Chemistry 2 (chemical changes)

Metal oxides

Metals react with oxygen to produce metal oxides.

eg  magnesium burns in oxygen to produce magnesium oxide

     2Mg(s)   +   O₂(g)  →   2MgO(s)

Metal oxides will be solids because they are giant ionic structures.

The reactions are oxidation reactions because the metals gain oxygen

Be able to explain reduction and oxidation in terms of loss or gain of oxygen.

Be careful about oxidation/reduction which can also be defined in terms of loss/gain of electrons (see later). In the example, magnesium gains oxygen (hence it is oxidised), and loses electrons to the oxygen (also oxidation because of loss of electrons). This is HT only.

Reactivity series

Metal atoms form positive ions when they react with other substances. The reactivity of a metal depends on how easily it forms positive ions. Metals can be arranged in order of their reactivity in a reactivity series.

The metals potassium, sodium, lithium, calcium, magnesium, zinc, iron and copper can be put in order of their reactivity from their reactions with water and dilute acids. (NB you only need to consider room temp. - not with steam)

The non-metals hydrogen and carbon are often included in the reactivity series (see over) which helps to explain reactions of metals (eg extraction/displacement).

A more reactive metal can displace a less reactive metal from a compound

Example:

zinc can displace copper from its compounds (because more reactive/higher in the series than copper), but it can't displace magnesium.

Extraction of metals and reduction

Unreactive metals such as gold are found in the Earth as the metal itself but most metals are found as compounds that require chemical reactions to extract the metal.

Metals less reactive than carbon can be extracted from their oxides by reduction with carbon eg iron but not aluminium.

Reduction involves the loss of oxygen.

Knowledge of the details of processes used in the extraction of metals is not required.

Be able to:

• interpret or evaluate specific metal extraction processes when given appropriate information
• identify the substances which are oxidised or reduced in terms of gain or loss of oxygen.

Oxidation and reduction in terms of electrons (HT only)

Oxidation is the loss of electrons and reduction is the gain of electrons (OILRIG)

When metals lose electrons to become positive ions, they are oxidised.

When metal ions gain electrons to become neutral atoms, they are reduced.

Oxidation and reduction must happen together (redox). Half-equations are used to show each half of the reaction.

Examples:

When magnesium burns in oxygen, it is oxidised because it gains oxygen, but also because it loses electrons to become positive ions   Mg - 2e^- → Mg²⁺

When copper II oxide is heated with carbon, the copper is reduced (it loses its oxygen and gains electrons to become a neutral atom) Cu²⁺ + 2e^- → Cu

Reactions of acids with metals

You should learn the formulae of hydrochloric acid (HCl) and sulphuric acid (H₂SO₄)

These acids react with some metals to produce salts (chlorides or sulphates) and hydrogen.

Example:

magnesium and sulphuric acid        Mg(s) + H₂SO₄(aq)  → MgSO₄(aq)+  H₂(g)

It would be dangerous to mix very reactive metals (eg sodium) with acids. Metals low down in the Reactivity Series (eg copper, silver and gold) do not react with acids in this way. .

(HT only) These are redox reactions. Metal atoms are being oxidised by losing electrons to become positive ions eg Mg²⁺ and H⁺ ions (from the acid) are being reduced by gaining those electrons to become hydrogen gas: 

                 Mg(s) - 2e^- →Mg²⁺(aq)            and            2H⁺+ 2e^-→ H₂(g)

Neutralisation of acids and salt production

As well as the reaction between reactive metals and acids, salts are formed by:

acid + alkali (soluble metal hydroxide) → salt + water

acid + base (insoluble metal oxide or hydroxide) → salt + water

acid + metal carbonate → salt + water + carbon dioxide

NB in each case, the acid is neutralised

The salt produced depends on:

• the acid (hydrochloric acid → chlorides, nitric acid → nitrates, sulphuric acid → sulphates)
• the metal in the alkali/base/carbonate
  

You can work out the formulae of simple ions (eg Cl^-) from the Periodic Table Group, but should learn the sulphate ion SO₄^2-

Soluble salts

Soluble salts can be made from acids by reacting them with solid insoluble substances, such as metals, metal oxides, hydroxides or carbonates. The solid is added to the acid until no more reacts (the acid is the limiting reactant) and the excess solid is filtered off to produce a solution of the salt.

