Edexcel International A Level Chemistry: Shapes of molecules and metallic bonding (YCH11)

This guide teaches Topic 3 parts 3C Shapes of molecules and 3D Metallic bonding of the Pearson Edexcel International Advanced Subsidiary/Advanced Level in Chemistry (XCH11/YCH11) specification, Issue 1, September 2017. It covers learning outcomes 3.16 to 3.22, the final outcomes of Topic 3: Bonding and Structure. All of it sits in Unit 1: Structure, Bonding and Introduction to Organic Chemistry, the IAS compulsory unit, so everything here is IAS content. The specification lists no core practical for 3C or 3D.

For fast recall, use the shapes and metallic bonding revision notes. To test yourself, try the shapes and metallic bonding practice questions. The wider Unit 1 picture is in the structure, bonding and organic chemistry study guide, and the earlier parts of Topic 3 (3A and 3B) are in the ionic and covalent bonding study guide and the bonding revision notes. See the Edexcel International A Level Chemistry course hub for every unit, and tick outcomes off on the printable checklist. To find gaps across the course, take one of the free 10-minute diagnostics.

What this unit covers

Outcome	What you must be able to do	Stage
3.16	Use the principles of electron-pair repulsion theory to interpret and predict shapes of simple molecules and ions	IAS
3.17	Explain the terms bond length and bond angle	IAS
3.18	Know and explain the shapes and bond angles of BeCl₂, BCl₃, CH₄, NH₃, NH₄⁺, H₂O, CO₂, gaseous PCl₅, SF₆ and C₂H₄	IAS
3.19	Predict shapes and bond angles of molecules and ions analogous to those in 3.18	IAS
3.20	Describe metals as giant lattices of metal ions in a sea of delocalised electrons	IAS
3.21	State that metallic bonding is the strong electrostatic attraction between metal ions and delocalised electrons	IAS
3.22	Use that model to interpret simple properties of metals, including electrical conductivity and high melting temperature	IAS

3.16 Electron-pair repulsion theory

The theory rests on four ideas.

1. The electron pairs in the outer shell of the central atom repel each other, because they are all negative.
2. The pairs move as far apart as they can, so repulsion is at a minimum. That arrangement fixes the shape.
3. A lone pair is held closer to the central nucleus than a bonding pair, so it spreads out more and repels more strongly. Two lone pairs repel each other most, a lone pair and a bonding pair less, two bonding pairs least.
4. A double bond holds both of its pairs in the same region between the same two atoms. For shape purposes it counts as one region of electron density, just like a single bond.

The electron-pair arrangement includes lone pairs; the shape describes only where the atoms are. NH₃ has a tetrahedral arrangement of four pairs but a pyramidal shape.

A counting method that works every time

1. Write down the number of outer-shell electrons on the central atom (for example 3 for boron, 5 for nitrogen and phosphorus, 6 for oxygen and sulfur).
2. Subtract the central atom’s own electrons used in bonding: one per single bond, two per double bond, three per triple bond.
3. For a negative ion add one electron per negative charge. For a positive ion subtract one per positive charge.
4. Divide what is left by 2. That gives the number of lone pairs.
5. Lone pairs plus bonded atoms gives the number of regions and the arrangement; the atom positions give the shape.

Worked example 1 (3.16, 3.18). Compare NH₃ with NH₄⁺.

• NH₃: nitrogen has 5 outer electrons. Three N-H bonds use 3. That leaves 2 electrons, which is 1 lone pair. Four regions in total (3 bonding, 1 lone).
• NH₄⁺: nitrogen has 5 outer electrons. Four N-H bonds use 4. The positive charge removes 1 more. That leaves 0, so there are no lone pairs. Four regions, all bonding.
• Both have four regions, so both have a tetrahedral arrangement.
• NH₃ is pyramidal, H-N-H about 107°. The lone pair repels the three bonding pairs more strongly than they repel each other, squeezing them together.
• NH₄⁺ is tetrahedral, H-N-H 109.5°. When NH₃ accepts H⁺, its lone pair becomes the dative bond. All four pairs are now bonding pairs and repel equally, so the angle opens out by about 2.5°.

