Chemistry Unit 2

Ions are electrically charged particles formed when atoms lose or gainelectrons. They have the same electronic structures as noble gases.

Metal atoms form positive ions, while non-metal atoms form negative ions. The strong electrostatic forces of attraction between oppositely charged i

Ions are electrically charged particles formed when atoms lose or gainelectrons. They have the same electronic structures as noble gases.

Metal atoms form positive ions, while non-metal atoms form negative ions. The strong electrostatic forces of attraction between oppositely charged ions are called ionic bonds.

Ions

How ions form

Ions are electrically charged particles formed when atoms lose or gain electrons. This loss or gain leaves a complete highest energy level, so the electronic structure of an ion is the same as that of a noble gas - such as a helium, neon or argon.

Metal atoms and non-metal atoms go in opposite directions when they ionise:

• Metal atoms lose the electron, or electrons, in their highest energy level and become positively charged ions
• Non-metal atoms gain an electron, or electrons, from another atom tobecome negatively charged ions

Positively charged sodium and aluminium ions

Negatively charged oxide and chloride ions

How many charges?

There is a quick way to work out what the charge on an ion should be:

• The number of charges on an ion formed by a metal is equal to the group number of the metal
• The number of charges on an ion formed by a non-metal is equal to the group number minus eight
• Hydrogen forms H⁺ ions.

	Group 1	Group 2	Group 3	Group 4	Group 5	Group 6	Group 7	Group 0
Example element	Na	Mg	Al	C	N	O	Cl	He
Charge	1+	2+	3+	Note 1	3-	2-	1-	Note 2
Symbol of ion	Na+	Mg2+	Al3+	Note 1	N3-	O2-	Cl-	Note 2

Note 1: carbon and silicon in Group 4 usually form covalent bonds by sharing electrons.

Note 2: the elements in Group 0 do not react with other elements to form ions.

Representing positive ions

You need to be able to show the electronic structure of some common metal ions, using diagrams like these:

Lithium, Li

Lithium is in Group 1. It has one electron in its highest energy level. When this electron is lost, a lithium ion Li⁺ is formed.

Sodium, Na

Sodium is also in Group 1. It has one electron in its highest energy level. When this electron is lost, a sodium ion Na⁺ is formed.

Neon atom

Note that a sodium ion has the same electronic structure as a neon atom (Ne).

But be careful - a sodium ion is not a neon atom. This is because the nucleus of a sodium ion is the nucleus of a sodium atom and has 11 protons - but the nucleus of a neon atom has only 10.

Magnesium, Mg

Magnesium is in Group 2. It has two electrons in its highest energy level. When these electrons are lost, a magnesium ion Mg²⁺ is formed.

A magnesium ion has the same electronic structure as a neon atom (Ne).

Calcium, Ca

Calcium is also in Group 2. It has two electrons in its highest energy level. When these electrons are lost, a calcium ion Ca²⁺ is formed.

A calcium ion has the same electronic structure as an argon atom (Ar).

Representing negative ions

You need to be able to show the electronic structure of some common non-metal ions, using diagrams like these:

Fluorine, F

Fluorine is in Group 7. It has seven electrons in its highest energy level. It gains an electron from another atom in reactions, forming a fluoride ion, F^-.

Note that the atom is called fluorine, but the ion is called fluoride.

Neon atom

Note that a fluoride ion has the same electronic structure as a neon atom (Ne).

Once again, a fluoride ion is not a neon atom, because thenucleus of a fluoride ion is the nucleus of a fluorine atom, with 9 protons, and not of a neon atom, with 10.

Chlorine, Cl

Chlorine is in Group 7. It has seven electrons in its highest energy level. It gains an electron from another atom in reactions, forming a chloride ion, Cl^-.

Oxygen, O

Oxygen is in Group 6. It has six electrons in its highest energy level. It gains two electrons from one or two other atoms in reactions, forming an oxide ion, O^2-.

Ionic bonding

When metals react with non-metals, electrons are transferred from the metal atoms to the non-metal atoms, forming ions. The resulting compound is called an ionic compound.

