Chemistry GCSE Notes

Structure and Bonding

Atomic number – number of protons in an atom. E.g. 11 in sodium-23.
Mass number – number of protons and neutrons in an atom. E.g. 23 in sodium-23.
Isotopes – atoms of the same element that have the same number of protons, but different numbers of neutrons. E.g. carbon-12 and carbon-13. Many isotopes are radioactive.
Electrons are arranged in shells according to the “octet rule” – 2,8,8.
Elements with full shells are very stable (non-reactive). Elements needing to gain or lose 1 or 2 electrons to achieve a full shell are very reactive (i.e. elements in Groups 1, 2, 6 and 7).
Ionic bonding – transfer of electrons from a metal to a non-metal. Ionic compounds have high melting and boiling points, are usually soluble in water and can conduct electricity when molten. The structure of an ionic compound is a giant lattice. Ionic compounds are often called salts. E.g. sodium chloride.
Covalent bonding – sharing of electrons between non-metals. Simple molecular covalent compounds have low boiling points, are usually insoluble in water, but soluble in organic solvents, and do not conduct electricity. E.g. carbon dioxide, methane. Giant structure covalent compounds have very high melting and boiling points, insoluble in most solvents and do not conduct electricity (Exception: graphite can conduct electricity). E.g. silicon dioxide, diamond.
Metallic bonding – regular lattice of positive ions, surrounded by a “sea” of free electrons. Metals can conduct electricity, are malleable and ductile, and generally have high melting and boiling points.
Allotropes – different forms of an element in the same physical state. E.g. graphite and diamond.

Quantitative Chemistry

The numbers of each atom and the total charge on each side of a chemical equation must balance.
Ionic equations do not include spectator ions, which are ions that do not take part in the reaction.
A mole contains Avagadro’s number of particles – 6.02 x 1023. A particle could be any elemental unit such as an atom, molecule, ion or formula unit. The symbol for moles is n.
The relative atomic or molar mass of an atom or molecule is its mass relative to 1/12 of carbon-12.
Moles = Mass / Molar Mass. In symbols: n = m / M.
Concentration = Moles / Volume. In symbols: c = n / V. The unit for concentration (also known as molarity) called molar, which is equivalent to mol / dm3 (or M).
1 mole of any gas occupies 24 dm3 (or 24 L) at room temperature and atmospheric pressure.
Empirical formula – the simplest formula of a substance. The molecular formula of a substance is either the empirical formula or some multiple of it.

Periodic Table

In the periodic table, vertical columns are called groups. Elements in the same group have similar properties. The horizontal rows are called periods.
The reactivity of Group 1 elements, the alkali metals, increases down the group because the outer electron becomes easier to remove due to the increasing distance from the nucleus.
The reactivity of Group 7 elements, the halogens, increases up the group because the electron to be gained is closer to the nucleus. Fluorine is the most reactive halogen.
Alkali metals are very reactive and have to be stored in oil to exclude air and water. They are less hard and have lower melting points and densities compared to other metals.
Alkali metals react with water to form the metal hydroxide and hydrogen gas.

- - 2Na(s) + 2H2O(l) → 2 NaOH(aq) + H2(g)

When heated, the alkali metals burn with a characteristic flame to form the oxides (see below).

- - 4Li(s) + O2(g) → 2Li2O(s)

Halogens are very reactive. Fluorine is a pale yellow gas, chlorine is a yellow/green gas, bromine is a red/brown volatile liquid and iodine is a dark grey solid.
The halogens are poorly soluble in water, and aqueous solutions of halogens are acidic and act as bleaches. The halogens dissolve readily in hexane, a non-polar solvent, to give colourless (chlorine), orange (bromine) or purple (iodine) solutions.
Halogens react with metals to form ionic compounds (salts).

- - 2K(s) + Cl2(g) → 2KCl(s)
  - 2Fe(s) + 3Cl2(g) → 2FeCl3(s)

A more reactive halogen will displace a less reactive halogen from one of its compounds.

- - 2KBr(aq) + Cl2(g) → 2KCl(aq) + Br2(g)

The Group 8 elements, the noble gases, are unreactive because they have a full outer shell of electrons.
The transition metals are a block of metals between Groups 2 and 3 of the periodic table. Besides common metallic properties, they can form more than one ion and their compounds are often coloured.

