(I once asked a class, "Why does sodium react so violently with chlorine, but the product — table salt — is so safe you can eat it?" No one could answer. Then I said, "Because when they bond, they become something entirely new. The violence is the price of stability." That is the whole story of chemical bonding. Two unstable things come together and create something stable, something useful, something completely different from either of them. And that is happening all around you, all the time.) Think about the last meal you ate. The salt on your chips is a compound of sodium and chlorine — one a violent metal, the other a poisonous gas. The water you drank is a compound of hydrogen and oxygen — one explosive, the other essential for combustion. The sugar in your tea is a compound of carbon, hydrogen, and oxygen — none of which, on their own, tastes sweet. Bonding is the great transformer of chemistry. It takes the elements of the periodic table — most of which are unstable on their own — and combines them into the thousands of stable compounds that make up our world. In this chapter, we're going to learn why atoms bond, how they bond, and what kind of compounds they form. By the end, you'll look at every substance around you and see the bonds that hold it together. Let's begin.
Atoms bond because they want to be stable. And what does "stable" mean for an atom? It means having a full outer shell of electrons — the same arrangement as a noble gas.
Noble gases (like helium, neon, argon) have full outer shells. They are unreactive because they don't need to gain or lose electrons. But most other atoms do not have full outer shells. They are unstable. They react with other atoms — they form bonds — to achieve the noble gas configuration.
There are three main ways atoms can achieve a full outer shell:
(Here's the simplest way to think about it. Atoms are like people at a party. Some are desperate to give away what they have — those are the metals. Some are desperate to receive — those are the non-metals. And some are happy to share — that's how covalent bonds work. The whole point is that everyone wants to end up with a full set of eight — the "octet rule.")
Most atoms are most stable when they have eight electrons in their outermost shell. This is called the octet rule. (Hydrogen and helium are exceptions — they are stable with two electrons in their first shell, which is full.)
Atoms achieve an octet by:
An ionic bond is formed when a metal transfers one or more electrons to a non-metal. The metal becomes a positive ion (cation), and the non-metal becomes a negative ion (anion). The oppositely charged ions are held together by strong electrostatic attraction — this attraction is the ionic bond.
Sodium (2,8,1) loses its one outer electron to form Na⁺ (2,8). Chlorine (2,8,7) gains that electron to form Cl⁻ (2,8,8). Both ions now have full outer shells — Na⁺ has the configuration of neon, and Cl⁻ has the configuration of argon. The oppositely charged ions attract each other, forming a giant ionic lattice.
Na → Na⁺ + e⁻
Cl + e⁻ → Cl⁻
Na⁺ + Cl⁻ → NaCl
Magnesium (2,8,2) loses two electrons to form Mg²⁺ (2,8). Oxygen (2,6) gains two electrons to form O²⁻ (2,8). The charges balance: one Mg²⁺ for every O²⁻.
Mg → Mg²⁺ + 2e⁻
O + 2e⁻ → O²⁻
Mg²⁺ + O²⁻ → MgO
| Property | Explanation |
|---|---|
| High melting and boiling points | Strong electrostatic forces between ions in the lattice require a lot of energy to break. |
| Solid at room temperature | The ions are locked in a giant lattice. |
| Brittle | When a force is applied, ions of the same charge are forced next to each other and repel, causing the crystal to shatter. |
| Do not conduct electricity when solid | Ions are fixed in the lattice and cannot move. |
| Conduct electricity when molten or dissolved | Ions are free to move and carry charge. |
| Usually soluble in water | Water molecules pull the ions away from the lattice. |
A covalent bond is formed when two non-metal atoms share a pair of electrons. Both atoms contribute one electron to the shared pair, and both achieve a full outer shell.
Each hydrogen atom has one electron. They share their electrons to form a single covalent bond. Both now have two electrons in their outer shell — a full first shell.
H + H → H₂ (with a shared pair of electrons)
Each chlorine atom has seven outer electrons. They share one pair, so both achieve eight — a full outer shell.
Oxygen has six outer electrons. It needs two more to complete its octet. Each hydrogen has one electron. Oxygen shares one pair with each hydrogen, forming two single covalent bonds.