Consider reactions between metals and non-metals, for example:

• sodium + chlorine → sodium chloride
• magnesium + oxygen → magnesium oxide
• calcium + chlorine → calcium chloride

In each of these reactions, the metal atoms give electrons to the non-metal atoms. The metal atoms become positive ions and the non-metal atoms become negative ions.

There is a strong electrostatic force of attraction between these oppositely charged ions, called an ionic bond. The animation shows ionic bonds being formed in sodium chloride, magnesium oxide and calcium chloride.

Ionic compounds

Group 1 and Group 7

The elements in Group 1 of the Periodic Table are called the alkali metals. They form ionic compounds when they react with non-metals. Their ions have a single positive charge. For example, sodium forms sodium ions, Na+.

The elements in Group 7 of the Periodic Table are called the halogens. They form ionic compounds when they react with metals. Their ions have a single negative charge. For example, chlorine forms chloride ions, Cl–.

Sodium chloride

Sodium chloride, NaCl, forms when sodium and chlorine react together. It contains oppositely charged ions held together by strongelectrostatic forces of attraction – the ionic bonds. The ions form a regular lattice in which the ionic bonds act in all directions.

Dot-and-cross diagrams

You need to be able to draw dot-and-cross diagrams to show the ions in some common ionic compounds.

Sodium chloride, NaCl

Sodium ions have the formula Na⁺, while chloride ions have the formula Cl^-. You need to show one sodium ion and one chloride ion. In the exam, make sure the dots and crosses are clear, but do not worry about colouring them.

Magnesium oxide, MgO

Magnesium ions have the formula Mg²⁺, while oxide ions have the formula O^2-. You need to show one magnesium ion and one oxide ion.

Calcium chloride, CaCl₂

Calcium ions have the formula Ca²⁺. Chloride ions have the formula Cl^-.

You need to show two chloride ions, because two chloride ions are needed to balance the charge on a calcium ion.

Formulae of ionic compounds

Ionic compounds are represented by formulae. Symbols and numbers show the atoms in the compound.

For example: ZnCO₃ is the formula for zinc carbonate.

One zinc atom (Zn) and one carbon atom (C) are chemically bonded with three oxygen atoms (O₃). Notice that we don't need to write a 1 next to the Zn or C.

Brackets in formulae

Sometimes brackets are used.

For example: Fe(OH)₃ is the formula for iron(III) hydroxide.

Iron(III) hydroxide consists of one iron atom joined with three oxygen and three hydrogen atoms. The formula is written like this because the oxygen and hydrogen atom often act together.

Constructing formulae

The formula of a compound can be worked out if the ions in it are known. For example, the compound formed from Na⁺ and SO₄^2- will consist of two Na⁺ ions to every one SO₄^2- ion so that the compound is neutral overall. The formula is therefore Na₂SO₄.

Here are the formulae of some common ions.

Positive ions (cations)

Name	Formula	Name	Formula	Name	Formula
ammonium	NH4+	magnesium	Mg2+	zinc	Zn2+
hydrogen	H+	calcium	Ca2+	lead	Pb2+
lithium	Li+	barium	Ba2+	iron(II)	Fe2+
sodium	Na+	silver	Ag+	iron(III)	Fe3+
potassium	K+	copper(II)	Cu2+	aluminium	Al3+

Negative ions (anions)

Name	Formula	Name	Formula	Name	Formula
fluoride	F-	hydrogen carbonate	HCO3-	sulfide	S2-
chloride	Cl-	hydroxide	OH-	sulfate	SO42-
bromide	Br-	nitrate	NO3-	carbonate	CO3-
iodide	I-	oxide	O2-

A covalent bond is a strong bond between two non-metal atoms. It consists of a shared pair of electrons. A covalent bond can be represented by a straight line or dot-and-cross diagram.

Hydrogen and chlorine can each form one covalent bond, oxygen two bonds, nitrogen three, while carbon can form four bonds.

A shared pair of electrons

You will need to understand what covalent bonding is, and to remember some of the properties of molecules that are formed in this way.

A covalent bond forms when two non-metal atoms share a pair of electrons. The electrons involved are in the highest occupied energy levels - or outer shells - of the atoms. An atom that shares one or more of its electrons will complete its highest occupied energy level.