Chemical Reactions – Types, Rates and Energy

Reactions can be classified into different types, including oxidation, reduction, thermal decomposition, neutralisation, displacement, precipitation, cracking, addition and polymerisation reactions (see below).
The rate of a reaction can be calculated by measuring the volume change, colour change, formation of precipitate, mass change, temperature change or pH change during the reaction.
Collision theory – for a reaction to occur, particles must collide and they must collide with sufficient energy. This energy is known as the activation energy.
Increasing the concentration of a solution, the pressure of a gas or the surface area of a solid increases the rate of reaction by increasing the frequency of collisions.
Increasing the temperature increases the rate of reaction by increasing both the frequency and the energy of collisions. A 10 °C increase in temperature approximately doubles the reaction rate.
Catalysts speed up the rate of reaction, but are not themselves consumed in the process. Catalysts are often transition metal compounds. Catalysts work by lowering the activation energy of a reaction.
Enzymes are proteins that act as biological catalysts. Enzymes are responsible for the controlled rate of nearly all chemical reactions in the body. Most enzymes work best at 37 °C and neutral pH. Enzymes can be used in many industrial processes, such as wine or cheese-making.
A reversible reaction is where complete conversion to product (100% yield) does not take place. At equilibrium, there is no change in the concentrations of the reactant and products because the forward and reverse reactions are happening at the same rate.
Exothermic reactions release energy. The enthalpy (or heat) of reaction (ΔH) has a negative value. Endothermic reactions absorb energy. The enthalpy (or heat) of reaction has a positive value.
A reaction is exothermic if the energy required to break the bonds in the reactants is less than the energy released when the bonds in the product are formed, and vice versa for an endothermic reaction.
Average bond enthalpies represent the energy of the bonds, and they can be used to calculate enthalpy.

Water

Water is a very importance substance as it is critical for many fundamental processes. Tap water can be produced by taking water from a clean river or reservoir, followed by filtering to remove solid particles and chlorination to kill bacteria. Fluoride is often added to improve dental health of consumers. Waste water can be recycled by filtering and treatment with bacteria to break down the waste.
The presence of water in a liquid can be determined using anhydrous copper(II) sulphate (white → blue) or cobalt(II) chloride paper (blue → pink). Only pure water freezes at 0 °C and boils at 100 °C.
Distilled water and rain water are soft water, they do not produce scum with soap solution. Hard water contains dissolved calcium and magnesium compounds, which produces scum with soap and forms scales or deposits in kettles and heating elements. Hard water supplies calcium which is good for health.
Temporary hardness, caused by calcium hydrogen carbonate, can be removed by boiling. Permanent hardness, caused by calcium and magnesium sulphate, can be removed by sodium carbonate (washing soda), which precipitates the insoluble carbonates. All hardness can be removed by distillation (expensive) or with an ion-exchange column which replaces calcium and magnesium with sodium.
Only two gases, hydrogen chloride and ammonia, are very soluble in water. The solubility of gases in water increases with pressure and decreases with temperature.
Water is a polar solvent. it has small positive and negative charges. Water is good at dissolving polar molecules and ionic compounds. Non-polar solvents such as hexane or tetrachloromethane (CCl4) are poor at dissolving ionic compounds, but good at dissolving covalent molecular compounds.
Saturated solution – a solution that contains as much solute as can be dissolved.
Solubility generally increases with temperature, and this can be plotted as a solubility curve on a graph.

Acids, Bases and Salts

Acids release protons (hydrogen ions) in water. Mineral acids include hydrochloric, sulphuric and nitric acid. Organic acids include ethanoic acid (vinegar) and citric acid (lemon). Only one proton in ethanoic acid (Ch2COOH) can be released. Dry acids do not show acidic properties, water must be present.
Acids react with fairly reactive metals (e.g. magnesium, zinc) to form a salt and hydrogen.

- - Mg(s) + H2SO4(aq) → MgSO4(aq) + H2(g)

Acids react with metal hydroxides or metal oxides to give a salt and water. Alternatively, acids can react with carbonates to give salt and carbon dioxide and water. Soluble metal hydroxides are called alkalis.

- - HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
  - HCl(aq) + ZnO(aq) → ZnCl2(aq) + H2O(l)
  - HCl(aq) + CaCO3(s) → CaCl2(s) + CO2(g) + H2O(l)

Strong acids such as hydrochloric acid ionise completely in water. Weak acids such as ethanoic acid do not ionise completely in water, only some protons are released and most molecules remain un-ionised.
Neutralisation involves reaction between an acid and alkali to produce a neutral substance. Acidic soil, stomach acid or acid rain can be treated by neutralisation with a weak alkali such as slaked lime, magnesium hydroxide or limestone. Neutralisation can be summarized by the ionic equation:

- - H+(aq) + OH-(aq) → H2O(l)

Soluble salts can be prepared by reaction of a metal, metal oxide, metal hydroxide or metal carbonate with an appropriate acid, followed by filtration and evaporation to a small volume to allow crystals to form (total evaporation does not produce crystals). All metal nitrates, chlorides (except silver, lead) and sulphates (except lead, barium) are soluble. Sodium, potassium and ammonium carbonate are soluble.
Insoluble salts can be prepared by precipitation by mixing two suitable solutions.

- - Ba(NO3)2 + H2SO4(aq) → BaSO4(s) + 2HNO3(aq)

Tests for Ions

Flame tests can be used to test for the presence of many metal ions, including sodium (orange-yellow), potassium (lilac-pink), lithium (red), calcium (brick red), copper (green), lead (blue) and barium (pale green). Some metal ions such as magnesium give no colour.
The ammonium ion (NH4+) can be tested with sodium hydroxide solution. Heating causes the release of ammonia gas which turns damp red litmus paper blue.
Treating a solution of a metal compound with sodium hydroxide solution gives a coloured precipitate (insoluble metal hydroxide) for iron(II) (green), iron(III) (red-brown), copper(II) (blue), calcium (white) and magnesium (white) compounds. Aluminium(III) initially gives a white precipitate, which re-dissolves upon addition of excess sodium hydroxide.
Carbonate ions – tested with dilute hydrochloric acid to produce carbon dioxide gas, which turns lime water milky.
Sulphate ions – tested with dilute hydrochloric acid and barium chloride solution to produce a white precipitate of barium sulphate.
Halide ions – tested with dilute nitric acid and silver nitrate to produce a precipitate of silver chloride (white), bromide (cream) or iodide (yellow).
Sulphite ions – heated with dilute hydrochloric acid produces colourless sulphur dioxide gas, which turns a piece of filter paper soaked with orange potassium dichromate(VI) green.

Collection and Tests for Gases

Gases that are insoluble or slightly soluble in water can be collected over water (inverted test tube).
Gases with a lower density than air can be collected by upward delivery (inverted test tube).
Gases with a higher density than air can be collected by downward delivery (upright test tube).
Oxygen – a relights a glowing splint.
Hydrogen – produces a “squeaky pop” with a lighted splint and the flame is extinguished.
Ammonia – turns damp red litmus paper blue.
Carbon dioxide – turns lime water milky.
Sulphur dioxide – turns orange potassium dichromate(VI) solution green.
Chlorine – turns damp blue litmus paper red then bleaches white.

Redox

Oxidation and reduction reactions occur together – a redox reaction.
Oxidation – gain of oxygen or loss of hydrogen, or more generally, the loss of electrons. Oxidising agents oxidise other substances, but are themselves reduced. Combustion is an oxidation reaction.
Reduction – loss of oxygen or gain of hydrogen, or more generally, the gain of electrons. Reducing agents reduce other substances, but are themselves oxidised.
OILRIG – “oxidation is loss, reduction is gain”.
Half-equations include electrons and show the oxidation and reduction steps separately. In this example, iron is the reducing agent and is itself oxidised. Chlorine is the oxidising agent and is itself reduced.

- - Fe2+ → Fe3+ + e– (oxidation half-equation)
  - Cl2 + 2e– → 2Cl– (reduction half-equation)
  - 2Fe2+ + Cl2 → F2e3+ + 2Cl– (full equation)

Chemical cells produce electricity through a chemical reaction. For example, magnesium and copper rods are dipped into salt solution and connected externally by a wire. Magnesium is oxidised and loses electrons which travel through the wire to the copper. The further apart the metals are in the reactivity series, the higher the voltage of the cell.
A hydrogen-oxygen fuel cell produces electricity efficiently, used in spaceships and some hybrid cars.

