The Covalent Model

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Covalent Bonds

A covalent bond is the electrostatic attraction between a shared pair of electrons and the positively charged nuclei of two atoms.
Covalent bonding occurs between two non-metals; electrons are shared, not transferred.
When a covalent bond forms, two atomic orbitals overlap to form a molecular orbital.
In a typical covalent bond, each atom contributes one electron to the shared pair.
A single covalent bond is represented by a short straight line, e.g. H–H, Cl–Cl, H–Br.
Shared electrons are constantly moving and are best represented as charge clouds.
By sharing electrons, both atoms achieve a noble gas configuration, making them more stable.
The octet rule refers to the tendency of atoms to gain a valence shell with a total of 8 electrons; exceptions include electron-deficient (e.g. Be, B) and expanded octet species.

A single covalent bond

Lewis Formulas

Lewis formulas (electron dot diagrams) show all valence electrons in a covalently bonded species, including bonding pairs and lone pairs.
Electron pairs can be represented by dots, crosses, a combination of both, or a line.
Steps to draw a Lewis formula: 1) Count total valence electrons; 2) Draw skeletal positions with the central atom (usually least electronegative) in the middle; 3) Add bonding pairs then lone pairs to complete octets (duets for H); 4) Check all valence electrons are used and octets/duets are satisfied.
For example, methane (CH₄) has 8 valence electrons, four C–H bonding pairs, and no lone pairs on carbon.
Ammonia (NH₃) has 8 valence electrons, three N–H bonding pairs, and one lone pair on nitrogen.
Water (H₂O) has 8 valence electrons, two O–H bonding pairs, and two lone pairs on oxygen.
Carbon dioxide (CO₂) has 16 valence electrons, two C=O double bonds, and two lone pairs on each oxygen.
Some molecules have incomplete octets, e.g. BeCl₂ (Be has 4 electrons) and BF₃ (B has 6 electrons), but they are still stable.

Multiple Bonds

Non-metal atoms can share more than one pair of electrons to form single, double, or triple covalent bonds.
A single bond (C–C) has 2 shared electrons; a double bond (C=C) has 4; a triple bond (C≡C) has 6.
Quadruple covalent bonds are not possible because the repulsion between 8 electrons between the two nuclei is too great.
Bond energy is the amount of energy required to break one mole of a covalent bond in the gaseous state, with units kJ mol⁻¹.
The higher the bond energy, the stronger the bond.
Bond length is the distance between the nuclei of two covalently bonded atoms.
As more electrons are shared, electron density between atoms increases, pulling atoms closer: bond length decreases and bond strength increases.
Triple bonds are the shortest and strongest covalent bonds; single bonds are the longest and weakest.

A double covalent bond

Coordinate Bonds

In a coordinate bond (dative covalent bond), both electrons in the shared pair come from the same atom.
This typically occurs when one atom has a lone pair and the other is electron-deficient (has an unfilled outer shell).
Example: in the ammonium ion (NH₄⁺), the nitrogen atom in ammonia donates its lone pair to an H⁺ ion.
The H⁺ ion is electron-deficient and accepts the pair of electrons, forming a coordinate bond.
Coordinate bonds are also found in transition metal complexes, e.g. [Al(H₂O)₆]³⁺, where water molecules donate lone pairs to the metal ion.
SL students should recognise that coordinate bonding can involve species beyond simple molecules like NH₄⁺.

Shapes of Molecules

VSEPR theory: electron pairs around a central atom repel each other and arrange as far apart as possible to minimise repulsion.
Rules: all electron pairs spread out; lone pairs repel more strongly than bonding pairs; multiple bonds behave like a single bond.
Each region of electron density is an electron domain; it may contain one, two, or three pairs of electrons.
Two domains → linear geometry, bond angle 180° (e.g. BeCl₂, CO₂, HC≡CH).
Three domains → trigonal planar geometry, bond angle 120° (e.g. BF₃, CH₂=CH₂, CH₂O); with one lone pair, shape is bent with angle ~118° (e.g. SO₂).
Four domains → tetrahedral geometry, bond angle 109.5° (e.g. CH₄, NH₄⁺); with one lone pair, shape is trigonal pyramidal with angle ~107° (e.g. NH₃); with two lone pairs, shape is bent with angle ~104.5° (e.g. H₂O).
Order of electron pair repulsion: lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair.
Always draw the Lewis structure first to identify lone pairs before determining shape and bond angles.