Covalent bonds are strong - a lot of energy is needed to break them. Substances with covalent bonds often form molecules with low melting and boiling points, such as hydrogen and water.

The animation shows a covalent bond being formed between a hydrogen atom and a chlorine atom, to form hydrogen chloride.

After bonding, the chlorine atom is now in contact with eight electrons in its highest energy level - so it is stable. The hydrogen atom is now in contact with two electrons in its highest energy level - so the hydrogen is also stable.

How many bonds?

Atoms may form multiple covalent bonds - that is, share not just one pair of electrons but two or more pairs. Atoms of different elements will form either one, two, three or four covalent bonds with other atoms.

There is a quick way to work out how many covalent bonds an element will form. The number of covalent bonds is equal to eight minus the group number (you can brush up on group numbers by reading through the section in AQA GCSE Science on thePeriodic Table). The table below gives more detail on this rule:

	Group 4	Group 5	Group 6	Group 7
Example	Carbon	Nitrogen	Oxygen	Chlorine
Number of bonds	8 - 4 = 4	8 - 5 = 3	8 - 6 = 2	8 - 7 = 1

Hydrogen forms one covalent bond. The noble gases in Group 0 do not form any.

Representing covalent bonds

Covalent bonds can be represented in several different ways.

Straight lines and models

Straight lines are the most common way to represent covalent bonds, with each line representing a shared pair of electrons. 2D or 3D molecular models are especially useful for showing the relationship between atoms in multiple covalent bonds. Below are some examples of straight lines and images of 3D models.

Models for covalent bonds

Element	Formula	Chemical structure	Ball-and-stick model
Hydrogen	H2
Water	H2O
Ammonia	NH3
Methane	CH4

Double and triple bonds

Note that molecules can have a double covalent bond - meaning they have two shared pairs of electrons - or a triple covalent bond - three shared pairs of electrons. A double covalent bond is shown by a double line, and a triple bond by a triple line.

A molecule of oxygen (O₂) consists of two oxygen atoms held together by a double bond, like this: 

A molecule of nitrogen (N₂) has two nitrogen atoms held together by a triple bond, like this: 

Dot-and-cross diagrams - elements

Dot-and-cross diagrams

Dot-and-cross diagrams are used to represent covalent bonds. The shared electron from one atom is shown as a dot, while the shared electron from the other atom is shown as a cross.

When drawing dot-and-cross diagrams for covalent bonds, you only need to show the electrons in the highest occupied energy level, as only these are involved.

The animations show covalent bonds represented by both displayed formulae (which use straight lines to represent bonds) and dot-and-cross diagrams:

Covalent bonding between two hydrogen atoms to form a molecule of hydrogen gas, H₂.

Covalent bonding between two oxygen atoms to form a molecule of oxygen gas, O₂.

Elements

For your examination, you need to be able to draw dot-and-cross diagrams for hydrogen, chlorine and oxygen.

You do not need to use colours in your answers.

Dot-and-cross diagrams - compounds

You will also need to be able to draw dot-and-cross diagrams representing thecovalent bonds in the molecules of some common compounds:

Hydrogen chloride, HCl

Hydrogen atoms and chlorine atoms can each form one covalent bond. One pair ofelectrons is shared in a hydrogen chloride molecule (HCl).

Water, H₂O

Hydrogen atoms can each form one covalent bond, while oxygen atoms can each form two covalent bonds. Two pairs of electrons are shared in a water molecule (H₂O).

Ammonia, NH₃

Hydrogen atoms can each form one covalent bond, while nitrogen atoms can each form three covalent bonds. Three pairs of electrons are shared in an ammonia molecule (NH₃).

Methane, CH₄

Hydrogen atoms can each form one covalent bond, while carbon atoms can each form four covalent bonds. Four pairs of electrons are shared in a methane molecule (CH₄).

Simple molecular substances consist of molecules in which the atoms are joined by strong covalent bonds. However, the molecules are held together by weak forces so these substances have low melting and boiling points. They do not conduct electricity.

Giant covalent structures contain many atoms joined together by covalent bonds to form a giant lattice. They have high melting and boiling points. Graphite and diamond have different properties because they have different structures. Graphite conducts heat and electricity well because it also has free electrons.

