The covalent model — IB Chemistry HL Revision Notes & Topical Questions

Section 1 of 75Covalent bonds and the octet rule

Covalent bond

Covalent bond the electrostatic attraction between a shared pair of electrons and the positively charged nuclei of two atoms.

Between non-metals

• Between non-metals no electrons are transferred; two atomic orbitals overlap to form a molecular orbital.

Normal bond

• Normal bond each atom contributes one electron to the shared pair, and both reach a noble gas configuration.

Stability

• Stability electrons are more stable shared between two nuclei than held by one atom.

Octet rule

• Octet rule atoms tend to gain a valence shell of 8 electrons.

Electron deficient

• Be in BeCl₂ has 4 electrons, and B in BF₃ and Al have 6.

Expanded octet

• Expanded octet some period 3 central atoms hold more than 8 electrons.

Section 2 of 75Lewis formulas

Lewis formula

• Lewis formula shows all valence electrons, as bonding pairs and lone pairs, using dots, crosses or lines.

Not a Lewis formula

• Cl–Cl, because it does not show the lone pairs.

1

Count the total valence electrons of all atoms, adjusting for any charge.

2

Draw the skeleton with the least electronegative atom in the centre.

3

Add a bonding pair between each pair of bonded atoms.

4

Add lone pairs to complete octets (duets for hydrogen), starting with the outer atoms.

5

Check that all electrons are used and every atom has a full shell.

Section 3 of 75Multiple and coordinate bonds

Single, double, triple

• Single, double, triple they share 2, 4 and 6 electrons.

No quadruple bonds

• No quadruple bonds repulsion between 8 electrons between the nuclei is too great.

Bond energy

• Bond energy the energy to break one mole of a covalent bond in the gas phase, in kJ mol⁻¹;
  • a higher value means a stronger bond.

Bond length

• Bond length the distance between the nuclei; more shared electrons pull the nuclei closer.

Trend

• Trend triple bonds are shortest and strongest, single bonds longest and weakest.

Coordinate bond

Coordinate bond both electrons of the shared pair come from the same atom.

• It forms when one atom has a lone pair and the other is electron deficient.
• NH₃ donates its lone pair to H⁺ to form NH₄⁺.
• At HL, [Al(H₂O)₆]³⁺ has six coordinate bonds from water oxygen atoms.

Section 4 of 75Shapes of molecules (VSEPR)

1Principle

• Principle electron domains repel and spread as far apart as possible.
  • (bonding pairs, lone pairs, or a multiple bond counted once)

2Repulsion order

• Repulsion order lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair.

3Lone pairs

• Lone pairs they are closer to the nucleus, so they repel more and squeeze bond angles.

4Lewis first

• Lewis first draw the Lewis formula first so hidden lone pairs are not missed.

5Domain geometry versus molecular shape

• Domain geometry versus molecular shape they are the same only when there are no lone pairs.

Bonding	Lone	Molecular shape	Angle	Example
2	0	Linear	180°	BeCl2, CO2
3	0	Trigonal planar	120°	BF3
2	1	Bent	about 118°	SO2
4	0	Tetrahedral	109.5°	CH4
3	1	Trigonal pyramidal	107°	NH3
2	2	Bent	104.5°	H2O

Section 5 of 75Bond and molecular polarity

Electronegativity

• Electronegativity the ability of an atom to attract the shared pair;
  • values are on the Pauling scale in the data booklet.

Nonpolar bond

• Nonpolar bond identical atoms share electrons equally, as in H₂ and Cl₂.

Polar bond

• Polar bond the more electronegative atom is δ− and the other δ+;
  • a larger electronegativity difference gives a more polar bond.

Dipole arrow

• Dipole arrow the crossed tail is at the δ+ atom and the head points to the δ− atom.

Nonpolar molecule

• Nonpolar molecule only nonpolar bonds, or polar bonds in a symmetrical arrangement so dipoles cancel, as in CO₂ and CCl₄.

Polar molecule

• Polar molecule polar bonds arranged asymmetrically so dipoles do not cancel, as in H₂O, NH₃ and CH₃Cl.

Section 6 of 75Giant covalent structures

Giant covalent structures

Giant covalent structure

• Giant covalent structure strong covalent bonds extend throughout the lattice, so there are no discrete molecules.

Silicon and SiO₂

• Silicon and SiO₂ tetrahedral networks; SiO₂ has each Si bonded to four O and each O to two Si.

Diamond

• Diamond each C bonded to four others, very hard, and non-conducting because all outer electrons are bonded.

Graphite

• Graphite each C bonded to three others in flat hexagonal layers;
  • delocalised electrons conduct and weak forces between layers make it soft.

Graphene

• Graphene a single layer of graphite, one atom thick, strong, flexible and conducting.

C60

• C60 a molecule, not a giant structure; a semiconductor.

Properties

• Properties very high melting points, mostly insoluble, and usually non-conducting except graphite and graphene.

Section 7 of 75Intermolecular forces

Intermolecular forces

Intermolecular versus covalent

• Intermolecular versus covalent forces between molecules are much weaker than covalent bonds within molecules.

London (dispersion) forces

• London (dispersion) forces present in all substances;
  • strength rises with the number of electrons and with surface area.

Dipole–dipole

• Dipole–dipole between permanent dipoles, in addition to London forces.

Dipole–induced dipole

• Dipole–induced dipole a polar molecule distorts the electron cloud of a nonpolar one.

Hydrogen bonding

• Hydrogen bonding the strongest;
  • needs H bonded to O, N or F and a lone pair on O, N or F of a neighbouring molecule.

Order

• London < dipole–dipole < hydrogen bonding.

Van der Waals

• Van der Waals collective name for London, dipole–dipole and dipole–induced dipole forces.

Section 8 of 75Physical properties and chromatography

Boiling point

• Boiling point stronger intermolecular forces give higher melting and boiling points;
  • weaker forces give volatile substances.

Solubility

• Solubility like dissolves like: polar in polar solvents, nonpolar in nonpolar solvents;
  • large hydrocarbon parts reduce solubility in water.

Conductivity

• Conductivity molecular substances do not conduct, though some polar ones ionise in water, such as HCl.

Chromatography

• Chromatography separates by different solubility in the mobile phase and attraction to the stationary phase;
  • paper and thin-layer chromatography are required.

Rf

• Rf distance moved by the component divided by the distance moved by the solvent front, with no units and always less than 1.

Section 9 of 75HL: resonance, expanded octets and formal charge

Resonance

• Resonance more than one valid Lewis structure exists, and the real structure is a hybrid with delocalised electrons.

Examples

• NO₃⁻, CO₃²⁻, O₃, RCOO⁻ and benzene all have equal bond lengths between single and double.

Benzene evidence

• Benzene evidence hydrogenation enthalpy of −208 instead of −360 kJ mol⁻¹, all C–C bonds 140 pm, no bromine water decolourisation, no C=C infrared peak.

Expanded octet

• Expanded octet period 3 and below can hold 10 or 12 electrons, as in PCl₅, SF₆, SF₄ and XeF₄.

Formal charge

• FC = V − ½B − N;
  • the preferred structure has formal charges closest to zero, with any negative on the most electronegative atom.

Section 10 of 75HL: sigma bonds, pi bonds and hybridisation

Sigma bond

• Sigma bond head-on overlap, present in all single bonds and in every multiple bond.

Pi bond

• Pi bond sideways overlap of p orbitals, present only in double and triple bonds.

Double bond

• Double bond one σ and one π; triple bond — one σ and two π.

sp3

• sp3 four domains, tetrahedral, 109.5°.

sp2

• sp2 three domains, trigonal planar, 120°, with one π bond.

sp

• sp two domains, linear, 180°, with two π bonds.

