IP Chemistry Upper Sec 03: Chemical Bonding and Structure

Study guide26 Nov 2025, 00:00 ZUpdated 17 Jul 2026

Reviewed by

Azmi·Senior Chemistry Specialist

Ionic, covalent, and metallic bonding, plus structure-property links for IP Sec 3-4 Chemistry.

For Integrated Programme students: Your current school materials, teacher instructions, and assessment scope take precedence because IP topic sequence and depth vary by school. This is an Eclat IP guide, not the O-Level / SEC G3 exam-track guide.

How this chapter applies

Eclat core: ionic, covalent, and metallic bonding; dot-and-cross diagrams; lattices; alloys; simple molecular, giant covalent, and macromolecular structures; and structure-to-property reasoning form the main route.
School-sensitive extension: electronegativity, molecular shape, polarity, and named intermolecular forces add explanatory depth where a school introduces them. Do not substitute these ideas for the bonding model requested by the current question.
2027 national comparison: K324 Topic C3 covers ion formation, the three bonding models, alloys, material structures, diamond and graphite, and deductions between structure and properties.
Check your school: diagram conventions and the depth of polarity or intermolecular-force treatment vary. Match the notation and explanatory level in the current school materials.
Exam-track route: use the separate O-Level and SEC G3 Chemistry notes for K324 topic ownership.

Choose the right Chemistry resource

Use this page for Topic 3 Chemical Bonding and Structure explanation practice. If your need is broader or narrower, use the route below.

Need	Best Eclat page
Official 6092 paper format and topic order	O-Level Chemistry 6092 syllabus guide
Sec 3 foundation overview before bonding	Sec 3 Chemistry notes
Mole, formula, and equation calculations after bonding	Chemical Calculations notes
Paper 2 answer layout and timing	O-Level Chemistry Paper 2 structured question strategy

The core idea is simple: Bonding questions are structure-to-property questions.

Use it as a working check: First name the particles and forces. Then explain melting point, boiling point, conductivity, hardness, or solubility from that structure.

Then go one layer deeper: Example: solid sodium chloride does not conduct because ions are fixed in the lattice, but molten sodium chloride conducts because ions can move.

Structure-to-property answer chain

For bonding explanations, build the answer in this order: identify the structure, name the particles, name the force or mobile charge carrier, then link that feature to the observed property.

Question cue	Structure clue	Explanation chain
High melting point for $\ce{NaCl}$	Giant ionic lattice	Many oppositely charged ions are held by strong electrostatic attractions, so a lot of energy is needed to overcome the lattice.
Solid does not conduct but molten liquid conducts	Ionic compound	Ions are fixed in position when solid, but ions are mobile when molten and can carry charge.
Low boiling point for $\ce{CH4}$	Simple molecular substance	Covalent bonds inside each molecule are strong, but only weak intermolecular forces act between molecules.
Graphite conducts but diamond does not	Different giant covalent structures	Graphite has delocalised electrons between layers; diamond has all outer electrons localised in covalent bonds.
Alloy is harder than pure metal	Metallic lattice with different atom sizes	Different-sized atoms disrupt regular layers, so the layers cannot slide over one another easily.

Common trap: do not say a simple molecule has a low boiling point because covalent bonds are weak. The covalent bonds inside the molecule stay intact; the weak forces being overcome are intermolecular forces between molecules.

What you must know

Ionic bonding: electron transfer to form ions with noble gas configuration; strong electrostatic attraction between oppositely charged ions in a giant ionic lattice.
Covalent bonding: electron sharing between non-metals; draw dot-and-cross for $\ce{H2}$ , $\ce{O2}$ , $\ce{H2O}$ , $\ce{CH4}$ , $\ce{CO2}$ , $\ce{NH3}$ . Distinguish simple molecular (low mp/bp, no conduction) vs giant covalent (diamond, graphite, $\ce{SiO2}$ ).
Metallic bonding: lattice of positive metal ions in a sea of delocalised electrons → good conductivity, malleability/ductility.
Structure-property links: high mp/bp from strong forces (ionic, metallic, giant covalent); low mp/bp for simple molecules because only weak intermolecular forces. Alloys disrupt layers to reduce slippage.
Allotropes: diamond (4 covalent bonds, very hard, no conduction) vs graphite (3 covalent bonds, layers with weak forces, conducts via delocalised electrons).