Simple molecules

Covalent bonds form between non-metal atoms. Each bond consists of a shared pair of electrons, and is very strong. Covalently bonded substances fall into two main types:

1. Simple molecules
2. Giant covalent structures

Simple molecules

A molecule of carbon dioxide

These contain only a few atoms held together by strong covalent bonds. An example is carbon dioxide (CO₂), the molecules of which contain one atom of carbon bonded with two atoms of oxygen.

Properties of simple molecular substances

• Low melting and boiling points - This is because the weak intermolecular forces break down easily.
• Non-conductive - Substances with a simple molecular structure do notconduct electricity. This is because they do not have any free electrons or an overall electric charge.

Higher tier only

Hydrogen, ammonia, methane and water are also simple molecules with covalent bonds. All have very strong bonds between the atoms, but much weaker forces holding the molecules together. When one of these substances melts or boils, it is these weak 'intermolecular forces' that break, not the strong covalent bonds. Simple molecular substances are gases, liquids or solids with low melting and boiling points.

Macromolecules

Macromolecules have giant covalent structures. They contain a lot of non-metal atoms, each joined to adjacent atoms by covalent bonds. Their atoms are arranged into giant lattices, which are strong structures because of the many bonds involved. Substances with giant covalent structures have very high melting points, because a lot of strong covalent bonds must be broken. Graphite, for example, has a melting point of more than 3,600°C.

Diamond

Diamond

Diamond is a form of carbon in which each carbon atom is joined to four other carbon atoms, forming a giant covalent structure. As a result, diamond is very hard and has a high melting point. It does not conduct electricity.

Graphite

Graphite

Graphite is a form of carbon in which the carbon atoms form layers. Each carbon atom in a layer is joined to only three other carbon atoms.The layers can slide over each other because there are no covalent bonds between them. This makes graphite much softer than diamond. It is used in pencils and as a lubricant. Graphite conducts electricity.

Silica

Silica

Silica, which is found in sand, has a similar structure to diamond. It is also hard and has a high melting point. However, it contains silicon and oxygen atoms instead of carbon atoms.

Polymers

Polymers have properties which depend on the chemicals they are made from, and the conditions in which they are made. For example, poly(ethene) can be low-density or high-density depending upon the catalyst and reaction condition used to make it. The table summarises some differences in their properties:

	LDPE low-density poly(ethene)	HDPE high-density poly(ethene)
Branches on polymer molecules	Many	Few
Relative strength	Weak	Strong
Maximum useable temperature	85°C	120°C

Thermosoftening polymers

Polymer with no cross-links

Thermosoftening polymers soften when heated and can be shaped when hot. The shape will harden when it is cooled, but can be reshaped when heated up again. Poly(ethene) is a thermosoftening polymer. Its tangled polymer chains can uncoil and slide past each other, making it a flexible material.

Thermosetting polymers

Polymer with cross-links

Thermosetting polymers have different properties to thermosoftening polymers. Once moulded, they do not soften when heated and they cannot be reshaped. Vulcanised rubber is athermoset used to make tyres. Its polymer chains are joined together by cross-links, so they cannot slide past each other easily.

Ionic compounds

Ionic bonds form when a metal reacts with a non-metal. Metals form positive ions, while non-metals form negative ions. Ionic bonds are the electrostatic forces of attraction between oppositely charged ions.

Melting points and boiling points

Ionic bonds are very strong so a lot of energy is needed to break them. Ionic compounds contain many of these strong bonds so they have high melting and boiling points.

Conduction of electricity

Ionic compounds conduct electricity when they are dissolved in water or when they are melted. This is because their ions are free to move and carry the current. However, ionic compounds do not conduct electricity when they are solid. This is because their ions cannot move around in their lattice structure.

Metals

Layers of atoms slide over each other when metals are bent or stretched

Metals are malleable - they can be bent and shaped. This is because they consist of layers ofatoms. These layers can slide over one another when the metal is bent, hammered or pressed.

Metals - Higher tier

Metals form giant structures in which electrons in the outer shells of the metal atoms are free to move. The metallic bond is the force of attraction between these free electrons and metal ions. Metallic bonds are strong, so metals can maintain a regular structure and usually have high melting and boiling points.