Salt solutions can be crystallised to produce solid salts. Be able to describe how to make pure, dry samples of named soluble salts from information provided. Revise the required practical to prepare a pure, dry sample of a soluble salt from an insoluble metal oxide or carbonate using warm dilute acid and a water bath or electric heater to evaporate the solution.

Example: Stir excess black copper II oxide solid with dilute sulphuric acid, warm and stir until no more will dissolve. Filter off the excess black solid and evaporate the blue solution (either completely, to give copper II sulphate powder, or leave to evaporate slowly if crystals are wanted).

                          CuO(s) + H₂SO₄(aq) → CuSO₄(aq) + H₂O(l)

The pH scale and neutralisation

Acids produce hydrogen ions (H⁺)in aqueous solutions.

Aqueous solutions of alkalis contain hydroxide ions (OH^-).

The pH scale, from 0 to 14, is a measure of the acidity or alkalinity of a solution, and can be measured using universal indicator or a pH probe.

A solution with pH 7 is neutral. Aqueous solutions of acids have pH values < 7 and aqueous solutions of alkalis have pH values > 7.

In neutralisation reactions between an acid and an alkali, hydrogen ions react with hydroxide ions to produce water.

This important reaction can be represented by the equation:

                                             H⁺ + OH^- ⇄  H₂O

Electrolysis

When an ionic compound is melted or dissolved in water, the ions are free to move about within the liquid or solution. These liquids and solutions are able to conduct electricity and are called electrolytes.

Passing an electric current through electrolytes causes the ions to move to the electrodes.

Positively charged ions move to the negative electrode (the cathode). They are sometimes called cations.

Negatively charged ions (anions) move to the positive electrode (the anode).

Ions are discharged at the electrodes producing neutral elements. This process is called electrolysis.

Be able to write half equations for the reactions occurring at the electrodes during electrolysis, and may be required to complete and balance supplied half equations.

Electrolysis of molten ionic compounds

NB It takes a lot of heat energy to melt ionic compounds.

When a simple ionic compound is melted and electrolysed, the positive (metal) ions will move to the cathode and be discharged as the metal element. The negative (non-metal) ions will move to the anode and be discharged as the non-metal element.

Example:

If lead II bromide (PbBr₂) is melted, the following ions are present:

                       PbBr₂ ⇄ Pb²⁺ + 2Br^-

When an electric current is passed through it, using inert electrodes, the lead (Pb²⁺) ions move to the cathode and are discharged as lead metal Pb²⁺ + 2e^- → Pb.

Bromide (Br^-) ions are discharged at the anode as bromine 2Br^- - 2e^-→ Br₂

Using electrolysis to extract metals

Metals can be extracted from molten compounds using electrolysis. Extraction of metals from their compounds (ores) always involves reduction.

Electrolysis is used if the metal is too reactive to be extracted by reduction with carbon or if the metal reacts with carbon. Large amounts of energy are used in the extraction process to melt the compounds and to produce the electrical current.

Aluminium is manufactured by the electrolysis of a molten mixture of aluminium oxide and cryolite using carbon as the positive electrode (anode). The cryolite is another aluminium compound added as a deliberate impurity, which lowers the mp of the electrolyte and reduces the energy costs.

Molten aluminium forms at the cathode (steel)   Al³⁺ + 3e^-  →   Al

Oxygen is formed at the carbon anodes 6O^- - 6e^- →   3O_{2 .} The anodes need to be replaced constantly as they burn to form carbon dioxide.

Electrolysis of aqueous solutions

The ions discharged when an aqueous solution is electrolysed using inert electrodes depend on the relative reactivity of the elements involved. The least reactive elements are discharged.

At the negative electrode (cathode), hydrogen is produced if the metal is more reactive than hydrogen.

At the positive electrode (anode), oxygen is produced unless the solution contains halide ions when the halogen is produced.

This happens because in the aqueous solution water molecules break down producing hydrogen ions and hydroxide ions that are discharged.

                                       H₂O  ⇄  H⁺+ OH^-

Revise the required practical to investigate electrolysis of aqueous solutions involves developing a hypothesis)

Representing reactions at electrodes as half equations (HT only)

During electrolysis, at the cathode (negative electrode), positively charged ions gain electrons and so the reactions are reductions.

At the anode (positive electrode), negatively charged ions lose electrons and so the reactions are oxidations.

