Ionic, Covalent and Metallic Bonding

• newEditor
• Chemistry
• Upper Secondary (Key Stage 4)
• GCSE

Why Atoms Bond

This GCSE chemistry lesson explores Ionic, Covalent and Metallic Bonding. Students compare how ionic, covalent and metallic bonds form and link bonding and structure to melting point, conductivity and other material properties. It supports KS4 classroom teaching, independent study and exam revision using clear, subject-specific vocabulary.

Atomic Bonding, Types of Bonding, KS4 Chemistry, GCSE Chemistry, Ionic Bonding, Covalent Bonding, Metallic Bonding, Chemical Bonding, Electrostatic Attraction, Shared Pair of Electrons, Delocalised Electrons, Lattice Structure, Giant Ionic Lattice, Giant Covalent, Simple Molecules, MCQs, Open Questions, Polls, Emoji Polls, Exit Tickets

Learning goal

I can explain how and why atoms bond to become more stable, and distinguish between ionic, covalent and metallic bonding.

What causes ions to bond?

Use this to surface prior ideas such as electrons, sharing, transfer, stability, ions, metals and molecules. Pick a few responses to connect back to the learning goal.

Bonding starts with outer electrons

Chemical bonding depends on outer-shell electrons.

Atoms bond because a full outer shell is more stable than an incomplete one.

To reach a more stable arrangement, atoms may:

• lose electrons

• gain electrons

• share electrons

Keep this at GCSE level: the main idea is stability through fuller outer shells. Avoid overcomplicating with advanced orbital language.

Three main ways atoms bond

• Ionic bonding: electrons are transferred

• Covalent bonding: electrons are shared

• Metallic bonding: outer electrons become delocalised

The type of atoms involved helps decide which bond forms.

Emphasise that these are different particle-level models, not just labels to memorise.

Stability and electron arrangement

A stable atom usually has a full outer shell.

For many of the elements students meet at this level, bonding can be understood as atoms changing their outer electrons until they reach a noble-gas-like arrangement.

This is why bonding is closely linked to electron arrangement.

You can mention that this is a useful model for GCSE and works well for the examples in this lesson.

Bond type depends on the atoms

• Metals + non-metals usually form ionic bonds

• Non-metals + non-metals usually form covalent bonds

• Metals usually bond metallically

These patterns help us predict structure and properties.

This slide prepares students to connect bonding type with examples and later with properties.

Ions form by electron transfer

In ionic bonding, electrons move from a metal atom to a non-metal atom.

• the metal atom loses electrons and becomes a positive ion

• the non-metal atom gains electrons and becomes a negative ion

The ions now have more stable electron arrangements.

Use sodium and chlorine as the first example, since students will meet sodium chloride later in the lesson.

Why do atoms bond in this GCSE model?

A

To fill outer shells

B

To empty the nucleus

C

To change element

D

To split protons

Tests the core idea from slides 2 to 5: bonding is explained through outer electrons and gaining a more stable full outer shell.

Complete the summary

Chemical bonding depends on outer-shell electrons. Atoms bond because a full outer shell is more stable. They may lose, gain or share electrons to reach this arrangement.

outer-shell electrons

outer shell

share

protons

inner shell

delocalise

nucleus

Echoes the exact phrasing from the opening teaching block so students secure the key vocabulary before moving into bond types.

How ionic bonding forms

metal loses electron(s) → positive ion forms

non-metal gains electron(s) → negative ion forms

Opposite charges attract, so the ions bond strongly.

This electrostatic attraction is an ionic bond.

Stress that the bond is the attraction between ions, not the electron transfer itself.

Sodium chloride as an example

• Sodium transfers one electron to chlorine

• Sodium becomes Na⁺

• Chlorine becomes Cl⁻

• The ions attract and form sodium chloride

This is a typical ionic compound made from a metal and a non-metal.

If useful, briefly sketch electron arrangements for Na and Cl to show why each ion is more stable.

Giant ionic lattices

Ionic compounds form a giant lattice of alternating positive and negative ions.

• the pattern repeats in all directions

• each ion is attracted to ions of opposite charge around it

• many strong attractions must be overcome to melt the solid

Make the link between structure and properties explicit: giant structure, many strong attractions, high melting point.

Properties of ionic compounds

Ionic compounds usually have:

• high melting points because of strong attractions in the lattice

• electrical conductivity when molten or dissolved because ions can move

• no conductivity when solid because ions are fixed in place

• often good solubility in water

Avoid saying all ionic compounds are soluble; 'often' is safer and still matches GCSE expectations.

What is an ionic bond?

A

A shared pair

B

A proton transfer

C

Electrostatic attraction

D

A moving electron

Targets the common misconception flagged on slide 7: the bond is the attraction between oppositely charged ions, not the electron transfer itself.

Ionic bonding and properties

In ionic bonding, a metal loses electrons and a non-metal gains them. This forms positive and negative ions. Ionic compounds have high melting points and conduct electricity when molten or dissolved because the ions can move.

loses

ions

melting points

shares

atoms

low melting points

molecules

Checks mechanism plus structure-property link from slides 6 to 10, using sodium chloride as the nearby model example.

Covalent bonding is sharing

In covalent bonding, non-metal atoms share pairs of electrons.

The shared electrons count towards the outer shells of both atoms, so each atom can reach a more stable arrangement.