Metals are good conductors of electricity and heat. This is because the free electrons can move throughout the metal.

Alloys

An alloy is a mixture of two or more elements, where at least one element is a metal. Many alloys are mixtures of two or more metals.

Layers

Alloys contain atoms of different sizes. These different sizes distort the regular arrangements of atoms. This makes it more difficult for the layers to slide over each other, so alloys are harder than the pure metal.

It is more difficult for layers of atoms to slide over each other in alloys

Copper, gold and aluminium are too soft for many uses. They are mixed with other metals to make them harder for everyday use. For example:

• Brass - used in electrical fittings - is 70 per cent copper and 30 per cent zinc
• 18-carat gold - used in jewellery - is 75 per cent gold and 25 per cent copper and other metals
• Duralumin - used in aircraft manufacture - is 96 per cent aluminium and 4 per cent copper and other metals

Shape memory alloys

Shape memory alloys can return to their original shape after being bent or twisted. Nitinol is a shape memory alloy made from nickel and titanium. It is used in dental braces and spectacle frames.

Nanoscience

Working with nanoparticles is called nanotechnology

A nanometre, 1 nm, is one billionth of a metre (or a millionth of a millimetre). Nanoparticles range in size from about 100 nm down to about 1 nm. They are typically the size of smallmolecules, and far too small to see with a microscope.

Properties and uses of nanoparticles

Nanoparticles have a very large surface area compared with their volume, so they are often able to react very quickly. This makes them useful as catalysts to speed up reactions. They can, for example, be used in self-cleaning ovens and windows.

Nanoparticles also have different properties to the same substance in normal-sized pieces. For example, titanium dioxide is a white solid used in house paint and certain sweet-coated chocolates. Titanium dioxide nanoparticles are so small that they do not reflect visible light, so cannot be seen. They are used in sun screens to block harmful ultraviolet light without appearing white on the skin.

In addition to new cosmetics such as sun screens and deodorants, nanoscience may lead to the development of:

• New catalysts
• New coatings
• New computers
• Stronger and lighter building materials
• Sensors that detect individual substances in tiny amounts

Graphite and fullerenes - Higher tier

Graphite

Graphite is soft and slippery because there are only weak intermolecular forces between its layers.

From left to right - graphite, diamond, silica

Graphite is a good conductor of heat and electricity. This is because, like metals, graphite contains delocalised electrons. These electrons are free to move through the structure of the graphite.

Fullerenes

Structure of a buckminsterfullerene molecule - a large ball of 60 atoms

Carbon exists as graphite and diamond, but it can also form fullerenes. These are cages and tubes with different number of carbon atoms. Buckminsterfullerene is one type of fullerene. Its molecules are spherical and contain 60 carbon atoms.

Fullerenes may be used for drug delivery systems in the body, in lubricants and as catalysts.

The tube fullerenes are called nanotubes. These are very strong. They are useful in reinforcing structures where lightness and strength are needed - for example, in tennis racket frames.

Atomic structure, analysis and quantitative chemistry


Mass number = number of protons and neutrons.

Isotopes are atoms of an element with different numbers of neutrons.

The relative formula mass = Relative atomic mass of all atoms in the formula of the compound.

The relative formula mass in grams is one mole of the substance.

Nucleus has protons and neutrons

Electrons are arranged around it in energy levels.

Relative mass of particle

Neutron 1

Proton 1

Electron very small 1/1836

Atomic number = number of protons.

Atoms of an element will have the same number of protons. Atomic number = same.

They can have different neutrons so the mass number changes (isotopes)

Isotopes have identical chemical properties but some are radioactive.

Mr is relative formula mass

Ar is relative atomic mass

One mole of the substance is the Mr in grams.

Method of analysis


Paper chromatography

Used to analyse coloured substances e.g. coloured pigments and artificial colours.

This work as coloured substances have different dissolving rates in  liquid and bonding with paper.

Substances that are better at dissolving than bonding with paper travel further.

Substances that travel the same distance with the same colour are likely to be the same.


