Bonding and structure

In a nutshell

Bonding is about electrostatic attraction: between oppositely charged ions (ionic), or between a shared pair of electrons and two nuclei (covalent).

This subtopic shows you how to draw those bonds, predict the shape of a molecule from its electron pairs, decide whether a bond or molecule is polar, and use intermolecular forces to explain melting points, boiling points, solubility and conductivity.

Assumed knowledge: Electron structure and ion charges and formulae (Compounds, formulae and equations).

Core content

Ionic bonding

Ionic bonding is the electrostatic attraction between positive and negative ions.

The ions form when a metal atom loses electrons and a non-metal atom gains them. The transfer is how the ions are made; the bond itself is the attraction between the ions.

Dot-and-cross diagrams for ionic compounds

• Draw each ion separately, inside square brackets, with its charge at the top right.
• Use crosses for electrons that started on one atom and dots for the other, so the examiner can see which electrons moved.
• Showing the outer shell only is normally enough.
• Get the ratio right: CaCl₂ needs two chloride ions (or a "2" in front of one).

Giant ionic lattices

In the solid, the ions pack into a giant ionic lattice: a regular 3D arrangement of oppositely charged ions strongly attracted in all directions.

In NaCl, each Na⁺ is surrounded by six Cl^- ions and each Cl^- by six Na⁺ ions. There are no separate "NaCl molecules".

Physical properties of ionic compounds

• Ions with larger charges (and smaller ions) attract each other more strongly, so the melting point is higher. MgO (Mg²⁺, O^2-) melts at over 2800 °C; NaCl (Na⁺, Cl^-) melts at about 800 °C.
• Not every ionic compound dissolves. When the attractions in the lattice are very strong, water cannot break the lattice down.
• In ionic compounds, ions carry the charge. Electrons do not move through an ionic compound.

Covalent bonding

A covalent bond is the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.

Covalent bonds form between non-metal atoms. Each shared pair is one bond.

Single and multiple bonds

• Single bond: one shared pair (C-H, O-H, N-H).
• Double bond: two shared pairs (O=O, C=O).
• Triple bond: three shared pairs (N≡N).

A lone pair is a pair of outer-shell electrons on an atom that is not used in bonding. Lone pairs matter for shape, polarity and hydrogen bonding, so always draw them.

More or fewer than eight electrons

The "full outer shell of eight" is a guide, not a law. The spec covers molecules with up to six electron pairs around the central atom.

• BF₃: boron has only 6 electrons (3 pairs) around it.
• SF₆: sulfur has 12 electrons (6 pairs) around it. Sulfur is in Period 3, so it can hold more than eight outer electrons.

Dative covalent (coordinate) bonds

In a dative covalent bond, both electrons in the shared pair come from one atom. That atom must have a lone pair; the other atom must have space to accept it.

$\mathrm{NH_3 + H^+ \longrightarrow NH_4^+}$

$\mathrm{H_2O + H^+ \longrightarrow H_3O^+}$

• In NH₄⁺, the nitrogen lone pair forms the fourth N-H bond. In a displayed formula, a dative bond is drawn as an arrow from the atom donating the pair.
• Once formed, a dative bond is identical to the other bonds. All four N-H bonds in NH₄⁺ are the same, so the ion is tetrahedral.

Average bond enthalpy

Average bond enthalpy measures covalent bond strength: the larger the value, the stronger the bond.

For this subtopic you only need that link. The definition and the calculations come in Enthalpy changes (OCR 3.2.1).

The shapes of simple molecules and ions

Electron pairs around a central atom repel each other and move as far apart as possible. Where they end up sets the shape.

The two rules that earn the marks:

1. Electron pairs repel and take up positions of minimum repulsion (as far apart as possible).
2. Lone pairs repel more than bonded pairs. Each lone pair squeezes the bond angle by about 2.5°.

The six shapes you must know

Drawing shapes in 3-D

• A plain line is a bond in the plane of the paper.
• A solid wedge comes towards you.
• A dashed wedge goes away from you.
• Draw any lone pairs on the central atom; the examiner checks for them.

Why lone pairs shrink the angle

CH₄, NH₃ and H₂O each have four electron pairs around the central atom. The pairs point to the corners of a tetrahedron in all three.

What changes is how many of those pairs are lone pairs:

• CH₄: 4 bonded, 0 lone: 109.5°
• NH₃: 3 bonded, 1 lone: 107°
• H₂O: 2 bonded, 2 lone: 104.5°

The shape is named from the atoms only, so NH₃ is pyramidal, not tetrahedral, even though its four pairs point tetrahedrally.

