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

Section 1 of 47Covalent bonding

Covalent bond

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

Occurrence

• Occurrence between two non-metals, with electrons shared rather than transferred.

Orbital overlap

• Orbital overlap two atomic orbitals overlap to form a molecular orbital, and sharing is more stable than holding electrons on one atom.

Octet rule

• Octet rule atoms share electrons to gain a noble-gas valence shell of eight electrons.
  • (a duet for hydrogen)

Exceptions

• Be in BeCl₂ has 4 electrons around it and B in BF₃ has 6, and these molecules are stable.

Multiple bonds

• Multiple bonds single, double and triple bonds share 2, 4 and 6 electrons, and quadruple bonds are impossible because repulsion is too great.

Section 2 of 47Drawing Lewis formulas

1

Count the total valence electrons, adding one per negative charge and subtracting one per positive charge.

2

Draw the skeleton with the least electronegative atom central (never hydrogen), joined by single bonds.

3

Complete the octets of outer atoms with lone pairs, then place any remaining electrons on the central atom.

4

If the central atom lacks an octet, convert lone pairs on outer atoms into double or triple bonds.

5

Check that all valence electrons are used and every atom has an octet (a duet for H).

Lewis formula

• Lewis formula shows all valence electrons as bonding pairs and lone pairs, so Cl–Cl alone is not a Lewis formula.

Representations

• Representations a pair can be dots, crosses, a mixture of both, or a line.

Section 3 of 47Bond length and strength; coordinate bonds

Bond strength

• Bond strength triple > double > single, because more shared electrons pull the nuclei closer.

Bond length

• Bond length single > double > triple, e.g. C–C 154 pm, C=C 134 pm and C≡C 120 pm.

Bond energy

• Bond energy the energy needed to break one mole of a bond in the gaseous state, in kJ mol⁻¹.

Coordinate bond

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

Formation

• Formation a lone pair is donated into an empty orbital, e.g. NH₃ + H⁺ → NH₄⁺.
• Identical bonds — once formed it is the same as other covalent bonds, so all four N–H bonds in NH₄⁺ are identical.
• Electron-deficient acceptor — BF₃ can accept a lone pair, e.g. forming F₃B–NH₃.

Section 4 of 47VSEPR shapes

Principle

• Principle electron domains repel and spread as far apart as possible.
  • (bonding pairs, lone pairs, each multiple bond counting as one)

Repulsion order

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

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

1Domain geometry versus molecular shape

• Domain geometry versus molecular shape they can differ, e.g. NH₃ has tetrahedral domains but a trigonal pyramidal shape.

2Lone pairs

• Lone pairs each reduces the bond angle by about 2–2.5°.

Section 5 of 47Bond polarity and molecular polarity

1Electronegativity

Electronegativity the ability of an atom to attract the shared pair in a bond, on the Pauling scale (F is highest at 4.0).

2Polar bond

• Polar bond unequal sharing gives δ+ on the less electronegative atom and δ− on the other, and a larger difference means a more polar bond.

3Non-polar bond

• Non-polar bond identical atoms such as H₂ and Cl₂ share equally.

4Non-polar molecule with polar bonds

• Non-polar molecule with polar bonds symmetrical geometry cancels the dipoles, as in CO₂ and CCl₄.

5Polar molecule

• Polar molecule asymmetrical geometry leaves a net dipole, as in H₂O, NH₃ and CHCl₃.

6Rule

• Rule a molecule with only non-polar bonds cannot be polar.

Section 6 of 47Giant covalent structures

Substance	Structure	Conductivity	Properties
Diamond	Each C bonded tetrahedrally to 4 others	None	Very hard, very high melting point
Graphite	Each C bonded to 3 in hexagonal layers; layers held by London forces	Conducts along layers (one delocalised electron per atom)	Soft and slippery
Graphene	A single layer of graphite	Excellent	Strong, flexible and one atom thick
C60 fullerene	Discrete spherical molecules with 20 hexagons and 12 pentagons	Poor (semiconductor)	Low melting point; not a giant structure
Silicon	Diamond-like network	Poor	High melting point
Silicon dioxide	Each Si bonded to 4 O and each O to 2 Si, formula SiO2	None	High melting point, hard, insoluble

• Melting points
  • covalent bonds, not intermolecular forces, must break in network solids.

Section 7 of 47Intermolecular forces

1Intermolecular forces

• Intermolecular forces act between molecules and are much weaker than covalent bonds, which are intramolecular.

2London (dispersion) forces

• London (dispersion) forces arise from instantaneous dipoles inducing dipoles in neighbours, are present in all substances and grow with more electrons and greater surface area.

3Dipole-induced dipole

• Dipole-induced dipole a permanent dipole distorts the electron cloud of a non-polar neighbour.