Instrumental method of analysis


This method uses machines they are:

Fast

Accurate (reliable)

Sensitive (can detect small amounts in a small sample)

Gas chromatography

A mixture of compounds can be separated.

It has a long glass tube called a column with a powdered solid material inside that is put in an oven.

The sample is dissolved in a solvent and then injected into one end of the column.

An unreactive gas usually nitrogen carries the sample through

The different substances travel at different speeds so they are separated.

The separated substances leave the column one after the other, they are detected by a detector.

Peaks show number of compounds

Position of peaks show the retention time for each compound



Mass spectrometry



It identifies substances very quickly and accurately in small amounts.

It provides the relative formula mass of the substances separated by gas chromatography.

The peak furthest to the right is the molecular ion peak. this is the relative formula mass of the substance.

GC is usually linked with MS so substances can be separated and identified

The rate of a reaction can be measured by the rate at which a reactant is used up, or the rate at which a product is formed.

The temperature, concentration, pressure of reacting gases, surface area of reacting solids, and the use of catalysts, are all factors which affect the rate of a reaction.

Chemical reactions can only happen if reactant particles collide with enough energy. The more frequently particles collide, and the greater the proportion of collisions with enough energy, the greater the rate of reaction.

Measuring rates

Different reactions can happen at different rates. Reactions that happen slowly have a low rate of reaction. Reactions that happen quickly have a high rate of reaction. For example, the chemical weathering of rocks is a very slow reaction: it has a low rate of reaction. Explosions are very fast reactions: they have a high rate of reaction.

Reactants and products

There are two ways to measure the rate of a reaction:

1. Measure the rate at which a reactant is used up
2. Measure the rate at which a product is formed

The method chosen depends on the reaction being studied. Sometimes it is easier to measure the change in the amount of a reactant that has been used up; sometimes it is easier to measure the change in the amount of product that has been produced.

Things to measure

The measurement itself depends on the nature of the reactant or product:

• The mass of a substance - solid, liquid or gas - is measured with a balance
• The volume of a gas is usually measured with a gas syringe, or sometimes an upside down measuring cylinder or burette

It is usual to record the mass or total volume at regular intervals and plot a graph. The readings go on the vertical axis, and the time goes on the horizontal axis.

For example, if 24 cm³ of hydrogen gas is produced in two minutes, the mean rate of reaction = 24 ÷ 2 = 12 cm³ hydrogen / min.

Factors affecting the rate

You will be expected to remember the factors that affect the rate of reactions, and to plot or interpret graphs from rate experiments.

How to increase the rate of a reaction

The rate of a reaction increases if:

• The temperature is increased
• The concentration of a dissolved reactant is increased
• The pressure of a reacting gas is increased
• Solid reactants are broken into smaller pieces
• A catalyst is used

Rate of reaction and changing conditions

The graph above summarises the differences in the rate of reaction at different temperatures, concentrations and size of pieces. The steeper the line, the greater the rate of reaction. Reactions are usually fastest at the beginning, when the concentration of reactants is greatest. When the line becomes horizontal, the reaction has stopped.

Collisions and reactions

You will be expected to explain, in terms of particles and their collisions, why changing the conditions of a reaction changes its rate.

Collisions

For a chemical reaction to occur, the reactant particles must collide. Collisions with too little energy do not produce a reaction.

The collision must have enough energy for the particles to react. The minimum energy needed for particles to react is called the activation energy.

Changing concentration or pressure

If the concentration of a dissolved reactant is increased, or the pressure of a reacting gas is increased:

• There are more reactant particles in the same volume
• There is a greater chance of the particles colliding
• The rate of reaction increases

Changing particle size

If a solid reactant is broken into small pieces or ground into a powder:

• Its surface area is increased
• More particles are exposed to the other reactant
• There is a greater chance of the particles colliding
• The rate of reaction increases

Changing the temperature

If the temperature is increased:

• The reactant particles move more quickly
• More particles have the activation energy or greater
• The particles collide more often, and more of the collisions result in a reaction
• The rate of reaction increases

Using a catalyst

Catalysts increase the rate of reaction without being used up. They do this by lowering the activation energy needed. With a catalyst, more collisions result in a reaction, so the rate of reaction increases. Different reactions need different catalysts.