Reactions at electrodes can be represented by half equations, for example:

at a cathode:      2H⁺+ 2e^-→ + H₂(g)                          (gain of electrons = reduction)

at an anode:       4OH^- - 4e^- → H₂O + O₂(g)              (loss of electrons = oxidation)

NB the electrons can be written on either side of these half-equations, but be careful about getting the + or - on the correct side.

see separate card in Word (2.14)

Factors which affect the rates of chemical reactions

Factors which affect the rates of chemical reactions include the:

• concentrations of reactants in solution (higher conc → faster)
• pressure of reacting gases (higher pressure → faster)
• surface area of solid reactants (greater SA → faster)
• temperature (higher temp → faster)
• presence of catalysts (faster)
  

Revise required practical to investigate how changes in concentration affect the rates of reactions by a method involving measuring the volume of a gas produced and a method involving a change in colour or turbidity. This should involve developing a hypothesis.

Collision theory explains how factors affect rates of reactions. Chemical reactions can occur only when reacting particles collide with each other and with sufficient energy. The minimum amount of energy that particles must have to react is called the activation energy.

Increasing the concentration, pressure of reacting gases, and the surface area of solid reactants increases the frequency of collisions and so increases the rate of reaction.

Increasing the temperature increases the frequency of collisions and makes the collisions more energetic, and so increases the rate of reaction.

Be able to:

• predict and explain using collision theory the effects of changing conditions of concentration, pressure and temperature on the rate of a reaction
• predict and explain the effects of changes in the size of pieces of a reacting solid in terms of surface area to volume ratio
• use simple ideas about proportionality when using collision theory to explain the effect of a factor on the rate of a reaction.

Catalysts

Catalysts change the rate of chemical reactions but are not used up during the reaction (they can be used over and over). Different reactions need different catalysts.

Enzymes act as catalysts in biological systems.

Catalysts increase the rate of reaction by providing a different pathway for the reaction that has a lower activation energy.

See over for a reaction profile for a catalysed reaction.

Be able to

• identify catalysts in reactions from their effect on the rate of reaction and because they are not included in the chemical equation for the reaction
• explain catalytic action in terms of activation energy.

You do not need to know the names of catalysts other than those specified in the subject content.

Crude oil, hydrocarbons and alkanes

Crude oil is a finite resource found in rocks. It is the remains of an ancient biomass consisting mainly of plankton that were buried in mud.

Crude oil is a mixture of a very large number of compounds. Most of the compounds in crude oil are hydrocarbons, which are molecules made up of hydrogen and carbon atoms only.

Most of the hydrocarbons in crude oil are hydrocarbons called alkanes. The general formula for the homologous series of alkanes is C_nH_2n+2

Structural formulae are shown overleaf. You should be able to recognise substances as alkanes given their formulae but do not need to know specific names other than the first four:

• methane  CH₄
• ethane    C₂H₆
• propane  C₃H₈
• butane    C₄H₁₀

The many hydrocarbons in crude oil may be separated into fractions, each of which contains molecules with a similar number of carbon atoms, by fractional distillation.

The fractions can be processed to produce fuels and feedstock for the petrochemical industry.

Many of the fuels on which we depend, such as petrol, diesel oil, kerosene, heavy fuel oil and liquefied petroleum gases, are produced from crude oil.

Many useful materials on which modern life depends are produced by the petrochemical industry, such as solvents, lubricants, polymers, detergents.

The huge array of natural and synthetic compounds occurs due to carbon's ability to form 4 strong covalent bonds in chains and rings, giving families of similar compounds.

You should be able to explain how fractional distillation works (evaporation and condensation).

You do not need to know the names of other specific fractions or fuels.

Some properties of hydrocarbons depend on the size of their molecules, including bp, viscosity and flammability. These properties influence how hydrocarbons are used as fuels.

With increasing molecular size:

• boiling point and viscosity increase (due to increasing intermolecular forces)
  
• flammability decreases (less volatile)
  

Combustion of hydrocarbons releases energy. During combustion, the carbon and hydrogen in the fuels are oxidised. Complete combustion of a hydrocarbon produces carbon dioxide and water eg for the complete combustion of propane (C₃H₈)

                                  C₃H₈ + 5O₂ → 3CO₂ + 4H₂O

Combustion in a limited supply of air/oxygen can produce carbon monoxide which can be lethal.

Large hydrocarbons can be broken down (cracked) to produce smaller, more useful molecules eg to meet the demand for fuels which ignite more easily.

Cracking can be done by passing the hydrocarbon over a very hot catalyst or by mixing with steam and heating for a few seconds at a very  high temperature.

The products of cracking include alkanes and alkenes.