4Dipole-dipole

• Dipole-dipole attraction between the δ+ and δ− ends of polar molecules.

5Hydrogen bonding

• H bonded to N, O or F is attracted to a lone pair on N, O or F in a neighbouring molecule, the strongest of these forces.

6Strength order

• London < dipole-dipole < hydrogen bonding for similar molar mass.

7Branching

• Branching straight-chain alkanes have higher boiling points than branched isomers.

8Hydrogen bonding examples

• Hydrogen bonding examples the anomalously high boiling points of H₂O, HF and NH₃, and ice being less dense than water.

Section 8 of 47Properties of molecular substances

Properties of molecular substances

Melting and boiling points

• Melting and boiling points depend on the strength of intermolecular forces, and low boiling points mean volatile.

Solubility

• "like dissolves like", so polar and hydrogen-bonding solutes dissolve in water and non-polar solutes in non-polar solvents such as hexane.

Conductivity

• Conductivity molecular substances have no mobile charged particles, so they do not conduct, except graphite and graphene.

Polar molecules that ionise

• Polar molecules that ionise such as HCl conduct in solution because ions form.

Small alcohols

• Small alcohols ethanol dissolves in water, but solubility falls as the hydrocarbon chain lengthens.

Section 9 of 47Paper chromatography and R_f

1

Draw a pencil baseline about 1 cm from the bottom of the paper.

2

Spot the samples on the baseline and let them dry.

3

Stand the paper in solvent with the level below the baseline and cover to saturate the atmosphere.

4

Remove the paper when the solvent is near the top and mark the solvent front immediately.

5

Locate colourless spots with a UV lamp or a locating agent such as ninhydrin.

6

Measure distances from the baseline and calculate R_f.

Rf

Rf distance moved by the spot ÷ distance moved by the solvent front, both from the baseline.

Properties of Rf

• Properties of Rf it has no unit and lies between 0 and 1.

Interpretation

• Interpretation components more soluble in the solvent and less attracted to the stationary phase travel further and have a higher R_f;
  • compare with references run under identical conditions.

Section 10 of 47More from the specification

More from the specification

Covalent bond nature

• Electrostatic attraction between a shared pair of valence electrons and two positively charged atomic nuclei.

Polar vs nonpolar molecules

• Symmetrical charge distribution produces nonpolar molecules with dispersion forces only;
  • asymmetrical distribution yields permanent dipole moments.

Electrical conductivity limitations

• Simple covalent molecular substances do not conduct electricity in any state because they lack mobile charged particles.
  • (ions or delocalised electrons)

Physical properties of simple molecules

• Low melting and boiling points due to weak intermolecular forces operating between intact covalent molecules.

Physical property

• A property observed or measured without changing the material.

Section 11 of 47Single, double and triple covalent bonds

• Sharing
  • non-metal atoms can share one, two or three pairs, so both atoms reach a noble-gas configuration and gain stability.

Bond	Shared pairs	Shared electrons	Length	Strength
Single	1	2	Longest	Weakest
Double	2	4	Intermediate	Intermediate
Triple	3	6	Shortest	Strongest

Quadruple bonds

• Quadruple bonds are not possible because repulsion between eight electrons is too great.

Reason

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

Carbon–carbon lengths

• C–C about 154 pm, C=C about 134 pm and C≡C about 120 pm, with bond enthalpies rising in the same order.

Section 12 of 47Coordinate (dative) bonds

Definition

Definition a covalent bond in which both electrons of the shared pair come from the same atom, which has a lone pair.

Acceptor

• Acceptor the other atom is electron-deficient with an empty orbital, such as H⁺.

Ammonium ion

• NH₃ donates its lone pair to H⁺ to form NH₄⁺, and similarly H₃O⁺ forms from H₂O.

After formation

• After formation a coordinate bond is identical to any other covalent bond.

Complex ions

• [Al(H₂O)₆]³⁺ has six coordinate bonds from the oxygen of each water, which is higher-level content, so SL needs only the idea of a donated lone pair.

Section 13 of 47VSEPR theory

Principle

• Principle bonding and non-bonding electron pairs behave like negative charge clouds that repel and spread as far apart as possible.
  • (electron domains)

Electron domain

• Electron domain each bonding pair, lone pair or multiple bond counts as one domain.

Multiple bonds

• Multiple bonds are treated as a single domain.

Repulsion order

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

Lone pairs

• Lone pairs are held closer to the nucleus, repel more strongly and reduce bond angles.

Naming the shape

• Naming the shape name the molecular shape from the positions of the atoms only, ignoring lone pairs.