Materials from Rocks and Air

Rocks such as limestone, sandstone and slate are used as building materials. Due to costs, new materials such as bricks (clay), mortar (calcium hydroxide, sand, water), cement (limestone, clay), concrete (cement, sand, stones), glass (limestone, sand, sodium carbonate).
Limestone (CaCO3) can be heated to produce quicklime (CaO), which can be mixed with water to produce slaked lime (Ca(OH)2).
Most metals are found in ore deposits. The method used to extract metals depends on their reactivities.
Very reactive metals, such as potassium, sodium, calcium, magnesium, and aluminium form very stable ores that can only be decomposed by electrolysis, an expensive process. Aluminium oxide (bauxite) is dissolved in molten cryolite (Na3AlF6) to lower the melting point and increase the solubility of the ore. Molten aluminium is produced at the cathode, and oxygen is produced at the anode. Oxygen reacts with the carbon of the anode so the anode has to be replaced frequently.
Electrolysis of molten sodium chloride produces sodium and chlorine. Electrolysis of aqueous sodium chloride (brine) produces sodium hydroxide, chlorine and hydrogen.
Metals in the middle of the reactivity series can be extracted by reduction, often with carbon. In a blast furnace, coke (carbon) burns in air to produce carbon monoxide (i), which reacts to produce carbon monoxide (ii). Carbon monoxide reacts with iron oxide (iron ore) to produce iron (iii). Calcium carbonate (limestone) decomposes into carbon dioxide and calcium oxide (iv), which reacts with silicon dioxide impurities to produce calcium silicate slag (v). Slag can be used as a fertiliser and for roads.
- - - C(s) + O2(g) → CO2(g)
    - CO2(g) + C(s) → CO2(g)
    - Fe2O3(s) + 3CO(g) → 2Fe(l) + 3CO2(g)
    - CaCO3(g) → CaO(s) + CO2(g)
    - CaO(s) + SiO2(s) → CaSiO3(l)
Metals low on the reactivity series, such as mercury, can be extracted by heating because the ores are unstable. Some very unreactive metals, such as gold are found uncombined on Earth.
Nitrogen is produced from the fractional distillation of air and hydrogen is produced as a by-product of cracking (see below). These are combined to make ammonia in the Haber process.

- - N2 + 3H2 ↔ 2Nh2

This reaction is reversible and 100% yield is not possible. A temperature of 450 °C and a pressure of 200 atmospheres gives the best balance between reaction rate and yield, which is about 10%. A finely divided iron catalyst speeds up the rate of reaction.
Plants require large quantities of nitrogen, phosphorus and potassium. These have been traditionally obtained from natural sources, but have now been largely replaced by cheaper, artificial fertilisers.
Ammonia can be converted into ammonium nitrate, ammonium sulphate or urea, which are nitrogen fertilisers. Ammonium phosphate provides phosphorus and potassium sulphate provides potassium.
Overuse of nitrogen fertilisers can lead to eutrophication. Nitrates washed into rivers or lakes causes the overgrowth of plants and algae, which eventually die and are decomposed by bacteria. Bacteria multiply and consume all the oxygen, leading to the death of fish and other lifeforms.

Metals

Rusting is the oxidation of iron into hydrated iron(III) oxide, Fe2O3.xH2O. Both oxygen and water are required for rusting. Rusting is accelerated by salt or acid, including carbon dioxide and sulphur dioxide.
Rusting can be prevented by painting, oiling/greasing or tin coating, which prevents oxygen and water from coming in contact with the iron, but only while the covering is intact.
Sacrificial protection for ships uses blocks of magnesium or zinc, which are more reactive than iron and will rust in preference to iron. As long as they are continually replaced the iron will not rust.
Electroplating can be used to coat steel with nickel or chromium, which protects the steel and also gives a decorative appearance. Stainless steel does not rust as easily as other types of steel.
Aluminium is lower than magnesium, but higher than zinc in the reactivity series. However, aluminium is unsurprisingly unreactive due to a thin layer of unreactive aluminium oxide rust coated on the surface. If this layer is removed aluminium becomes more reactive. Anodising is an electrolytic process where the oxygen generated at the anode reacts with the aluminium to form a thicker layer of aluminium oxide.
Metals such as copper can be purified by electrophoresis. Pure copper is connected at the cathode and the impure copper is connected at the anode. Copper from the anode goes into solution as copper ions which are deposited at the cathode. Impurities either remain in solution or collect in the anode mud.
Alloys are harder than pure metals because the atoms are of different sizes and cannot slide easily over one other. Alloys have many applications because they have better properties for many uses.
Steel is an alloy of iron with a small amount of carbon. Brass is an alloy of copper and zinc. Bronze is an alloy of copper and tin. Solder is an alloy of tin and lead.
Most iron is converted into steel, which involves blowing pure oxygen onto the surface of molten iron to remove carbon and phosphorous impurities, which escape as gases. Limestone is added to remove other impurities as slag. Finally, carbon and other elements are added to give the desired type of steel.

