Bonding

In a nutshell

Bonding is how atoms and ions are held together, and it is the single best predictor of a substance's physical properties.

There are three types of strong (chemical) bond, ionic, covalent and metallic, plus weaker forces that act between molecules. Which type a substance has, and the structure it forms, decides its melting point, its conductivity and whether it dissolves.

Assumed knowledge: Atomic structure, Amount of substance.

Core content

Ionic bonding

Ionic bonding is the electrostatic attraction between oppositely charged ions, held in a giant lattice.

Ions form by electron transfer: a metal atom loses electrons to form a positive ion, and a non-metal atom gains them to form a negative ion.

You can predict the charge on a simple ion from the element's position in the Periodic Table:

• Group 1 forms 1+, Group 2 forms 2+, Group 3 forms 3+.
• Group 5 forms 3−, Group 6 forms 2−, Group 7 forms 1−.

To construct the formula of an ionic compound, balance the total positive and negative charge so the compound is neutral (see Worked examples).

You must know the formulae of these compound (polyatomic) ions:

Ionic bonding is stronger (and the melting point higher) when the ions are smaller and/or have a higher charge.

For example, magnesium oxide has a higher melting point than sodium chloride: Mg²⁺ and O²⁻ are smaller and more highly charged than Na⁺ and Cl⁻, so the electrostatic attraction between the oppositely charged ions is stronger and more energy is needed to overcome it.

Covalent and dative covalent bonds

A single covalent bond is a shared pair of electrons between two atoms.

• The shared electrons are attracted to the nuclei of both atoms, holding them together.
• Multiple bonds contain more than one shared pair: a double bond is two shared pairs, a triple bond is three.

A co-ordinate (dative covalent) bond is a shared pair of electrons in which both electrons come from the same atom. Once formed, it behaves exactly like an ordinary covalent bond.

You must be able to represent bonds correctly:

• a covalent bond is drawn as a line,
• a co-ordinate bond is drawn as an arrow, pointing from the atom that supplies the lone pair to the atom that is electron deficient.

Common species with a dative bond are the ammonium ion NH₄⁺, the oxonium ion H₃O⁺ and the ammonia-boron trifluoride adduct H₃N→BF₃.

Metallic bonding

Metallic bonding is the electrostatic attraction between the delocalised electrons and the positive metal ions arranged in a giant lattice.

Each metal atom releases its outer-shell electrons into a shared "sea" of delocalised electrons; the atoms become positive ions fixed in a regular lattice.

Three factors make metallic bonding stronger (and the melting point higher):

1. more protons in the nucleus (greater nuclear charge),
2. more delocalised electrons per atom (from the outer shell),
3. a smaller ion.

For example, magnesium has stronger metallic bonding than sodium (and a higher melting point): each Mg atom delocalises two outer electrons rather than one, the Mg²⁺ ion is smaller, and it has one more proton. So there is a stronger attraction between the positive ions and the delocalised electrons.

The four types of crystal structure

Every solid in this topic is one of four crystal types. The spec names six crystals as examples, two of them (diamond and graphite) macromolecular:

A single warning that examiners repeat: use the words molecule and intermolecular forces only for molecular (simple covalent) substances. There are no molecules in sodium chloride, in a metal, or in diamond.

Melting point, conductivity and changes of state

You must be able to relate melting point and conductivity to the structure and bonding.

Two cases worth learning precisely:

• Graphite conducts electricity because each carbon bonds to only three others, leaving one delocalised electron per atom that can move along the layers. Diamond does not conduct because all four outer electrons are used in covalent bonds, so there are no free electrons.
• Metals are malleable because the layers of positive ions can slide over one another without breaking the metallic bonding (the attraction is the same whichever ions are adjacent).

Energy changes on changing state. Melting or boiling means overcoming the forces between particles, so energy must be put in.

• To melt or boil a simple molecular substance you break only the weak intermolecular forces between the molecules, not the strong covalent bonds inside them. This is why these substances have low melting and boiling points.
• To melt an ionic, metallic or macromolecular solid you must break the strong bonding throughout a giant structure, so far more energy is needed.

Still don't get it? · why melting iodine does not break any covalent bonds

Think of a box of paperclips, where each paperclip is made of two balls welded together, and the clips are only lightly tangled with their neighbours. Tip the box and the clips come apart from each other easily, but every weld stays intact. Nothing about separating the clips breaks a weld.

