Sec 3 Chemistry — Chemical Bonding and Structure
Prepared by Miss Clarissa Ng · www.clartutors.com
Part A · How atoms bond
1 Why Do Atoms Bond at All?
Chemical bonding happens in both elements and compounds. Most atoms form bonds with other atoms to reach a stable (usually duplet or octet) electronic configuration with a completely filled valence shell — the arrangement the noble gases already have.
Whatever the type of bond, one thing is always true: every kind of chemical bonding is an electrostatic force of attraction between oppositely charged particles. The three types differ only in what those particles are:
| Type of bonding | Between | The attraction is between |
|---|---|---|
| Ionic | A metal and a non-metal | Oppositely charged ions — one has lost electrons, the other has gained them. |
| Covalent | Two non-metal atoms | The positively charged nuclei and the shared pair(s) of electrons between them. |
| Metallic | Metal atoms | The metal cations and the sea of delocalised valence electrons around them. |
Exam habit: never write "atoms stick together". Say what is attracted to what — the marks are for the charged particles and the word electrostatic.
2 Ionic Bonding — Electrons Are Transferred
Ionic bonding is the electrostatic forces of attraction between oppositely charged ions in an ionic compound. It generally occurs between cations and anions.
Ionic bonds form when metal atoms transfer their valence electrons to non-metal atoms. Both atoms then reach a stable duplet or octet configuration with a completely filled valence shell. The metal becomes a cation and the non-metal becomes an anion.
Worked example — sodium oxide
4Na + O2 → 2Na2O
- Each sodium atom (2.8.1) loses its single valence electron and becomes a Na+ ion, now 2.8.
- The oxygen atom (2.6) gains two electrons and becomes an O2− ion, now 2.8.
- Two sodium atoms are needed for every one oxygen atom, so the formula is Na2O.
- Oxygen exists as diatomic O2 molecules, so the balanced equation uses four sodium atoms and gives two formula units: 4Na + O2 → 2Na2O.
- The compound is electrically neutral overall — the electrons lost by the metal equal the electrons gained by the non-metal.
The balance rule: in an ionic compound the total number of electrons lost by the metal atoms equals the total number gained by the non-metal atoms. Use it to check any formula you are asked to write.
3 Covalent Bonding — Electrons Are Shared
Covalent bonding is the electrostatic forces of attraction between the positively charged nuclei of atoms and their shared pair(s) of electrons. It generally occurs between atoms of non-metals.
A covalent bond forms when two atoms share at least one pair of valence electrons, so that both reach a stable duplet or octet configuration. The shared electrons sit in the overlapping region of the valence shells of both atoms.
The number of pairs shared decides the name of the bond:
| Pairs of electrons shared | Bond | Where you meet it |
|---|---|---|
| One pair | Single covalent bond | Hydrogen, H₂ — one shared pair holding the two atoms together. |
| Two pairs | Double covalent bond | Oxygen, O₂ — two shared pairs, which is why oxygen is written with a double bond. |
| Three pairs | Triple covalent bond | Nitrogen, N₂ — three shared pairs, the strongest of the three. |
Exam habit: count the pairs, then name the bond. One shared pair is a single bond — never "one bond", because the mark is for the pair.
Part B · What the structure looks like
4 Giant Ionic Lattice
Ionic compounds do not exist as single molecules. Their ions are arranged in a giant ionic (crystal) lattice structure — a three-dimensional network in which a very large number of ions are packed in a regular pattern.
- Oppositely charged ions are arranged to be as near to each other as possible.
- Similarly charged ions are arranged as far apart as possible.
- This maximises attraction and minimises repulsion, giving a highly stable structure. The ions are held in place by strong electrostatic forces of attraction.
5 Simple Molecular Structure
Simple covalent substances have a simple molecular structure: they exist as many simple (discrete) molecules. Two very different strengths of attraction are at work in the same substance —
| What is holding what | |
|---|---|
| Inside a molecule | The atoms are held together by strong covalent bonds — the shared pairs of electrons. |
| Between molecules | The molecules are held to one another only by weak intermolecular forces of attraction. |
This single difference explains almost every property question: melting or boiling a simple covalent substance overcomes only the weak forces between molecules — the covalent bonds inside each molecule are not broken. That is why the melting and boiling points are low.
6 Giant Covalent Structures
Covalent substances fall into two groups, and the difference is entirely structural:
| Simple covalent | Giant covalent | |
|---|---|---|
| Examples | Oxygen, water, ammonia | Diamond, graphite, silicon dioxide |
| Structure | A simple molecular structure — separate molecules, held to each other by weak intermolecular forces. | A giant covalent structure — an extensive network of atoms held by strong covalent bonds. |
Diamond — the network in three dimensions
- Diamond has a giant three-dimensional tetrahedral network of carbon atoms held by strong carbon–carbon (C–C) single covalent bonds.
- Each carbon atom is bonded to four other carbon atoms, reaching a stable octet with a completely filled valence shell.
Graphite — the network in layers
- Graphite consists of flat layers of carbon atoms, with weak intermolecular forces of attraction (van der Waals forces) between the layers.