A covalent bond is the shared pair of electrons.

Students often confuse transfer and sharing. Contrast covalent bonding clearly with ionic bonding.

Simple molecules

Some covalently bonded substances form simple molecules.

Examples include:

• carbon dioxide

• water

• oxygen
  

Within each molecule, the covalent bonds are strong.

You can mention that the forces between molecules are much weaker than the covalent bonds inside a molecule.

Why simple molecules have low melting points

Simple molecular substances often have low melting and boiling points.

This is because melting usually does not break the covalent bonds inside the molecules.

Instead, it overcomes the weaker forces between molecules.

Use carbon dioxide as the example named in the prompt. Keep the focus on structure causing property.

Giant covalent structures

Some covalent substances do not form small molecules.

They form giant covalent structures, where atoms are joined by a network of strong covalent bonds throughout the whole solid.

This is an important contrast with simple molecules: same bonding type, different structure, different properties.

Diamond and graphite

Both diamond and graphite are forms of carbon with giant covalent structures, but their arrangements are different.

• Diamond: each carbon bonds to four others

• Graphite: each carbon bonds to three others in layers

Different structures lead to different properties.

Flag that students should not assume same element means same properties; structure matters.

Properties of diamond and graphite

• Diamond has a very high melting point and does not conduct electricity

• Graphite has a high melting point and does conduct electricity

Diamond is hard because of its strong 3D network.

Graphite conducts because it has delocalised electrons within its layers.

Students often expect all covalent substances not to conduct. Graphite is the key exception in this lesson.

Why does carbon dioxide have a low melting point?

A

Weak forces between molecules

B

Weak bonds inside molecules

C

Free-moving ions

D

Delocalised electrons

Uses the taught example of carbon dioxide. Distractors reflect confusion with ionic and metallic bonding, and the idea that covalent bonds themselves are weak.

Covalent structures

In covalent bonding, non-metal atoms share pairs of electrons. Some substances form simple molecules, while others form giant covalent structures such as diamond and graphite.

share

molecules

giant covalent structures

transfer

lattices

metallic bonds

ions

Reinforces the key contrast across slides 11 to 16: same bonding type, different structures, different properties.

Why does graphite conduct but diamond does not?

Expected elements: graphite conducts because it has delocalised electrons within its layers; diamond does not because it has no free electrons to move. Accept concise explanations in students' own words.

Metallic bonding

In a metal, atoms lose control of their outer electrons.

The metal forms a lattice of positive ions surrounded by delocalised electrons.

The attraction between the positive ions and the delocalised electrons is metallic bonding.

Use the phrase 'sea of delocalised electrons' only if helpful; keep the formal definition on the slide.

Copper as an example

Copper shows typical metallic bonding.

• it has a giant lattice structure

• its electrons are delocalised

• these mobile electrons help copper conduct electricity

• this is why copper is widely used in wires

Link the model directly to a familiar use so students see why metallic bonding matters.

Properties of metals

Metals usually have:

• good electrical conductivity because delocalised electrons can move

• high melting points because metallic bonds are strong

• good thermal conductivity

• low solubility in water

Stay close to the requested property set, but thermal conductivity is a useful linked property of the same model.

Bond type and properties

Ionic - giant lattice of positive and negative ions

Covalent - shared pairs of electrons; may be simple molecules or giant covalent structures

Metallic - positive ions in a lattice with delocalised electrons

Ionic - high melting point; conduct when molten or dissolved; often soluble

Simple molecular - low melting point; do not conduct

Giant covalent - very high melting point; usually do not conduct, except graphite

Metallic - conduct electricity; usually high melting point; insoluble

Structure and bonding

Typical properties

This comparison slide is the landing point of the lesson: students should now be able to predict properties from bonding and structure.

Why does copper conduct electricity?

A

Ions can dissolve

B

Electrons are delocalised

C

Atoms share pairs

D

Molecules are weak

Checks the metallic model from slides 17 to 20 and links directly to the copper wire example.

Metallic bonding summary

In metals, positive ions are arranged in a lattice and surrounded by delocalised electrons. This explains why metals conduct electricity and usually have high melting points.

ions

delocalised electrons

electricity

molecules

shared pairs

water

negative ions

Anchors the load-bearing terms from the final teaching block and the comparison slide.

Which bond type do you now feel most confident about?

1

Ionic

2

Covalent

3

Metallic

4

Still unsure

How confident are you with the different types of bonding now?

Glossary

outer shell - the electrons in the highest occupied energy level of an atom

stable arrangement - an electron arrangement with a full outer shell in this model

ion - a charged particle formed when an atom gains or loses electrons

ionic bond - the electrostatic attraction between oppositely charged ions

covalent bond - a shared pair of electrons between atoms

metallic bond - the attraction between positive metal ions and delocalised electrons

delocalised electrons - electrons that are free to move through a metal or within graphite layers

Glossary

giant lattice - a repeating three-dimensional arrangement of particles

simple molecule - a small group of covalently bonded atoms

giant covalent structure - a large network of atoms joined by covalent bonds

electrical conductivity - how easily charge flows through a substance

solubility - how well a substance dissolves in a solvent such as water

graphite - a form of carbon with layers and delocalised electrons

diamond - a form of carbon with a rigid giant covalent structure