Iodine is the same. Each I2 molecule is two iodine atoms held by a strong covalent bond (the weld). Separate I2 molecules are only held to each other by weak van der Waals forces (the light tangling). When iodine melts or boils, you pull the I2 molecules apart from one another, so you overcome the weak van der Waals forces between molecules. The strong covalent bond inside each I2 molecule is untouched.

In the exam: the energy to melt or boil a simple molecular substance overcomes the weak intermolecular forces between molecules. Writing that "the covalent bonds break" scores zero, and so does saying the forces are "disturbed" or "weakened" rather than broken or overcome.

Shapes of simple molecules and ions

The shape of a molecule or ion is set by electron pair repulsion (the VSEPR idea):

• bonding pairs and lone (non-bonding) pairs both behave as charge clouds that repel each other,
• the pairs in the outer shell of the central atom arrange themselves as far apart as possible, to minimise repulsion,
• repulsion is not equal: lone pair-lone pair > lone pair-bond pair > bond pair-bond pair.

A lone pair repels more than a bonding pair, so each lone pair squeezes the remaining bonds closer and reduces the bond angle by about 2.5°.

This table covers shapes with up to six electron pairs. The first five rows are the shapes with no lone pairs; the lower rows show what happens as bonding pairs are replaced by lone pairs.

To deduce a shape (the method the marks are for): count the bonding pairs and lone pairs on the central atom, state that the pairs repel to get as far apart as possible, allow for lone pairs repelling more than bonding pairs, then give the shape and bond angle. Worked examples below apply this to species with up to six pairs, including lone pairs.

Electronegativity and bond polarity

Electronegativity is the power of an atom to attract the pair of electrons in a covalent bond. It is measured on the Pauling scale; fluorine is the most electronegative element (4.0), and the most electronegative atoms are F, O, N and Cl.

When two atoms of different electronegativity share a pair of electrons, the electrons are pulled towards the more electronegative atom. The electron distribution in the bond is unsymmetrical, so the bond is a polar covalent bond with a permanent dipole.

Show polarity with partial charges: the more electronegative atom is δ−, the other is δ+.

For example, in hydrogen chloride chlorine is more electronegative than hydrogen, so the bond is polar, with δ⁺ on the H and δ⁻ on the Cl.

Covalent and ionic bonding are the two extremes of a continuum. A bond between atoms of similar electronegativity is essentially pure covalent; a very large electronegativity difference gives ionic bonding; an intermediate difference gives a polar covalent bond.

Why some molecules with polar bonds are not polar

A molecule only has an overall (net) permanent dipole if the individual bond dipoles do not cancel.

If a molecule is symmetrical (identical bonds arranged symmetrically, with no lone pairs on the central atom), the bond dipoles point outwards equally and cancel, so there is no net dipole and the molecule is non-polar, even though each bond is polar.

• Non-polar despite polar bonds: CO₂ (linear, symmetrical), CCl₄ and CF₄ (tetrahedral, symmetrical).
• Polar: CH₃Cl and CHCl₃ are not symmetrical, so their dipoles do not cancel and the molecule has a permanent dipole.

Forces between molecules

There are three kinds of force between molecules. They are much weaker than covalent, ionic or metallic bonding, and they decide the melting and boiling points of molecular substances.

1. Van der Waals forces (also called induced dipole-dipole, London or dispersion forces) act between all molecules.

They arise in three steps (this is the full mark-scheme chain):

1. electrons are moving constantly, so at any instant the electron distribution in a molecule is uneven, creating a temporary (instantaneous) dipole,
2. this induces a dipole in a neighbouring molecule,
3. the δ+ of one molecule is attracted to the δ− of the adjacent molecule.

Van der Waals forces get stronger when a molecule has more electrons (so bigger molecules have higher boiling points).

Down Group 7, boiling point rises from F₂ (a gas) to I₂ (a solid) because each larger molecule has more electrons, so the van der Waals forces between molecules are stronger.

2. Permanent dipole-dipole forces act between polar molecules (those with a permanent dipole), in addition to van der Waals forces. The δ+ of one molecule attracts the δ− of another. They are stronger than van der Waals forces alone, so polar molecules boil higher than non-polar molecules of similar size.

3. Hydrogen bonding is the strongest of the three. It occurs when a hydrogen atom is bonded to N, O or F (the three most electronegative atoms, each with an available lone pair). The δ+ hydrogen of one molecule is attracted to a lone pair on the N, O or F of a neighbouring molecule.

When you draw a hydrogen bond you must show: all lone pairs on the N, O or F; the δ+ and δ− partial charges; and the hydrogen bond going from the δ+ H to a lone pair, drawn straight (180°).

Hydrogen bonding, ice and anomalous boiling points

Hydrogen bonding explains two famous anomalies.