Section 14 of 47Shapes with two and three electron domains

Shapes with two and three electron domains

Two domains

• Two domains arrange on opposite sides at 180°, giving linear molecules such as BeCl₂, CO₂ and ethyne.

Three domains, no lone pair

• Three domains, no lone pair trigonal planar at 120°, as in BF₃, ethene and methanal.

Three domains, one lone pair

• Three domains, one lone pair bent (V-shaped), with the angle reduced to below 120° (about 118°).

SO₂

• SO₂ has two bonding domains and one lone pair on sulfur, so it is bent, and its sulfur has 10 valence electrons.

Accepted names for bent

• Accepted names for bent non-linear, angular and V-shaped.

Section 15 of 47Shapes with four electron domains

• Four domains
  • arrange tetrahedrally with an ideal angle of 109.5°.

Bonding pairs	Lone pairs	Shape	Angle	Example
4	0	Tetrahedral	109.5°	CH4, NH4+
3	1	Trigonal pyramidal	about 107°	NH3
2	2	Bent	about 104.5°	H2O

• Angle reduction
  • one lone pair reduces the angle slightly, and two lone pairs reduce it further because of lone pair–lone pair repulsion.

Section 16 of 47Predicting shape step by step

1

Draw the Lewis structure and count the electron domains around the central atom.

2

Arrange the domains as far apart as possible: 2 linear, 3 trigonal planar, 4 tetrahedral.

3

Name the molecular shape from the atoms only.

4

Reduce the angles slightly for each lone pair.

Species	Domains	Lone pairs	Shape	Angle
CO2	2	0	Linear	180°
BF3	3	0	Trigonal planar	120°
CH4	4	0	Tetrahedral	109.5°
NH3	4	1	Trigonal pyramidal	107°
H2O	4	2	Bent	104.5°

1Domain geometry versus molecular geometry

• Domain geometry versus molecular geometry they are the same when there are no lone pairs (CH₄) and differ when there are (NH₃).

2Hidden lone pairs

• Hidden lone pairs always draw the Lewis structure first so none are missed.

Section 17 of 47Resonance and delocalisation

Resonance

• Resonance where more than one valid Lewis structure can be drawn, as in ozone, carbonate and benzene, the real structure is a resonance hybrid.

Bond lengths

• Bond lengths in a hybrid they lie between those of single and double bonds because the electrons are delocalised.

Carbonate

• Carbonate all three C–O bonds are equal in length.

Section 18 of 47More from the specification

1

Molecular Polarity and Permanent Dipoles

• A covalent molecule is polar if bond dipoles do not cancel due to asymmetrical geometry,
• whereas symmetrical molecules are nonpolar because opposing dipoles cancel out completely.

2

London Dispersion Forces

• London dispersion forces are weak, temporary intermolecular forces present between all covalent substances, arising from instantaneous fluctuations in electron density that induce dipoles in neighboring atoms.

3

Hydrogen Bonding in Polar Solvents

• Hydrogen bonding is a strong specific intermolecular dipole-dipole attraction occurring when hydrogen is covalently bonded to highly electronegative nitrogen, oxygen,
• or fluorine, enabling polar substances to dissolve in a polar solvent like water.

4

Giant Covalent Network Substances

• In contrast to simple molecular substances held by weak intermolecular forces, network solids like silicon and silicon dioxide form rigid tetrahedral covalent lattices with extremely high melting temperatures.

5

Dispersion force

• Another name for induced dipole–induced dipole force.

6

Intermolecular forces

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

7

Giant covalent

• Network of atoms joined by covalent bonds.

8

Lewis formula

• A diagram showing all bonding pairs and lone pairs.

9

Stationary phase

• The solid or liquid phase fixed in place in chromatography.

10

Physical property

• A property observed or measured without changing the material.

11

Dipole

• A separation of charge in a bond or molecule.

12

Hydrogen

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

13

Substance

• A material with only one type of particle.

14

Solvent

• Substance that a solute dissolves in.

15

Boiling

• Liquid to gas throughout the liquid at the boiling point.

16

Chromatography

• Method separating substances by their different solubilities in a solvent.

Section 19 of 47Electronegativity and bond polarity

Electronegativity

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

Scale

• Pauling values are in Section 9 of the IB Data Booklet, and a higher value means a more electronegative element.

Non-polar bond

• Non-polar bond identical atoms, as in Cl₂, share the electrons equally.

Polar bond

• Polar bond different electronegativities draw the electrons towards the more electronegative atom, which becomes δ− while the other becomes δ+.

Charge centres

• Charge centres the centres of positive and negative charge no longer coincide.

Difference and character

• Difference and character a difference of about 0.4 or less is mostly covalent, intermediate values are polar covalent and above about 1.8 gives ionic character.

