Bonding and Structure

What you'll learn

This guide covers the fundamental principles of chemical bonding and structure required for WJEC GCSE Chemistry. You'll understand how atoms combine through ionic, covalent and metallic bonding, and how these different bond types determine the physical and chemical properties of substances. This topic is critical for answering questions worth 15-20% of your exam marks.

Key terms and definitions

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

Ionic bonding — the electrostatic attraction between oppositely charged ions

Covalent bonding — the sharing of pairs of electrons between atoms

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

Giant structure — a three-dimensional lattice containing millions of atoms or ions bonded together

Simple molecular structure — small groups of atoms held together by covalent bonds, with weak forces between molecules

Delocalised electrons — electrons that are free to move throughout a structure, not fixed to one atom

Electrostatic attraction — the force of attraction between oppositely charged particles

Core concepts

Ionic bonding and ionic compounds

Ionic bonding occurs between metals and non-metals. Metal atoms lose electrons to form positive ions (cations), while non-metal atoms gain electrons to form negative ions (anions). The strong electrostatic attraction between these oppositely charged ions creates the ionic bond.

Formation of ions:

• Group 1 metals lose 1 electron to form +1 ions (e.g. Na⁺, K⁺)
• Group 2 metals lose 2 electrons to form +2 ions (e.g. Mg²⁺, Ca²⁺)
• Group 7 non-metals gain 1 electron to form -1 ions (e.g. Cl⁻, Br⁻)
• Group 6 non-metals gain 2 electrons to form -2 ions (e.g. O²⁻, S²⁻)

Electronic structure representation:

For sodium chloride (NaCl):

• Sodium atom (2,8,1) → Sodium ion (2,8)⁺
• Chlorine atom (2,8,7) → Chloride ion (2,8,8)⁻

Both ions achieve stable electronic configurations similar to noble gases.

Dot and cross diagrams:

In dot and cross diagrams for ionic compounds:

• Show the outer shell electrons only
• Use dots for electrons from one atom, crosses for the other
• Draw square brackets around each ion with the charge outside
• For sodium chloride: [2,8]⁺ and [2,8,8]⁻

Properties of ionic compounds:

Ionic compounds form giant ionic lattices — regular three-dimensional arrangements of alternating positive and negative ions.

• High melting and boiling points: Strong electrostatic forces between ions require significant energy to overcome
• Electrical conductivity: Do not conduct when solid (ions fixed in position), but conduct when molten or dissolved in water (ions free to move and carry charge)
• Solubility: Many ionic compounds dissolve in water as water molecules can separate the ions
• Brittleness: When struck, layers of ions may shift, causing like charges to align and repel, shattering the structure

Covalent bonding and molecular substances

Covalent bonding occurs between non-metal atoms. Atoms share pairs of electrons to achieve stable electronic configurations. Each shared pair constitutes one covalent bond.

Simple covalent molecules:

Common examples include:

• Hydrogen (H₂): single bond, each H shares 1 electron
• Chlorine (Cl₂): single bond, each Cl shares 1 electron
• Water (H₂O): two single bonds, oxygen shares electrons with two hydrogen atoms
• Ammonia (NH₃): three single bonds, nitrogen shares electrons with three hydrogen atoms
• Methane (CH₄): four single bonds, carbon shares electrons with four hydrogen atoms
• Carbon dioxide (CO₂): two double bonds, carbon shares two pairs of electrons with each oxygen
• Oxygen (O₂): double bond, each oxygen shares 2 electrons

Dot and cross diagrams for covalent molecules:

Show only outer shell electrons. Shared pairs appear in the overlap between atoms. For water (H₂O):

• Oxygen has 6 outer electrons (shown as dots)
• Each hydrogen has 1 electron (shown as crosses)
• Two pairs are shared, giving oxygen 8 outer electrons and each hydrogen 2

Properties of simple molecular substances:

• Low melting and boiling points: Weak intermolecular forces between molecules require little energy to overcome. The covalent bonds within molecules are strong, but these don't break during melting or boiling
• Poor electrical conductivity: No charged particles free to move
• Solubility: Variable — some dissolve in water (e.g. ammonia), others dissolve in organic solvents

Giant covalent structures

Some non-metals form giant covalent structures (also called macromolecular structures) where millions of atoms are joined by covalent bonds in continuous networks.

