and the lattice enthalpy is large.
Lattice enthalpy increases with higher ionic charges and smaller ionic radii (∝ q₁q₂/r): MgO > NaF > NaCl > NaI.
The Born–Haber cycle applies Hess's law (sublimation, ionisation, dissociation, electron gain, lattice enthalpy) to find lattice enthalpy.
Properties: hard, brittle, high melting and boiling points, conduct only molten or in solution, mostly soluble in water (ion–dipole attraction beats the lattice).
Fajans' rules: covalent character increases with a small, highly charged cation, a large, highly charged anion, and a cation without a noble-gas configuration (Cu⁺ is more polarising than Na⁺). So LiCl > NaCl > KCl, AlCl₃ > MgCl₂ > NaCl, and AlI₃ is the most covalent aluminium halide.
Isoelectronic ions shrink as nuclear charge rises: O²⁻ > F⁻ > Na⁺ > Mg²⁺ > Al³⁺. A cation is smaller than its atom, an anion larger.

Try Na + Cl, Mg + O, Al + O (criss-cross to Al₂O₃) and Fe + O (iron(III) oxide).

Class 9–12 — The covalent bond, Lewis structures, formal charge, resonance

Main points

An atom forms as many covalent bonds as it needs to complete its octet: H 1 · O 2 · N 3 · C 4. Each shared pair is one bond: single (σ), double (σ + π), triple (σ + 2π).
Drawing a Lewis structure: count the total valence electrons (adjust for charge), join atoms with single bonds, complete octets of outer atoms, put leftover pairs on the central atom, and form multiple bonds if the central atom is short of an octet.
Formal charge = V − N − B/2 (valence electrons − non-bonding electrons − half the bonding electrons). The best structure has formal charges close to zero, with any negative charge on the more electronegative atom.
Resonance: when one Lewis structure is not enough, the real molecule is a hybrid, more stable than any single structure. O₃ bond order 1.5; CO₃²⁻ and NO₃⁻ bond order 1.33; benzene C–C bond order 1.5.
Exceptions to the octet: incomplete (BeCl₂, BF₃, AlCl₃: Lewis acids), expanded (PCl₅, SF₆, using d orbitals from period 3), and odd-electron (NO, NO₂).
Coordinate (dative) bond: both electrons come from one atom: NH₄⁺, H₃O⁺, CO, F₃B←NH₃, Al₂Cl₆. Once formed it is identical to an ordinary covalent bond.

Try H + O (H₂O), N + H (NH₃), C + H (CH₄), C + O (CO₂), N + N (triple bond), then Be + Cl and Al + Cl for the incomplete octet.

Class 11–12 — Polarity, dipole moment and bond parameters

Main points

A polar bond carries partial charges δ⁺ / δ⁻. Dipole moment μ = q × d (a vector pointing from the δ⁺ end towards the δ⁻ end; unit debye, 1 D = 3.34 × 10⁻³⁰ C·m).
A molecule is polar only if the bond dipoles do not cancel. Symmetric shapes cancel: CO₂, BF₃, CH₄, CCl₄, SF₆, XeF₄, BeCl₂ have μ = 0. Bent or pyramidal shapes do not: H₂O 1.85 D, NH₃ 1.47 D.
NH₃ (1.47 D) vs NF₃ (0.23 D): in NH₃ the N–H bond dipoles point the same way as the lone-pair dipole; in NF₃ they oppose it.
Isomers: trans-1,2-dichloroethene and p-dichlorobenzene have μ = 0; the cis, ortho and meta forms are polar.
Dipole moment of the hydrogen halides falls HF > HCl > HBr > HI as ΔEN falls.
Percentage ionic character ≈ 16ΔEN + 3.5ΔEN² (Hannay–Smyth): HCl about 19 %, NaCl about 53 %.
Bond parameters: bond length (C–C 154, C=C 134, C≡C 120 pm), bond enthalpy (346, 614, 839 kJ/mol), bond angle, bond order. Higher bond order → shorter, stronger bond.
Odd one out: F₂ (159 kJ/mol) is weaker than Cl₂ (243) because of lone pair–lone pair repulsion on the small F atoms.