Section 20 of 47Dipoles

Dipole

Dipole a separation of charge across a bond.

Dipole moment

• Dipole moment measures how polar a bond is and increases with the electronegativity difference.

Arrow

• Arrow the dipole arrow has a crossed tail at δ+ and points towards the more electronegative atom.

Example

• Example in HCl, chlorine is more electronegative, so Cl is δ− and H is δ+.

Section 21 of 47Polarity of molecules

Polarity of molecules

Two factors

• Two factors polarity depends on the polarity of each bond and the arrangement of the bonds in space.

Only non-polar bonds

• Only non-polar bonds the molecule is non-polar, e.g. H₂ and Cl₂.

Polar bonds, symmetrical

• Polar bonds, symmetrical dipoles cancel and the molecule is non-polar, as in linear CO₂ and tetrahedral CCl₄.

Polar bonds, asymmetrical

• Polar bonds, asymmetrical dipoles do not cancel, leaving a net dipole, as in H₂O, NH₃, CH₃Cl and CHCl₃.

Rule

• Rule a molecule is polar only if it has polar bonds and its shape does not cancel the dipoles.

Section 22 of 47Giant covalent (network) structures

Giant covalent (network) structures

Structure

• Structure large networks of atoms joined by covalent bonds throughout the lattice, with no discrete molecules.

Silicon

• Silicon each atom forms four covalent bonds in a tetrahedral arrangement with angles of about 109.5°.

Silicon(IV) oxide

• Silicon(IV) oxide each Si bonds to four O and each O to two Si in a tetrahedral network,
  • and the empirical formula SiO₂ shows the ratio, not a molecule.

Melting points

• Melting points very high because many strong covalent bonds must break.

Solubility

• Solubility insoluble in water and most solvents.

Section 23 of 47Diamond and graphite

Diamond

• Diamond each carbon forms four covalent bonds tetrahedrally, giving a rigid, very hard, 3D network.

Conductivity

• Conductivity none, because all four outer electrons of each carbon are used in bonding.

Graphite

• Graphite each carbon forms three bonds in hexagonal layers with angles of 120°.

Conductivity

• Conductivity the fourth electron is delocalised between the layers, so graphite conducts.
• Softness — layers are held by weak intermolecular forces and slide over each other, so it is soft and slippery.

Section 24 of 47Graphene and fullerene

Graphene and fullerene

Graphene

• Graphene a single layer of carbon atoms in a hexagonal lattice, one atom thick, strong and flexible, and it conducts through delocalised electrons.

Buckminsterfullerene (C60)

• Buckminsterfullerene (C60) a molecular form of carbon with 60 atoms in a sphere of 20 hexagons and 12 pentagons, each carbon bonded to three others.

C60 conductivity

• C60 conductivity the delocalised fourth electrons allow limited movement, so it is a poor conductor or semiconductor.

Not giant

• C₆₀ has a fixed formula and is a molecule, so its melting point is low.

Allotropes

• Allotropes different structural forms of the same element, such as diamond, graphite, graphene and C₆₀.

Section 25 of 47More from the specification

1

Electron Domain Geometry and Trigonal Arrangements

• A central atom surrounded by three electron domains adopts a trigonal planar electron domain geometry with 120° bond angles,
• whereas three bonding pairs and one lone pair adopt a trigonal pyramidal molecular geometry.

2

Intermolecular Forces and Boiling Points

• Covalent substances exhibit boiling points governed by intermolecular forces: nonpolar molecules rely solely on weak London dispersion forces,
• while polar molecules with hydrogen bonding possess significantly higher boiling points.

3

Molecular Polarity and Symmetrical Domains

• Molecules with symmetrical arrangements of identical polar bonds are nonpolar overall because individual bond dipoles cancel out completely.
  • (such as trigonal planar BF₃)

4

Separation of Substances by Chromatography

• In paper and thin-layer chromatography, components of a mixture are separated based on their relative affinities for the polar stationary phase and nonpolar mobile solvent, governed by intermolecular forces.

5

Dispersion force

• Another name for induced dipole–induced dipole force.

6

Lewis formula

• A diagram showing all bonding pairs and lone pairs.

7

Electron domain

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

8

Stationary phase

• The solid or liquid phase fixed in place in chromatography.

9

Physical property

• A property observed or measured without changing the material.

10

Hydrogen

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

11

Substance

• A material with only one type of particle.

12

Boiling

• Liquid to gas throughout the liquid at the boiling point.

13

Chromatography

• Method separating substances by their different solubilities in a solvent.

14

Bonding pair

• A shared pair of electrons forming a covalent bond.

Section 26 of 47Conductivity of covalent solids

Conductivity of covalent solids

Graphene

• Graphene a single layer of carbon atoms in a hexagonal lattice, each bonded to three others at 120°, conducts because the fourth outer electron of each carbon is delocalised.