3.17 Bond length and bond angle

Bond length is the internuclear distance for two atoms joined by a covalent bond. It is the separation at which the attraction of both nuclei for the shared electrons balances the repulsion between the two nuclei.

Bond angle is the angle between two bonds that meet at the same atom. It is written with three atoms, the middle one being the atom at the vertex, for example the H-O-H angle in water.

More shared electrons between the same two atoms pull the nuclei closer together. Using typical values of 0.154 nm for C-C, 0.134 nm for C=C and 0.120 nm for C≡C, the double bond is about 13% shorter than the single bond and the triple bond about 22% shorter.

3.18 The ten named shapes

Species	Bonding regions	Lone pairs	Shape	Bond angle
BeCl₂ (gas)	2	0	linear	180°
CO₂	2 (both C=O)	0	linear	180°
BCl₃	3	0	trigonal planar	120°
C₂H₄ (each carbon)	3 (one is C=C)	0	trigonal planar at each C; whole molecule planar	about 120°
CH₄	4	0	tetrahedral	109.5°
NH₄⁺	4	0	tetrahedral	109.5°
NH₃	3	1	pyramidal	about 107°
H₂O	2	2	bent (V-shaped, non-linear)	about 104.5°
PCl₅ (gas)	5	0	trigonal bipyramidal	90° (axial) and 120° (equatorial)
SF₆	6	0	octahedral	90°

Points the explanations need:

• BeCl₂ and BCl₃: beryllium has only 2 outer electrons and boron 3, so there are two or three bonding pairs, no lone pairs, and fewer than eight electrons round the central atom.
• CO₂ has two double bonds. Each counts as one region, so two regions point opposite each other.
• C₂H₄: each carbon has three regions (two C-H and one C=C), so each carbon is trigonal planar. All six atoms lie in one plane. The double bond region repels a little more than a single bond, so the measured H-C-H angle is slightly below 120°; quote “about 120°” unless a question gives data.
• H₂O has two lone pairs. Lone pair-lone pair repulsion is the strongest of all, so the angle is squeezed further than in NH₃. The trend 109.5° → 107° → 104.5° for CH₄, NH₃, H₂O is the key comparison.
• PCl₅ is named as gaseous because the solid has a different, ionic structure. In the gas, three chlorine atoms lie in a triangle round the phosphorus (equatorial, 120° apart) and two sit above and below (axial, 90° to the triangle). All five outer electrons of phosphorus are used in bonds.
• SF₆: sulfur uses all six outer electrons, giving six bonding pairs; adjacent F-S-F angles are 90°.

Drawing shapes

Use the standard 3D convention:

• an ordinary line shows a bond lying flat on the page;
• a solid wedge shows a bond pointing out towards you;
• a dashed (hashed) wedge shows a bond pointing back, away from you.

Show lone pairs on the central atom and label the bond angle. A tetrahedral species is two ordinary lines, one wedge and one dashed wedge.

3.19 Predicting analogous shapes

An analogous species has the same number of bonding regions and lone pairs as one of the named ten, so it has the same shape. Count first; never guess from the formula.

Worked example 2. Predict the shape of the oxonium ion, H₃O⁺.

• Oxygen outer electrons: 6.
• Used in three O-H bonds: 3. Positive charge: take away 1.
• Left over: 6 - 3 - 1 = 2, which is 1 lone pair.
• Regions: 3 bonding + 1 lone = 4.

Four regions with one lone pair, like NH₃. H₃O⁺ is pyramidal with H-O-H about 107°.

Worked example 3. Predict the shape of the hexafluorosilicate ion, SiF₆²⁻.