Hybridisation

• Hybridisation mixing of s and p orbitals to give equal hybrid orbitals, explaining equal bonds in methane.

Section 11 of 75Covalent bonds

Definition

• Definition the electrostatic attraction between a shared pair of electrons and two positively charged nuclei.

Formation

• Formation two atomic orbitals overlap to form a molecular orbital between non-metals.

Contribution

• Contribution each atom normally gives one electron and both reach a noble gas configuration.

Charge cloud

• Charge cloud shared electrons are best pictured as a charge cloud.

Electron deficient

• Be (4 electrons in BeCl₂), B and Al (6 electrons in BF₃) do not complete an octet.

Expansion

• Expansion some period 3 central atoms hold more than 8 electrons.

Section 12 of 75Drawing Lewis formulas

Lewis formula

• Lewis formula shows all valence electrons as bonding pairs and lone pairs.

Cl–Cl

• Cl–Cl not a Lewis formula because it omits the lone pairs.

1

Count the total valence electrons of all atoms, adjusting for charge.

2

Draw the skeleton with the least electronegative atom in the centre.

3

Add a bonding pair between each pair of bonded atoms.

4

Add lone pairs to complete octets (duets for H), starting with the outer atoms.

5

Check that all electrons are used and each atom has a full octet.

1Pair counts

• CH₄ has 4 bonding pairs, NH₃ has 3 bonding pairs and 1 lone pair, and H₂O has 2 bonding and 2 lone pairs.

2Multiple bonds in formulas

• CO₂ has two C=O bonds and HCN has a C–H single bond and a C≡N triple bond.

Section 13 of 75Multiple bonds and bond properties

Bond	Shared electrons	Length	Strength
Single	2	Longest	Weakest
Double	4	Intermediate	Intermediate
Triple	6	Shortest	Strongest

No quadruple bond

• No quadruple bond repulsion of 8 electrons between the nuclei is too great.

Bond energy

• Bond energy the energy needed to break one mole of a covalent bond in the gas phase, in kJ mol⁻¹.

Bond length

• Bond length the distance between the two nuclei.

Reason

• Reason more shared electrons increase electron density between the nuclei, so attraction rises and atoms are pulled closer.

Section 14 of 75Coordinate bonds

Coordinate (dative) bond

Coordinate (dative) bond a covalent bond in which both electrons come from the same atom.

Conditions

• Conditions one atom has a lone pair and the other is electron deficient.

Ammonium ion

• Ammonium ion the lone pair on N in NH₃ is donated to an H⁺ ion with an empty orbital, so NH₃ + H⁺ → NH₄⁺.

HL complex ions

• [Al(H₂O)₆]³⁺ has six coordinate bonds, each from an oxygen lone pair of water.

SL scope

• SL scope coordinate bonding is not limited to NH₄⁺.

Section 15 of 75VSEPR theory

Electron clouds

• Electron clouds bonding and non-bonding pairs behave like negative charge clouds that repel.

Rule 1

• Rule 1 all electron pairs spread as far apart as possible.

Rule 2

• Rule 2 lone pairs repel more strongly than bonding pairs, because they are held closer to the nucleus.

Rule 3

• Rule 3 a multiple bond counts as one domain.

Repulsion order

• Repulsion order lone pair–lone pair > lone pair–bonding pair > bonding pair–bonding pair.

Electron domain

• Electron domain each region of electron density around the central atom.

Section 16 of 75Shapes with two, three and four domains

Bonding	Lone	Domain geometry	Molecular shape	Angle	Examples
2	0	Linear	Linear	180°	BeCl2, CO2, HC≡CH
3	0	Trigonal planar	Trigonal planar	120°	BF3, CH2=CH2, CH2O
2	1	Trigonal planar	Bent	about 118°	SO2
4	0	Tetrahedral	Tetrahedral	109.5°	CH4, NH4+
3	1	Tetrahedral	Trigonal pyramidal	107°	NH3
2	2	Tetrahedral	Bent	104.5°	H2O

Names for bent

• Names for bent non-linear, angular and V-shaped are accepted.

SO₂

• SO₂ a double bond counts as one domain, so S has two bonding domains and one lone pair;
  • it is an expanded octet with 10 electrons.

Section 17 of 75Using VSEPR correctly

1Draw first

• Draw first always draw the Lewis formula first to find lone pairs.

2Do not confuse

• Do not confuse electron domain geometry and molecular geometry are the same only when there are no lone pairs.

3Example

• CH₄ is tetrahedral in both, but NH₃ has tetrahedral domains and a trigonal pyramidal shape.

Section 18 of 75More from the specification

More from the specification

Covalent bonding and sigma bonds

• A single covalent bond,
• such as between carbon and hydrogen atoms in hydrocarbons, consists of an axial end-on overlap of atomic orbitals forming a stable sigma bond.

Molecular dipole and polar bonds

• When bonded atoms differ in electronegativity, electron density shifts toward the more electronegative atom creating a bond dipole;
  • if molecular geometry is asymmetrical, individual dipoles do not cancel,
• producing a polar molecule.

Resonance structures and resonance hybrid

• Molecules such as benzene (C₆H₆) or carbonate ions cannot be described by a single Lewis structure,
• so chemists represent them as resonance structures contributing to a delocalized resonance hybrid structure with equal bond lengths and intermediate bond orders.

Intermolecular forces in molecular structures

• The physical properties of covalent substances depend on intermolecular forces (London dispersion forces, dipole-dipole interactions,
• and hydrogen bonding) acting between molecular structures.

Dispersion force

• Another name for induced dipole–induced dipole force.

Giant covalent

• Network of atoms joined by covalent bonds.

Hydrogen

• A gas that gives a squeaky pop with a lighted splint.

Dipole

• A separation of charge in a bond or molecule.

Resonance

• The existence of several valid Lewis structures for one species.

Sigma bond

• A bond from end-on overlap of orbitals along the line between two nuclei.

Formal charge

• V − ½B − N for an atom in a Lewis structure.

Intermolecular forces

• Weak attractions between separate molecules, overcome on melting or boiling.

Section 19 of 75Electronegativity and bond polarity

Electronegativity

Electronegativity the ability of an atom to attract the shared pair of electrons in a covalent bond.

Pauling scale

• Pauling scale values are in the data booklet (section 9);
  • the higher the value, the more electronegative the element.

Nonpolar bond

• Nonpolar bond identical atoms in H₂, O₂ and Cl₂ share electrons equally.

Polar bond

• Polar bond the electrons are drawn towards the more electronegative atom, which becomes δ− while the other becomes δ+.

Example

• Example in HCl, Cl is δ− and H is δ+.

Size of difference

• Size of difference a larger electronegativity difference gives a more polar bond and a larger dipole moment.

Dipole arrow

• Dipole arrow the crossed tail sits at the δ+ atom and the head points to the δ− atom.

Section 20 of 75Molecular polarity

Case	Result	Examples
Only nonpolar bonds	No bond dipoles, so the molecule is nonpolar	H2, Cl2
Polar bonds arranged symmetrically	Dipoles cancel, so the molecule is nonpolar	CO2, CCl4
Polar bonds arranged asymmetrically	Dipoles do not cancel, so the molecule is polar	H2O, NH3, CH3Cl

Impossible case

• Impossible case a molecule with only nonpolar bonds cannot be polar.

Method

• Method consider both the bond polarity and the molecular geometry.

Section 21 of 75Expanded octets (HL)

Expansion

• Expansion elements in period 3 and below can hold more than eight electrons using vacant d orbitals.

Five domains

• Five domains trigonal bipyramidal domain geometry.

Six domains

• Six domains octahedral domain geometry.

Molecular shape

• Molecular shape it depends on how many domains are lone pairs.