Graphite

• Graphite conducts for the same reason, through delocalised electrons.

Most network solids

• Most network solids do not conduct because all electrons are used in bonding.

Buckminsterfullerene (C60)

• Buckminsterfullerene (C60) is a semiconductor.

Solubility

• Solubility most covalent network structures are insoluble in water.

Section 27 of 47Intermolecular forces overview

Intermolecular forces overview

Molecular substances

• Molecular substances have covalent bonds within molecules but only intermolecular forces between them.

Strength

• Strength these forces are much weaker than covalent or ionic bonds.

Properties controlled

• Properties controlled melting and boiling point, volatility and solubility.

Types

• London (dispersion), dipole-induced dipole, dipole-dipole and hydrogen bonding.

Van der Waals forces

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

Intramolecular forces

• Intramolecular forces covalent bonds within a molecule, not intermolecular forces.

Section 28 of 47London (dispersion) forces

Origin

• Origin electrons move constantly,
• so the distribution is momentarily uneven,
• giving a temporary dipole that induces a dipole in a neighbour,
• leading to attraction.

Occurrence

• Occurrence present between all atoms and molecules and the only intermolecular force in non-polar substances.

Strength

• Strength typically 1–50 kJ mol⁻¹.

Number of electrons

• Number of electrons more electrons make a larger, more polarisable cloud, so the forces and boiling points increase, as down group 17 from F₂ to I₂.

Surface area

• Surface area a long straight-chain isomer has more contact area and a higher boiling point than a compact branched isomer with the same number of electrons.

Terminology

• Terminology always call these London (dispersion) forces.

Section 29 of 47Dipole-dipole and dipole-induced dipole forces

Dipole-dipole

• Dipole-dipole molecules with a permanent dipole, from an electronegativity difference and an asymmetric shape, attract through opposite ends.

Effect

• Effect polar compounds usually have higher boiling points than similar-sized non-polar molecules,
• because dipole-dipole forces act in addition to London forces.

Comparison

• Comparison butane and propanone have the same number of electrons, but only propanone has a permanent dipole, so it has the higher boiling point.

Dipole-induced dipole

• Dipole-induced dipole a polar molecule such as HCl distorts the electron cloud of a non-polar molecule such as Cl₂, creating a weak attraction.

Section 30 of 47Hydrogen bonding

Definition

Definition an attraction between a hydrogen atom bonded to N, O or F and a lone pair on an N, O or F atom of a neighbouring molecule.

Strength

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

Requirements

• H attached directly to N, O or F, and a lone pair on a neighbouring N, O or F.

Not a covalent bond

• Not a covalent bond it is an intermolecular attraction, shown by dotted lines.

Section 31 of 47Ranking intermolecular forces

Force	Present in	Relative strength
London (dispersion)	All substances	Weakest
Dipole-dipole	Polar molecules	Intermediate
Hydrogen bonding	Molecules with N–H, O–H or F–H	Strongest

Melting and boiling

• Melting and boiling only intermolecular forces are overcome, and covalent bonds stay intact.

Polarisability

• London forces increase with the number of electrons because a larger cloud is more easily distorted.

Section 32 of 47More from the specification

1

Electron domains and tetrahedral geometry

• In molecules with four shared electron pairs around a central atom, the four electron domains adopt a tetrahedral geometry with bond angles of 109.5 degrees.
  • (such as methane or silicon tetrachloride, SiCl₄)

2

Molecular symmetry and nonpolar molecules

• When identical terminal atoms symmetrically surround a central atom in tetrahedral geometry, individual bond dipoles cancel,
• resulting in a nonpolar molecule.

3

Giant covalent structures of silicon

• Silicon and silicon dioxide form tetrahedral network covalent lattices with shared electron pairs extending across three dimensions,
• resulting in high melting points and insolubility.

4

Solvent interactions and chromatography

• In paper chromatography, separation of solutes depends on differential partitioning between the polar cellulose stationary phase and the nonpolar organic mobile phase solvent.

5

Electron domains and tetrahedral geometry

• In molecules with four shared electron pairs around a central atom, the four electron domains adopt a tetrahedral geometry with bond angles of 109.5 degrees.
  • (such as methane or silicon tetrachloride, SiCl₄)

6

Molecular symmetry and nonpolar molecules

• When identical terminal atoms symmetrically surround a central atom in tetrahedral geometry, individual bond dipoles cancel,
• resulting in a nonpolar molecule.

7

Giant covalent structures of silicon

• Silicon and silicon dioxide form tetrahedral network covalent lattices with shared electron pairs extending across three dimensions,
• resulting in high melting points and insolubility.

