Notes · ChemistryUK · A-Levels

Bonding

How atoms hold together decides everything about a substance's properties. This chapter develops the three chemical bonds - ionic, covalent (including dative) and metallic - and the four types of structure they build, then uses electron-pair repulsion to predict the shapes of molecules and ions, and finishes with electronegativity, polarity and the intermolecular forces that explain melting points, solubility and the strange behaviour of water.

5 sections~17 min reading time3 competenciesLevel Foundation 1 · Standard 3 · Advanced 1

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Exam profile

AO1 · Describe ionic, covalent, dative and metallic bonding, the four crystal structures and the three intermolecular forcesAO2 · Predict the shapes and bond angles of molecules and ions using electron-pair repulsion theoryAO3 · Relate melting point, conductivity and solubility to structure and bonding, and evaluate the role of hydrogen bonding

Operators:statedescribeexplainpredictdeducejustify

basic level

AS-Level requires the three bond types, the four structures with their properties, electron-pair repulsion shapes and angles, electronegativity, polarity and the three intermolecular forces.

higher level

The full A-Level applies these ideas synoptically - explaining boiling-point trends, solubility, the shapes of transition-metal complexes and the behaviour of organic molecules - so a secure grasp of structure and bonding underpins much of the rest of the course.

Depth

Text

Contents · 5 sections

Bonding

Ionic bonding and giant ionic lattices#

FoundationAQA 7405 3.1.3.1DfE GCE Chemistry - bonding

Electron transfer forming sodium chloride

Fig. 1Sodium loses an electron to chlorine; the resulting Na+ and Cl- ions attract electrostatically in a giant lattice.

Ionic bonding is the electrostatic attraction between oppositely charged ions, formed when a metal atom transfers one or more electrons to a non-metal atom. Sodium loses its single outer electron to form

\text{Na}^+

and chlorine gains one to form

\text{Cl}^-

; each ion now has a full outer shell (a noble-gas configuration). The bond is not directional - each ion attracts all its oppositely charged neighbours - so the ions build a regular, repeating three-dimensional giant ionic lattice.

The strength of ionic bonding, and hence properties such as melting point, depends on the charges on the ions and their sizes. Higher charges attract more strongly (magnesium oxide, with

2+

and

2-

ions, has a much higher melting point than sodium chloride with

1+

and

1-

), and smaller ions can pack closer, increasing the attraction. This is quantified later as lattice enthalpy, but the qualitative rule 'higher charge and smaller ions means stronger bonding' explains most comparisons at this stage.

The properties of ionic compounds follow directly from the lattice. They have high melting and boiling points because a great deal of energy is needed to overcome the strong electrostatic forces throughout the giant structure. They are hard but brittle, because a blow that shifts one layer brings like charges next to each other, and the repulsion splits the crystal. Many dissolve in water, whose polar molecules surround and stabilise the separated ions.

Electrical conductivity is the diagnostic test. Ionic solids do not conduct because the ions are locked in place, but when molten or dissolved the ions become free to move and carry charge, so the liquid or solution conducts electricity (and is decomposed by it). Being able to explain conductivity 'when molten or aqueous but not when solid' in terms of mobile ions is a standard exam requirement.

Worked example

Explaining a melting-point difference

Explain why magnesium oxide (mp 2852 C) has a much higher melting point than sodium chloride (mp 801 C).

Compare the charges
MgO contains Mg2+ and O2- ions (charges 2+/2-), whereas NaCl contains Na+ and Cl- (charges 1+/1-).
Relate to attraction
The larger charges in MgO create stronger electrostatic attraction between the ions; Mg2+ and O2- are also smaller than Na+ and Cl-, so the ions pack closer.
Link to melting point
Stronger attractions throughout the lattice need more energy to overcome, so MgO melts at a far higher temperature.

Result: The higher ionic charges (and smaller ions) in MgO give much stronger lattice attractions and a higher melting point.

Typical mistakes

Saying an ionic compound conducts when solid - the ions are fixed until it is molten or dissolved.
Describing an ionic 'molecule' - ionic compounds are giant lattices with an empirical formula, not discrete molecules.

Active revision

Magnesium oxide and sodium chloride are both ionic. Explain, using ionic charges and sizes, why magnesium oxide has the higher melting point.

Active recall

Recall the key points — then reveal.

Sources: GCE AS and A level subject content for the sciences (Department for Education) · AQA A-level Chemistry 7405 specification (AQA)

Covalent and dative covalent bonding#

StandardAQA 7405 3.1.3.2DfE GCE Chemistry - covalent bonding

The dative bond in the ammonium ion

Fig. 2The nitrogen lone pair forms a dative bond to H+, giving NH4+ with four identical N-H bonds and an overall 1+ charge.