Section 22 of 75Shapes with five electron domains (HL)

Bonding	Lone	Molecular shape	Example	Polarity
5	0	Trigonal bipyramidal	PCl5	Nonpolar (symmetrical)
4	1	Seesaw	SF4	Polar
3	2	T-shaped	ClF3	Polar
2	3	Linear (180°)	I3−, XeF2	Nonpolar

Section 23 of 75Shapes with six electron domains (HL)

Bonding	Lone	Molecular shape	Example	Polarity
6	0	Octahedral	SF6	Nonpolar (symmetrical)
5	1	Square pyramidal	BrF5	Polar
4	2	Square planar	XeF4	Nonpolar (lone pairs above and below the plane)

Section 24 of 75Predicting polarity step by step

1

Draw the Lewis structure.

2

Decide the shape with VSEPR.

3

Mark the bond dipoles with arrows pointing to the more electronegative atom.

4

Decide whether the dipoles cancel by symmetry.

1Nonpolar despite polar bonds

• CO₂ and BF₃.

2Polar

• NH₃, H₂O and CH₃Cl.

Section 25 of 75Using polarity to predict properties

Boiling point

• Boiling point polar molecules attract each other by dipole–dipole forces in addition to London forces,
• so they usually boil higher than nonpolar molecules of similar mass.

Solubility

• Solubility polar molecules dissolve in polar solvents such as water;
  • nonpolar molecules dissolve mainly in nonpolar solvents.

Charged rod test

• Charged rod test a thin stream of polar liquid is deflected by a charged rod, but a nonpolar liquid is not.

Section 26 of 75More from the specification

1Sigma and Pi Bonding Overlap

• A single covalent bond consists of one axial sigma bond formed by end-to-end orbital overlap,
• whereas a double covalent bond comprises one sigma bond and one lateral pi bond formed by parallel p-orbital overlap.

2Valence Octet in Carbon Compounds

• In organic molecules such as methane, the central carbon atom forms four single covalent bonds with surrounding hydrogen atoms, completing a stable valence electron octet.

3Orbital Hybridization

• Carbon 2s and 2p atomic orbitals undergo hybridization to generate equivalent hybrid orbitals that maximize bonding overlap.
  • (sp³ tetrahedral in alkanes, sp² trigonal planar in alkenes)

4Resonance Hybrids and Delocalization

• Molecules that cannot be represented accurately by a single fixed double or single bond structure exist as resonance hybrids with delocalized pi electron density distributed across adjacent atoms.

5Dispersion force

• Another name for induced dipole–induced dipole force.

6Formal charge

• V − ½B − N for an atom in a Lewis structure.

7Intermolecular forces

• Weak attractions between separate molecules, overcome on melting or boiling.

8Lewis formula

• A diagram showing all bonding pairs and lone pairs.

9Giant covalent

• Network of atoms joined by covalent bonds.

10Hydrogen

• A gas that gives a squeaky pop with a lighted splint.

11Resonance

• The existence of several valid Lewis structures for one species.

12Oxygen

• A gas that relights a glowing splint.

13Tetrahedral

• Having each atom bonded to four others arranged around it.

14Sigma bond

• A bond from end-on overlap of orbitals along the line between two nuclei.

Section 27 of 75Giant covalent structures

Giant covalent structures

Definition

• Definition strong covalent bonds extend throughout the lattice, so there are no individual molecules.

Examples

• Examples silicon, silicon dioxide, diamond, graphite and graphene.

Silicon

• Silicon each atom bonded to four others, tetrahedral, 109.5°.

Silicon(IV) oxide

• Silicon(IV) oxide each Si bonded to four O and each O to two Si;
  • empirical formula SiO₂, the main component of sand.

Section 28 of 75Carbon allotropes

Diamond

• Diamond each C bonded to four others in a tetrahedral network at 109.5°.

Diamond properties

• Diamond properties extremely rigid and hard because many strong bonds must be broken.
• It does not conduct because all four outer electrons are used in bonding.

Graphite

• Graphite each C bonded to three others in flat hexagonal layers at 120°.
• The fourth electron is delocalised, so graphite conducts.
• Weak forces between layers let them slide, so it is soft and slippery.

Graphene

• Graphene a single layer of graphite, about a million times thinner than paper, strong and flexible, conducting through delocalised electrons.

Buckminsterfullerene C60

• Buckminsterfullerene C60 a football-shaped molecule of 20 hexagons and 12 pentagons, with each C bonded to three others.
• Some electron delocalisation makes it a semiconductor.
• It is a molecule with a fixed formula, not a giant structure.

Section 29 of 75Properties of giant covalent substances

	Diamond	Graphite	Graphene	C60	Silicon	SiO2
Melting and boiling points	Very high	Very high	Very high	Low	High	Very high
Electrical conductivity	None	Good	Very good	Poor	Semiconductor	None
Other	Hardest natural substance	Soft, slippery	Thinnest, strongest material	Light, strong	Good thermal conductor	Hard

High melting points

• High melting points many strong covalent bonds must be broken.

Solubility

• Solubility most network solids are insoluble in water.

Section 30 of 75London (dispersion) forces

Intermolecular forces

• Intermolecular forces act between molecules and are much weaker than covalent or ionic (intramolecular) bonds.

Roles

• Roles they determine melting and boiling points, volatility and solubility.

Cause

• Cause moving electrons give a temporary dipole that induces a dipole in a neighbour, causing weak attraction.

Presence

• Presence they exist between all atoms and molecules and are the only forces in nonpolar substances, typically 1–50 kJ mol⁻¹.

Number of electrons

• Number of electrons more electrons give stronger forces, so noble gas boiling points rise with atomic number.

Surface area

• Surface area unbranched, extended isomers have stronger forces than compact ones.

Van der Waals forces

• Van der Waals forces collective term for London, dipole–dipole and dipole–induced dipole attractions.

Section 31 of 75Dipole–dipole and dipole–induced dipole attractions

Dipole–dipole

• Dipole–dipole between the δ+ end of one polar molecule and the δ− end of another, in addition to London forces.

Permanent dipole

• Permanent dipole needs polar bonds in an asymmetric shape.

Effect

• Effect polar compounds usually have higher boiling points than nonpolar compounds of similar size.

Butane and propanone

• Butane and propanone they have the same number of electrons, but only propanone is polar, so propanone boils higher.

Dipole–induced dipole

• Dipole–induced dipole a polar molecule such as HCl distorts the electron cloud of a nonpolar molecule such as Cl₂, inducing a temporary dipole.

Section 32 of 75Hydrogen bonding

Hydrogen bonding

Strength

• Strength the strongest intermolecular force, a special case of dipole–dipole attraction.

Condition 1

• Condition 1 a hydrogen atom covalently bonded to O, N or F.

Condition 2

• Condition 2 a lone pair on an O, N or F atom of a neighbouring molecule.

Cause

• Cause the bond is strongly polarised, so the hydrogen is very δ+ and is attracted to the lone pair.

Notation

• Notation hydrogen bonds are drawn as dotted lines.

Number formed

• Number formed it depends on the H atoms on O, N or F and the lone pairs available.

Ammonia

• Ammonia one hydrogen bond per molecule.
  • (one lone pair on N)

Water

• Water two hydrogen bonds per molecule.
  • (two lone pairs and two H atoms on O)

Strength order

• London < dipole–dipole < hydrogen bonding.

Section 33 of 75More from the specification

More from the specification

Lewis structures and octet rule

• Lewis structures and octet rule valence electron diagrams display bonding pairs and non-bonding electron pairs distributed around central atoms to achieve octets.

Hybrid atomic orbitals and sigma bonds

• Hybrid atomic orbitals and sigma bonds central carbon atoms form localized sigma bonds through coaxial overlap of sp hybrid orbitals with neighbouring atoms.