8

Solvent interactions and chromatography

• In paper chromatography, separation of solutes depends on differential partitioning between the polar cellulose stationary phase and the nonpolar organic mobile phase solvent.

9

Lewis formula

• A diagram showing all bonding pairs and lone pairs.

10

Electron domain

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

11

Stationary phase

• The solid or liquid phase fixed in place in chromatography.

12

Physical property

• A property observed or measured without changing the material.

13

Solvent

• Substance that a solute dissolves in.

14

Tetrahedral

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

15

Chromatography

• Method separating substances by their different solubilities in a solvent.

16

Bonding pair

• A shared pair of electrons forming a covalent bond.

17

Giant covalent

• Network of atoms joined by covalent bonds.

Section 33 of 47Hydrogen bonding in ammonia and water

Ammonia

• Ammonia each molecule forms on average one hydrogen bond per molecule, because it has one N–H hydrogen set and one lone pair on nitrogen.

Water

• Water each molecule can form two hydrogen bonds, from two H atoms on oxygen and two lone pairs on oxygen.

Result

• Result extensive hydrogen bonding explains the unusually high boiling point of water.

Section 34 of 47Ranking boiling points

1

Compare molar masses or numbers of electrons to judge the London forces.

2

Decide whether each substance is polar.

3

Decide whether it has H bonded to N, O or F.

4

Rank by the strongest type of force present, using London forces to settle ties.

Similar molar mass

• Similar molar mass propan-1-ol (M = 60) has hydrogen bonding, propanone (M = 58) is polar with dipole-dipole forces,
  • and butane (M = 58) has only London forces.

Order

• Order butane < propanone < propan-1-ol.

Hydrogen bonding effect

• Hydrogen bonding effect ethanol boils higher than propane although their molar masses are similar.

Common error

• Common error basing boiling points on molar mass alone and ignoring hydrogen bonding.

Section 35 of 47Solubility

Solubility

Rule

• Rule like dissolves like.

Polar and hydrogen-bonding solutes

• Polar and hydrogen-bonding solutes such as ethanol and sugar dissolve in water because they form new hydrogen bonds with water.

Small alcohols

• Small alcohols ethanol is highly soluble because its –OH group forms hydrogen bonds with water.

Longer chains

• Longer chains hexanol is much less soluble because its long non-polar chain disrupts the hydrogen bonding of water.

Non-polar substances

• Non-polar substances such as hexane dissolve in non-polar solvents through London forces between solute and solvent.

Polar substances in non-polar solvents

• Polar substances in non-polar solvents dissolve poorly because their dipoles cannot interact effectively.

Giant covalent substances

• Giant covalent substances are insoluble in polar and non-polar solvents because the covalent bonds are too strong to break.

Section 36 of 47Properties of covalent substances compared

Property	Non-polar molecular	Polar molecular	Giant covalent
Melting and boiling point	Low	Low to moderate	Very high
Volatility	High	Moderate to high	Low
Solubility in polar solvents	Insoluble	Sometimes soluble	Insoluble
Solubility in non-polar solvents	Soluble	Sometimes soluble	Insoluble
Electrical conductivity	None	None (except those that ionise in solution)	None (except graphite and graphene)

• Ionising polar molecules
  • HCl forms H⁺ and Cl⁻ in water and conducts in solution.

Section 37 of 47Principle of chromatography

1Purpose

• Purpose separates and analyses the components of a mixture.

2Phases

• Phases a liquid mobile phase moves over a solid stationary phase.

3Rate

• Rate components move at different rates depending on their solubility in the mobile phase and their attraction to the stationary phase.

4Distance

• Distance substances more soluble in the solvent travel further, and non-polar substances are more soluble in a non-polar solvent.

5Polar component

• Polar component is held more strongly by a polar stationary phase and moves a shorter distance.

Section 38 of 47Paper chromatography method

• Stationary phase
  • paper made of cellulose, whose polar –OH groups form hydrogen bonds with polar molecules.

1

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

2

Place small spots of the sample on the baseline.

3

Suspend the paper in solvent with the baseline above the solvent level.

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 read the chromatogram.

Paper chromatography method

Insoluble component

• Insoluble component if a component does not travel with the solvent, it is insoluble in it, so another solvent should be tried.

Several solvents

• Several solvents may be needed for clear separation of all the components.

Section 39 of 47More from the specification

More from the specification

VSEPR Theory and Electron Domains

• According to valence shell electron pair repulsion (VSEPR) theory, electron domains around a central atom repel each other to adopt an electron domain geometry that minimizes electrostatic repulsion.

Tetrahedral Geometry of Group 14 Elements

• When a central carbon or silicon atom forms four single covalent bonds, the four shared electron pairs orient towards the corners of a regular tetrahedron with bond angles of 109.5°.

