What you'll learn
Chemical bonding explains how atoms combine to form compounds and molecules. This topic is fundamental to understanding the properties of materials you encounter daily—from table salt used in Caribbean cooking to the aluminium in bauxite mined in Jamaica and Guyana. You'll explore three main bond types and learn to predict compound properties based on bonding.
Key terms and definitions
Ion — an atom or group of atoms that has gained or lost electrons, giving it a positive or negative charge
Ionic bond — the electrostatic force of attraction between oppositely charged ions, formed when electrons transfer from metal atoms to non-metal atoms
Covalent bond — a chemical bond formed when two non-metal atoms share one or more pairs of electrons
Metallic bond — the electrostatic attraction between positive metal ions and a 'sea' of delocalised electrons
Electrostatic attraction — the force of attraction between opposite charges (positive and negative)
Valency — the combining power of an element, determined by the number of electrons in its outer shell that can be lost, gained or shared
Molecule — a group of two or more atoms held together by covalent bonds
Lattice structure — a regular, repeating three-dimensional arrangement of particles (atoms, ions or molecules) in a solid
Core concepts
Electron arrangement and bonding
Atoms bond to achieve a stable electron arrangement. The most stable arrangement is a full outer shell, typically containing eight electrons (the octet rule), or two electrons for hydrogen and helium.
Elements in Group 1 have one electron in their outer shell and readily lose this electron to form 1+ ions. Group 2 elements have two outer electrons and form 2+ ions. Group 7 elements need one more electron to complete their outer shell and form 1- ions, while Group 6 elements form 2- ions.
The periodic table helps predict bonding behaviour:
- Metals (Groups 1, 2, 3 and transition metals) lose electrons to form positive ions
- Non-metals (Groups 5, 6, 7) gain or share electrons to form negative ions or covalent bonds
- Noble gases (Group 0/8) already have full outer shells and rarely form bonds
Ionic bonding
Ionic bonding occurs between metals and non-metals. The metal atom loses electrons to become a positive ion (cation), while the non-metal atom gains electrons to become a negative ion (anion). The strong electrostatic attraction between oppositely charged ions holds the compound together.
Formation of sodium chloride (NaCl):
- Sodium (2,8,1) loses one electron → Na⁺ (2,8)
- Chlorine (2,8,7) gains one electron → Cl⁻ (2,8,8)
- The Na⁺ and Cl⁻ ions attract to form NaCl
Formation of magnesium oxide (MgO):
- Magnesium (2,8,2) loses two electrons → Mg²⁺ (2,8)
- Oxygen (2,6) gains two electrons → O²⁻ (2,8)
- The Mg²⁺ and O²⁻ ions attract to form MgO
Ionic compounds form giant ionic lattice structures where each ion is surrounded by oppositely charged ions in a regular, repeating pattern. This gives ionic compounds characteristic properties:
Properties of ionic compounds:
- High melting and boiling points (strong electrostatic forces require much energy to break)
- Conduct electricity when molten or dissolved in water (ions are free to move and carry charge)
- Do not conduct electricity when solid (ions are fixed in position)
- Usually soluble in water
- Hard but brittle (layers of ions can slide, causing like charges to repel and the crystal to shatter)
Caribbean relevance: Common salt (sodium chloride) is produced in the Caribbean through solar evaporation of seawater in countries like Turks and Caicos, The Bahamas, and Bonaire. The ionic nature of NaCl makes it soluble in water and gives it the high melting point needed for the evaporation process.
Covalent bonding
Covalent bonding occurs between non-metal atoms that share pairs of electrons. Each shared pair constitutes one covalent bond. Atoms share electrons to achieve full outer shells.
Types of covalent bonds:
Single bond — one pair of shared electrons (e.g., H-H in hydrogen, Cl-Cl in chlorine)
Double bond — two pairs of shared electrons (e.g., O=O in oxygen, CO₂)
Triple bond — three pairs of shared electrons (e.g., N≡N in nitrogen)
Examples of covalent molecules:
Hydrogen (H₂):
- Each H atom has 1 electron
- They share 1 pair of electrons
- Both atoms achieve 2 electrons (full outer shell)
Water (H₂O):
- Oxygen has 6 outer electrons, needs 2 more
- Each hydrogen has 1 electron, needs 1 more
- Oxygen forms 2 single bonds with 2 hydrogen atoms
Carbon dioxide (CO₂):
- Carbon has 4 outer electrons
- Each oxygen has 6 outer electrons
- Carbon forms double bonds with 2 oxygen atoms: O=C=O
Ammonia (NH₃):
- Nitrogen has 5 outer electrons
- Forms 3 single bonds with 3 hydrogen atoms
Covalent substances exist as either simple molecular structures or giant covalent structures.