- Within a layer each carbon atom is bonded to three other carbon atoms, arranged in regular hexagonal rings.
Silicon dioxide — the same idea with two elements
Silicon dioxide is found in sand, and its structure is very like diamond's. Each silicon atom is bonded to four oxygen atoms, and each oxygen atom is bonded to two silicon atoms. That gives a ratio of 1:2, which is why the formula is SiO2.
Diamond and graphite are both pure carbon — yet one is the hardest natural substance and the other is soft enough to write with. They are allotropes: different forms of the same element with a different structural arrangement of atoms. Structure, not the element, decides the properties.
Part C · Metallic bonding
7 Metallic Bonding and the Structure of Metals
Metallic bonding is the electrostatic forces of attraction between the metal cations and the sea of delocalised valence electrons. It generally occurs in metals.
The atoms of a metal lose their valence electrons to form positively charged ions — cations. Those cations are arranged regularly in layers in a three-dimensional lattice, and the lost electrons are free to move throughout the whole structure as a sea of delocalised valence electrons. The particles together form a giant metallic lattice structure.
Exam habit: the electrons are delocalised, not "free electrons floating around". Name the two things being attracted — the cations and the sea of delocalised valence electrons — and the mark is yours.
Part D · Explaining the properties
8 Ionic Compounds Compared with Simple Covalent Substances
Every property question in this topic is answered the same way: name the structure, then say which attraction is being overcome. Put the two side by side and the pattern is obvious.
Group 1 — structure, melting point and state
| Ionic compound | Simple covalent substance | Why | |
|---|---|---|---|
| Structure | Giant ionic lattice | Simple molecular | Ions pack into a network; molecules stay as separate units. |
| Melting and boiling point | High | Low | Ionic: a large amount of thermal energy is needed to overcome the strong electrostatic forces between ions. Covalent: only the weak forces between molecules must be overcome — the covalent bonds inside each molecule are not broken. |
| State at room temperature | Solid | Liquid or gas | Follows directly from the melting point. |
Group 2 — conductivity and solubility
| Ionic compound | Simple covalent substance | Why | |
|---|---|---|---|
| Electrical conductivity | Only when molten or aqueous | Does not conduct in any state (except acids and alkalis) | Ionic: in the solid state the ions are held in fixed positions and cannot move, so there are no mobile charge carriers; molten or dissolved, the forces holding them are overcome and the ions become free-moving. Covalent: the substance is electrically neutral with no mobile charge carriers at all. |
| Solubility in water | Usually soluble | Usually insoluble | Ionic lattices are broken up by water; neutral molecules are not. |
| Solubility in organic solvents | Insoluble | Soluble | The reverse — hexane is the organic solvent normally used to tell the two apart. |
Exam habit: the conductivity answer is always about mobile charge carriers. Say whether charged particles can move, not just whether the substance "has ions" — a solid ionic compound has plenty of ions and still cannot conduct.
9 Giant Covalent Substances Compared
Diamond, graphite and silicon dioxide all have giant covalent structures, yet their properties differ. The table below is the one to learn, because it separates what is the same from what is different.
| Diamond | Graphite | Silicon dioxide | |
|---|---|---|---|
| Bonding | Each C bonded to 4 others | Each C bonded to 3 others in a layer | Each Si bonded to 4 O; each O bonded to 2 Si |
| Hardness | Hardest — atoms held rigidly in fixed positions by strong covalent bonds that cannot be broken easily | Soft and slippery — the layers slide over one another easily because the forces between them are weak | Hard — the network is held rigidly in fixed positions |
| Melting point | Extremely high | Extremely high | High |
| Conducts electricity? | No — all four valence electrons are used for bonding, so there are no delocalised electrons | Yes — only three of the four valence electrons are used for bonding, so one delocalised electron per atom can move along the layers | No — all valence electrons are used for covalent bonding |
| Solubility | Insoluble in water and organic solvents | Insoluble in water and organic solvents | Insoluble in water and organic solvents |
| Typical use | Tool tips for drills and saws — the hardest naturally occurring substance | Lubricant in machinery — it is slippery | Making glass — it is hard with a high melting point |
The one question that catches everyone: why does graphite conduct when diamond does not? Both are carbon. Graphite's atoms use only three of their four valence electrons for bonding, leaving one delocalised electron per atom; diamond's atoms use all four, leaving nothing to carry a current.
10 Metals, and Why We Use Alloys
| Property | Explanation |
|---|---|
| High melting and boiling points | A large amount of thermal energy is needed to overcome the strong electrostatic forces of attraction between the metal cations and the sea of delocalised valence electrons. (Exception: Group 1 metals have relatively low melting points — sodium melts at 98 °C.) |
| Solid at room temperature | Most metals are solids. (Exception: mercury, which is a liquid.) |
| Conducts electricity | In both the solid and molten states, because the delocalised valence electrons move freely within the lattice as mobile charge carriers. Comparing two metals with the same number of atoms, the one with more delocalised valence electrons conducts better. |
| Conducts heat | The same free-moving delocalised electrons transfer thermal energy through the metal. |
| High density | The metal cations are packed closely together in the giant metallic lattice. (Exception: Group 1 metals.) |
| Malleable and ductile | Malleable: they can be bent or flattened into thin sheets when hammered. Ductile: they can be pulled into wires without breaking. When a force is applied, the layers of atoms slide over each other and the atoms take up new positions without the structure breaking. |
Alloys — why pure metals are rarely used
Pure metals are not widely used in industry: they are softer than alloys, and they corrode easily because they react readily with water and/or oxygen.