Detailed notes

Ionic properties: strong lattice → high mp/bp; soluble in water due to ion-dipole interactions; brittle because same-charge layers align on impact and repel. Conduct when molten/aqueous (mobile ions), not when solid.
Covalent polarity and intermolecular forces: electronegativity differences create polar bonds; molecule shape decides overall polarity. London forces scale with electron count and surface area; hydrogen bonding raises bp ( $\ce{H2O}$ , $\ce{HF}$ , $\ce{NH3}$ ).
Giant covalent vs simple molecular: diamond/ $\ce{SiO2}$ hard, high mp, non-conducting; graphite conducts along layers and is soft due to weak forces between layers.
Metallic bonding: strength increases with more delocalised electrons and smaller, higher-charged cations ( $\ce{Mg}$ > $\ce{Na}$ ). Delocalised electrons carry heat and electricity; layers slide unless disrupted by different atom sizes (alloys).
Shapes and angles (qualitative): $\ce{CO2}$ linear 180°, $\ce{H2O}$ bent ~104.5°, $\ce{NH3}$ trigonal pyramidal ~107°, $\ce{CH4}$ tetrahedral 109.5°; lone pairs increase repulsion.

Data-to-structure checkpoint

When a question gives melting point, boiling point, conductivity, and solubility data, identify the structure before writing the explanation.

Data pattern	Likely structure	First explanation to write	Common trap
High melting point, conducts when molten or aqueous but not solid	Ionic lattice	Ions are fixed in the solid but mobile when molten or dissolved.	Saying electrons move through the ionic solid.
Low melting and boiling points, no electrical conductivity	Simple molecular	Weak intermolecular forces are overcome; molecules have no mobile ions or electrons.	Saying covalent bonds are weak.
High melting point, conducts as solid and molten	Metallic or graphite-like giant structure	Mobile delocalised electrons carry charge through the structure.	Calling every solid conductor a metal without checking other data.
Very high melting point, does not conduct	Giant covalent network such as diamond or $\ce{SiO2}$	Many strong covalent bonds extend through the structure.	Treating it like a simple molecule because it has covalent bonds.

Worked check: a solid has melting point above $\pu{800^\circ C}$ , does not conduct as a solid, but conducts when molten. The change in conductivity shows the charge carriers are ions that become mobile on melting, so the substance is likely ionic.

Conductivity carrier checkpoint

For conductivity questions, first identify the mobile charged particle. A substance conducts only when charged particles can move through the structure.

Does it contain mobile charged particles?
  ions fixed in solid lattice       -> no conduction
  ions mobile when molten/aqueous   -> conducts
  delocalised electrons in lattice  -> conducts as a solid
  simple molecules only             -> no conduction

Substance or state	Possible charge carrier	Can the carrier move?	Conductivity conclusion
Solid $\ce{NaCl}$	Ions	No, ions are fixed in the giant ionic lattice.	Does not conduct.
Molten $\ce{NaCl}$	Ions	Yes, ions can move through the liquid.	Conducts.
Aqueous $\ce{NaCl}$	Ions	Yes, ions can move through the solution.	Conducts.
Copper metal	Delocalised electrons	Yes, electrons move through the metallic lattice.	Conducts as a solid.
Liquid $\ce{CH4}$	None	No mobile ions or delocalised electrons are present.	Does not conduct.

Misconception check: melting does not make every substance conduct. It helps ionic compounds because ions become mobile. Simple molecular substances still lack mobile charged particles when liquid.

Ionic brittleness checkpoint

For ionic solids, do not explain brittleness by saying "it is hard, so it breaks". Hardness and brittleness come from related but different structure ideas.

Step in the explanation	What to say	Why it matters	Common trap
Before force is applied	Oppositely charged ions alternate in a regular giant ionic lattice.	Strong attractions hold the lattice rigid, giving a high melting point and hardness.	Saying the ions are free to move in the solid.
When a force shifts one layer	Ions in one layer are pushed sideways relative to the next layer.	The regular + and - pattern is disturbed.	Drawing the layers as if every ion stays opposite an unlike charge.
After the shift	Like-charged ions may become aligned next to each other.	Repulsion between like charges causes the crystal to split.	Saying covalent bonds break in $\ce{NaCl}$ .

Worked check: in a crystal of $\ce{NaCl}$ , a sharp force can shift an ion layer. If $\ce{Na+}$ ions line up beside $\ce{Na+}$ ions and $\ce{Cl-}$ ions line up beside $\ce{Cl-}$ ions, strong repulsion acts across the shifted plane, so the crystal fractures.

Misconception check: ionic solids are hard because many strong electrostatic attractions resist deformation, but they are brittle because a layer shift can line up like charges and create repulsion.