A covalent bond is a shared pair of electrons, held between two nuclei by the electrostatic attraction of both nuclei for the shared pair. It forms between non-metal atoms, and each shared pair gives both atoms a share in a more stable, usually full, outer shell. Atoms can share more than one pair: oxygen forms a double bond (

\text{O}=\text{O}

, two shared pairs) and nitrogen a triple bond (

\text{N}{\equiv}\text{N}

, three shared pairs), which are correspondingly shorter and stronger than single bonds.

A dative covalent (coordinate) bond is a covalent bond in which both shared electrons come from the same atom - one atom donates a lone pair into an empty orbital of the other. Once formed it is identical to any other covalent bond; only its origin differs. It is drawn as an arrow pointing from the donor to the acceptor. The classic examples are the ammonium ion

\text{NH}_4^+

(the nitrogen lone pair bonds to

\text{H}^+

) and the dimer

\text{Al}_2\text{Cl}_6

(a chlorine lone pair bonds into aluminium's empty orbital).

Covalent bonding produces two very different kinds of structure. Simple molecular substances (such as

\text{Cl}_2

\text{CO}_2

and

\text{H}_2\text{O}

) consist of small molecules with strong covalent bonds inside each molecule but only weak intermolecular forces between molecules, so they have low melting points and do not conduct. Giant covalent (macromolecular) substances (such as diamond, graphite and silicon dioxide) have a continuous network of covalent bonds, so they have very high melting points.

Because it is important not to over-generalise, note that 'covalent means low melting point' is only true for simple molecular substances - it is the weak forces between molecules, not the covalent bonds, that break on melting. In diamond, melting requires breaking strong covalent bonds throughout the giant lattice, which is why it is one of the highest-melting substances known. Always distinguish what is being broken: intermolecular forces in a molecular solid, covalent bonds in a macromolecular one.

Worked example

Identifying a dative bond

Ammonia reacts with a hydrogen ion to form the ammonium ion, NH3 + H+ -> NH4+. Explain how the fourth N-H bond forms and why all four bonds are then identical.

Locate the lone pair
Nitrogen in NH3 has one lone pair of electrons not used in bonding.
Form the dative bond
This lone pair is donated into the empty 1s orbital of H+, forming a dative covalent bond (both electrons from N).
Explain the equivalence
Once formed, the dative bond is an ordinary shared pair, so the ion has four identical N-H bonds and the 1+ charge is spread over the whole ion.

Result: The nitrogen lone pair forms a dative bond to H+; all four N-H bonds are then equivalent.

Typical mistakes

Claiming covalent bonds break when a simple molecular solid melts - only the intermolecular forces do.
Forgetting that a dative bond, once formed, is identical to an ordinary covalent bond.

Active revision

Explain, in terms of what is broken, why silicon dioxide (a giant covalent solid) has a much higher melting point than carbon dioxide (a simple molecular solid).

Active recall

Recall the key points — then reveal.

Sources: AQA A-level Chemistry 7405 specification (AQA)

Metallic bonding and the four structures#

StandardAQA 7405 3.1.3.3DfE GCE Chemistry - metallic bonding and structures

The four types of structure

Fig. 3Melting point, conductivity and solubility form a fingerprint that identifies each structure.

Metallic bonding is the electrostatic attraction between a lattice of positive metal ions and a 'sea' of delocalised outer-shell electrons that are free to move throughout the whole structure. The strength of the bonding increases with the charge on the ion (more delocalised electrons per atom) and with a smaller ionic radius; this is why magnesium, with

2+

ions, is harder and higher-melting than sodium with

1+

ions.

The delocalised electrons explain the characteristic metallic properties. Metals conduct electricity (and heat) in both the solid and liquid states, because the mobile electrons carry charge; they are malleable and ductile, because layers of ions can slide over one another without breaking the bonding (the electron sea simply flows with them); and they have generally high melting points because the electrostatic attraction between ions and electrons is strong.

It is worth holding the four structure types side by side, because a large fraction of structure-and-bonding questions ask you to identify a structure from its properties, or predict properties from a structure. Ionic (giant lattice of ions), giant covalent (network of covalent bonds), simple molecular (small molecules with weak forces between them) and metallic (ions in an electron sea) each have a signature pattern of melting point, conductivity and solubility.

Reading properties as evidence is the skill being tested. A high melting point with conduction only when molten points to ionic; a very high melting point with no conduction (except graphite) points to giant covalent; a low melting point with no conduction points to simple molecular; a high melting point with conduction as a solid points to metallic. Graphite is the informative exception - a giant covalent solid that conducts because each carbon has a delocalised electron in layers that can slide, making it a lubricant too.