Resonance hybrids and electron delocalization

• Resonance hybrids and electron delocalization equivalent Lewis resonance structures depict molecular geometry where delocalised pi electrons stabilize formal charge distributions across bonds.

Orbital hybridization in carbon

• Carbon atom 2s and 2p atomic orbitals hybridize into four equivalent sp3 hybrid orbitals, three sp2 orbitals (trigonal planar, 120°),
  • (tetrahedral, 109.5°)
• or two sp orbitals (linear, 180°).

Sigma and pi covalent bonding

• Sigma bonds form by direct head-on overlap of atomic or hybrid orbitals along the internuclear axis;
  • pi bonds form by lateral side-on overlap of unhybridized p orbitals.

VSEPR molecular geometry prediction

• Valence Shell Electron Pair Repulsion theory predicts molecular geometry based on minimizing electrostatic repulsion between bonding pairs and lone pairs of electrons.

Dispersion force

• Another name for induced dipole–induced dipole force.

Oxygen

• A gas that relights a glowing splint.

Molecular geometry

• Arrangement of the atoms only.

Sigma bond

• A bond from end-on overlap of orbitals along the line between two nuclei.

Electron domain

• A bonding pair, lone pair or multiple bond region around an atom.

Formal charge

• V − ½B − N for an atom in a Lewis structure.

Orbital

• A region of space around the nucleus that holds up to two electrons of opposite spin.

Charge

• Q = It, measured in coulombs.
  What the symbols mean: Q = charge (C) · I = current (A) · t = time (s)

Resonance

• The existence of several valid Lewis structures for one species.

Section 34 of 75Melting point, boiling point and volatility

What is overcome

• What is overcome when a molecular substance changes state, the intermolecular forces are overcome, not the covalent bonds.

Room temperature

• Room temperature these forces are weak, so many covalent substances are gases or liquids.

Trend

• Trend stronger forces give higher melting and boiling points.

Volatile

• Volatile substances with low boiling points are described as volatile.

Factors

• Factors molar mass, polarity and the type of force, in the order London < dipole–dipole < hydrogen bonding.

Equal mass

• Equal mass for molecules of similar M_r, hydrogen bonding raises the boiling point most, then dipole–dipole forces, then London forces alone.

Section 35 of 75Solubility

Rule

• Rule like dissolves like.

Nonpolar in nonpolar

• Nonpolar in nonpolar dissolves through London forces between solute and solvent.

Polar in polar

• Polar in polar dissolves through dipole–dipole attractions or hydrogen bonding.

Ethanol in water

• Ethanol in water soluble because it forms hydrogen bonds with water.

Longer chains

• Longer chains the nonpolar hydrocarbon part dominates, so hexanol is much less soluble than ethanol.

Polar in nonpolar

• Polar in nonpolar polar substances do not dissolve well in nonpolar solvents.

Giant covalent

• Giant covalent insoluble in both, because too much energy is needed to break the lattice.

Section 36 of 75Conductivity

Conductivity

Molecular covalent substances

• Molecular covalent substances do not conduct in any state, because they contain neutral molecules with no mobile ions or delocalised electrons.

Exception in solution

• Exception in solution some polar covalent substances ionise in water, for example HCl gives H⁺ and Cl⁻.

Giant covalent

• Giant covalent most do not conduct; graphite and graphene conduct through delocalised electrons.

Section 37 of 75Comparing covalent substances

	Nonpolar molecular	Polar molecular	Giant covalent
Melting and boiling points	Low	Low to moderate	Very high
Volatility	High	Moderate	Very low
Solubility in polar solvents	Insoluble	Often soluble	Insoluble
Solubility in nonpolar solvents	Soluble	Poor	Insoluble
Conductivity	None	None (ionic in solution for some)	None (except graphite and graphene)

Section 38 of 75Chromatography

Principle

• Principle a liquid mobile phase passes over a solid stationary phase and components move at different rates.

Basis

• Basis rates depend on solubility in the mobile phase and attraction to the stationary phase, which are differences in intermolecular forces.

IB scope

• IB scope paper chromatography and thin-layer chromatography (TLC);
  • column and gas chromatography are not required.

Paper chromatography

• Paper chromatography the paper is the stationary phase and the solvent is the mobile phase.
  • (water or ethanol)
• Paper is cellulose with polar –OH groups.
• Polar substances interact more with the paper, travel more slowly and do not go as far.
• Less polar substances are more soluble in the solvent and travel further.

TLC

• TLC the same principle, with a thin layer of inert silica or alumina on glass or plastic as the stationary phase.
• The –OH groups on its surface adsorb components to different extents.

Section 39 of 75Running a chromatogram

1

Draw a pencil baseline near the bottom, because ink would run.

2

Place small spots of sample on the line.

3

Suspend the paper with the baseline above the solvent.

4

Let the solvent rise by capillary action.

5

Remove the paper before the solvent reaches the top and mark the solvent front.

6

Dry the paper and examine the chromatogram.

• Component not moving
  • it is insoluble in that solvent, so a different solvent should be tried.

Section 40 of 75R_f values

Definition

• R_f = distance moved by the component ÷ distance moved by the solvent front, both measured from the baseline.

No units

• No units it is a ratio of two lengths, quoted as a decimal.

Always below 1

• Always below 1 a component cannot travel further than the solvent front.

Higher Rf

• Higher Rf the component is more soluble in the mobile phase and interacts less with the stationary phase.

Use

• R_f values help identify the components of a mixture.

Section 41 of 75More from the specification

1VSEPR theory and electron pairs

• Valence Shell Electron Pair Repulsion (VSEPR) theory dictates that bonding pairs and lone pairs of electrons around a central atom repel each other to adopt an electron domain geometry that minimizes electrostatic repulsion.

2Molecular shape vs electron geometry

• The molecular shape is defined by the positions of bonded atoms around the central atom;
  • lone pairs exert greater repulsive force than bonding pairs,
• reducing bond angles from ideal geometric values.

3Octet rule and formal charge

• In Lewis structures, outer shell electrons are arranged to satisfy the octet rule for second-row atoms, with formal charge calculations used to identify the most stable arrangement.
  • (eight valence electrons)

4Orbital hybridisation and sigma frameworks

• Mixing atomic orbitals generates equivalent hybrid orbitals (sp, sp2, sp3) that form directional sigma bonds with overlapping orbitals of adjacent atoms.

5Dispersion force

• Another name for induced dipole–induced dipole force.

6Oxygen

• A gas that relights a glowing splint.

7Tetrahedral

• Having each atom bonded to four others arranged around it.

8Bonding pair

• A shared pair of electrons forming a covalent bond.

9Molecular geometry

• Arrangement of the atoms only.

10Sigma bond

• A bond from end-on overlap of orbitals along the line between two nuclei.

11Electron domain

• A bonding pair, lone pair or multiple bond region around an atom.

12Formal charge

• V − ½B − N for an atom in a Lewis structure.

13Lewis formula

• A diagram showing all bonding pairs and lone pairs.

14Orbital

• A region of space around the nucleus that holds up to two electrons of opposite spin.

15Charge

• Q = It, measured in coulombs.
  What the symbols mean: Q = charge (C) · I = current (A) · t = time (s)

16Resonance

• The existence of several valid Lewis structures for one species.

Section 42 of 75Resonance structures

Resonance structures

Resonance

• Resonance occurs when more than one valid Lewis structure can be drawn because of delocalised electrons.

Resonance hybrid

• Resonance hybrid the real structure does not switch between forms;
  • electron density is spread over several atoms.

Explains

• Explains bond lengths, bond strengths and electron distribution.

Requirement

• Requirement a π bond that can occupy more than one position, with adjacent atoms of similar electronegativity so lone pairs or double bonds shift without breaking octet rules.

Notation

• Notation dotted lines show delocalised electrons in a hybrid diagram.