Example:   Decane can be cracked by passing the vapour over a very hot catalyst

                      C₁₀H₂₂ → C₈H₁₈ + C₂H₄    (structural formulae overleaf)

to form more useful products - octane (petrol) and ethene (eg to make polyethene)

Alkenes have a double covalent bond and are more reactive than alkanes. The test for alkenes is to react with bromine water (turns from brown to colourless).

Alkenes are used to produce polymers and as starting materials (feedstock) for the production of many other chemicals.

Structure and formulae of alkenes

Alkenes are hydrocarbons with a double carbon-carbon bond. The general formula for the homologous series of alkenes is C_nH_2n

Alkene molecules are unsaturated because they contain two fewer hydrogen atoms than the alkane with the same number of carbon atoms.

The first four members of the homologous series of alkenes are ethene, propene, butene and pentene (you don't need to know any other specific names).

Ethene has the formula C₂H₄

The structural formula for an alkene molecule is shown overleaf.

Alkenes are hydrocarbons with the functional group C=C. The reactivity of functional groups determine the reactions of organic compounds.

Alkenes react with oxygen in combustion reactions in the same way as other hydrocarbons, but they tend to burn in air with smoky flames because of incomplete combustion.

Alkenes react with hydrogen, water and the halogens, by the addition of atoms across the carbon=carbon double bond so that the double bond becomes a single carbon-carbon bond.

Example:                CH₃-CH=CH₂ + Br₂  → CH₃-CHBr-CH₂Br

See overleaf for structural formulae.

You should be able to draw fully displayed structural formulae of the first four members of the alkenes and the products of their addition reactions with hydrogen, water, chlorine, bromine and iodine.

Addition polymerisation

Alkenes can be used to make polymers such as poly(ethene) and poly(propene) by addition polymerisation. 

In addition polymerisation reactions, many small molecules (monomers) join together to form very large molecules (polymers). For example individual ethene (CH₂=CH₂) molecules will join together in the presence of a catalyst to form polyethene (CH₂-CH₂)_n 

In addition polymers the repeating unit has the same atoms as the monomer because no other molecule is formed in the reaction.

Be able to:

• recognise addition polymers and monomers from diagrams in the forms shown and from the presence of the functional group C=C in the monomers
• draw diagrams to represent the formation of a polymer from a given alkene monomer
• relate the repeating unit to the monomer.

Pure substances

In chemistry, a pure substance is a single element or compound, not mixed with any other substance.

Pure elements and compounds melt and boil at specific temperatures.

Melting point and boiling point data can be used to distinguish pure substances from mixtures.

In everyday language, a pure substance can mean a substance that has had nothing added to it, so it is unadulterated and in its natural state, eg pure milk.

You should be able to use melting point and boiling point data to distinguish pure from impure substances.

A formulation is a mixture that has been designed as a useful product.

Many products are complex mixtures in which each chemical has a particular purpose. Formulations are made by mixing the components in carefully measured quantities to ensure that the product has the required properties. Formulations include fuels, cleaning agents, paints, medicines, alloys, fertilisers and foods.

You should be able to identify formulations given appropriate information.

You do not need to know the names of components in proprietary products.

Chromatography can be used to separate mixtures and identify substances. Chromatography involves a stationary phase (the paper or other material) and a mobile phase (the solvent).

The ratio of the distance moved by a compound (centre of spot from origin) to the distance moved by the solvent can be expressed as its Rf value:

Rf = distance moved by substance/distance moved by solvent

Rf will be between 0-1. The bigger the number, the faster the substance has moved and the more soluble it must be in that solvent. Different compounds have different Rf values in different solvents, which can help identify the compounds. The substances in a mixture may separate into different spots but a pure compound will produce a single spot in all solvents.

Be able to (required practical):

• explain how paper chromatography separates mixtures and suggest how chromatographic methods can be used for distinguishing pure substances from impure substances
• interpret chromatograms and determine Rf values from them
  

Tests for:

hydrogen- a burning splint held at the open end of a test tube of the gas. Hydrogen burns rapidly with a pop sound.

oxygen - a glowing splint inserted into a test tube of the gas. The splint relights in oxygen.

carbon dioxide - an aqueous solution of calcium hydroxide (lime water). When carbon dioxide is shaken with or bubbled through limewater the limewater turns milky (cloudy).

chlorine  - damp litmus paper. When damp litmus paper is put into chlorine gas the litmus paper is bleached and turns white. NB chlorine also smells like bleach.