Diamond (carbon):

• Each carbon atom forms four covalent bonds with other carbon atoms
• Tetrahedral arrangement creates a very rigid three-dimensional structure
• Extremely hard — used in cutting tools and drill bits
• Very high melting point (>3500°C) — millions of strong covalent bonds must break
• Does not conduct electricity — no delocalised electrons

Graphite (carbon):

• Each carbon atom forms three covalent bonds, creating layers of hexagonal rings
• Fourth outer electron from each carbon is delocalised between layers
• Layers held together by weak intermolecular forces
• Soft and slippery — layers slide over each other easily, used in pencils and lubricants
• High melting point — strong covalent bonds within layers
• Conducts electricity — delocalised electrons can move through the structure
• Used in electrodes and battery terminals

Silicon dioxide (silica):

• Each silicon atom bonds to four oxygen atoms, each oxygen bonds to two silicon atoms
• Giant three-dimensional structure similar to diamond
• Very hard with high melting point
• Found in sand and quartz
• Does not conduct electricity

Metallic bonding and properties of metals

Metallic bonding involves the electrostatic attraction between positive metal ions and a 'sea' of delocalised electrons. Metal atoms release their outer shell electrons, which become free to move throughout the entire structure.

Structure of metals:

• Positive metal ions arranged in regular layers
• Delocalised electrons move freely between ions
• Strong electrostatic attraction holds the structure together

Properties of metals:

• Good electrical conductivity: Delocalised electrons are free to move and carry charge
• Good thermal conductivity: Energy transferred rapidly through delocalised electrons and vibrating ions
• Malleable and ductile: Layers of ions can slide over each other without breaking the metallic bonding, as delocalised electrons adjust position. Metals can be hammered into shape (malleable) or drawn into wires (ductile)
• High melting and boiling points: Strong electrostatic attraction between ions and delocalised electrons requires significant energy to overcome (though lower than giant covalent structures)
• Lustrous (shiny): Delocalised electrons reflect light

Alloys:

Alloys are mixtures of metals, or metals with small amounts of other elements. Examples include:

• Steel (iron + carbon)
• Brass (copper + zinc)
• Bronze (copper + tin)

Alloys are harder than pure metals because atoms of different sizes disrupt the regular layer arrangement, making it more difficult for layers to slide over each other.

Nanoparticles and nanoscience

Nanoparticles are structures with dimensions between 1 and 100 nanometres (1 nm = 1 × 10⁻⁹ m). They contain only a few hundred atoms.

Properties of nanoparticles:

• Very high surface area to volume ratio compared to bulk materials
• Different properties from bulk materials — may have different melting points, electrical conductivity, or chemical reactivity
• Surface atoms constitute a significant proportion of total atoms

Applications:

• Medicine: Drug delivery systems, cancer treatment targeting specific cells
• Electronics: Smaller, faster computer components
• Materials: Stronger, lighter sports equipment; self-cleaning windows
• Catalysis: More effective catalysts due to high surface area
• Sun creams: Nanoparticle titanium dioxide and zinc oxide provide UV protection without leaving white marks

Risks and concerns:

• Unknown long-term health effects if nanoparticles enter the body or environment
• May cross biological barriers (skin, blood-brain barrier)
• Difficult to predict environmental impact
• Require careful regulation and testing

Polymers

Polymers are very large molecules formed when many small molecules (monomers) join together through covalent bonds in a process called polymerisation.

Formation:

Monomers contain double bonds (C=C). During polymerisation, these double bonds break, allowing monomers to link into long chains. For example:

• Ethene (C₂H₄) polymerises to form poly(ethene)
• Propene polymerises to form poly(propene)
• Styrene polymerises to form poly(styrene)

Structure:

Polymer molecules consist of thousands of atoms in long chains. Individual molecules are not covalently bonded to each other, but held together by weak intermolecular forces.