Molecular Shape and Symmetry in Nonpolar Compounds

• Symmetrical tetrahedral molecules,
• such as methane (CH₄) or tetrachloromethane (CCl₄), have a nonpolar overall structure because identical bond polarities cancel out symmetrically.

Dispersion Forces in Nonpolar Hydrocarbons

• Nonpolar covalent substances composed of carbon and hydrogen rely solely on weak London dispersion forces between molecules, with intermolecular strength increasing as total electron count increases.

Dispersion force

• Another name for induced dipole–induced dipole force.

Lewis formula

• A diagram showing all bonding pairs and lone pairs.

Electron domain

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

Physical property

• A property observed or measured without changing the material.

Tetrahedral

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

Section 40 of 47Principle of chromatography

Distribution

• Distribution each component distributes between a stationary phase and a mobile phase.

Movement

• Movement components more soluble in the mobile phase and less attracted to the stationary phase move further up the plate.

Basis

• Basis separation depends on differences in intermolecular forces between the sample and each phase.

Course focus

• Course focus the IB course covers paper chromatography and TLC, and column and gas chromatography work on similar principles but are not required.

Section 41 of 47Thin layer chromatography (TLC)

Stationary phase

• Stationary phase a thin layer of an inert solid such as silica or alumina on a glass or plastic plate.

Surface

• Surface the surface has –OH groups that form hydrogen bonds with components, which are adsorbed to different extents.

Mobile phase

• Mobile phase a solvent, as in paper chromatography.

Advantages

• Advantages faster than paper chromatography, gives better separation and uses very small samples.

Section 42 of 47TLC method

1

Draw a pencil baseline near the bottom of the plate, since ink would dissolve and run.

2

Apply small spots of the samples and known references on the baseline.

3

Stand the plate in a covered container with solvent below the baseline, so the air is saturated with vapour.

4

Let the solvent rise until near the top and mark the solvent front immediately.

5

Dry the plate.

6

Make colourless spots visible with UV light or a locating agent such as iodine.

7

Measure the distances moved by the spots and by the solvent front.

Section 43 of 47Calculating Rf

Formula

R_f = distance moved by the spot ÷ distance moved by the solvent front.

Measurement

• Measurement both distances are measured from the baseline, to the centre of the spot.

Form

• R_f values are quoted as decimals.

Units

• Units there are none, because the lengths are in the same unit and cancel.

Range

• Range always less than 1, because a spot cannot move further than the solvent.

Section 44 of 47Using Rf values

Using Rf values

Identification

• Identification a spot from an unknown mixture with the same R_f as a reference, run under identical conditions, probably contains that substance.

Number of components

• Number of components the number of spots shows the number of components.

Conditions

• R_f depends on the solvent and stationary phase, which must be stated, so values from different solvents cannot be compared.

Confidence

• Confidence agreement in several different solvents makes the identification stronger.

Observation	Interpretation
Spot stays on the baseline	Insoluble in that solvent; try another solvent
High Rf (near 1)	Very soluble in the mobile phase
Same Rf as a standard	Likely the same compound

Section 45 of 47More from the specification

More from the specification

Nature of Covalent Bonding

• A covalent bond consists of the electrostatic attraction between positively charged nuclei and shared pairs of valence electrons localized between the bonded atoms.

Electronegativity Differences and Bond Dipoles

• When two bonded atoms possess differing electronegativities, bonding electrons are pulled closer to the more electronegative atom, generating a permanent bond dipole with partial charges (δ+ and δ-).

Molecular Polarity

• Polar vs. Nonpolar Molecules — A molecule containing polar covalent bonds is polar if its geometric structure is asymmetrical,
• while symmetrical molecules are nonpolar because opposing dipoles cancel out.
  • (such as tetrahedral carbon tetrachloride, CCl₄)

Octet Rule and Electron Sharing in Carbon Compounds

• In organic carbon compounds, a carbon atom shares four pairs of electrons to achieve a stable octet configuration,
• forming four single covalent bonds, double bonds,
• or triple bonds.

Dispersion force

• Another name for induced dipole–induced dipole force.

Boiling

• Liquid to gas throughout the liquid at the boiling point.

Covalent bond

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

Bonding pair

• A shared pair of electrons forming a covalent bond.

Valence electrons

• Electrons in the outer shell.

Giant covalent

• Network of atoms joined by covalent bonds.

Lewis formula

• A diagram showing all bonding pairs and lone pairs.

Electron domain

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

Physical property

• A property observed or measured without changing the material.

Electron

• Negative particle moving around the nucleus.

Dipole

• A separation of charge in a bond or molecule.