The proportions of different gases in the atmosphere

For 200 million years, the proportions of different gases in the atmosphere have been much the same as they are today:

• about four-fifths (approximately 80%) nitrogen
• about one-fifth (approximately 20%) oxygen
• small proportions of various other gases, including carbon dioxide, water vapour and noble gases

Theories about what was in the Earth’s early atmosphere and how the atmosphere was formed have changed and developed over time. Evidence is limited because of the time scale of 4.6 billion years.

One theory suggests that during the first billion years of the Earth’s existence there was intense volcanic activity that released gases that formed the early atmosphere and water vapour that condensed to form the oceans. At the start of this period the Earth’s atmosphere may have been like the atmospheres of Mars and Venus today, consisting of mainly carbon dioxide with little or no oxygen gas.

Volcanoes also produced nitrogen which gradually built up in the atmosphere and there may have been small proportions of methane and ammonia.

When the oceans formed carbon dioxide dissolved in the water and carbonates were precipitated producing sediments, reducing the amount of carbon dioxide in the atmosphere.

You should be able to, given appropriate information, interpret evidence and evaluate different theories about the Earth’s early atmosphere.

Algae and plants produced the oxygen that is now in the atmosphere by photosynthesis, which can be represented by the equation:

                                                       light

              6CO₂    +    6H₂O                  →                    C₆H₁₂O₆  +   6O₂

 carbon dioxide  +   water                →               glucose     +   oxygen

Algae first produced oxygen about 2.7 billion years ago and soon after this oxygen appeared in the atmosphere.

Over the next billion years plants evolved and the percentage of oxygen gradually increased to a level that enabled animals to evolve.

Algae and plants reduced the proportion of carbon dioxide in the atmosphere by photosynthesis.

Carbon dioxide was also reduced by the formation of sedimentary rocks and fossil fuels that contain carbon.

Be able to:

• describe the main changes in the atmosphere over time and some of the likely causes of these changes
• describe and explain the formation of deposits of limestone, coal, crude oil and natural gas.

Greenhouse gases in the atmosphere maintain temperatures on Earth high enough to support life.

Water vapour, carbon dioxide and methane are greenhouse gases.

Electromagnetic radiation at most wavelengths from the Sun passes through the Earth’s atmosphere.

The Earth absorbs electromagnetic radiation with short wavelengths and so warms up. Heat is radiated from the Earth as longer wavelength infrared radiation.

Some of this infrared radiation is absorbed (trapped) by greenhouse gases in the atmosphere and so it warms up.

Some human activities increase the amounts of greenhouse gases in the atmosphere. These include carbon dioxide and methane.

Be able to recall two human activities that increase the amounts of each of these greenhouse gases (eg burning fossil fuels and deforestation due to increasing consumption of meat)

Based on peer-reviewed evidence, most scientists believe that human activities are causing the temperature of the atmosphere to rise and that this is resulting in global climate change. However, global climate change is a complex issue. This leads to simplified models, speculation and opinions presented in the media that may be biased.

Be able to:

• evaluate the quality of evidence in a report about global climate change given appropriate information
• describe uncertainties in the evidence base
• recognise the importance of peer review of results and of communicating results to a wide range of audience

An increase in average global temperature is a major cause of climate change.

There are several potential effects of global climate change.

You should be able to:

• describe briefly four potential effects of global climate change (eg flooding due to rise in sea levels, melting polar ice caps threatening animal life, migration of populations away from increasing desert regions because of difficulty in growing food crops, rising sea temperatures killing coral reefs and threatening other sea life)
• discuss the scale, risk and environmental implications of global climate change.

The carbon footprint is the total amount of carbon dioxide and other greenhouse gases emitted over the full life cycle of a product, service or event.

The carbon footprint can be reduced by reducing emissions of carbon dioxide and methane.

You should be able to:

• describe actions to reduce emissions of carbon dioxide and methane (eg reducing use of fossil fuels/increasing use of renewable energy sources; reducing consumption of meat)
  
• give reasons why actions may be limited (eg denial of the severity of the threat by some, pressure by interested profit making organisations, unwillingness to change lifestyle eg reduced use of cars, perceived cost of 'greener' initiatives)
  

Atmospheric pollutants from fuels

The combustion of fuels is a major source of atmospheric pollutants.

Most fuels, including coal, contain carbon and/or hydrogen and may also contain some sulphur.