Catalysts are important in industry because they reduce costs.

Exothermic reactions transfer energy to the surroundings. Endothermic reactions take in energy from the surroundings.

Reversible reactions are where the products can react to remake the original reactants. If the forward reaction is exothermic, the reverse reaction is endothermic.

Exothermic reactions

When a chemical reaction occurs, energy is transferred to or from the surroundings - and there is often a temperature change.

Exothermic reactions transfer energy to the surroundings. The energy is usually transferred as heat energy, causing the reaction mixture and its surroundings to become hotter. The temperature increase can be detected using a thermometer. Some examples of exothermic reactions are:

• Combustion (burning)
• Many oxidation reactions, for example rusting
• Neutralisation reactions between acids and alkalis

When a flame burns it transfers heat to its surroundings.

Exothermic reactions can be used for everyday purposes. For example, hand warmers and self-heating cans for drinks (such as coffee) use exothermic reactions.

Endothermic reactions

These are reactions that take in energy from the surroundings. The energy is usually transferred as heat energy, causing the reaction mixture and its surroundings to get colder. The temperature decrease can also be detected using a thermometer.

Some examples of endothermic reactions are:

• Electrolysis
• The reaction between ethanoic acid and sodium carbonate
• The thermal decomposition of calcium carbonate in a blast furnace

Endothermic reactions can be used for everyday purposes. For example, certain sports injury cold packs use endothermic reactions.

The animation shows an exothermic reaction between sodium hydroxide and hydrochloric acid, and an endothermic reaction between sodium carbonate and ethanoic acid.

Reversible reactions

In reversible reactions, the reaction in one direction will be exothermic and the reaction in the other direction will be endothermic.

The decomposition of ammonium chloride is a reversible reaction:

ammonium chlorideammonia + hydrogen chloride

Ammonium chloride decomposes when it is heated, so the forward reaction is endothermic - energy must be transferred from the surroundings for it to happen. The backward reaction is exothermic - energy is transferred to the surroundings when it happens.

Copper sulfate

The reaction between anhydrous copper sulfate and water is reversible:

hydrated copper sulfate (blue)anhydrous copper sulfate (white) + water

Water is driven off from hydrated copper sulfate when it is heated, so the forward reaction is endothermic - energy must be transferred from the surroundings for it to happen. The backward reaction is exothermic - energy is transferred to the surroundings when it happens. This is easily observed. When water is added to anhydrous copper sulfate, enough heat is released to make the water bubble and boil.

Acids have a pH of less than 7. Bases have a pH of more than 7. When bases are dissolved in water, they are known as alkalis. Salts are made when an acid reacts with a base, carbonate or metal. The name of the salt formed depends on the metal in the base and the acid used. For example, salts made using hydrochloric acid are called chlorides.

Acids and bases

Diagram of pH scale and universal indicator colours

Acids

Substances with a pH of less than 7 are acids. The more strongly acidic the solution, the lower its pH number. Acidic solutions turn blue litmus paper red. They turn universal indicator paper red if they are strongly acidic, and orange or yellow if they are weakly acidic.

Bases

Substances that can react with acids and neutralise them to make a salt and water are called bases. They are usually metal oxides or metal hydroxides. For example, copper oxide and sodium hydroxide are bases.

Alkalis

Bases that dissolve in water are called alkalis. Copper oxide is not an alkali because it does not dissolve in water. Sodium hydroxide is an alkali because it does dissolve in water.

Alkaline solutions have a pH of more than 7. The stronger the alkali, the higher the pH number. Alkalis turn red litmus paper blue. They turn universal indicator paperdark blue or purple if they are strongly alkaline, and blue-green if they areweakly alkaline.

Neutral solutions

Neutral solutions have a pH of 7. They do not change the colour of litmus paper, but they turn universal indicator paper green. Water is neutral.

Neutralisation reactions

Ions are charged particles which are formed when atoms, or groups of atoms, lose or gain electrons. For the examination, you need to know which ions are produced by acids, and which are produced by alkalis. You will also need to know the ionic equation for neutralisation.