• Silicon outer electrons: 4.
• Used in six Si-F bonds: 6. The 2- charge: add 2.
• Left over: 4 - 6 + 2 = 0, so no lone pairs.
• Regions: 6, all bonding.

Six bonding pairs and no lone pairs, like SF₆. SiF₆²⁻ is octahedral, F-Si-F 90°.

Worked example 4. Predict the shapes of carbon disulfide, CS₂, and tetrafluoroethene, C₂F₄.

• CS₂ is S=C=S. Carbon uses all 4 outer electrons in two double bonds, so it has no lone pairs and two regions. Like CO₂, it is linear, 180°.
• C₂F₄ is F₂C=CF₂. Each carbon has one C=C and two C-F bonds: three regions and no lone pairs. Like C₂H₄, each carbon is trigonal planar, F-C-F about 120°, and the molecule is planar.

To “explain” a shape, give the numbers of bonding pairs and lone pairs, say the pairs repel to be as far apart as possible (lone pairs more strongly), then state the shape and angle.

3D Metallic bonding

3.20 The model

Every metal consists of a giant lattice of positive ions surrounded by a sea of delocalised electrons. Each atom gives its outer electrons to the whole structure. Those electrons do not belong to any one ion; they can move through the lattice. Sodium gives one electron per atom and forms Na⁺ ions; magnesium gives two and forms Mg²⁺ ions; aluminium gives three and forms Al³⁺ ions.

3.21 The bond

Metallic bonding is the strong electrostatic attraction between the metal ions and the delocalised electrons. It acts in all directions throughout the lattice. Do not describe it as atoms sharing electrons.

3.22 Interpreting properties

Electrical conductivity. The delocalised electrons are free to move. When a potential difference is applied, they drift through the lattice towards the positive terminal, carrying charge. Metals conduct as solids and when molten, because the delocalised electrons are present in both states.

High melting temperature. To melt a metal, the ions must gain enough energy to move past one another. That means overcoming much of the strong attraction between the ions and the delocalised electrons, which acts throughout a giant lattice, so a lot of energy is needed.

The strength of metallic bonding increases when:

• the ions carry a larger charge;
• there are more delocalised electrons per atom;
• the ions are smaller, so the delocalised electrons are closer to the centre of positive charge.

Worked example 5. Sodium melts at about 98 °C and magnesium at about 650 °C. Explain the difference.

• Magnesium forms Mg²⁺ ions; sodium forms Na⁺ ions. The larger charge on the ion attracts the electron sea more strongly.
• Magnesium releases two delocalised electrons per atom; sodium releases one. The electron sea is denser.
• The Mg²⁺ ion is smaller than the Na⁺ ion, so the charges are closer together.
• Overall, magnesium ions attract their electron sea more strongly, so melting magnesium takes more energy. Its melting temperature is about 552 °C higher.

Other simple properties. Metals are malleable: when a force is applied, rows of ions shift to new positions and the electron sea still holds them, because the bonding is not fixed between particular atoms.

Common errors

• Counting only bonded atoms and forgetting lone pairs. Water has two bonded atoms but four regions.
• Treating a double bond as two regions. CO₂ would come out as tetrahedral, which is wrong.
• Giving “tetrahedral” for NH₃. That is the electron-pair arrangement; the shape is pyramidal.
• Writing 90° only for PCl₅. The equatorial angle of 120° is needed as well.
• Leaving out the charge when counting electrons for an ion.
• Saying metals conduct because “ions move”. In a solid metal the ions stay in the lattice; the delocalised electrons carry the charge.

Official syllabus

Pearson Edexcel International Advanced Subsidiary/Advanced Level in Chemistry (XCH11/YCH11) specification, Issue 1, September 2017, published by Pearson Education Limited. Topic 3: Bonding and Structure, parts 3C Shapes of molecules and 3D Metallic bonding, outcomes 3.16 to 3.22.