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Chemistry & Biology · NCEA Level 1

92023 · Properties of materials

Understand how the physical properties of materials inform their use.

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Key concepts

  • Structure explains properties; properties explain uses

    This chain is the whole standard. Take copper wiring: copper is a metallic lattice with delocalised electrons (structure), so those electrons are free to move and carry charge (property), so it is used to carry electricity around a house (use). Every strong answer walks all three steps in order. If you find yourself writing "copper is used for wires because it conducts electricity", you have skipped the structure step and given only half an explanation - the half that does not show you understand any chemistry.

  • Ionic bonding and the ionic lattice

    When a metal reacts with a non-metal, electrons transfer: the metal atom becomes a positive ion and the non-metal atom a negative ion. Those oppositely charged ions attract one another strongly in every direction, so they pack into a repeating three-dimensional lattice - not into molecules. Sodium chloride is a lattice of alternating Na⁺ and Cl⁻ ions, which is why the formula NaCl means "one sodium ion for every chloride ion", not "one NaCl molecule". The strength of that attraction is behind almost every property an ionic solid has.

  • Why ionic solids are high-melting, hard and brittle

    Melting an ionic solid means giving the ions enough energy to break free of a lattice of strong electrostatic attractions pulling in all directions, so melting points are high - sodium chloride melts at 801 °C and magnesium oxide at about 2 850 °C. The same attractions make the solid hard. But the lattice is also brittle: hit it hard enough and a layer of ions shifts by one position, which suddenly lines up ions of the same charge. They repel, and the crystal splits cleanly along that plane rather than bending.

  • Why ionic solids conduct only when molten or dissolved

    Electrical conduction needs charged particles that are free to move. In a solid ionic lattice the ions are charged but locked in place, so nothing flows and the solid is an insulator. Melt it, or dissolve it in water, and the lattice breaks up: the ions can now move, and the liquid conducts. This is not a laboratory curiosity - it is the reason aluminium is extracted at Tiwai Point from aluminium oxide dissolved in a molten bath rather than from the solid powder. No mobile ions, no current, no metal.

  • Metallic bonding: cations in a sea of delocalised electrons

    In a metal, each atom releases its outer electrons into a shared pool. What is left is a lattice of positive ions surrounded by electrons that belong to no particular atom - a "sea" of delocalised electrons. The attraction between the positive lattice and that negative sea is the metallic bond. It is strong, which is why most metals are dense solids with high melting points: copper melts at 1 085 °C and tungsten at about 3 420 °C, which is what lets a tungsten filament glow white-hot without melting.

  • Why metals conduct, bend and can be alloyed

    Delocalised electrons are free to drift through the whole lattice, so metals conduct electricity as solids - unlike ionic compounds - and conduct heat well for the same reason. They are malleable and ductile because the layers of cations can slide over one another while the electron sea flows with them and keeps the attraction intact, so the metal changes shape instead of shattering. Alloying disrupts that sliding on purpose: steel is iron with carbon atoms of a different size wedged into the lattice, so the layers no longer slip past each other easily and the metal is harder and stronger than pure iron. Most of the steel used in New Zealand construction begins as ironsand mined off the west coast of the North Island.

  • Covalent bonding: molecular substances

    Two non-metal atoms can share a pair of electrons instead of transferring them - a covalent bond. In a molecular substance the atoms within each molecule are held by those strong covalent bonds, but separate molecules are attracted to each other only weakly. Melting or boiling only has to overcome those weak forces between molecules, so molecular substances melt and boil at low temperatures: iodine melts at 114 °C and candle wax at around 55 °C. They also do not conduct, because the electrons are held in fixed bonds and there are no free charged particles.