Carbon has four outer electrons. It needs four more. Each hydrogen has one. Carbon shares one pair with each of four hydrogens, forming four single covalent bonds.
Carbon has four outer electrons. Each oxygen has six. Carbon shares two pairs with each oxygen, forming two double covalent bonds. This gives carbon eight electrons and each oxygen eight electrons.
| Property | Explanation |
|---|---|
| Low melting and boiling points | Weak intermolecular forces between molecules — little energy needed to separate them. |
| Often gases or liquids at room temperature | Because of the weak forces between molecules. |
| Do not conduct electricity | No free ions or electrons to carry charge. |
| Usually insoluble in water (but soluble in organic solvents) | Non-polar molecules do not interact strongly with polar water. |
(Important distinction: in a covalent compound, the covalent bonds within the molecule are strong. But the forces between molecules are weak. That's why covalent compounds have low melting and boiling points — you're not breaking the bonds, you're just separating the molecules from each other. Think of it like a group of friends holding hands. The hand-holding is strong, but the friends can still be separated from other groups.)
Some covalent substances don't exist as small molecules. Instead, they form giant covalent structures — networks of atoms held together by covalent bonds extending throughout the entire material. These substances have very different properties from simple covalent compounds.
Diamond is a form of carbon where each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral arrangement. This creates a rigid, three-dimensional network.
| Property | Explanation |
|---|---|
| Extremely hard | All atoms are held by strong covalent bonds — no easy way to break the structure. |
| Very high melting point | Strong covalent bonds throughout the structure require enormous energy to break. |
| Does not conduct electricity | All outer electrons are locked in covalent bonds — no free electrons. |
| Used in cutting tools and jewellery | Because of its hardness and brilliance. |
Graphite is another form of carbon. Each carbon atom is covalently bonded to three other carbon atoms in flat hexagonal layers. The fourth electron is delocalised — free to move between the layers.
| Property | Explanation |
|---|---|
| Soft and slippery | Layers are held by weak forces and can slide over each other. |
| Conducts electricity | Delocalised electrons are free to move and carry charge. |
| High melting point | Covalent bonds within layers are strong. |
| Used in pencils and as a lubricant | Because of its softness and slipperiness. |
| Used as electrodes | Because it conducts electricity and is unreactive. |
| Feature | Diamond | Graphite |
|---|---|---|
| Bonding | Each C bonded to 4 others (tetrahedral) | Each C bonded to 3 others (hexagonal layers) |
| Structure | Rigid 3D network | Layers held by weak forces |
| Hardness | Extremely hard | Soft and slippery |
| Electrical conductivity | Does not conduct | Conducts (delocalised electrons) |
| Uses | Cutting tools, jewellery | Pencils, lubricants, electrodes |
| Property | Ionic Compounds | Simple Covalent Compounds |
|---|---|---|
| Formed between | Metals and non-metals | Non-metals and non-metals |
| Bond type | Transfer of electrons | Sharing of electrons |
| Particles present | Ions | Molecules |
| Melting/boiling point | High | Low |
| State at room temp | Solid | Gas, liquid, or low-melting solid |
| Electrical conductivity (solid) | Does not conduct | Does not conduct |
| Electrical conductivity (molten/aqueous) | Conducts | Does not conduct |
| Solubility in water | Usually soluble | Usually insoluble |
| Examples | NaCl, MgO, CaCO₃ | H₂O, CO₂, CH₄, Cl₂ |
Metals also bond — but in a different way. In a metal, the outer electrons are delocalised — they are free to move throughout the entire structure. The metal consists of a lattice of positive ions immersed in a "sea" of delocalised electrons. The attraction between the positive ions and the delocalised electrons is the metallic bond.
This explains the properties of metals:
(Picture a metal as a swimming pool full of positive ions, with a "sea" of electrons flowing around them. That's metallic bonding. It's why metals conduct, why they bend instead of shatter, and why they feel cool to the touch. It's a beautiful picture once you see it.)
Question: Predict the type of bond formed between: (a) sodium and chlorine, (b) hydrogen and oxygen, (c) magnesium and oxygen.