The gases released into the atmosphere when a fuel is burned may include carbon dioxide, water vapour, carbon monoxide, sulphur dioxide and oxides of nitrogen. Solid particles and unburned hydrocarbons may also be released that form particulates in the atmosphere.

You should be able to:

• describe how carbon monoxide, soot (carbon particles), sulphur dioxide and oxides of nitrogen are produced by burning fuels
• predict the products of combustion of a fuel given appropriate information about the composition of the fuel and the conditions in which it is used (eg incomplete combustion, forming carbon monoxide).

Carbon monoxide is a toxic gas. It is colourless and odourless and so is not easily detected.

Sulphur dioxide and oxides of nitrogen cause respiratory problems in humans and cause acid rain.

Particulates cause global dimming and health problems for humans.

You should be able to describe and explain the problems caused by increased amounts of these pollutants in the air.

Sustainable development

Humans use the Earth’s resources to provide warmth, shelter, food and transport. Natural resources, supplemented by agriculture, provide food, timber, clothing and fuels.

Finite resources from the Earth, oceans and atmosphere provide energy and materials.

Chemistry plays an important role in improving agricultural and industrial processes to provide new products and in sustainable development. This meets the needs of current generations without compromising the ability of future generations to meet their own needs.

You should be able to:

• state examples of natural products that are supplemented or replaced by agricultural and synthetic products
• distinguish between finite and renewable resources given appropriate information.
• extract and interpret information about resources from charts, graphs and tables
• use orders of magnitude to evaluate the significance of data.

The Earth’s resources of metal ores are limited.

Copper ores are becoming scarce and new ways of extracting copper from low-grade ores include phytomining, and bioleaching. These methods avoid traditional mining methods of digging, moving and disposing of large amounts of rock.

Phytomining uses plants to absorb metal compounds. The plants are harvested and then burned to produce ash that contains metal compounds.

Bioleaching uses bacteria to produce leachate solutions that contain metal compounds.

The metals can be extracted from these compounds.

For example, copper can be obtained from solutions of copper compounds by displacement using scrap iron or by electrolysis.

You should be able to evaluate alternative biological methods of metal extraction, given appropriate information.

Corrosion and its prevention

Corrosion is the destruction of materials by chemical reactions with substances in the environment. Rusting is an example of corrosion.

Both air and water are necessary for iron to rust.

Corrosion can be prevented by applying a coating that acts as a barrier, such as greasing, painting or electroplating. Aluminium develops a thin oxide coating that protects the metal from further corrosion.

Some coatings are reactive and contain a more reactive metal to provide sacrificial protection, eg zinc is used to galvanise iron (because zinc is more reactive than iron).

You should be able to:

• describe experiments and interpret results to show that both air and water are necessary for rusting
• explain sacrificial protection in terms of relative reactivity.

Most metals in everyday use are alloys.

Bronze is an alloy of copper and tin. Brass is an alloy of copper and zinc.

Gold used as jewellery is usually an alloy with silver, copper and zinc. The proportion of gold in the alloy is measured in carats. 24 carats being 100% (pure gold), which is too soft for most jewellery purposes, and 18 carats being 75% gold.

Steels are alloys of iron that contain specific amounts of carbon and other metals. High carbon steel is strong but brittle. Low carbon steel is softer and more easily shaped. Steels containing chromium and nickel (stainless steels) are hard and resistant to corrosion.

Aluminium alloys are low density (useful where weight is an issue eg aircraft bodies).

You should be able to:

• recall a use of each of the alloys specified
• interpret and evaluate the composition and uses of alloys other than those specified given appropriate information.

Most glass we use is soda-lime glass (from sand, sodium carbonate and limestone). Borosilicate glass (eg Pyrex), made from sand and boron trioxide, has a higher mp.

Ceramics (pottery/bricks) are made by shaping wet clay and heating in a furnace.

Polymer properties depend on the monomers and the reacting conditions/catalysts used eg low density (LD) and high density (HD) poly(ethene) are both produced from ethene.

Thermosoftening polymers melt when heated (their long molecules are jumbled and slide easily over each other when heated). Thermosetting polymers have strong cross-linked molecules which do not allow them to melt when heated.

Most composites are made of two materials, a matrix or binder surrounding/binding together fibres or fragments of the other material (the reinforcement) eg fibreglass.

You should be able to, given appropriate information, compare the physical properties of glass and clay ceramics, polymers, composites and metals. You should also be able to explain how the properties of materials are related to their uses and select appropriate materials.