Chemical bonding: ionic, covalent and metallic bonding

What you'll learn

Chemical bonding explains how atoms join together to form compounds and why different substances have different properties. You'll study three types of bonding: ionic bonding (between metals and non-metals), covalent bonding (between non-metals), and metallic bonding (in metals). Understanding bonding is essential for predicting properties, drawing structures and explaining the behaviour of materials in everyday life and industrial processes.

Key terms and definitions

Ionic bond — the electrostatic force of attraction between oppositely charged ions, formed when electrons are transferred from metal atoms to non-metal atoms

Covalent bond — a shared pair of electrons between two non-metal atoms

Metallic bond — the electrostatic attraction between positive metal ions and delocalised electrons

Ion — a charged particle formed when an atom loses or gains electrons

Electron transfer — the process by which electrons move from one atom to another, creating ions in ionic bonding

Delocalised electrons — electrons that are not associated with a particular atom or bond but are free to move throughout a structure

Giant structure — a three-dimensional network of atoms or ions extending throughout the substance, containing millions of particles

Simple molecular structure — a structure consisting of small molecules held together by weak intermolecular forces

Core concepts

Ionic bonding

Ionic bonding occurs between metal and non-metal atoms. During ionic bond formation, metal atoms lose electrons to form positive ions (cations), while non-metal atoms gain electrons to form negative ions (anions). Both types of ions achieve stable electron configurations, usually with full outer shells like the nearest noble gas.

Formation of ionic compounds:

• Metal atoms have few electrons in their outer shell and lose these electrons easily
• Non-metal atoms have nearly full outer shells and gain electrons to complete them
• The electrostatic attraction between the oppositely charged ions forms the ionic bond
• For example, sodium (electronic structure 2,8,1) loses one electron to form Na⁺ (2,8)
• Chlorine (electronic structure 2,8,7) gains one electron to form Cl⁻ (2,8,8)
• These ions attract to form sodium chloride, NaCl

Dot and cross diagrams for ionic compounds:

You must be able to draw dot and cross diagrams showing:

• The electronic structure of the atoms before bonding
• The transfer of electrons (shown as crosses from one atom, dots from another)
• The charges on the resulting ions
• Square brackets around each ion with the charge outside

Common ionic compounds to know:

• Sodium chloride (NaCl): Na⁺ and Cl⁻
• Magnesium oxide (MgO): Mg²⁺ and O²⁻
• Calcium chloride (CaCl₂): Ca²⁺ and two Cl⁻

Properties of ionic compounds:

• High melting and boiling points due to strong electrostatic forces between ions requiring large amounts of energy to overcome
• Solid at room temperature with giant ionic lattice structures
• Conduct electricity when molten or dissolved in water because ions are free to move and carry charge
• Do not conduct electricity when solid because ions are fixed in position
• Often soluble in water because water molecules can attract and separate the ions
• Brittle — layers of ions can slide and like charges repel, causing the structure to break

Covalent bonding

Covalent bonding occurs between non-metal atoms. Atoms share pairs of electrons to achieve full outer shells. Each shared pair of electrons forms one covalent bond.

Formation of covalent bonds:

• Non-metal atoms need to gain electrons to fill their outer shell
• Instead of transferring electrons, non-metals share electrons
• Each atom contributes one electron to the shared pair
• The shared electrons count towards the outer shell of both atoms
• The positive nuclei of both atoms are attracted to the shared electrons, holding the atoms together

Dot and cross diagrams for covalent molecules:

You must be able to draw dot and cross diagrams for simple molecules:

• Hydrogen (H₂): single covalent bond, each H shares one electron
• Chlorine (Cl₂): single covalent bond, each Cl shares one electron
• Water (H₂O): oxygen forms two single bonds with two hydrogen atoms
• Ammonia (NH₃): nitrogen forms three single bonds with three hydrogen atoms
• Methane (CH₄): carbon forms four single bonds with four hydrogen atoms
• Oxygen (O₂): double covalent bond (two shared pairs)
• Carbon dioxide (CO₂): carbon forms two double bonds with two oxygen atoms

Simple molecular structures:

Most covalent substances exist as simple molecules with these properties:

• Low melting and boiling points because weak intermolecular forces between molecules require little energy to overcome
• Often gases or liquids at room temperature
• Do not conduct electricity because molecules are neutral with no free charges
• Insoluble in water (usually) but may dissolve in organic solvents

Giant covalent structures:

Some covalent substances form giant structures where millions of atoms are joined by covalent bonds:

Diamond:

• Each carbon atom forms four covalent bonds in a tetrahedral arrangement
• Extremely hard — all bonds must be broken to break the structure
• Very high melting point
• Does not conduct electricity — no free electrons
• Used in cutting tools and drill bits

Graphite:

• Each carbon atom forms three covalent bonds in layers
• Soft and slippery — layers can slide over each other with only weak forces between them
• Very high melting point — strong covalent bonds within layers
• Conducts electricity — one electron per carbon atom is delocalised and free to move
• Used in pencils and as a lubricant

Silicon dioxide (silica):

• Each silicon atom bonds to four oxygen atoms, each oxygen bonds to two silicon atoms
• Giant structure similar to diamond
• Very hard with a very high melting point
• Main component of sand and glass

Metallic bonding

Metallic bonding occurs in metallic elements and alloys. Metal atoms lose their outer shell electrons to form positive ions. These electrons become delocalised and move freely throughout the structure, forming a "sea of electrons" around the positive metal ions.