Worked example

Deducing structure from properties

A white solid melts at 993 C, does not conduct when solid, conducts when molten, and dissolves in water. Deduce and justify its structure and bonding.

Melting point
The high melting point rules out simple molecular; it suggests a giant structure.
Conductivity
Conducting only when molten (not solid) is the signature of mobile ions freed on melting - characteristic of an ionic lattice.
Solubility
Dissolving in water fits ionic (polar water surrounds the ions), confirming the deduction.

Result: The substance is a giant ionic solid: high melting point, conducts only when molten or aqueous, soluble in water.

Typical mistakes

Saying metals conduct 'because ions move' - it is the delocalised electrons that carry the charge.
Forgetting graphite as the giant covalent solid that conducts electricity.

Active revision

A solid has a very high melting point, does not dissolve in water and conducts electricity as a solid but not when its bonds are considered as localised. Identify and justify the type of structure. (Hint: consider graphite.)

Active recall

Recall the key points — then reveal.

Sources: AQA A-level Chemistry 7405 specification (AQA)

Shapes of molecules and ions (electron-pair repulsion)#

StandardAQA 7405 3.1.3.4DfE GCE Chemistry - shapes of molecules and ions

Tetrahedral methane

Fig. 4Four bonding pairs repel to the corners of a tetrahedron, giving H-C-H angles of 109.5 degrees.

The shape of a simple molecule or ion is set by electron-pair repulsion theory: the pairs of electrons in the outer (valence) shell of the central atom arrange themselves as far apart as possible, because like charges repel, and this minimises the repulsion. You count the number of electron pairs around the central atom (bonding pairs plus lone pairs), and that number fixes the basic geometry: two pairs give linear, three give trigonal planar, four give tetrahedral, five give trigonal bipyramidal and six give octahedral.

Lone pairs change the picture because they repel more strongly than bonding pairs. A lone pair is held closer to the central atom and occupies more space, so the order of repulsion is lone pair-lone pair

>

lone pair-bonding pair

>

bonding pair-bonding pair. Each lone pair therefore squeezes the bonding pairs together, reducing the bond angle by roughly

2.5^{\circ}

per lone pair. Methane (

\text{CH}_4

, four bonding pairs) is tetrahedral at

109.5^{\circ}

; ammonia (

\text{NH}_3

, three bonding and one lone pair) is trigonal pyramidal at

107^{\circ}

; water (

\text{H}_2\text{O}

, two bonding and two lone pairs) is bent at

104.5^{\circ}

To predict a shape, draw the dot-and-cross diagram to count the bonding and lone pairs, arrange them for minimum repulsion, then describe only the arrangement of the atoms (lone pairs are invisible in the named shape). A double or triple bond counts as a single region of electron density for this purpose - carbon dioxide,

\text{O}=\text{C}=\text{O}

, has two bonding regions and is linear at

180^{\circ}

. For ions, adjust the electron count for the charge:

\text{NH}_4^+

has four bonding pairs and is tetrahedral, while

\text{H}_3\text{O}^+

has three bonding and one lone pair and is pyramidal.

The standard angles are worth memorising because they are frequently asked: linear

180^{\circ}

, trigonal planar

120^{\circ}

, tetrahedral

109.5^{\circ}

, trigonal pyramidal

107^{\circ}

, bent (from four pairs)

104.5^{\circ}

, octahedral

90^{\circ}

, and trigonal bipyramidal

120^{\circ}

(equatorial) and

90^{\circ}

(axial). Always justify a shape by referring to the number of bonding and lone pairs and the fact that electron pairs repel to be as far apart as possible.

Shapes from electron-pair repulsion

Fig. 5The number of bonding and lone pairs fixes the shape and bond angle.

Worked example

Shape of the sulfate ion region and of water

Predict and explain the shape and bond angle of a water molecule, H2O.

Count the electron pairs
Oxygen has 6 outer electrons; two form bonds to H, leaving two lone pairs, so there are 4 electron pairs (2 bonding, 2 lone).
Arrange for minimum repulsion
Four pairs point to the corners of a tetrahedron, but only the two O-H bonds are 'seen', so the shape is bent (V-shaped).
Adjust the angle for lone pairs
The two lone pairs repel more strongly than the bonding pairs, squeezing the H-O-H angle from 109.5 down to 104.5.
$109.5^{\circ} - 2 \times 2.5^{\circ} \approx 104.5^{\circ}$

Result: Water is bent with a bond angle of 104.5 degrees, because two lone pairs on oxygen repel the bonding pairs closer together.

Typical mistakes

Including lone pairs when naming the shape - the name describes only the atoms' arrangement.
Forgetting that lone pairs repel more than bonding pairs, so each reduces the bond angle.