Bonding

• The joining of atoms.

Tetrahedral

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

Section 46 of 47Key terms

Term	Meaning
Covalent bond	The attraction between a shared pair of electrons and the two nuclei.
Coordinate bond	A covalent bond in which both electrons come from the same atom.
VSEPR theory	Electron domains around a central atom repel to be as far apart as possible, with lone pairs repelling most.
Electronegativity	The ability of an atom to attract the shared pair in a covalent bond.
Allotropes	Different structural forms of the same element in the same state.
London (dispersion) forces	Weak attractions between instantaneous and induced dipoles.
Hydrogen bond	Attraction between H bonded to N, O or F and a lone pair on N, O or F of another molecule.

Term	Meaning
Covalent bond	A shared pair of electrons between two atoms.
Coordinate bond	A covalent bond in which both electrons come from one atom.
Electron domain	A region of electron density: a bonding pair, lone pair or multiple bond.
VSEPR	Valence shell electron pair repulsion theory for predicting shapes.
Lone pair	A pair of valence electrons not involved in bonding.
Bond angle	The angle between two bonds at a central atom.

Term	Meaning
Electronegativity	The ability of an atom to attract the shared electron pair in a bond.
Polar covalent bond	A covalent bond with unequal sharing of electrons.
Dipole	A separation of positive and negative charge.
Covalent network structure	A giant lattice of atoms joined by covalent bonds.
Allotrope	A different structural form of the same element.
Delocalised electron	An electron free to move through a structure.

Term	Meaning
Intermolecular force	An attraction between molecules, weaker than a covalent bond.
London (dispersion) force	An attraction between temporary and induced dipoles.
Dipole-dipole force	An attraction between permanent dipoles in polar molecules.
Hydrogen bond	An attraction between H on N, O or F and a lone pair on N, O or F of another molecule.
Graphene	A single layer of carbon atoms in a hexagonal lattice.
Polarisability	How easily an electron cloud is distorted.

Term	Meaning
Hydrogen bond	An attraction between H on N, O or F and a lone pair on another N, O or F.
Volatility	The tendency of a substance to evaporate easily.
Mobile phase	The solvent that moves through a chromatography system.
Stationary phase	The solid, such as paper, that stays fixed.
Solubility	The ability of a substance to dissolve in a solvent.
Non-polar	Having no net dipole.

Term	Meaning
TLC	Thin layer chromatography using a thin adsorbent layer on a plate.
Rf	The distance moved by the spot divided by the distance moved by the solvent front.
Solvent front	The furthest point the solvent reaches.
Baseline	The line on which the sample is spotted.
Mobile phase	The solvent that carries the components.
Stationary phase	The adsorbent layer that stays fixed.

Section 47 of 47Exam technique

1

Saying covalent molecules melt by breaking covalent bonds — only intermolecular forces are overcome (except in network solids).

2

Assuming polar bonds always make a polar molecule — symmetrical shapes such as CO₂ cancel the dipoles.

3

Forgetting lone pairs when predicting shape — draw the Lewis formula first.

4

Quoting R_f with units or values above 1 — it is a ratio less than 1.

1

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

2

Naming the shape from the domain geometry — name it from the atoms only.

3

Giving 109.5° for NH₃ and H₂O — lone pairs reduce the angle.

4

Thinking a coordinate bond differs from a normal covalent bond once formed — they are identical.

1

Saying any molecule with polar bonds is polar — symmetry may cancel the dipoles.

2

Calling SiO₂ a molecule — it is a giant covalent network.

3

Saying graphite is soft because of weak covalent bonds — the weak forces are between the layers.

4

Forgetting diamond is an insulator — all four outer electrons are used in bonding.

1

Saying intermolecular forces are broken when a covalent substance boils — only intermolecular forces are overcome, covalent bonds stay intact.

2

Stating London forces exist only in non-polar molecules — they exist between all molecules.

3

Ignoring molecular shape when comparing isomers — branching lowers the boiling point.

4

Calling hydrogen bonding a covalent bond — it is an intermolecular force.

1

Basing boiling points on molar mass alone — hydrogen bonding can outweigh the effect of mass.

2

Thinking hydrogen bonds occur in every compound that contains hydrogen — H must be bonded to N, O or F.

3

Using ink for the chromatography baseline — it dissolves and runs, so use pencil.

4

Saying long-chain alcohols dissolve well because they contain OH — the non-polar chain reduces solubility.

1

Measuring distances from the bottom of the plate — measure from the baseline.

2

Quoting R_f with units or as a value greater than 1 — it is a unitless ratio less than 1.

3

Not marking the solvent front before it dries — mark it immediately.

4

Comparing R_f values from different solvents — the solvent changes the values.