Solution:
(a) Sodium (metal) + chlorine (non-metal) → ionic bond (NaCl).
(b) Hydrogen (non-metal) + oxygen (non-metal) → covalent bond (H₂O).
(c) Magnesium (metal) + oxygen (non-metal) → ionic bond (MgO).
Question: Draw a dot-and-cross diagram for sodium chloride, showing the transfer of electrons.
Solution: Sodium has one electron in its outer shell (shown as a dot). Chlorine has seven (shown as crosses). The sodium dot is transferred to the chlorine, so chlorine now has eight electrons in its outer shell. The diagram shows Na⁺ with no outer electrons and Cl⁻ with eight (seven crosses plus one dot). The charges are shown as + and −.
Question: Explain why sodium chloride conducts electricity when molten but not when solid.
Solution: In solid sodium chloride, the ions are held tightly in a giant lattice and cannot move. There are no free charge carriers, so it does not conduct. When molten, the lattice breaks down and the ions are free to move. The charged ions carry the current, so molten sodium chloride conducts electricity.
1. Why do atoms form chemical bonds?
Atoms form bonds to achieve a stable, full outer shell of electrons — usually the same configuration as a noble gas.
2. Describe the formation of an ionic bond.
An ionic bond forms when a metal transfers one or more electrons to a non-metal. The metal becomes a positive ion and the non-metal becomes a negative ion. The oppositely charged ions attract each other strongly.
3. Describe the formation of a covalent bond.
A covalent bond forms when two non-metal atoms share a pair of electrons. Both atoms contribute one electron to the shared pair and both achieve a full outer shell.
4. Why do ionic compounds have high melting points?
Ionic compounds have high melting points because the strong electrostatic forces between the ions in the giant lattice require a lot of energy to break.
5. Why do simple covalent compounds have low boiling points?
Simple covalent compounds have weak intermolecular forces between molecules. Only a small amount of energy is needed to separate the molecules, so boiling points are low.
6. Why does graphite conduct electricity but diamond does not?
In graphite, each carbon atom is bonded to only three others, leaving one delocalised electron per atom. These delocalised electrons are free to move and carry charge. In diamond, all four outer electrons are locked in covalent bonds, so there are no free electrons to carry charge.
7. Describe metallic bonding and state one property it explains.
Metallic bonding consists of a lattice of positive ions immersed in a sea of delocalised electrons. The attraction between the ions and the electrons is the metallic bond. This explains why metals conduct electricity — the delocalised electrons are free to move.
Question 1 [10 marks]
(a) Explain why atoms form chemical bonds. [2]
(b) Describe the formation of an ionic bond between sodium and chlorine, using electron transfer. [4]
(c) State two properties of ionic compounds and explain each in terms of their structure. [4]
(a) Atoms form bonds to achieve a stable, full outer shell of electrons — usually the same configuration as a noble gas. (2 marks: 1 for "full outer shell", 1 for "stability".)
(b) Sodium (2,8,1) loses its one outer electron to form Na⁺ (2,8). Chlorine (2,8,7) gains that electron to form Cl⁻ (2,8,8). Both ions now have full outer shells. The oppositely charged ions attract each other strongly, forming an ionic bond. (4 marks: 1 for electron loss, 1 for electron gain, 1 for full shells, 1 for electrostatic attraction.)
(c) Two of: High melting point — strong electrostatic forces in the lattice require lots of energy to break. Brittle — when a force is applied, ions of the same charge are forced together and repel. Does not conduct when solid — ions are fixed in the lattice. Conducts when molten or aqueous — ions are free to move. (4 marks: 2 for each property + explanation.)
Question 2 [10 marks]
(a) Describe the formation of a covalent bond, using hydrogen as an example. [3]
(b) Draw a dot-and-cross diagram for a water molecule (H₂O). [3]
(c) Explain why simple covalent compounds have low melting and boiling points. [4]
(a) A covalent bond forms when two non-metal atoms share a pair of electrons. In hydrogen, each H atom has one electron. They share their electrons, so both achieve a full first shell of two electrons. (3 marks: 1 for sharing, 1 for each H contributing one electron, 1 for full shell.)