Still don't get it? · why lone pairs repel more

Tie four balloons together at their knots. They spread out into a 3-D star with every balloon as far from the others as it can get. Now swap one balloon for a fatter one. It takes up more room, so it shoves the other three closer together.

Step by step:

1. Count the electron pairs around the central atom. In NH₃ there are four: three shared with H atoms and one lone pair.
2. Four pairs spread out tetrahedrally, which on its own would give 109.5°.
3. A bonded pair is pulled out between two nuclei. A lone pair belongs to the central atom only, so it sits closer to it and spreads out more. That is the "fat balloon".
4. The lone pair therefore repels the bonded pairs more strongly than they repel each other, pushing the N-H bonds closer together: 109.5° falls to 107°.
5. Water has two lone pairs, so the squeeze happens twice: 104.5°.

Exam wording: "there are 3 bonded pairs and 1 lone pair; electron pairs repel as far apart as possible; lone pairs repel more than bonded pairs; so the shape is pyramidal and the bond angle is 107°."

Multiple bonds count as one bonding region

A double or triple bond points in one direction, so for shape it counts like a single bonded pair.

CO₂ has two C=O double bonds and no lone pairs on carbon. That is two bonding regions, repelling equally, so CO₂ is linear, 180°.

Ions: count the electrons first

For a central atom joined by single bonds to H or halogen atoms, count the electrons around the central atom, then halve:

$\text{electrons around centre} = \text{central atom's outer electrons} + \text{one per bonded atom} + \text{negative charge} - \text{positive charge}$

Then subtract the bonded pairs to find the lone pairs. Worked examples 1 and 2 below use this.

Electronegativity and bond polarity

Electronegativity is the ability of an atom to attract the bonding electrons in a covalent bond.

It is measured on the Pauling scale. You do not need to learn the values; they are given in the question. You do need to know that electronegativity increases towards fluorine (across a period and up a group).

Hydrogen has a Pauling value of 2.20. Worked example 3 uses the chart values plus this one.

Using the difference to predict bond type

Think of it as a scale from equal sharing to complete transfer, not three sealed boxes.

Polar bonds and permanent dipoles

When two bonded atoms have different electronegativities, the bonding pair is pulled towards the more electronegative atom.

• The more electronegative atom gets a small negative charge, δ-; the other gets δ+.
• This charge separation is a permanent dipole, and the bond is a polar bond.
• Example: in H-Cl, chlorine (3.16) is more electronegative than hydrogen (2.20), so H is δ+ and Cl is δ-.

Polar molecules: the shape decides

A molecule with polar bonds is only polar overall if the dipoles do not cancel. That depends on shape.

• H₂O: two polar O-H bonds in a non-linear shape. The dipoles do not cancel, so water has an overall dipole.
• CO₂: two polar C=O bonds, but the molecule is linear and symmetrical. The dipoles point in opposite directions and cancel, so CO₂ is non-polar.

Intermolecular forces

Intermolecular forces are the weak attractions between molecules. Covalent bonds are the strong attractions within a molecule.

When a simple molecular substance melts or boils, it is the intermolecular forces that are overcome. The covalent bonds stay intact.

Both dipole-dipole types can be called van der Waals' forces. A polar molecule has London forces as well as its permanent dipole-dipole interactions.

Induced dipole-dipole interactions (London forces)

1. Electrons in a molecule are always moving, so at any instant they may be unevenly spread. This creates an instantaneous dipole.
2. That dipole induces a dipole in a neighbouring molecule, by repelling or attracting its electrons.
3. The two dipoles attract: δ+ of one to δ- of the next.

More electrons mean larger instantaneous and induced dipoles, so stronger London forces and higher boiling points.

Still don't get it? · induced dipole-dipole interactions

Picture a crowd of people milling around inside a room with no fixed places. At any moment, just by chance, more of them might drift to the left-hand side. For that instant, the room is lopsided.

Now put a second room next door, separated by a thin wall, and suppose the people in each room do not like being near the people next door. When the first room's crowd bunches up by the shared wall, the crowd next door shuffles away from that wall. The two rooms are now lopsided in step with each other.