Properties:

• Solid at room temperature
• Lower melting points than metals or ionic compounds (weak forces between chains)
• Do not conduct electricity
• Variable strength and flexibility depending on structure
• Some are thermosoftening (soften when heated, harden when cooled)
• Thermosetting polymers have cross-links between chains, making them rigid and heat-resistant

Worked examples

Example 1: Explaining properties based on structure

Question: Sodium chloride has a melting point of 801°C, but chlorine gas has a melting point of -102°C. Explain this difference in terms of structure and bonding. [4 marks]

Mark scheme answer:

Sodium chloride is an ionic compound [1] with a giant ionic lattice structure [1]. Strong electrostatic forces of attraction between oppositely charged ions require significant energy to overcome [1].

Chlorine exists as simple molecules (Cl₂) [1] with weak intermolecular forces between molecules that require little energy to overcome [1]. (Award any 4 points)

Key point: Link structure to bonding type, then bonding to properties. Don't confuse strong covalent bonds within molecules with weak forces between molecules.

Example 2: Drawing bonding diagrams

Question: Draw a dot and cross diagram to show the bonding in a molecule of water (H₂O). Show outer shell electrons only. [3 marks]

Mark scheme answer:

• Oxygen atom shown with 6 outer electrons [1]
• Two hydrogen atoms each shown with 1 electron, with different symbols (dots/crosses) used for electrons from different atoms [1]
• Two shared pairs shown in overlap regions, with unshared pairs on oxygen clearly shown [1]

Key point: Use dots for one element, crosses for another. Show all outer electrons, including lone pairs that aren't involved in bonding.

Example 3: Comparing graphite and diamond

Question: Graphite and diamond are both forms of carbon. Explain why graphite conducts electricity but diamond does not. [3 marks]

Mark scheme answer:

In graphite, each carbon forms three covalent bonds [1], leaving one outer electron delocalised/free to move [1]. These delocalised electrons can carry charge through the structure [1].

In diamond, each carbon forms four covalent bonds [1], so all electrons are held in bonds/there are no delocalised electrons [1]. (Award any 3 points)

Key point: Link the number of bonds formed by each carbon to the availability of delocalised electrons, then to electrical conductivity.

Common mistakes and how to avoid them

• Confusing ionic and covalent bonding: Remember ionic bonding involves transfer of electrons (between metals and non-metals), covalent bonding involves sharing electrons (between non-metals). Don't say ionic compounds "share" electrons.

• Stating molecules have high melting points: Simple covalent molecules have low melting points because weak intermolecular forces break, not the strong covalent bonds within molecules. Only giant structures have high melting points.

• Forgetting charge and square brackets on ions: In dot and cross diagrams for ionic compounds, always draw square brackets around each ion with the charge shown outside, e.g. [2,8,8]⁻.

• Confusing properties of graphite and diamond: Learn specific differences: diamond is hard (used in cutting tools), graphite is soft and slippery (used in pencils); diamond doesn't conduct electricity, graphite does (delocalised electrons).

• Saying metals conduct "because they have electrons": All atoms have electrons. Metals conduct because they have delocalised electrons that are free to move throughout the structure.

• Not linking structure to properties in exam answers: Always explain the mechanism: identify the structure type → describe the bonding/forces present → explain why this gives the property observed.

Exam technique for "Bonding and Structure"

• Command words matter: "Describe" requires you to state features; "Explain" requires you to give reasons using scientific principles. For a 3-mark "explain" question, give three distinct scientific points linking cause and effect.

• Draw diagrams carefully: Use a ruler for dot and cross diagrams. Clearly distinguish dots from crosses. Show all outer electrons in the correct positions. Don't forget charges on ions.

• Use correct scientific terminology: Write "electrostatic attraction between oppositely charged ions" rather than "ions stick together." Write "weak intermolecular forces between molecules" not "weak bonds between molecules."

• Structure extended answers logically: For 4-6 mark questions, organise your answer into clear points: identify the structure type, describe the bonding, explain how this affects properties. Use link words like "therefore," "because," and "this means that."

Quick revision summary

Ionic bonding (metal + non-metal) involves electron transfer creating oppositely charged ions with giant lattice structures, high melting points, and conducting when molten. Covalent bonding (non-metal + non-metal) involves electron sharing; simple molecules have low melting points, giant covalent structures have very high melting points. Metallic bonding involves positive ions in a sea of delocalised electrons, giving good conductivity and malleability. Match structure type to properties in exam questions by explaining the bonding and forces present.