Active revision

Predict the shape and bond angle of the $\text{H}_3\text{O}^+$ ion, justifying your answer in terms of electron pairs.

Active recall

Recall the key points — then reveal.

Sources: AQA A-level Chemistry 7405 specification (AQA)

Electronegativity, polarity and intermolecular forces#

AdvancedAQA 7405 3.1.3.5AQA 7405 3.1.3.6DfE GCE Chemistry - polarity and intermolecular forces

Hydrogen bonding between water molecules

Fig. 6A hydrogen bond (dashed) links the delta-positive H of one water molecule to a lone pair on the delta-negative O of another.

Electronegativity is the power of an atom to attract the pair of electrons in a covalent bond towards itself, measured on the Pauling scale (fluorine, the most electronegative, is

4.0

). It increases across a period (rising nuclear charge, similar shielding, so a stronger pull) and up a group (the bonding pair is closer to the nucleus). When two atoms of different electronegativity bond, the shared pair is pulled towards the more electronegative atom, making the bond polar - one end slightly negative (

\delta-

), the other slightly positive (

\delta+

) - a permanent dipole.

A molecule can contain polar bonds yet be non-polar overall if its shape is symmetrical, because the bond dipoles cancel. Carbon dioxide has two polar

\text{C}=\text{O}

bonds, but because it is linear the dipoles point in exactly opposite directions and cancel, so the molecule has no overall dipole. Water has two polar

\text{O-H}

bonds and a bent shape, so the dipoles add to give a net dipole - water is a polar molecule. Deciding molecular polarity therefore requires both the bond polarities and the shape.

Three kinds of force act between molecules. Van der Waals (London, or induced dipole-dipole) forces exist between all molecules: the constant motion of electrons creates instantaneous dipoles that induce dipoles in neighbours; they strengthen with more electrons (larger

M_r

) and with greater surface contact, which is why boiling points rise down a homologous series and are lower for branched isomers. Permanent dipole-dipole forces act between polar molecules in addition to van der Waals forces. Hydrogen bonding, the strongest, occurs when hydrogen is bonded directly to the small, highly electronegative atoms nitrogen, oxygen or fluorine, and is attracted to a lone pair on the N, O or F of a neighbouring molecule.

Hydrogen bonding explains several anomalies that are examination favourites. Water, ammonia and hydrogen fluoride have far higher boiling points than the trend of their groups would predict, because extra energy is needed to break the hydrogen bonds. Ice is less dense than liquid water because each molecule forms hydrogen bonds to four others in an open, tetrahedral lattice that holds the molecules further apart than in the liquid - so ice floats, a property with profound consequences for life. When you explain a physical property, always name the specific intermolecular force being overcome, and never say that covalent bonds break on boiling a molecular liquid.

Worked example

Explaining an anomalous boiling point

Ammonia (NH3) boils at -33 C, whereas phosphine (PH3) boils at -88 C, even though PH3 has more electrons. Explain this.

Identify the forces
PH3 molecules are held only by weak van der Waals and small dipole forces; NH3 molecules can hydrogen bond, because H is bonded to the small, very electronegative N with a lone pair.
Compare their strength
Hydrogen bonds are much stronger than the van der Waals forces in PH3, despite PH3 having more electrons.
Link to boiling point
More energy is needed to overcome the hydrogen bonds in ammonia, so it boils at a higher temperature.

Result: Ammonia boils higher because it forms hydrogen bonds, which are stronger than the van der Waals forces holding phosphine together.

Typical mistakes

Assuming a molecule with polar bonds must be polar overall (symmetry can cancel the dipoles, as in CO2).
Saying covalent bonds break on boiling water - only the hydrogen bonds (and other intermolecular forces) are overcome.

Active revision

Explain why the boiling point of water (100 C) is much higher than that of hydrogen sulfide (-60 C), even though both are Group 6 hydrides.

Active recall

Recall the key points — then reveal.

Sources: AQA A-level Chemistry 7405 specification (AQA)

References & sources

Sources

Department for Education

GCE AS and A level subject content for the sciences

AQA

AQA A-level Chemistry 7405 specification

Bonding

5 sections~17 min reading time3 competenciesLevel Foundation 1 · Standard 3 · Advanced 1

Bonding

Key points

Explaining a melting-point difference

Key points

Identifying a dative bond

Key points

Deducing structure from properties

Key points

Shape of the sulfate ion region and of water

Key points

Explaining an anomalous boiling point

Sources

Bonding

Key points

Explaining a melting-point difference

Key points

Identifying a dative bond

Key points

Deducing structure from properties

Key points

Shape of the sulfate ion region and of water

Key points

Explaining an anomalous boiling point

Sources