Intermolecular-force selection checkpoint

For simple molecular substances, identify the strongest intermolecular force present before comparing melting or boiling points. Do not discuss the covalent bonds inside the molecules unless the question asks about chemical reactions.

Molecule type	Strongest force to name first	What to check	Common trap
Non-polar molecules such as $\ce{CH4}$	London dispersion forces	Larger electron cloud or larger surface area usually gives stronger London forces.	Saying there are no forces between non-polar molecules.
Polar molecules without $\ce{H-N}$ , $\ce{H-O}$ , or $\ce{H-F}$	Permanent dipole-dipole attractions plus London forces	Check whether bond dipoles cancel because of molecular shape.	Treating every polar bond as proof of a polar molecule.
Molecules with $\ce{H}$ bonded to $\ce{N}$ , $\ce{O}$ , or $\ce{F}$	Hydrogen bonding, plus weaker forces	Check for a highly polar $\ce{H-N}$ , $\ce{H-O}$ , or $\ce{H-F}$ bond and a lone pair on another molecule.	Calling every molecule that contains hydrogen a hydrogen-bonding molecule.
Giant covalent or ionic substances	Do not use this intermolecular-force checklist	Explain the lattice or network forces instead.	Comparing $\ce{NaCl}$ with $\ce{CH4}$ using only molecular forces.

Worked check: $\ce{CH3CH2OH}$ has an $\ce{O-H}$ bond, so ethanol molecules can form hydrogen bonds with one another. $\ce{CH3OCH3}$ has polar bonds and lone pairs on oxygen but no $\ce{O-H}$ bond, so it cannot donate hydrogen bonds to another identical molecule. This helps explain why ethanol has the stronger intermolecular attractions.

Misconception check: hydrogen bonding is an intermolecular attraction, not a covalent bond inside one molecule. Breaking it during boiling separates molecules; it does not split the molecule into atoms.

Worked walkthroughs

Draw dot-and-cross and state structure/force: $\ce{MgO}$ ionic lattice → high mp, conducts when molten. $\ce{CH4}$ simple molecular → low mp, no conduction. $\ce{SiO2}$ giant covalent → very high mp, no conduction.
Boiling point comparison: explain $\ce{CH4} < \ce{CH3Cl} < \ce{CH3CH2OH}$ by London < dipole-dipole < hydrogen bonding strength.
Graphite vs diamond: bonding (3 vs 4 per carbon), delocalised electrons in graphite layers → conduction and lubrication; diamond has no mobile electrons → insulator, very hard.
Alloy reasoning: Brass (Cu + Zn) harder than Cu because different atom sizes disrupt layer sliding, reducing malleability.

Pitfalls and fixes

Calling intermolecular forces “covalent bonds”-clarify they are attractions between molecules.
Saying ionic solids conduct-only molten/aqueous conduct due to mobile ions.
Forgetting to link properties to structure (e.g., high mp because many strong bonds/forces to overcome).
Mislabeling hydrogen bonding-only when H bonded to $\ce{N}$ , $\ce{O}$ , or $\ce{F}$ and a lone pair present.

Practice drills

Predict shape and polarity of $\ce{NH3}$ vs $\ce{BF3}$ ; link to solubility in water.
Explain why $\ce{Mg}$ has a higher mp than $\ce{Na}$ , but $\ce{Al}$ is higher still.
Describe why ice floats on water using hydrogen bonding and open lattice structure.
Compare conduction in molten $\ce{NaCl}$ , solid graphite, and solid $\ce{SiO2}$ .

Quick applications

Draw dot-and-cross for $\ce{NaCl}$ and $\ce{MgCl2}$ and explain why molten/aqueous forms conduct but solid does not (ions fixed vs mobile).
Explain why $\ce{CH4}$ boils at low temperature (weak intermolecular forces) while diamond needs high energy to break covalent network.
Deduce likely structure from data: high mp + conducts when molten = ionic; low mp + no conduction = simple molecular; high mp + conducts as solid = metallic or graphite.
Use alloy example: brass (Cu + Zn) is harder than copper because different atom sizes prevent layer sliding.

Exam cues

When describing any lattice, name the particles and the electrostatic attraction between them.
Simple molecules do not conduct because there are no mobile ions/electrons in the structure (even if molten).
For graphite conduction, emphasise delocalised electrons along layers; for diamond, all electrons localised in bonds.
Alloys: “different atom sizes disrupt regular layers, so layers cannot slide easily,” giving strength and lower ductility.

Next steps

Revise the full hub: https://eclatinstitute.sg/blog/ip-chemistry-upper-sec-notes

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Sources

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