Step by step:

1. The "people" are electrons. They move constantly, so their spread in a molecule is uneven at any instant: one end is briefly δ- and the other δ+.
2. Electrons repel electrons. The δ- end of the first molecule pushes the electrons of its neighbour away, so the neighbour's near end becomes δ+.
3. Opposite charges attract, so the two molecules attract. A moment later the electrons move again and a new pair of dipoles forms. The attraction is always flickering, which is why it is weak.
4. A molecule with more electrons can become far more lopsided, so the dipoles, and the attractions, are bigger.

Exam wording: "uneven distribution of electrons causes an instantaneous dipole, which induces a dipole in a neighbouring molecule; the dipoles attract. More electrons means stronger induced dipole-dipole interactions, so more energy is needed to overcome them."

The halogens show the trend clearly (you will explain it again in The halogens, OCR 3.1.3):

Permanent dipole-dipole interactions

These act between polar molecules: the δ+ end of one molecule attracts the δ- end of a neighbour. They are present in addition to London forces.

Compare three substances that all have 18 electrons, so their London forces are similar. Only HCl is polar:

HCl's much higher boiling point comes from its permanent dipole-dipole interactions, on top of London forces.

Hydrogen bonding

A hydrogen bond is an intermolecular attraction between a lone pair on an N, O or F atom of one molecule and a δ+ hydrogen atom bonded to N, O or F (in -NH, -OH or HF) on a neighbouring molecule.

Why these three elements:

• N, O and F are very electronegative, so an N-H, O-H or F-H bond is very polar and the H is strongly δ+.
• N, O and F are small and carry lone pairs, which the δ+ H is attracted to.

How to draw it so you get the marks:

• Show the δ+ on H and δ- on O, N or F.
• Show at least one lone pair on the O, N or F receiving the hydrogen bond.
• Draw the hydrogen bond as a dashed line from the H to the lone pair. The three atoms O-H···O lie in a straight line (about 180°).

Anomalous properties of water

Water has relatively high melting and boiling points. Hydrogen bonds are stronger than the other intermolecular forces, so more energy is needed to overcome them.

The hydrides of Groups 14 to 17 show how unusual this is:

• From Period 3 to Period 5, boiling points rise in every group: more electrons, stronger London forces.
• In Period 2, NH₃, H₂O and HF are far higher than the trend predicts, because of hydrogen bonding.
• CH₄ has no lone pairs and C is not electronegative enough, so it has no hydrogen bonding and sits at the bottom of its line.

Ice is less dense than liquid water.

• In ice, each water molecule is hydrogen bonded to four others in a regular, open lattice.
• The hydrogen bonds hold the molecules further apart than in the liquid.
• When ice melts, some hydrogen bonds break, the open lattice collapses and the molecules move closer together.
• So ice (about 0.92 g cm^-3) is less dense than liquid water (about 1.00 g cm^-3), and ice floats.

Simple molecular lattices

A simple molecular lattice is a regular arrangement of covalently bonded molecules attracted by intermolecular forces.

• Iodine, I₂: I₂ molecules (each held together by a strong covalent bond) attracted to each other by induced dipole-dipole interactions.
• Ice: H₂O molecules attracted to each other by hydrogen bonds (plus London forces).

Physical properties of simple molecular substances

Polar simple molecular substances that can hydrogen bond with water (such as ethanol) can dissolve in water.

Worked examples

Example 1: Predict and explain the shape of NCl₃.

Count the electrons around nitrogen (Group 15, so 5 outer electrons; each Cl forms one single bond and adds one):

$\text{electrons around N} = 5 + 3 = 8$

$\text{electron pairs} = 8 \div 2 = 4$

$\text{lone pairs} = 4 - 3 \text{ bonded pairs} = 1$

The model answer, one mark-earning point per line:

1. N has 3 bonded pairs and 1 lone pair.
2. Electron pairs repel and get as far apart as possible.
3. Lone pairs repel more than bonded pairs.
4. So NCl₃ is pyramidal, with a bond angle of about 107°.

Where marks go: writing "tetrahedral" (that names the electron-pair arrangement, not the shape), or forgetting point 3, which is the reason the angle is 107° and not 109.5°.

Example 2: Predict the shapes of the ions NH₂^-, PCl₄⁺, and the molecule SF₆.

NH₂^-: N has 5 outer electrons, 2 bonded H, and the 1- charge adds one electron.

$5 + 2 + 1 = 8 \text{ electrons} = 4 \text{ pairs: } 2 \text{ bonded}, \ 2 \text{ lone}$

Shape: non-linear, about 104.5°.