Class 11–12 — VSEPR shapes and hybridisation

Main points

Count the electron domains on the central atom (bonded atoms + lone pairs; a double or triple bond counts as one). They spread as far apart as possible.
2 linear (sp) · 3 trigonal planar (sp²) · 4 tetrahedral (sp³) · 5 trigonal bipyramidal (sp³d) · 6 octahedral (sp³d²).
Repulsion order: lone pair–lone pair > lone pair–bond pair > bond pair–bond pair. So CH₄ 109.5° > NH₃ 107° > H₂O 104.5°.
With lone pairs: NH₃ pyramidal · H₂O bent · SO₂ bent · SF₄ see-saw · ClF₃ T-shaped · XeF₂ linear · BrF₅ square pyramidal · XeF₄ square planar. On a trigonal bipyramid the lone pairs go equatorial.
In PCl₅ the axial bonds (two 90° neighbours more) are longer than the equatorial bonds.
Hybridisation mixes atomic orbitals into equivalent hybrid orbitals: sp (50 % s, 180°), sp² (33 % s, 120°), sp³ (25 % s, 109.5°). More s-character → shorter bond, higher electronegativity of the carbon, more acidic C–H (ethyne > ethene > ethane).
σ bond: head-on overlap (strongest, free rotation); π bond: sideways overlap of p orbitals (weaker, stops rotation). Count: σ = number of bonds, π = extra bonds in multiple bonds.
Examples to remember: NH₄⁺ tetrahedral, H₃O⁺ pyramidal, NO₃⁻ trigonal planar, SO₄²⁻ tetrahedral, I₃⁻ linear (sp³d), ClO₃⁻ pyramidal.

Use the Molecule shapes tab: set bonding and lone pairs, or click BeCl₂ … XeF₄, to see all thirteen common shapes.

Class 11–12 — Molecular orbital theory (MOT)

Main points

Atomic orbitals combine (LCAO) into bonding molecular orbitals (lower energy, electron density between the nuclei) and antibonding orbitals (higher energy, marked *, a node between the nuclei).
Electrons fill the lowest orbitals first (Aufbau), two per orbital with opposite spins (Pauli), and singly into equal-energy orbitals before pairing (Hund).
Bond order = ½ (N_b − N_a). Bond order 0: no molecule (He₂, Be₂, Ne₂). Higher bond order: shorter, stronger bond.
Energy order: for B₂, C₂, N₂ and lighter: σ1s < σ*1s < σ2s < σ*2s < π2p < σ2pz < π*2p < σ*2pz. For O₂, F₂, Ne₂ σ2pz drops below π2p.
Magnetism: unpaired electrons → paramagnetic (B₂, O₂, O₂⁺, O₂⁻, NO); all paired → diamagnetic (N₂, C₂, F₂, O₂²⁻).
Standard results: N₂ 3 · O₂ 2 · F₂ 1 · O₂⁺ 2.5 · O₂⁻ 1.5 · O₂²⁻ 1 · N₂⁺ 2.5 · NO 2.5 · H₂⁺ 0.5. Bond length order: O₂⁺ < O₂ < O₂⁻ < O₂²⁻.
Lewis theory cannot explain why O₂ is paramagnetic; MOT does (two unpaired π* electrons).

Open the Orbitals (MOT) tab and compare O₂, O₂⁺, O₂⁻ and O₂²⁻ to see the bond order fall as electrons enter π*.

Class 11–12 — Hydrogen bonding and intermolecular forces

Main points

Strength order (typical): ion–dipole > hydrogen bond (10–40 kJ/mol) > dipole–dipole > London dispersion. All are far weaker than a covalent bond (hundreds of kJ/mol).
A hydrogen bond X–H···Y needs H bonded to a small, very electronegative atom (N, O, F) and a lone pair on another N, O or F.
Anomalous boiling points: H₂O, HF and NH₃ boil far above their group-mates (H₂S, HCl, PH₃) because of hydrogen bonding.
Ice floats: hydrogen bonds hold ice in an open hexagonal network, so it is less dense than liquid water. Water has maximum density at 4 °C.
Intramolecular H-bond (o-nitrophenol, salicylaldehyde) makes the compound steam-volatile and less soluble in water; intermolecular (p-nitrophenol) raises boiling point and viscosity.
London forces grow with the number of electrons (polarisability): boiling points rise down group 18 and for halogens (F₂ gas, Br₂ liquid, I₂ solid).
Carboxylic acids exist as hydrogen-bonded dimers in non-polar solvents.

Class 9–12 — The metallic bond

Main points

Metal atoms release their valence electrons into a shared sea of delocalised electrons around a lattice of positive ions.
Because the bond is non-directional, layers slide without breaking: malleable and ductile. Free electrons give electrical and thermal conductivity and lustre.
Strength rises with more valence electrons and smaller atoms: Na (98 °C) < Mg (650 °C) < Al (660 °C); tungsten melts at about 3400 °C.
An alloy is a mixture of a metal with other elements; it is often harder and more resistant to corrosion (brass, steel, bronze).