Section 43 of 75Examples of resonance

Nitrate(V), NO₃⁻

• 5 + (3 × 6) + 1 = 24 electrons.
• One Lewis diagram shows one N=O and two N–O bonds, but three equivalent structures exist.
• In the hybrid the charge and π bond are delocalised over all three oxygens.
• All three N–O bonds are equal in length, between single and double.

Carbonate, CO₃²⁻

• Carbonate, CO₃²⁻ three equivalent structures and three equal C–O bonds.

Ozone, O₃

• Ozone, O₃ bent, with two equivalent structures and two equal O–O bonds.

Carboxylate, RCOO−

• Carboxylate, RCOO− the negative charge and π bond are shared between the two oxygens, giving two equal C–O bonds.

Section 44 of 75Structure of benzene

Structure of benzene

Formula

• Formula benzene is a ring of six carbons, C₆H₆.

Kekulé model

• Kekulé model alternating single and double bonds, which would make it react like an alkene; it does not.

Shape

• Shape a planar regular hexagon with 120° angles.

Hybridisation

• Hybridisation each carbon is sp² and forms three σ bonds, two to carbon neighbours and one to hydrogen.

Delocalised π system

• Delocalised π system the unhybridised p orbitals perpendicular to the ring overlap sideways, above and below the ring.

Electron sharing

• Electron sharing the six π electrons are shared equally across all six carbons.

Result

• Result unusual stability, and all C–C bonds equal in length and between single and double.

Diagram

• Diagram a circle inside the hexagon represents the delocalised system.

Section 45 of 75Evidence from enthalpy and bond length

Evidence	Kekulé model predicts	Observed
Enthalpy of hydrogenation	Three C=C bonds would give 3 × (−120) = −360 kJ mol−1 (cyclohexene is −120)	−208 kJ mol−1, 152 kJ mol−1 less exothermic
C–C bond lengths (X-ray)	Alternating 154 pm (single) and 134 pm (double)	All 140 pm, between single and double

• Delocalisation energy
  • the 152 kJ mol⁻¹ difference, showing benzene is more stable than the Kekulé model suggests.

Section 46 of 75Evidence from reactions and infrared

Evidence	Kekulé model predicts	Observed
Saturation test	Would decolourise bromine water	Does not decolourise it, so no isolated C=C bonds
Infrared spectroscopy	Strong C=C absorption at 1620–1680 cm−1 as in cyclohexene	No absorption there; weaker peaks near 1450, 1500 and 1580 cm−1 typical of arenes

Section 47 of 75Consequences of delocalisation

1Stability

• Stability delocalisation lowers energy, so delocalised species are more stable than one Lewis structure suggests.

2Carbonate

• Carbonate the C–O bond length (about 129 pm) lies between C–O (143 pm) and C=O (122 pm).

3Ozone

• Ozone the two O–O bonds are equal, longer than a double bond but shorter than a single bond.

4Carboxylate

• Carboxylate it is more stable than the alkoxide ion, so carboxylic acids are stronger acids than alcohols.

5Benzene reactions

• Benzene reactions it undergoes substitution, which keeps the delocalised ring, rather than the addition typical of alkenes.

Section 48 of 75More from the specification

More from the specification

Sigma and Pi Covalent Bonds

• A single covalent bond consists of an end-to-end orbital overlap called a sigma (σ) bond,
• whereas double and triple bonds contain one σ bond and one or two lateral overlap pi (π) bonds.

Electronegativity and Bond Dipoles

• Differences in electronegativity between bonded atoms cause asymmetric electron distribution, creating a polar covalent bond with partial delta charges (δ+ and δ-) and a permanent bond dipole.

Molecular Geometry and Net Dipole Moments

• A molecule containing polar bonds will have an overall net molecular dipole if its three-dimensional geometry is asymmetrical,
• whereas symmetrical geometries cancel dipoles out.
  • (e.g. linear CO₂, tetrahedral CCl₄)

Intermolecular Forces Between Molecules

• Polar molecules experience permanent dipole-dipole attractions,
• whereas all molecules exhibit London dispersion forces arising from instantaneous induced temporary dipoles.

Dispersion force

• Another name for induced dipole–induced dipole force.

Tetrahedral

• Having each atom bonded to four others arranged around it.

Molecular geometry

• Arrangement of the atoms only.

Sigma bond

• A bond from end-on overlap of orbitals along the line between two nuclei.

Electron domain

• A bonding pair, lone pair or multiple bond region around an atom.

Formal charge

• V − ½B − N for an atom in a Lewis structure.

Intermolecular forces

• Weak attractions between separate molecules, overcome on melting or boiling.

Giant covalent

• Network of atoms joined by covalent bonds.

Bonding

• The joining of atoms.

Molecule

• Two or more atoms held together by covalent bonds.

Dipole

• A separation of charge in a bond or molecule.

Section 49 of 75Expansion of the octet

Elements

• Elements those in period 3 and below can have more than eight electrons around the central atom.

Reason

• Reason they have vacant d orbitals in the valence shell to hold extra bonding pairs.

Examples

• PCl₅ has five bonding pairs and SF₆ has six.

Fluorine

• F cannot expand its octet, so each F has 3 lone pairs.

Section 50 of 75Lewis formulas with five domains

Species	Valence electrons	Lewis formula
PCl5	5 + 35 = 40	5 bonding pairs; 15 lone pairs, 3 on each Cl; P has 10 electrons
SF4	6 + 28 = 34	4 bonding pairs; 3 lone pairs on each F; 1 lone pair on S, which has 10 electrons
I3−	21 + 1 = 22	2 bonding pairs; 3 lone pairs on each outer I and on the central I (10 electrons)
ClF3	7 + 21 = 28	3 bonding pairs; 3 lone pairs on each F; 2 lone pairs on Cl

Section 51 of 75Lewis formulas with six domains

Species	Valence electrons	Lewis formula
SF6	6 + 42 = 48	6 bonding pairs; 3 lone pairs on each F; S has 12 electrons
BrF5	7 + 35 = 42	5 bonding pairs; 3 lone pairs on each F; 1 lone pair on Br (12 electrons)
XeF4	8 + 28 = 36	4 bonding pairs; 3 lone pairs on each F; 2 lone pairs on Xe

Section 52 of 75Geometry of expanded-octet species

Five domains

• Five domains trigonal bipyramidal domain geometry, giving PCl₅ , SF₄ (seesaw), ClF₃ (T-shaped) and I₃⁻ (linear).
  • (trigonal bipyramidal)

Six domains

• Six domains octahedral domain geometry, giving SF₆ (octahedral), BrF₅ and XeF₄ (square planar).
  • (square pyramidal)

Electron domain geometry

• Electron domain geometry the arrangement of all electron pairs.

Molecular geometry

• Molecular geometry the arrangement of the atoms only.

Multiple bonds

• Multiple bonds they count as one domain.

Section 53 of 75Formal charge

Purpose

• Purpose formal charge (FC) decides which Lewis formula is most likely when several are valid.

Assumption

• Assumption bonding electrons are shared equally regardless of electronegativity.

Formula

• FC = V − ½B − N.

V

V valence electrons of the free atom.

B

B number of bonding electrons.

N

N number of non-bonding (lone pair) electrons.

Check

• Check the formal charges add up to the overall charge of the species.

Section 54 of 75Choosing between structures

Choosing between structures

Criterion 1

• Criterion 1 the preferred formula has the difference between largest and smallest formal charge, ΔFC = FC_max − FC_min, closest to zero.

Criterion 2

• Criterion 2 any negative formal charge sits on the most electronegative atom.

SO₂

• SO₂ two S=O bonds give formal charges of 0 on all atoms, so ΔFC = 0.
  • (S has 10 electrons)

SO₂ alternative

• SO₂ alternative one S=O and one S–O gives S +1 and the single-bonded O −1, so ΔFC = 2 and it is not preferred.