Bond Polarity

Electronegativity is the ability of an atom to attract the shared pair of electrons in a covalent bond.
Electronegativity values are on the Pauling scale and are listed in Section 9 of the IB Chemistry Data Booklet.
In diatomic molecules like H₂, O₂, Cl₂, the atoms have equal electronegativity, so electrons are shared equally → nonpolar covalent bond.
When atoms have different electronegativities, bonding electrons are drawn closer to the more electronegative atom, creating a polar covalent bond.
The less electronegative atom gains a partial positive charge (δ⁺); the more electronegative atom gains a partial negative charge (δ⁻).
A dipole forms when there is a separation of charge across a covalent bond; the dipole moment measures how polar the bond is.
The greater the difference in electronegativity, the greater the dipole moment and the more polar the bond.
Dipole direction is shown by an arrow with a crossed tail near δ⁺ and arrowhead pointing toward δ⁻.

Molecular Polarity

To determine if a molecule with more than two atoms is polar, consider both bond polarity and molecular geometry.
Molecules with only nonpolar bonds cannot be polar because there is no bond dipole.
Molecules with polar bonds can be nonpolar overall if the bond dipoles are arranged symmetrically and cancel out (e.g. CO₂, CCl₄).
Molecules with polar bonds are polar overall if the bond dipoles are arranged asymmetrically and do not cancel (e.g. H₂O, NH₃, CH₃Cl).
CO₂ has two C=O dipoles pulling in opposite directions, so they cancel → nonpolar.
CCl₄ has four polar C–Cl dipoles arranged symmetrically in a tetrahedral shape → nonpolar.
H₂O has a bent shape causing dipoles to reinforce → polar; NH₃ has a trigonal pyramidal shape creating a net dipole → polar.
CH₃Cl has four polar bonds but asymmetrical arrangement, so dipoles do not cancel → polar.

Giant Covalent Structures

Giant covalent structures (covalent network lattices) are large networks of atoms joined by strong covalent bonds extending throughout the lattice.
Examples include silicon, silicon dioxide, diamond, graphite, buckminsterfullerene (C₆₀), and graphene.
In diamond, each carbon is covalently bonded to four others in a tetrahedral arrangement (109.5°); it is extremely hard and does not conduct electricity.
In graphite, each carbon is bonded to three others in hexagonal rings (120°); layers are held by weak intermolecular forces, and delocalised electrons make it a good conductor.
Graphene is a single layer of graphite; each carbon bonded to three others (120°), with delocalised electrons making it a very good conductor; it is strong, thin, and flexible.
Buckminsterfullerene (C₆₀) is a molecular form of carbon with 60 atoms in a spherical structure of hexagons and pentagons; each carbon bonds to three others; it is a semiconductor.
Silicon dioxide has each silicon bonded to four oxygen atoms and each oxygen bonded to two silicon atoms in a tetrahedral network; empirical formula SiO₂.
Giant covalent structures have very high melting and boiling points, are generally hard (except graphite), and are usually insoluble in water.

Intermolecular Forces

Intermolecular forces act between molecules and are much weaker than covalent or ionic bonds; they determine melting/boiling point, volatility, and solubility.
London (dispersion) forces arise from temporary dipoles and induced dipoles; they are present between all atoms and molecules and are the only intermolecular forces in nonpolar substances.
Strength of London forces increases with more electrons and larger surface area; they range from 1–50 kJ mol⁻¹.
Dipole–dipole attractions occur between polar molecules with permanent dipoles, between the δ⁺ end of one and the δ⁻ end of another.
Dipole–induced dipole attractions occur when a polar molecule distorts the electron cloud of a nonpolar molecule, creating a temporary dipole.
Hydrogen bonding is the strongest intermolecular force; it requires a H atom covalently bonded to O, N, or F, and a lone pair on an O, N, or F atom in a neighbouring molecule.
Each water molecule can form two hydrogen bonds (two lone pairs, two H atoms); each ammonia molecule can form one hydrogen bond (one lone pair, one N–H bond).
Van der Waals forces is a collective term for London (dispersion) forces, dipole–dipole attractions, and dipole–induced dipole attractions.