Simple molecular structures (like H₂O, CO₂, NH₃):
- Low melting and boiling points (weak forces between molecules, not within them)
- Do not conduct electricity (no charged particles)
- May be gases, liquids or soft solids at room temperature
- Often soluble in organic solvents, insoluble in water (except some like ammonia)
Giant covalent structures (like diamond, graphite, silicon dioxide):
- Very high melting and boiling points (many strong covalent bonds throughout structure)
- Usually hard and strong (diamond is the hardest natural substance)
- Do not conduct electricity (except graphite, which has delocalised electrons)
- Insoluble in water and organic solvents
Caribbean relevance: Natural gas (mainly methane, CH₄) is extracted in Trinidad and Tobago. Methane is a simple covalent molecule with low intermolecular forces, which is why it exists as a gas at room temperature and must be cooled or pressurised for transport.
Metallic bonding
Metallic bonding occurs in metals and alloys. Metal atoms lose their outer electrons, forming positive ions. The lost electrons form a "sea" of delocalised electrons that are free to move throughout the structure. The metallic bond is the electrostatic attraction between the positive metal ions and the negative electron sea.
Properties of metals:
- High melting and boiling points (strong metallic bonds)
- Conduct electricity (delocalised electrons move and carry charge)
- Conduct heat (electrons transfer kinetic energy efficiently)
- Malleable (can be hammered into shapes — layers of ions slide without breaking bonds)
- Ductile (can be drawn into wires)
- Lustrous (shiny appearance when polished)
Alloys are mixtures of metals (or metals with small amounts of non-metals). Atoms of different sizes disrupt the regular arrangement, making it harder for layers to slide. This makes alloys stronger and harder than pure metals.
Examples:
- Steel (iron + carbon) — stronger than pure iron
- Bronze (copper + tin) — used historically in Caribbean region
- Brass (copper + zinc) — used in musical instruments
Caribbean relevance: Bauxite ore mined in Jamaica, Guyana and Suriname is processed to extract aluminium. The metallic bonding in aluminium gives it excellent electrical conductivity and malleability, making it valuable for electrical wiring and aircraft construction. Trinidad's iron and steel industry also relies on understanding metallic bonding properties.
Comparing bond types
| Property | Ionic | Simple Covalent | Giant Covalent | Metallic |
|---|---|---|---|---|
| Elements involved | Metal + non-metal | Non-metal + non-metal | Non-metals | Metals only |
| Structure | Giant lattice | Small molecules | Giant network | Giant lattice |
| Melting/boiling point | High | Low | Very high | High |
| Electrical conductivity (solid) | No | No | No (except graphite) | Yes |
| Electrical conductivity (liquid) | Yes | No | N/A (doesn't melt easily) | Yes |
| Solubility in water | Usually soluble | Variable | Insoluble | Insoluble |
Formulae and valency
The chemical formula shows the ratio of atoms in a compound. For ionic compounds, this ratio balances the charges.
Rules for writing ionic formulae:
- Write the symbols with the metal first
- Balance charges so the total is zero
- Use subscripts to show the number of each ion
Examples:
- Sodium chloride: Na⁺ and Cl⁻ → NaCl (1:1 ratio)
- Magnesium chloride: Mg²⁺ and Cl⁻ → MgCl₂ (1:2 ratio to balance charges)
- Calcium oxide: Ca²⁺ and O²⁻ → CaO (1:1 ratio)
- Aluminium oxide: Al³⁺ and O²⁻ → Al₂O₃ (2:3 ratio gives 6+ and 6-)
Common ion charges to memorize:
- Group 1: 1+ (Na⁺, K⁺)
- Group 2: 2+ (Mg²⁺, Ca²⁺)
- Group 3: 3+ (Al³⁺)
- Group 6: 2- (O²⁻)
- Group 7: 1- (Cl⁻, Br⁻, I⁻)
- Common polyatomic ions: SO₄²⁻ (sulfate), NO₃⁻ (nitrate), CO₃²⁻ (carbonate), OH⁻ (hydroxide), NH₄⁺ (ammonium)
Worked examples
Example 1: Describing ionic bonding (4 marks)
Question: Describe how magnesium and chlorine atoms bond to form magnesium chloride. Include the electron arrangement of the atoms and ions formed.