Class 11–12 — Bonding in organic molecules

Main points

Carbon makes four bonds: sp³ (alkanes, tetrahedral 109.5°), sp² (alkenes, carbonyls, benzene, trigonal planar 120°), sp (alkynes, nitriles, linear 180°).
C–C 154 pm / 346 kJ · C=C 134 pm / 614 kJ · C≡C 120 pm / 839 kJ; benzene C–C 139 pm (bond order 1.5, six π electrons delocalised).
Counting bonds: σ = number of bonds; π = bonds beyond the first in each multiple bond. Ethene 5σ + 1π, ethyne 3σ + 2π, benzene 12σ + 3π, acetaldehyde 6σ + 1π.
Allene CH₂=C=CH₂: sp², sp, sp²; the two CH₂ planes are at 90°.
The π bond prevents rotation about C=C, which is why cis–trans isomerism exists.
C–H bond lengths: 109 pm (sp³), 108 pm (sp²), 106 pm (sp): shorter with more s-character.

Open the Organic molecules tab, pick propene or allene and click each carbon to read its hybridisation.

The three bond types at a glance

	Ionic	Covalent	Metallic
Formed between	metal + non-metal	non-metal + non-metal	metal atoms
Electrons	transferred	shared in pairs	delocalised sea
Attraction	between oppositely charged ions (non-directional)	between nuclei and shared pair (directional)	between ions and electron sea (non-directional)
Melting point	high	low (molecular) / very high (network: diamond, SiO₂)	moderate to very high
Electrical conduction	molten or dissolved only	poor (graphite is an exception)	good in solid and liquid
Hardness	hard, brittle	soft (molecular) / very hard (network)	malleable, ductile
Example	NaCl, MgO	H₂O, CH₄, diamond	Cu, Fe, brass

Quick formulas

Type of bondΔEN ≥ 1.7 (metal + non-metal): ionic
0.4 ≤ ΔEN < 1.7: polar covalent
ΔEN < 0.4: non-polar covalent

Ionic formula (criss-cross)A^m+ + Bⁿ⁻ → A_nB_m (divide by the common factor)

Bonds an atom forms8 − (outer electrons), or 1 for hydrogen. Shared pairs in total = lcm of the two atoms' bond counts.

Ionic character% = 16ΔEN + 3.5ΔEN²

Formal chargeFC = V − N − B/2

Dipole momentμ = q × d (1 D = 3.34 × 10⁻³⁰ C·m)

Lone pairs on the central atom(valence electrons − electrons used in bonds) ÷ 2

Electron domainsbonding pairs + lone pairs: 2 sp, 3 sp², 4 sp³, 5 sp³d, 6 sp³d²

Bond order (MOT)½ (N_b − N_a)

σ and π countsσ = number of bonds; π = Σ(bond order − 1)

Questions people ask

How do I know if a bond is ionic or covalent?

Look at the two elements. Metal + non-metal with a large electronegativity difference (about 1.7 or more) is ionic: electrons are transferred. Two non-metals share electrons, which is covalent, and it is polar if the difference is about 0.4 or more. Small, highly charged cations (Be²⁺, Al³⁺) give compounds with a lot of covalent character (Fajans' rules).

What is electronegativity?

How strongly an atom pulls the shared electrons in a bond towards itself. Fluorine is highest (4.0) and caesium is lowest (about 0.7).

Why does CO₂ have polar bonds but is not a polar molecule?

It is linear, so the two C=O dipoles point opposite ways and cancel. Water is bent, so its dipoles add up and the molecule is polar.

Why is O₂ paramagnetic although its Lewis structure shows no unpaired electrons?

The Lewis structure is too simple. In molecular orbital theory the last two electrons of O₂ occupy the two degenerate π* orbitals singly, so there are two unpaired electrons. Open the Orbitals tab to see it.

Are the shapes here exact?

The shapes and angles come from VSEPR theory, which is an excellent model for classes 11 and 12. Real angles differ a little (for example H₂S is 92°). The lab uses the simple octet rule for the Two atoms tab, so molecules with an expanded octet (SO₂, SO₃, SF₆) and molecules with several atoms of both elements (like N₂O₃) are best explored in the Molecule shapes tab or the lessons.