Physical Properties of Covalent Substances

When molecular covalent substances change state, intermolecular forces are overcome, not covalent bonds.
Stronger intermolecular forces → higher melting and boiling points; substances with low boiling points are volatile.
Strength of intermolecular forces depends on molecular size (molar mass), polarity, and type: London < dipole–dipole < hydrogen bonding.
Solubility rule: 'like dissolves like' — nonpolar substances dissolve in nonpolar solvents; polar substances dissolve in polar solvents.
Small alcohols (e.g. ethanol) are highly soluble in water due to hydrogen bonding; solubility decreases as the nonpolar hydrocarbon chain lengthens (e.g. hexanol is much less soluble).
Most covalent substances do not conduct electricity because they lack free-moving charged particles.
Some polar covalent substances conduct in solution because they ionise to produce ions (e.g. HCl forms H⁺ and Cl⁻ in water).
Giant covalent structures are generally insoluble and non-conducting, except graphite and graphene which have delocalised electrons.

Chromatography

Chromatography separates components of a mixture by passing a liquid mobile phase over a solid stationary phase.
Separation depends on differences in solubility in the mobile phase and attraction to the stationary phase.
In paper chromatography, the stationary phase is paper (cellulose with polar –OH groups) and the mobile phase is a solvent (e.g. water or ethanol).
Polar substances interact strongly with the paper, travel slower, and do not travel as far; nonpolar substances interact less, are more soluble in the solvent, and travel further.
Thin-layer chromatography (TLC) uses a thin layer of an inert solid (silica or alumina) on a flat surface as the stationary phase; the mobile phase is a solvent.
The retardation factor (Rf) is calculated as: Rf = distance travelled by component / distance travelled by solvent front.
Rf values are decimals, have no units, and are always less than 1.
If a component does not travel with the solvent, it is insoluble in that solvent and a different solvent should be used.

Paper chromatography

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Practice questions

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1.Which statement best defines a covalent bond?
Easy
- AThe electrostatic attraction between a shared pair of electrons and the positively charged nuclei of two atoms
- BThe transfer of electrons from a metal atom to a non-metal atom
- CThe electrostatic attraction between oppositely charged ions in a lattice
- DThe attraction between a metal cation and a sea of delocalised electrons
2.In a typical covalent bond, each atom contributes one electron to the shared pair.
Easy
True or false?
3.Which of the following are exceptions to the octet rule? (select all that apply)
Medium
- ABeryllium chloride, BeCl₂, where beryllium has only 4 electrons around it
- BBoron trifluoride, BF₃, where boron has only 6 electrons around it
- CSulfur dioxide, SO₂, where sulfur has 10 electrons in its valence shell
- DMethane, CH₄, where carbon has 8 electrons around it
- EWater, H₂O, where oxygen has 8 electrons around it
4.Which of the following molecules has a linear shape?
Medium
- ABeCl₂
- BBF₃
- CCH₄
- DH₂O
5.According to VSEPR theory, which statement about electron pair repulsion is correct?
Medium
- ALone pairs repel more strongly than bonding pairs
- BBonding pairs repel more strongly than lone pairs
- CLone pairs and bonding pairs repel equally
- DMultiple bonds repel more strongly than single bonds
6.In a molecule with polar bonds arranged symmetrically, the bond dipoles cancel and the molecule is nonpolar overall.
Easy
True or false?
7.Match each molecule to its correct molecular geometry.
Medium
- CH₄
- NH₃
- H₂O
- BF₃
- tetrahedral
- trigonal pyramidal
- bent
- trigonal planar
8.Place the following carbon–carbon bonds in order of increasing bond length.
Medium
- C≡C (triple bond)
- C=C (double bond)
- C–C (single bond)

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