CO₂

• O=C=O has all formal charges zero, so it is preferred over a C≡O/C–O structure with ΔFC = 2.

Section 55 of 75More from the specification

1Sigma (σ) Bonds and Orbital Overlap

• Formed by the direct, head-on overlap of atomic or hybrid orbitals along the internuclear axis, such as the C-H bonds in methane.

2Hybridisation in Carbon Chemistry

• Mixing of atomic s and p orbitals produces hybrid orbitals (sp³, sp²,
• or sp) with directional geometries suitable for forming stable covalent bonds.

3Resonance and Pi Delocalisation

• Conjugated planar molecules exhibit resonance where pi electrons are delocalised over multiple atoms rather than localized in fixed double bonds.

4Bond Polarity and Molecular Dipole Moments

• When polar bonds do not cancel out geometrically, the molecule possesses a permanent molecular dipole governed by intermolecular dipole-dipole forces.

5Sigma (σ) Bonds and Orbital Overlap

• Formed by the direct, head-on overlap of atomic or hybrid orbitals along the internuclear axis, such as the C-H bonds in methane.

6Hybridisation in Carbon Chemistry

• Mixing of atomic s and p orbitals produces hybrid orbitals (sp³, sp²,
• or sp) with directional geometries suitable for forming stable covalent bonds.

7Resonance and Pi Delocalisation

• Conjugated planar molecules exhibit resonance where pi electrons are delocalised over multiple atoms rather than localized in fixed double bonds.

8Bond Polarity and Molecular Dipole Moments

• When polar bonds do not cancel out geometrically, the molecule possesses a permanent molecular dipole governed by intermolecular dipole-dipole forces.

9Dispersion force

• Another name for induced dipole–induced dipole force.

10Covalent bond

• A shared pair of electrons between two atoms, giving noble gas configurations.

11Sigma bond

• A bond from end-on overlap of orbitals along the line between two nuclei.

12Intermolecular forces

• Weak attractions between separate molecules, overcome on melting or boiling.

13Giant covalent

• Network of atoms joined by covalent bonds.

14Hydrogen

• A gas that gives a squeaky pop with a lighted splint.

15Dipole

• A separation of charge in a bond or molecule.

16Resonance

• The existence of several valid Lewis structures for one species.

17Oxygen

• A gas that relights a glowing splint.

18Tetrahedral

• Having each atom bonded to four others arranged around it.

Section 56 of 75Orbital overlap and bond types

Orbital overlap and bond types

Formation

• Formation a single covalent bond forms when two non-metal atoms each contribute an unpaired electron in orbitals that overlap to form a molecular orbital.

Overlap

• Overlap the greater the overlap, the stronger the bond.

Types

• Types sigma (σ) bonds and pi (π) bonds.

Section 57 of 75Sigma bonds

Overlap

• Overlap head-on (axial) overlap of atomic or hybrid orbitals.

Electron density

• Electron density concentrated along the bond axis between the nuclei.

Strength

• Strength the strongest type of covalent bond.

Presence

• Presence in all single bonds and in every double and triple bond.

s–s overlap

• s–s overlap the 1s orbitals of two hydrogen atoms in H₂.

s–p overlap

• H 1s with F 2p in HF.

p–p overlap

• p–p overlap the 2p orbitals of two fluorine atoms in F₂.

Section 58 of 75Pi bonds

Overlap

• Overlap sideways overlap of adjacent unhybridised p orbitals.

Electron density

• Electron density above and below the plane of the σ bond, on opposite sides of the bond axis.

Occurrence

• Occurrence only in double and triple bonds.

Strength

• Strength weaker than σ bonds because sideways overlap is less effective.

Section 59 of 75Sigma and pi bonds in molecules

Molecule	σ bonds	π bonds	Description
Methane	4	0	Four head-on overlaps of carbon hybrid orbitals with H 1s
Ethene	5	1	C=C is one σ plus one π; π from sideways overlap of leftover p orbitals
Ethyne	3	2	C≡C is one σ plus two perpendicular π bonds
Nitrogen, N2	1	2	Triple bond of one σ and two π
HCN	2	2	C–H is σ; C≡N is one σ and two π

• Counting
  • every bond contains exactly one σ bond, and each extra bond in a double or triple bond is a π bond.

Section 60 of 75Effects of pi bonds

Reactivity

• π electron density lies outside the bond axis, exposed to electrophiles, so alkenes and alkynes are more reactive than alkanes.

No rotation

• No rotation a π bond prevents rotation about a double bond, which causes cis–trans isomerism.

Free rotation

• Free rotation single σ bonds rotate freely.

Bond strength

• C=C (about 614 kJ mol⁻¹) is less than twice C–C (about 346 kJ mol⁻¹), so the π component is weaker than the σ bond.

Nitrogen

• Nitrogen the triple bond (one σ, two π) makes N₂ unreactive at room temperature.

Section 61 of 75More from the specification

More from the specification

Octet Rule and Lewis Structures

• In covalent Lewis structures, outer atoms share electron pairs with central atoms to complete an octet of eight valence electrons, balancing total formal charge across the molecule.
  • (or a duet for hydrogen)

Bonding Pairs vs. Lone Pairs in Molecular Geometry

• Bonding pairs of shared electrons form covalent bonds,
• while non-bonding lone pairs exert greater electrostatic repulsion, compressing bond angles and altering overall molecular geometry from ideal electron domain arrangements.

Net Molecular Dipole and Polar Structures

• A molecule possesses a permanent net dipole moment if its 3D geometry is asymmetrical,
• causing bond polarities to sum vectorially rather than cancel,
• resulting in a polar structure.

Influence of Molecular Polarity on Intermolecular Forces

• Polar molecules with permanent dipoles experience dipole-dipole attractions in addition to dispersion forces,
• resulting in higher boiling points compared to nonpolar structures of similar molar mass.

Dispersion force

• Another name for induced dipole–induced dipole force.

Giant covalent

• Network of atoms joined by covalent bonds.

Oxygen

• A gas that relights a glowing splint.

Tetrahedral

• Having each atom bonded to four others arranged around it.

Bonding pair

• A shared pair of electrons forming a covalent bond.

Molecular geometry

• Arrangement of the atoms only.

Electron domain

• A bonding pair, lone pair or multiple bond region around an atom.

Formal charge

• V − ½B − N for an atom in a Lewis structure.

Intermolecular forces

• Weak attractions between separate molecules, overcome on melting or boiling.

Lewis formula

• A diagram showing all bonding pairs and lone pairs.

Bonding

• The joining of atoms.

Dipole

• A separation of charge in a bond or molecule.

Charge

• Q = It, measured in coulombs.
  What the symbols mean: Q = charge (C) · I = current (A) · t = time (s)

Resonance

• The existence of several valid Lewis structures for one species.

Section 62 of 75Why atoms hybridise

Hybridisation

Hybridisation the mixing of atomic orbitals to form new hybrid orbitals used in covalent bonding.

Explains

• Explains why atoms such as carbon form bonds of equal strength and direction.

Carbon ground state

• 1s² 2s² 2p², with only two unpaired electrons.

Promotion

• Promotion one 2s electron moves to the empty 2p orbital, giving four unpaired electrons whose orbitals mix.

Orbital shapes

• Orbital shapes s is spherical and the three p orbitals (p_x, p_y, p_z) are dumbbells at right angles.

Hybrid orbitals

• Hybrid orbitals asymmetric with one large lobe, like a p orbital.

Period 3

• 3s and 3p orbitals hybridise in the same way, and 3d may be used for expanded octets.

Section 63 of 75sp³ hybridisation

Formation

• Formation one s and three p orbitals of the same shell give four sp³ orbitals with ¼ s and ¾ p character.