PCl₄⁺: P has 5 outer electrons, 4 bonded Cl, and the 1+ charge removes one electron.

$5 + 4 - 1 = 8 \text{ electrons} = 4 \text{ pairs: } 4 \text{ bonded}, \ 0 \text{ lone}$

Shape: tetrahedral, 109.5°.

SF₆: S has 6 outer electrons and 6 bonded F; no charge.

$6 + 6 = 12 \text{ electrons} = 6 \text{ pairs: } 6 \text{ bonded}, \ 0 \text{ lone}$

Shape: octahedral, 90°.

Where marks go: adding the charge the wrong way round. A negative charge means extra electrons; a positive charge means fewer.

Example 3: Use the Pauling values in the chart to decide which bonds are polar, then whether CH₃Cl and CCl₄ are polar molecules.

Electronegativity differences (C = 2.55, Cl = 3.16, H = 2.20):

$\text{C-Cl: } 3.16 - 2.55 = 0.61 \quad \Rightarrow \text{polar, C is } \delta+ \text{, Cl is } \delta-$

$\text{C-H: } 2.55 - 2.20 = 0.35 \quad \Rightarrow \text{only slightly polar, usually treated as non-polar}$

Now the shape: both molecules have 4 bonded pairs and no lone pairs on carbon, so both are tetrahedral.

• CCl₄: four identical C-Cl dipoles arranged symmetrically. They cancel, so CCl₄ is non-polar.
• CH₃Cl: one C-Cl dipole is not balanced by the three C-H bonds. The dipoles do not cancel, so CH₃Cl is polar.

Where marks go: saying "CCl₄ has polar bonds, so it is polar". The mark needs the shape and the word cancel.

Example 4: Explain why water (boiling point 373 K) has a much higher boiling point than hydrogen sulfide (213 K).

The model answer:

1. Water molecules form hydrogen bonds between molecules (O is very electronegative and has lone pairs).
2. H₂S has only permanent dipole-dipole and induced dipole-dipole interactions; S is not electronegative enough for hydrogen bonding.
3. Hydrogen bonds are stronger than these other intermolecular forces.
4. So more energy is needed to overcome the intermolecular forces in water.

Where marks go: writing that "the O-H bonds are stronger than the S-H bonds" or that "bonds break when water boils". Boiling overcomes intermolecular forces, not covalent bonds.

Example 5: Explain why sodium chloride has a high melting point and conducts electricity when molten but not when solid.

1. NaCl has a giant ionic lattice.
2. There are strong electrostatic attractions between oppositely charged ions in all directions, so a lot of energy is needed to overcome them: high melting point.
3. In the solid, the ions are fixed in the lattice and cannot move.
4. When molten, the ions are free to move and carry charge, so it conducts.

Where marks go: saying "electrons move" in molten NaCl. In an ionic compound, the ions are the moving charges.

Common exam mistakes

• Ionic bonding as "transfer of electrons". Transfer is how the ions form. The bond is the electrostatic attraction between oppositely charged ions.
• Mixing structure words. Writing "intermolecular forces" or "molecules" for NaCl, or "ions" for iodine, contradicts the answer and loses the mark. Ionic: ions and a giant lattice. Simple molecular: molecules and intermolecular forces.
• Saying covalent bonds break on melting or boiling a simple molecular substance. Only the intermolecular forces are overcome.
• Electrons carrying the current in molten or aqueous ionic compounds. It is the ions that move.
• Dot-and-cross slips: missing square brackets or charges on ions, wrong ratio of ions (one Cl^- for CaCl₂), or leaving off lone pairs on N or O in a molecule.
• Dative bond arrow the wrong way round. The arrow points from the atom that donates the lone pair.
• Leaving out "lone pairs repel more than bonded pairs" when explaining NH₃ or H₂O. Without it there is no reason for 107° or 104.5°.
• Naming the electron-pair arrangement instead of the shape: calling NH₃ "tetrahedral" or H₂O "tetrahedral". Use OCR's names: pyramidal and non-linear.
• Swapping the angles: 107° belongs to NH₃, 104.5° to H₂O.
• Counting each double bond in CO₂ as two pairs and predicting a bent or tetrahedral shape. Each C=O is one bonding region, so CO₂ is linear.
• Electronegativity defined vaguely: "how much an element wants electrons". The definition needs an atom attracting the bonding electrons in a covalent bond.
• "Polar bonds, so a polar molecule." CO₂ and CCl₄ have polar bonds but are non-polar, because the dipoles cancel in a symmetrical shape.
• Explaining London forces with "bigger mass" or "bigger M_r". The cause is more electrons.
• Hydrogen-bond diagrams missing a lone pair or the δ+/δ- charges, or drawing the hydrogen bond within one molecule (the O-H covalent bond) instead of between molecules.
• Ice "less dense because the molecules move less". The mark needs hydrogen bonds holding the molecules further apart in an open lattice.
• Reversing bond enthalpy: a larger average bond enthalpy means a stronger bond.