Answer:
- Magnesium has electron arrangement 2,8,2 and loses 2 electrons to form Mg²⁺ with arrangement 2,8 ✓
- Chlorine has electron arrangement 2,8,7 and gains 1 electron to form Cl⁻ with arrangement 2,8,8 ✓
- Two chlorine atoms are needed for each magnesium atom to balance the charges / form MgCl₂ ✓
- The ionic bond is the electrostatic attraction between the oppositely charged Mg²⁺ and Cl⁻ ions ✓
Mark scheme notes: Award 1 mark for each bullet point. Accept alternative wording that conveys the same meaning. The charge balance must be mentioned explicitly.
Example 2: Predicting properties (3 marks)
Question: Sodium chloride has a melting point of 801°C, while methane has a melting point of -182°C. Explain this difference in terms of bonding and structure.
Answer:
- Sodium chloride has ionic bonding / is an ionic compound ✓
- It has a giant ionic lattice structure with strong electrostatic forces between oppositely charged ions ✓
- Methane has covalent bonding / is a simple molecular substance with weak intermolecular forces, so less energy is needed to separate the molecules ✓
Mark scheme notes: Must mention both types of bonding and relate to the strength of forces. Reference to structure type earns the mark.
Example 3: Drawing dot-and-cross diagrams (3 marks)
Question: Draw a dot-and-cross diagram to show the bonding in a molecule of ammonia (NH₃). Show only the outer shell electrons.
Answer:
- Nitrogen shown with 5 outer electrons (use dots or crosses consistently) ✓
- Three hydrogen atoms each shown with 1 electron (use opposite symbol from nitrogen) ✓
- Three pairs of electrons shared between nitrogen and hydrogen atoms, with nitrogen having one unshared pair ✓
Mark scheme notes: Electrons must be clearly shown as pairs between atoms. The unshared pair on nitrogen is essential for full marks.
Common mistakes and how to avoid them
Confusing atoms and ions: Remember that atoms are neutral, while ions have a charge. Always show charges on ions (e.g., Na⁺, not just Na). In ionic compounds, particles exist as ions, not atoms.
Thinking ionic compounds conduct electricity when solid: Ionic compounds only conduct when molten or dissolved because ions must be free to move. In solid form, ions are fixed in the lattice.
Mixing up "melting" and "boiling" in bonding explanations: Be specific—melting breaks the lattice structure but molecules may remain intact (for ionic/metallic substances). For covalent molecules, melting separates molecules but doesn't break covalent bonds within them.
Incorrectly balancing formulae: Always balance charges in ionic compounds. Check your formula: in MgCl₂, the 2+ charge on Mg²⁺ is balanced by two Cl⁻ ions (total 2-). Practice with different combinations.
Confusing the strength of covalent bonds with intermolecular forces: Covalent bonds (within molecules) are strong, but forces between simple covalent molecules are weak. This explains why water boils easily (weak intermolecular forces) but doesn't decompose into hydrogen and oxygen (strong O-H covalent bonds).
Failing to explain properties using bonding: When asked to "explain" a property, you must link it to the type of bonding and structure. Simply naming the bond type earns minimal marks. For example: "high melting point because ionic bonding creates strong electrostatic forces throughout the giant lattice structure."
Exam technique for "Chemical Bonding"
Command words matter: "State" requires a simple fact (1 mark). "Describe" needs a statement about what happens. "Explain" requires reasoning using bonding theory—always link properties to bonding type and structure (2-3 marks typically).
Drawing diagrams: For dot-and-cross diagrams, use dots for one element and crosses for another. Be consistent. Show only outer shell electrons unless specified. Shared pairs must be clearly shown between atoms. Brackets and charges are essential for ionic diagrams.
Multi-step explanations: For properties of compounds, use a three-step approach: (1) identify bond type, (2) describe structure, (3) link to property. Example: "NaCl has ionic bonding → forms giant ionic lattice → strong forces require high temperature to break → high melting point."
Allocate time by marks: Each mark typically requires one distinct point. A 4-mark question needs four separate facts or explanations. Don't write paragraphs for 1-mark questions, but do provide sufficient detail for higher-mark questions.
Quick revision summary
Chemical bonds form when atoms achieve stable electron arrangements. Ionic bonding (metal + non-metal) involves electron transfer and creates high-melting-point compounds that conduct when molten. Covalent bonding (non-metal + non-metal) involves electron sharing; simple molecular compounds have low melting points, while giant covalent structures have very high melting points. Metallic bonding creates malleable, conductive materials through delocalised electrons. Always link properties to bonding type and structure. Practice writing formulae by balancing ion charges and drawing dot-and-cross diagrams showing only outer electrons.