Geometry

• Geometry arranged tetrahedrally.

Methane

• Methane each sp³ orbital overlaps head-on with a hydrogen 1s orbital.
• Four identical σ bonds, tetrahedral domain and molecular geometry, 109.5°.

Ammonia

• Ammonia nitrogen has three bonding pairs and one lone pair in sp³ orbitals.
• Three σ bonds and a lone pair, tetrahedral domains, trigonal pyramidal shape, 107°.
• Hybrid orbitals can hold lone pairs as well as bonding pairs.

Section 64 of 75sp² hybridisation

1Formation

• Formation one s and two p orbitals give three sp² orbitals with ⅓ s and ⅔ p character.

2Geometry

• Geometry trigonal planar, 120°.

3Ethene

• Ethene each carbon uses three sp² orbitals for σ bonds, two to H and one to C.
• The unhybridised p orbital forms a π bond with the other carbon.
• C=C is one σ plus one π bond.

4Carbonyl groups

• Carbonyl groups both C and O are sp² in C=O.

Section 65 of 75sp hybridisation

Formation

• Formation one s and one p orbital give two sp orbitals with ½ s and ½ p character.

Geometry

• Geometry linear, 180°.

Ethyne

• Ethyne each carbon uses two sp orbitals, one to H and one to C.
• Two unhybridised p orbitals form two perpendicular π bonds.
• C≡C is one σ plus two π bonds.

Section 66 of 75Quick method for carbon

Hybridisation	p orbitals used in hybrid	π bonds	Bond type	Geometry	Angle
sp3	3	0	Single bonds	Tetrahedral	109.5°
sp2	2	1	Double bond	Trigonal planar	120°
sp	1	2	Triple bond	Linear	180°

Rule

• Rule carbon forms four bonds; the number of p orbitals left unhybridised equals the number of π bonds.

Period 3

• Period 3 elements such as P and S hybridise 3s and 3p orbitals, sometimes with 3d.

Section 67 of 75Predicting hybridisation from domains

Counting

• Counting count electron domains.
  • (bonded atoms plus lone pairs, multiple bonds counted once)

Two domains

• Two domains sp, linear, as in the carbon of CO₂.

Three domains

• Three domains sp², trigonal planar, as in the carbon of methanal.

Four domains

• Four domains sp³, tetrahedral, as in the oxygen of H₂O with two bonding pairs and two lone pairs.

HCN

• HCN carbon is sp; nitrogen is sp with a lone pair in an sp orbital.
  • (a σ bond to H, and a σ and two π bonds to N)

Section 68 of 75More from the specification

More from the specification

Lewis structure drawing rules

• To construct a valid Lewis structure, calculate total valence electrons, connect terminal atoms to the central atom via sigma bonds, distribute remaining electron pairs to complete octets,
• and verify minimal formal charge.

Molecular polarity and dipole moment

• A molecule is polar and possesses an overall net dipole moment if it contains polar bonds and has an asymmetrical structure where bond dipoles fail to cancel vectorially.

Resonance structures and delocalisation

• In molecules or polyatomic ions with delocalized pi systems, multiple valid resonance structures exist, sharing formal charge across equivalent peripheral atoms.
  • (e.g. ozone or nitrate)

Influence of intermolecular forces

• Net molecular dipole and charge distribution govern the strength of intermolecular forces, determining boiling points and solubility.
  • (dipole-dipole attractions and London dispersion forces)

Dispersion force

• Another name for induced dipole–induced dipole force.

Lewis formula

• A diagram showing all bonding pairs and lone pairs.

Giant covalent

• Network of atoms joined by covalent bonds.

Dipole

• A separation of charge in a bond or molecule.

Charge

• Q = It, measured in coulombs.
  What the symbols mean: Q = charge (C) · I = current (A) · t = time (s)

Resonance

• The existence of several valid Lewis structures for one species.

Sigma bond

• A bond from end-on overlap of orbitals along the line between two nuclei.

Formal charge

• V − ½B − N for an atom in a Lewis structure.

Intermolecular forces

• Weak attractions between separate molecules, overcome on melting or boiling.

Section 69 of 75Sigma and pi bonds

1Sigma bond

• Sigma bond formed by overlap along the line joining the two nuclei.

2Pi bond

• Pi bond formed by side-on overlap of unhybridised p orbitals.

3Double bond

• Double bond one σ and one π bond.

4Triple bond

• Triple bond one σ and two π bonds.

5First bond

• First bond the first bond between two atoms is always a σ bond, and extra bonds are π bonds.

Section 70 of 75sp² hybridisation in ethene

1Hybrid orbitals

• Hybrid orbitals each carbon mixes one s orbital with two p orbitals to form three sp² orbitals.

2σ bonds

• σ bonds two C–H bonds and one C–C bond.

3π bond

• π bond the unhybridised p orbital on each carbon overlaps side-on to form the π bond of C=C.

4Geometry

• Geometry three electron domains around each carbon, trigonal planar, about 120°.

Section 71 of 75sp hybridisation in ethyne

Hybrid orbitals

• Hybrid orbitals each carbon mixes one s orbital with one p orbital to form two sp orbitals.

σ bonds

• σ bonds one C–H bond and one C–C bond.

π bonds

• π bonds two unhybridised p orbitals on each carbon form two perpendicular π bonds.

Triple bond

• C≡C is one σ and two π bonds, not three σ bonds.

Geometry

• Geometry linear, 180°.

Section 72 of 75Comparing carbon centres

Carbon with	Hybridisation	Geometry	Angle
Four single bonds	sp3	Tetrahedral	109.5°
A double bond	sp2	Trigonal planar	120°
A triple bond	sp	Linear	180°

Section 73 of 75More from the specification

More from the specification

Covalent bonding and shared electron pairs

• Covalent bonding involves electrostatic attraction between positively charged atomic nuclei and shared pairs of bonding electrons,
• forming discrete molecular structures.

Polar covalent bonds and molecular dipoles

• A covalent bond between atoms of differing electronegativity creates an unequal distribution of electron density and a bond dipole;
  • asymmetrical arrangement of dipoles makes the overall molecule polar.

Intermolecular forces between covalent molecules

• The physical properties of molecular substances are governed by intermolecular forces,
• including London dispersion forces, dipole-dipole attractions,
• and intermolecular hydrogen bonding.

Hydrogen bonding in molecular structures

• Hydrogen bonding occurs when a hydrogen atom bonded to highly electronegative nitrogen, oxygen,
• or fluorine interacts with a lone pair on an adjacent electronegative atom in neighbouring molecules.

Covalent bonding and shared electron pairs

• Covalent bonding involves electrostatic attraction between positively charged atomic nuclei and shared pairs of bonding electrons,
• forming discrete molecular structures.

Polar covalent bonds and molecular dipoles

• A covalent bond between atoms of differing electronegativity creates an unequal distribution of electron density and a bond dipole;
  • asymmetrical arrangement of dipoles makes the overall molecule polar.

Intermolecular forces between covalent molecules

• The physical properties of molecular substances are governed by intermolecular forces,
• including London dispersion forces, dipole-dipole attractions,
• and intermolecular hydrogen bonding.

Hydrogen bonding in molecular structures

• Hydrogen bonding occurs when a hydrogen atom bonded to highly electronegative nitrogen, oxygen,
• or fluorine interacts with a lone pair on an adjacent electronegative atom in neighbouring molecules.

Dispersion force

• Another name for induced dipole–induced dipole force.

Resonance

• The existence of several valid Lewis structures for one species.

Oxygen

• A gas that relights a glowing splint.

Bonding pair

• A shared pair of electrons forming a covalent bond.

Formal charge

• V − ½B − N for an atom in a Lewis structure.

Intermolecular forces

• Weak attractions between separate molecules, overcome on melting or boiling.