Key definitions

• Ionic bonding: electrostatic attraction between positive and negative ions.
• Giant ionic lattice: a regular 3D arrangement of oppositely charged ions strongly attracted in all directions.
• Covalent bond: the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms.
• Dative covalent (coordinate) bond: a covalent bond in which the shared pair of electrons is supplied by one of the bonding atoms only.
• Lone pair: a pair of outer-shell electrons that is not involved in bonding.
• Average bond enthalpy: a measure of covalent bond strength; the larger the value, the stronger the bond (a formal definition is not required until 3.2.1).
• Electronegativity: the ability of an atom to attract the bonding electrons in a covalent bond.
• Polar bond (permanent dipole): a covalent bond between atoms of different electronegativities, giving a small charge separation (δ+ and δ-) across the bond.
• Polar molecule: a molecule with polar bonds whose dipoles do not cancel because of the molecule's shape, giving an overall dipole.
• Induced dipole-dipole interactions (London forces): attractions between an instantaneous dipole in one molecule and the dipole it induces in a neighbouring molecule.
• Permanent dipole-dipole interactions: attractions between the δ+ end of one polar molecule and the δ- end of a neighbouring polar molecule.
• Hydrogen bond: an intermolecular attraction between a lone pair on an N, O or F atom and a δ+ hydrogen atom bonded to N, O or F in a neighbouring molecule.
• Simple molecular lattice: a regular arrangement of covalently bonded molecules attracted by intermolecular forces.

Specification

• I can describe ionic bonding as electrostatic attraction between positive and negative ions, and draw dot-and-cross diagrams for ionic compounds. (2.2.2(a))
• I can explain the giant ionic lattice structure of a compound such as NaCl. (2.2.2(b))
• I can explain the melting and boiling points, solubility and electrical conductivity (solid, liquid and aqueous) of ionic compounds from their structure and bonding. (2.2.2(c))
• I can define a covalent bond as the strong electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms. (2.2.2(d))
• I can draw dot-and-cross diagrams of molecules and ions with single, multiple and dative covalent bonds, with up to six electron pairs around the central atom. (2.2.2(e))
• I can use average bond enthalpy as a measure of covalent bond strength (larger value, stronger bond). (2.2.2(f))
• I can predict the shapes and bond angles of molecules and ions with up to six electron pairs, using electron pair repulsion and the greater repulsion of lone pairs, and draw them in 3-D. (2.2.2(g))
• I can explain linear, non-linear, trigonal planar, pyramidal, tetrahedral and octahedral shapes, including the angles in CH₄ (109.5°), NH₃ (107°) and H₂O (104.5°). (2.2.2(h))
• I can define electronegativity, interpret Pauling values, and state that electronegativity increases towards fluorine. (2.2.2(i))
• I can explain a polar bond and permanent dipole, and explain whether a molecule is polar from its dipoles and its shape (e.g. H₂O polar, CO₂ non-polar). (2.2.2(j))
• I can explain permanent dipole-dipole and induced dipole-dipole (London) interactions. (2.2.2(k))
• I can describe and draw hydrogen bonding between molecules containing N, O or F and the H of -NH, -OH or HF, including the role of lone pairs. (2.2.2(l))
• I can explain why ice is less dense than water and why water has relatively high melting and boiling points. (2.2.2(m))
• I can explain simple molecular lattices such as I₂ and ice. (2.2.2(n))
• I can explain the melting and boiling points, solubility and electrical conductivity of simple molecular substances. (2.2.2(o))

Related notes

• Electron structure
• Compounds, formulae and equations
• Periodicity: metallic bonding, giant covalent lattices and melting points across Periods 2 and 3
• The halogens: boiling points of Cl₂, Br₂ and I₂
• Enthalpy changes: average bond enthalpies and calculations
• Alkanes: boiling points, chain length and branching