An alloy is a solid mixture made up of a metal and one or a few other elements. The other elements may be metals and/or non-metals (for example carbon). The physical properties of the alloy are different from those of its constituent pure metal(s).
- In an alloy the constituent atoms have different sizes. That disrupts the regular arrangement of the atoms in the pure metal, giving an irregular arrangement.
- A larger force is therefore needed to make the layers slide over each other, compared with the layers in the pure metal.
- So an alloy is stronger and harder than the pure metal — alloying is how a soft metal is made useful.
11 Put It Together — Exam-Style Question
A student is given three solids: sodium chloride, iodine (a simple covalent substance) and graphite.
[3](a) State the type of bonding present in each of the three solids.
[3](b) Sodium chloride does not conduct electricity when solid, but it does when molten. Explain why.
[2](c) Iodine has a low melting point. Explain why, in terms of its structure and bonding.
[3](d) Graphite conducts electricity but diamond does not, even though both are made of carbon. Explain why.
[2](e) Brass is an alloy of copper and zinc. Explain why brass is harder than pure copper.
Model answers.
(a) Sodium chloride: ionic bonding. Iodine: covalent bonding in a simple molecular structure. Graphite: covalent bonding in a giant covalent structure.
(b) In the solid state the ions are held in fixed positions in the giant ionic lattice, so there are no mobile charge carriers. When molten, the electrostatic forces of attraction are overcome, so the ions become free-moving and can carry charge.
(c) Iodine has a simple molecular structure. Only the weak intermolecular forces of attraction between the molecules need to be overcome, so only a small amount of thermal energy is needed. The strong covalent bonds within each molecule are not broken.
(d) In graphite each carbon atom uses only three of its four valence electrons for bonding, so there is one delocalised electron per atom which can move freely along the layers as a mobile charge carrier. In diamond each carbon atom bonds to four other atoms and uses all four valence electrons, so there are no delocalised electrons to conduct.
(e) The zinc atoms are of a different size from the copper atoms, which disrupts the regular arrangement of the layers. The layers therefore cannot slide over one another easily, so a larger force is needed — brass is harder and stronger than pure copper.
(a) Sodium chloride: ionic bonding. Iodine: covalent bonding in a simple molecular structure. Graphite: covalent bonding in a giant covalent structure.
(b) In the solid state the ions are held in fixed positions in the giant ionic lattice, so there are no mobile charge carriers. When molten, the electrostatic forces of attraction are overcome, so the ions become free-moving and can carry charge.
(c) Iodine has a simple molecular structure. Only the weak intermolecular forces of attraction between the molecules need to be overcome, so only a small amount of thermal energy is needed. The strong covalent bonds within each molecule are not broken.
(d) In graphite each carbon atom uses only three of its four valence electrons for bonding, so there is one delocalised electron per atom which can move freely along the layers as a mobile charge carrier. In diamond each carbon atom bonds to four other atoms and uses all four valence electrons, so there are no delocalised electrons to conduct.
(e) The zinc atoms are of a different size from the copper atoms, which disrupts the regular arrangement of the layers. The layers therefore cannot slide over one another easily, so a larger force is needed — brass is harder and stronger than pure copper.
★ Chapter Concept Map
Chemical Bonding and Structure — structure decides properties
Why atoms bondto reach a stable duplet or octet configuration with a filled valence shell, like the noble gases
ALWAYS AN
ATTRACTION
ATTRACTION
Between oppositely charged particlesionic: between ions · covalent: between nuclei and shared pairs · metallic: between cations and the sea of delocalised electrons
The three structures: giant ionic lattice · simple molecular (weak forces between molecules) · giant covalent (extensive network of strong bonds)
Ionic compoundhigh melting point, conducts only when molten or aqueous, usually soluble in water
WHAT IS BEING
OVERCOME?
OVERCOME?
Simple covalent substancelow melting point: only the weak forces between molecules are overcome, not the bonds inside them
Diamond and silicon dioxideall valence electrons used for bonding → hard, extremely high melting point, does not conduct
GIANT
COVALENT
COVALENT
Graphitethree of four valence electrons used → one delocalised electron per atom → conducts; weak forces between layers → slippery
Metalcations in layers inside a sea of delocalised electrons → conducts when solid or molten, and the layers slide so it bends
ALLOYS
ARE HARDER
ARE HARDER
Alloyatoms of different sizes disrupt the regular layers, so they cannot slide easily — harder and stronger than the pure metal