The anomalously high boiling points of H₂O, NH₃ and HF. Across Groups 4 to 7, boiling points generally rise down a group as molecules gain electrons and van der Waals forces strengthen. But the Period 2 hydrides H₂O, NH₃ and HF sit far above this trend because they also have hydrogen bonding, which is stronger and takes more energy to overcome.

Read the chart at Period 2: the Group 4 hydride (CH₄, no hydrogen bonding) is low and on-trend, while NH₃ (240 K), H₂O (373 K) and HF (293 K) all spike above their groups. Group 4 has no N, O or F bonded to hydrogen, so it shows no anomaly.

The low density of ice. In ice, each water molecule forms hydrogen bonds that hold the molecules in an open, regular lattice, so the molecules are held further apart than in liquid water. Because the same mass takes up more volume, ice is less dense than water and floats.

Still don't get it? · why ice is less dense than liquid water

Picture a crowd of people milling about in a room (liquid water): they can pack in close and shuffle past one another. Now ask everyone to hold hands at arm's length in a fixed, repeating pattern (ice): they lock into a tidy grid, but holding each other at arm's length leaves gaps, so the same number of people take up more floor space.

Water molecules do the same. In liquid water the molecules are close and constantly moving. When water freezes, each molecule forms hydrogen bonds to its neighbours in a fixed, open arrangement that holds the molecules further apart, leaving spaces in the structure.

In the exam: in ice the water molecules are held further apart (the structure is more open) than in liquid water, so ice is less dense. Say molecules, not atoms, and do not say the gaps are filled with air.

Worked examples

1. Construct the formula of an ionic compound.

Aluminium oxide. Aluminium is in Group 3, so it forms Al³⁺; oxygen is in Group 6, so it forms O²⁻.

• Balance the charge: two Al³⁺ give 6+, three O²⁻ give 6−.
• Formula: Al₂O₃.

Ammonium sulfate. The ions are NH₄⁺ and SO₄²⁻.

• Two NH₄⁺ (2+) balance one SO₄²⁻ (2−).
• Keep the compound ion together in brackets: (NH₄)₂SO₄.

2. Deduce the shape and bond angle (no lone pairs): SF₆.

1. Count the electron pairs on the central atom: sulfur has 6 outer electrons, and the 6 fluorine atoms each share 1, giving

$\text{electron pairs} = \frac{6 + 6}{2} = 6$

2. That is 6 bonding pairs and 0 lone pairs.
3. The 6 bonding pairs repel equally and get as far apart as possible.
4. Shape: octahedral, bond angle 90°.

3. Deduce the shape and bond angle (with a lone pair): H₃O⁺.

1. Oxygen has 6 outer electrons, 3 hydrogens share 1 each (+3), and the + charge removes 1 electron:

$\text{electron pairs} = \frac{6 + 3 - 1}{2} = 4$

2. That is 3 bonding pairs and 1 lone pair.
3. The 4 pairs repel to get as far apart as possible, but the lone pair repels more than the bonding pairs, squeezing the bonds together.
4. Shape: trigonal pyramidal, bond angle 107° (reduced from 109.5° by the lone pair).

4. Deduce the shape of a species with lone pairs and up to six pairs: ClF₃.

1. Chlorine has 7 outer electrons, 3 fluorines share 1 each (+3):

$\text{electron pairs} = \frac{7 + 3}{2} = 5$

2. That is 3 bonding pairs and 2 lone pairs (a variation of the 5-pair trigonal bipyramidal arrangement).
3. The 2 lone pairs repel most and sit as far apart as possible, so the three Cl-F bonds are pushed into a T-shape.
4. Bond angle: slightly less than 90° (about 87°), because the lone pairs repel the bonding pairs more strongly.

5. Model answer: explain why CCl₄ is non-polar but CHCl₃ is polar.

1. In both molecules each C-Cl bond is polar, because chlorine is more electronegative than carbon.
2. CCl₄ is tetrahedral and symmetrical, so the four bond dipoles are arranged symmetrically and cancel, giving no net dipole (non-polar).
3. In CHCl₃ one position is a C-H bond, so the molecule is not symmetrical; the dipoles do not cancel, so the molecule has a permanent dipole (polar).