Lewis formula

• A diagram showing all bonding pairs and lone pairs.

Giant covalent

• Network of atoms joined by covalent bonds.

Bonding

• The joining of atoms.

Molecule

• Two or more atoms held together by covalent bonds.

Hydrogen

• A gas that gives a squeaky pop with a lighted splint.

Dipole

• A separation of charge in a bond or molecule.

Charge

• Q = It, measured in coulombs.
  What the symbols mean: Q = charge (C) · I = current (A) · t = time (s)

Section 74 of 75Key terms

Term	Meaning
Covalent bond	Attraction between a shared electron pair and two nuclei.
Electron domain	A bonding pair, lone pair or multiple bond region around an atom.
Electronegativity	Ability of an atom to attract the shared pair in a bond.
Hydrogen bond	Strong intermolecular attraction involving H bonded to O, N or F.
Resonance hybrid	The real structure formed by delocalised electrons.
Sigma bond	Head-on overlap bond.
Pi bond	Sideways overlap bond.
Hybridisation	Mixing of atomic orbitals to form equal hybrid orbitals.

Term	Meaning
Lewis formula	A diagram showing all bonding pairs and lone pairs.
Coordinate bond	A covalent bond with both electrons from one atom.
VSEPR	Valence shell electron pair repulsion theory predicting shapes.
Bond energy	Energy to break one mole of a gaseous covalent bond.
Bond length	Distance between the nuclei of bonded atoms.
Electron domain	A region of electron density around an atom.

Term	Meaning
Polar bond	A covalent bond with unequal sharing of electrons.
Dipole	A separation of charge across a bond or molecule.
Dipole moment	A measure of how polar a bond is.
Nonpolar molecule	A molecule with no net dipole.
Seesaw	The shape with 4 bonding pairs and 1 lone pair on 5 domains.
Square planar	The shape with 4 bonding pairs and 2 lone pairs on 6 domains.

Term	Meaning
Giant covalent structure	A lattice of atoms held by covalent bonds throughout.
Allotropes	Different structural forms of the same element.
London forces	Weak attractions from temporary and induced dipoles.
Dipole–dipole attraction	Attraction between permanent dipoles.
Hydrogen bond	Attraction between H on O, N or F and a lone pair on a neighbouring O, N or F.
Delocalised electrons	Electrons free to move through a structure.

Term	Meaning
Volatile	Having a low boiling point and evaporating easily.
Mobile phase	The solvent that moves through the chromatography medium.
Stationary phase	The solid that stays in place, such as paper or silica.
Rf value	Component distance divided by solvent front distance.
Like dissolves like	Substances dissolve best in solvents of similar polarity.
Capillary action	Rise of liquid through paper without external force.

Term	Meaning
Resonance	The existence of several valid Lewis structures for one species.
Resonance hybrid	The real structure with delocalised electrons.
Delocalised electrons	Electrons shared over more than two atoms.
Delocalisation energy	The extra stability of benzene compared with the Kekulé model.
Kekulé structure	A benzene model with alternating single and double bonds.
Arene	A hydrocarbon containing a benzene ring.

Term	Meaning
Expanded octet	A central atom with more than 8 valence electrons.
Formal charge	V − ½B − N for an atom in a Lewis structure.
Seesaw	Molecular shape with 4 bonding and 1 lone pair on 5 domains.
T-shaped	Molecular shape with 3 bonding and 2 lone pairs on 5 domains.
Square pyramidal	Molecular shape with 5 bonding and 1 lone pair on 6 domains.
Square planar	Molecular shape with 4 bonding and 2 lone pairs on 6 domains.

Term	Meaning
Sigma bond	Bond formed by head-on orbital overlap.
Pi bond	Bond formed by sideways overlap of p orbitals.
Molecular orbital	The orbital formed when atomic orbitals overlap.
Bond axis	The line joining the two nuclei.
Unhybridised p orbital	A p orbital not used in hybridisation, available for a π bond.

Term	Meaning
Hybridisation	Mixing of atomic orbitals to give equivalent hybrid orbitals.
sp3 orbital	Hybrid of one s and three p orbitals, tetrahedral.
sp2 orbital	Hybrid of one s and two p orbitals, trigonal planar.
sp orbital	Hybrid of one s and one p orbital, linear.
Promotion	Moving an electron to a higher empty orbital before hybridising.
Unhybridised p orbital	A leftover p orbital that forms a π bond.

Term	Meaning
Sigma bond	Bond formed by overlap along the internuclear axis.
Pi bond	Bond formed by side-on overlap of p orbitals.
sp2 hybrid	One s and two p orbitals mixed, trigonal planar.
sp hybrid	One s and one p orbital mixed, linear.

Section 75 of 75Exam technique

1

Confusing electron domain geometry with molecular shape — lone pairs change the shape.

2

Saying a molecule with polar bonds must be polar — symmetrical molecules such as CO₂ and CCl₄ are nonpolar.

3

Breaking covalent bonds when a molecular substance melts — only intermolecular forces are overcome.

4

Counting a double bond as two domains in VSEPR — a multiple bond is one domain.

5

Giving an R_f value with units or a value above 1.

1

Using Cl–Cl as a Lewis formula — the lone pairs must be shown.

2

Counting a double bond as two electron domains.

3

Using 109.5° for ammonia and water — lone pairs reduce the angles to 107° and 104.5°.

4

Assuming every coordinate bond uses ammonia — any lone pair donor can form one.

1

Saying polar bonds always make a polar molecule — symmetrical arrangements cancel the dipoles.

2

Pointing the dipole arrow towards the less electronegative atom.

3

Calling XeF₄ polar because it has lone pairs — they sit above and below the plane and the molecule is symmetrical.

4

Using the electronegativity difference alone without checking the shape.

1

Calling C₆₀ a giant structure — it is a molecule with a fixed formula.

2

Saying graphite conducts because of ions — it conducts through delocalised electrons.

3

Saying covalent bonds break when a molecular substance melts — only intermolecular forces are overcome.

4

Describing London forces as "induced dipole" only — use the term London (dispersion) forces.

1

Saying covalent bonds break when a molecular substance boils — intermolecular forces are overcome.

2

Saying all covalent substances conduct in solution — only those that ionise do.

3

Measuring R_f from the bottom of the paper instead of the baseline.

4

Drawing the baseline in ink, which runs with the solvent.

1

Saying the nitrate ion flips between structures — it is one hybrid with equal bonds.

2

Quoting alternating bond lengths for benzene — all C–C bonds are 140 pm.

3

Predicting benzene decolourises bromine water — it does not.

4

Forgetting that the π electrons in benzene are above and below the ring.

1

Letting fluorine expand its octet — only central atoms in period 3 or below can.

2

Counting valence electrons without including the charge of an ion.

3

Using total bonding electrons B instead of ½B in the formal charge formula.

4

Choosing the structure with the largest charge separation instead of the smallest ΔFC.

1

Saying a double bond is two σ bonds — it is one σ and one π.

2

Saying π bonds form by head-on overlap — they form by sideways overlap.

3

Counting the C–H bonds as π bonds — all single bonds are σ.

4

Saying π bonds are stronger than σ bonds — σ bonds are stronger.

1

Giving the hybridisation of a double-bonded carbon as sp³ — it is sp².

2

Forgetting that lone pairs count as electron domains when assigning hybridisation.

3

Counting the unhybridised p orbital as part of the hybrid set — it forms the π bond.

4

Saying sp³ orbitals are spherical — they have one large lobe.

1

Describing a triple bond as three σ bonds.

2

Giving the ethene carbon bond angle as 109.5° instead of about 120°.

3

Forgetting that the π bond is formed from unhybridised p orbitals, not hybrid orbitals.

4

Saying ethyne carbons are sp² — they are sp.