(b) Oxygen (six outer electrons) shares one pair with each hydrogen (one outer electron). The diagram shows oxygen with two shared pairs (one to each H) and two lone pairs. Each hydrogen has one shared pair. (3 marks: 1 for correct bonds, 1 for lone pairs on oxygen, 1 for correct electron count.)
(c) Simple covalent compounds consist of small molecules held together by weak intermolecular forces. These forces are much weaker than covalent bonds, so only a small amount of energy is needed to separate the molecules. This is why melting and boiling points are low. The covalent bonds within the molecules remain intact — it is the forces between molecules that break. (4 marks: 1 for "weak intermolecular forces", 1 for "little energy needed", 1 for "bonds within molecule stay intact", 1 for clarity.)
Question 3 [10 marks]
(a) Compare the structures of diamond and graphite. [4]
(b) Explain why graphite conducts electricity but diamond does not. [3]
(c) Describe metallic bonding and explain why metals are malleable. [3]
(a) Diamond: each carbon atom is covalently bonded to four others in a rigid three-dimensional tetrahedral network. Graphite: each carbon atom is covalently bonded to three others in flat hexagonal layers, with weak forces between the layers. (4 marks: 2 for diamond, 2 for graphite.)
(b) In graphite, each carbon atom has one delocalised electron that is free to move between the layers. These delocalised electrons carry charge, so graphite conducts. In diamond, all four outer electrons are locked in covalent bonds, so there are no free electrons to carry charge. (3 marks: 1 for delocalised electrons in graphite, 1 for no free electrons in diamond, 1 for link to conductivity.)
(c) Metallic bonding consists of a lattice of positive ions immersed in a sea of delocalised electrons. The attraction between the ions and the electrons holds the structure together. Metals are malleable because the layers of positive ions can slide over each other without breaking the metallic bond — the delocalised electrons continue to hold the structure together. (3 marks: 1 for structure, 1 for sliding layers, 1 for "bond not broken".)
Challenge 1: Magnesium oxide (MgO) and sodium chloride (NaCl) are both ionic compounds. Explain why MgO has a much higher melting point than NaCl.
MgO has a higher melting point because the ions in MgO carry higher charges: Mg²⁺ and O²⁻ compared to Na⁺ and Cl⁻. The electrostatic force of attraction between ions is proportional to the product of their charges. Since the charges in MgO are double those in NaCl, the attraction between Mg²⁺ and O²⁻ is much stronger than between Na⁺ and Cl⁻. More energy is needed to break these stronger forces, so MgO has a higher melting point. Additionally, the ions in MgO are smaller, which brings them closer together and further strengthens the attraction.
Challenge 2: Explain why carbon dioxide (CO₂) is a gas at room temperature, while silicon dioxide (SiO₂) is a solid with a very high melting point, even though both are covalent compounds.
CO₂ is a simple covalent molecule. It consists of small, discrete molecules held together by weak intermolecular forces. Only a small amount of energy is needed to separate these molecules, so CO₂ is a gas at room temperature. SiO₂, on the other hand, is a giant covalent structure. Each silicon atom is covalently bonded to four oxygen atoms, and each oxygen is bonded to two silicons, forming a continuous three-dimensional network. To melt SiO₂, you must break strong covalent bonds throughout the entire structure — this requires an enormous amount of energy, so SiO₂ has a very high melting point. The key difference is simple molecules vs giant covalent structure.
Challenge 3: A student says, "Ionic compounds conduct electricity because they contain ions." Is this statement complete? Explain, referring to the states in which ionic compounds do and do not conduct.
The statement is incomplete. While it is true that ionic compounds contain ions, this alone does not mean they conduct electricity. In the solid state, the ions are held tightly in a giant lattice and cannot move. Without moving charge carriers, there is no conduction. Ionic compounds only conduct electricity when molten or dissolved in water, because in these states the lattice breaks down and the ions are free to move. So the correct statement is: ionic compounds conduct electricity when molten or aqueous because their ions are free to move, but they do not conduct when solid because the ions are fixed in the lattice.
"Bonding is the bridge between atoms and the world. Learn how atoms connect, and you will understand why matter behaves the way it does."