Structure of metals:

• Metal atoms are arranged in a regular giant structure
• Outer electrons are delocalised and free to move throughout the structure
• The metallic bond is the strong electrostatic attraction between positive metal ions and the sea of delocalised electrons
• This bonding extends throughout the metal

Properties of metals explained by metallic bonding:

• Good electrical conductors — delocalised electrons are free to move and carry charge through the structure
• Good thermal conductors — delocalised electrons transfer kinetic energy quickly through the structure
• High melting and boiling points — strong metallic bonds require large amounts of energy to overcome (though lower than giant ionic or covalent structures)
• Malleable and ductile — layers of metal ions can slide over each other while maintaining metallic bonding, so metals can be hammered into shape or drawn into wires without breaking
• Shiny appearance — delocalised electrons reflect light

Alloys:

Alloys are mixtures of metals or metals with other elements:

• Different sized atoms disrupt the regular arrangement
• Layers cannot slide over each other as easily
• Alloys are harder than pure metals
• Examples: steel (iron with carbon), brass (copper with zinc), bronze (copper with tin)

Comparing the three types of bonding

Predicting bonding type

Use the periodic table to predict bonding:

• Metal + non-metal → ionic bonding (e.g., NaCl, MgO, CaCl₂)
• Non-metal + non-metal → covalent bonding (e.g., H₂O, CO₂, CH₄)
• Metal only → metallic bonding (e.g., Cu, Fe, Al)

Worked examples

Example 1: Ionic bonding in magnesium oxide

Question: Magnesium reacts with oxygen to form magnesium oxide, MgO. (a) Describe what happens to the electrons when magnesium reacts with oxygen. [2 marks] (b) Explain why magnesium oxide has a high melting point. [2 marks]

Answer: (a)

• Magnesium atoms lose two electrons to form Mg²⁺ ions [1]
• Oxygen atoms gain two electrons to form O²⁻ ions [1]

(b)

• Magnesium oxide has a giant ionic lattice structure [1]
• Strong electrostatic forces between oppositely charged ions require large amounts of energy to overcome [1]

Examiner tip: Always mention both the structure AND the strength of forces when explaining melting points.

Example 2: Covalent bonding in water

Question: Water, H₂O, contains covalent bonds. (a) What is a covalent bond? [1 mark] (b) Draw a dot and cross diagram to show the bonding in a water molecule. Show only the outer shell electrons. [2 marks] (c) Explain why water has a low boiling point. [2 marks]

Answer: (a) A shared pair of electrons (between two non-metal atoms) [1]

(b)

• Correct diagram showing oxygen with 6 outer electrons (as dots), two hydrogen atoms each with 1 electron (as crosses) [1]
• Two shared pairs correctly shown between O and each H [1]

(c)

• Water consists of simple molecules [1]
• Weak intermolecular forces between molecules require little energy to overcome [1]

Examiner tip: Don't confuse the strong covalent bonds within molecules with the weak forces between molecules.

Example 3: Properties of metals

Question: Copper is used for electrical wiring. (a) Describe the structure and bonding in copper. [3 marks] (b) Explain why copper is a good electrical conductor. [2 marks]

Answer: (a)

• Giant structure of copper atoms/ions arranged regularly [1]
• Outer electrons are delocalised/free to move [1]
• Strong electrostatic attraction between positive ions and delocalised electrons [1]

(b)

• Copper contains delocalised electrons [1]
• These electrons are free to move through the structure and carry charge [1]

Examiner tip: Always link the property to the structure — explain HOW the structure causes the property.

Common mistakes and how to avoid them

• Confusing ionic and covalent bonding — remember that ionic bonding involves electron transfer and occurs between metals and non-metals, while covalent bonding involves electron sharing between non-metals

• Saying "molecules" for ionic compounds — ionic compounds do not exist as molecules; they form giant lattice structures. Never refer to "molecules of sodium chloride"

• Confusing intermolecular forces with covalent bonds — the bonds within simple covalent molecules are strong, but the forces between molecules are weak. This explains why simple molecular substances have low melting points despite having strong bonds

• Incomplete explanations of properties — always explain properties in terms of structure AND bonding. For example, don't just say "high melting point because of strong bonds" — specify what type of structure and what forces need to be overcome

• Poor dot and cross diagrams — ensure you show only outer shell electrons, use dots for one atom and crosses for the other consistently, and show square brackets with charges for ions

• Forgetting that graphite conducts electricity — graphite is the exception among giant covalent structures because it has delocalised electrons. This is a common exam question

Exam technique for "Chemical bonding: ionic, covalent and metallic bonding"

• "Describe" questions require you to state observations or facts without explanation. For bonding, state what happens to electrons (transferred, shared, delocalised) and what particles form (ions, molecules)

• "Explain" questions need both a statement and a reason linked together. For properties, always link the structure type to the forces present and how much energy is needed. Use "because" or "therefore" to connect your points clearly

• Drawing dot and cross diagrams usually awards 1 mark for correct number of electrons and 1 mark for correct sharing/transfer shown. Show only outer shell electrons unless specified otherwise. For ionic compounds, use square brackets and show charges

• Command word "compare" means you must make direct comparisons between two things, not describe them separately. Use comparative language: "higher than," "whereas," "in contrast to"

Quick revision summary

Chemical bonding explains how atoms join together. Ionic bonding occurs between metals and non-metals through electron transfer, forming ions held by electrostatic attraction in giant structures with high melting points. Covalent bonding occurs between non-metals through electron sharing; simple molecular substances have low melting points while giant covalent structures have very high melting points. Metallic bonding involves positive metal ions in a sea of delocalised electrons, explaining why metals conduct electricity and are malleable. Predict bonding type using the periodic table: metal + non-metal = ionic, non-metal + non-metal = covalent, metal only = metallic.