Carbon Chemistry

Crude oil is a mixture of hydrocarbons, which are compounds of carbon and hydrogen only. Alkanes form a homologous series with formula CnH2n+2. Alkanes are saturated hydrocarbons (only single carbon-carbon bonds), burn in air or oxygen, and have few other reactions.
Crude oil is separated into separate in an oil refinery by fractional distillation. The lighter fractions are used for fuel and the lower fractions are used for heating and road tar.
Complete combustion occurs in excess air or oxygen to give carbon dioxide and water.

- - CH4 + 2O2 → CO2 + 2H2O

Incomplete combustion occurs in a limited supply of air to produce water and carbon monoxide (toxic).

- - 2CH4 + 3O2 → 2CO + 4H2O

Higher boiling point fractions are usually in lower demand, so these long chains are broken into shorter molecules such as ethene (C2H4) by passing over a catalyst at high pressure, which is known as cracking.
Alkenes are unsaturated hydrocarbons (one double carbon-carbon bond) that form a homologous series with formula CnH2n. Alkenes decolourise (red-brown) bromine water, an example of an addition reaction.

- - C2H4 + Br2 → C2H4Br2

Unsaturated oils (alkenes) such as sunflower oil can be hardened by hydrogenation by heating over a nickel catalyst at 170 °C, producing margarine.
Alkanes with four or more carbons can exist as isomers, which are compounds with the same molecular formulae, but different structures. Isomerism becomes more common with higher alkanes, and can also occur with compounds in different homologous series.
Alcohols are a homologous series with formula CnH2n+1OH. Ethanol (C2H5OH) can be produced from the hydration of ethene by passage with steam over a phosphoric acid catalyst at 600 °C, a continuous process. Ethanol can also be produced from the fermentation of sugar using enzymes in yeast (a batch process). The method of production depends on the abundance of raw materials. Hydration is fast and produces pure ethanol, but uses a non-renewable energy resource and requires a lot of energy. Fermentation uses renewable resources, but is slow and requires large reaction volumes and vessels, and produces only dilute ethanol, which can be concentrated by fractional distillation.

- - C2H4 + H2O → C2H5OH (hydration)
  - C6H12O6 → 2C2H5OH + 2CO2 (fermentation)

Ethanol is used a a solvent, as a fuel, to make other organic chemicals, and to make alcoholic drinks. Th toxic methanol is often added to pure ethanol to make it undrinkable. Consuming high amounts of ethanol can lead to harmful health effects.
Carboxylic acids such as ethanoic acid are weak acids, and react with bases in the same way as mineral acids. Vinegar is a dilute solution of ethanoic acid.
The reaction of an alcohol with a carboxylic acid produces an ester and water. Esters are sweet-smelling liquids that are as flavouring agents. Natural fats and oils are esters of long chain acids and glycerol.

- - Ch2COOH + C2H5OH → Ch2COOC2H5 + H2O
  - ethanoic acid + ethanol → ethyl ethanoate + water

Ethene (the monomer unit) can be polymerized by passing over a heated catalyst to make poly(ethene), an addition polymer. Addition polymers such as poly(ethene) and poly(vinyl) chloride have many uses.
Thermoplastic polymers such as poly(ethene) are composed of unlinked polymer chains, and so they melt easily when heated. Thermosetting polymers such as Bakelite are composed of strongly cross-linked polymer chains, and they do not melt when heated. On further heating they decompose.
Polymers are non-biodegradable, and recycling is often not-economical compared to making new polymers. Polymers can be incinerated to produce carbon dioxide, water and potentially toxic gases.