  • Covalent network solids are a different animal

    In a network (or giant covalent) solid there are no separate molecules at all - every atom is covalently bonded into one continuous structure. In diamond each carbon atom is bonded to four others in a rigid three-dimensional framework, which is why it is the hardest natural material and does not melt below about 3 500 °C. Silicon dioxide - the mineral quartz, and the main component of most sand - is a network of silicon and oxygen atoms and melts near 1 710 °C, which is why quartz sand survives in a furnace. Pounamu (nephrite) is another silicate, and its interlocking fibrous crystals are what make it so tough to break and able to hold a fine edge. Same kind of bond as candle wax, wildly different properties - because the structure is different.

  • A melting point tells you what is being broken

    This is the single idea most often got backwards. A low melting point does not mean weak covalent bonds; it means the forces between molecules are weak. When wax melts, not one covalent bond inside a wax molecule is broken - the molecules simply come apart from each other. When quartz melts, actual covalent bonds must break, which is why it takes over a thousand degrees more. So when you read a melting point, ask "what has to be overcome for this to melt?" - forces between molecules, an ionic lattice, a metallic lattice, or a covalent network.

  • Reading property data and justifying a choice

    Given a table of melting points, conductivities and solubilities, you can usually name the structure: conducts as a solid → metallic; does not conduct as a solid but conducts when molten, high melting point, often soluble in water → ionic; low melting point and no conduction in any state → molecular; very high melting point, very hard, no conduction → covalent network (with graphite as the deliberate exception, since each carbon bonds to only three others and the spare electron is delocalised, so it conducts and its layers slide). Choosing a material then means matching required properties to that structure, comparing at least one realistic alternative, and being honest about the limitations - cost, mass, brittleness, corrosion - before you state your judgement.

Assessment

External · grade-score marked.

This is an external achievement standard. Whetū offers original exam-style questions marked on the NCEA grade-score scale. Grades are estimates, not official results.

A mock paper is available in Whetū: Mock paper · Properties of materials.

Learn

2 authored Learn units for this standard.

  • Structure and bonding

    A material's behaviour is not a list of adjectives. It follows from which particles sit where, and from the force that holds them. This page owns that first half of the argument: the five permitted types of material, the arrangement inside each, and the force you have to name. The types are molecular substances, metallic solids, ionic materials, natural and synthetic polymers, and covalent networks. Alloys sit inside the metallic type as a packing variant, not as a sixth family.

  • Physical properties and their uses

    A physical property is a characteristic you can observe or measure without changing the substance's chemical composition. The exam's list is density, thermal and electrical conductivity, melting and boiling points, solubility in water, malleability and hardness. A periodic table and an example list sit in the resource booklet, so you read the values you are given rather than dumping a memorised data book. Melting point and boiling point also fix the physical state at a stated temperature: below the melting point the sample is solid; between melting and boiling it is liquid; above boiling it is gas.

Practise

34 Practise questions in “Materials — structure, properties and uses”. Feedback here is formative and is not an official NCEA grade.

  • The bonding families

    Tell ionic, metallic and covalent structures apart.

  • Ionic structures

    Explain an ionic lattice and the properties it produces.

  • Metallic structures

    Use the cation lattice and electron sea to explain metals.

  • Covalent structures

    Separate small molecules from giant covalent networks.

  • Structure explains property

    Never skip the middle step between structure and use.

  • Melting point and conductivity

    Say what is being overcome, and what is free to move.

  • Choosing a material

    Match the properties a real job needs to a real material.

  • Reading property data

    Work out a structure from melting point, conduction and solubility.

  • Explaining a use

    Carry one thread from structure through property to use.

  • Evaluating a choice

    Compare an alternative and judge with the limitations named.

Sample questions

  1. Which three kinds of bonding do you use to explain the properties of materials in this standard?
  2. 92023 is an external standard. How is your response graded?
  3. Match each structure to the particles and forces inside it.
  4. What holds an ionic lattice such as sodium chloride together?
  5. What best explains the difference?
  6. Four students explain why. Whose interpretation is correct?
  7. Which description best matches the structure of a metal?
  8. Why are metals malleable?

Practise 92023 in Whetū

Exam-style questions

211 original exam-style questions in Whetū. These are practice papers, not official NZQA assessments.

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