Common exam mistakes

• Defining electronegativity as the power of an element (or molecule) rather than an atom, or leaving out that it attracts the pair of electrons in a covalent bond. Both wordings are needed.
• Comparing the "electronegativity of molecules" when discussing polarity or hydrogen bonding; electronegativity is a property of atoms, so compare the atoms (for example O versus S).
• Describing ionic bonding as attraction "between ions" without saying oppositely charged ions, or naming the wrong ions. The attraction is between the oppositely charged ions.
• Saying that melting or boiling a simple molecular substance (iodine, methane, methanol) breaks the covalent bonds, or that "the O-H bond breaks". You overcome the weak intermolecular forces between the molecules; the covalent bonds stay intact.
• Writing that the intermolecular forces are "disturbed" or "weakened". The mark needs them broken or overcome.
• Explaining van der Waals forces by "electron movement" alone, and forgetting to say a dipole is induced in a neighbouring molecule. All three steps are needed, and the forces must be stated to act between molecules.
• Saying hydrogen bonds (or any intermolecular force) are "strong" without the comparison. The mark needs "hydrogen bonds are stronger than van der Waals forces".
• Hydrogen-bond diagrams that miss the lone pair on N, O or F, omit the δ+/δ− charges, put a lone pair on hydrogen, or draw the bond to the wrong atom. The hydrogen bond goes from the δ+ H to a lone pair on the N, O or F of another molecule.
• Deducing a shape but forgetting that a lone pair reduces the bond angle: giving NH₃ or H₃O⁺ a bond angle of 109.5° instead of 107°.
• Explaining repulsion in terms of "the atoms repelling" rather than the electron pairs repelling.
• Confusing metallic and ionic bonding, or, for metals, forgetting to say the ions are positive and the electrons are delocalised.
• Explaining conductivity of solid sodium chloride as "no free electrons" instead of the ions being fixed in the lattice and unable to move; explaining a metal's malleability by "weak intermolecular forces" instead of layers of ions sliding over one another.
• For ice, saying the atoms (not molecules) are further apart, or that the gaps are filled with air. It is the molecules that are held further apart in a more open structure.
• For a dative bond, giving a generic description instead of naming which atom supplies the lone pair and which accepts it; drawing the arrow the wrong way, or (for NH₄⁺) thinking H⁺ donates the lone pair when it is the N that donates and H⁺ that accepts.

Key definitions

• Ionic bonding: the electrostatic attraction between oppositely charged ions (in a lattice).
• Covalent bond: a shared pair of electrons.
• Co-ordinate (dative covalent) bond: a shared pair of electrons in which both electrons are supplied by one atom.
• Metallic bonding: the electrostatic attraction between delocalised electrons and positive ions arranged in a lattice.
• Electronegativity: the power of an atom to attract the pair of electrons in a covalent bond.
• Polar covalent bond: a covalent bond between atoms of different electronegativity, in which the electron distribution is unsymmetrical, giving a permanent dipole.
• Van der Waals (induced dipole-dipole) forces: weak intermolecular forces caused by temporary dipoles, which arise from the movement of electrons, inducing dipoles in neighbouring molecules.
• Permanent dipole-dipole forces: forces of attraction between the permanent dipoles of polar molecules.
• Hydrogen bond: an intermolecular attraction between a hydrogen atom bonded to N, O or F and a lone pair on the N, O or F of another molecule.

Specification

• I can state that ionic bonding is the electrostatic attraction between oppositely charged ions in a lattice.
• I can predict the charge on a simple ion from its position in the Periodic Table and construct formulae for ionic compounds, including the compound ions sulfate, hydroxide, nitrate, carbonate and ammonium.
• I can describe a single covalent bond as a shared pair of electrons, and multiple bonds as multiple shared pairs.
• I can describe a co-ordinate (dative covalent) bond as a shared pair of electrons supplied by one atom, and represent a covalent bond with a line and a co-ordinate bond with an arrow.
• I can state that metallic bonding is the attraction between delocalised electrons and positive ions in a lattice.
• I can name the four crystal types (ionic, metallic, macromolecular, molecular) and the structures of diamond, graphite, ice, iodine, magnesium and sodium chloride.
• I can relate the melting point and conductivity of a material to its structure and bonding, and explain the energy changes associated with changes of state.
• I can explain the shapes of, and bond angles in, simple molecules and ions with up to six electron pairs (including lone pairs) around the central atom, using electron pair repulsion.
• I can define electronegativity and use partial charges to show that a bond is polar.
• I can explain why some molecules with polar bonds have no permanent dipole.
• I can explain the existence of van der Waals forces, permanent dipole-dipole forces and hydrogen bonding, and how they influence melting and boiling points.
• I can explain the importance of hydrogen bonding in the low density of ice and the anomalous boiling points of compounds.

Related notes

• Atomic structure
• Amount of substance
• Periodicity