State symbols

State symbols are used in symbol equations:

• (s) means solid
• (l) means liquid (not the same as dissolved in water - see below)
• (g) means gas
• (aq) means aqueous (dissolved in water)

Acids

When acids dissolve in water they produce aqueous hydrogen ions, H⁺(aq). For example, looking at hydrochloric acid:

HCl(aq) → H⁺(aq) + Cl^–(aq)

Alkalis

When alkalis dissolve in water they produce aqueous hydroxide ions, OH^–(aq). For example, looking at sodium hydroxide:

NaOH(aq) → Na⁺(aq) + OH^–(aq)

Ammonia is slightly different. This is the equation for ammonia in solution:

NH₃(aq) + H₂O(l) → NH₄⁺(aq) + OH^–(aq)

Be careful to write OH^– and not Oh^– or oh^–.

Neutralisation reaction

When the H⁺(aq) ions from an acid react with the OH^–(aq) ions from an alkali, a neutralisation reaction occurs to form water. This is the equation for the reaction:

H⁺(aq) + OH^–(aq) → H²O(l)

For example, hydrochloric acid and sodium hydroxide solution react together to form water and sodium chloride solution. The acid contains H⁺ ions and Cl^– ions, and the alkali contains Na⁺ ions and OH^– ions. The H⁺ ions and OH^– ions produce the water, and the Na⁺ ions and Cl^– ions produce the sodium chloride, NaCl(aq).

Making soluble salts

You need to be able to describe the reactions of acids with bases and metals. You should be able to work out the particular salt formed in the reaction.

Acids and bases

When acids react with bases, a salt and water are made:

• acid + metal oxide → salt + water
• acid + metal hydroxide → salt + water

Remember that most bases do not dissolve in water. But if a base can dissolve in water, it is also called an alkali.

Reactive metals

Acids will react with reactive metals, such as magnesium and zinc, to make a salt and hydrogen:

• acid + metal → salt + hydrogen

The hydrogen causes bubbling during the reaction, and can be detected using a lighted splint.

Naming salts

The name of the salt produced in a neutralisation reaction can be predicted. The first part of the name is 'ammonium' if the base used is ammonia. Otherwise, it is the name of the metal in the base. The second part of the name comes from the acidused:

• Chloride (if hydrochloric acid is used)
• Nitrate (if nitric acid is used)
• Sulfate (if sulfuric acid is used)

The table shows some examples:

Acid	+	Base	→	Salt + Water
Hydrochloric acid	+	Copper oxide	→	Copper chloride + water
Sulfuric acid	+	Sodium hydroxide	→	Sodium sulfate + water
Nitric acid	+	Calcium hydroxide	→	Calcium nitrate + water

Ammonium salts

Many artificial fertilisers are ammonium salts, made by the reaction of an acid with ammonia solution. For example:

Acid	Alkali	Fertiliser
Nitric acid	Ammonia solution	Ammonium nitrate
Phosphoric acid	Ammonia solution	Ammonium phosphate
Sulfuric acid	Ammonia solution	Ammonium sulfate

Crystallising salt solutions

You may be asked to describe how to make a soluble salt.

If the base dissolves in water, you need to add just enough acid to make a neutral solution. Check a small sample with universal indicator paper. If ammonia solution is used, you can add a little more than needed to get a neutral solution.

Warm the salt solution to evaporate the water. You get larger crystals if you evaporate the water slowly.

Copper oxide, and other transition metal oxides or hydroxides, do not dissolve in water. If the base does not dissolve in water, you need an extra step. You add the base to the acid until no more will dissolve and you have some base left over (called an excess). You filter the mixture to remove the excess base, then evaporate the water in the filtrate to leave the salt behind.

Making insoluble salts

Insoluble salts do not dissolve in water. They can be made by mixing appropriate solutions of ions together.

Soluble and insoluble salts

Soluble	Insoluble
All nitrates	None
Most sulfates	Lead sulfate, barium sulfate
Most chlorides, bromides and iodides	Silver chloride, silver bromide, silver iodide, lead chloride, lead bromide, lead iodide
Sodium carbonate, potassium carbonate	Most other carbonates
Sodium hydroxide, potassium hydroxide	Most other hydroxides

Notice that all nitrates and most chlorides are soluble. This is why many of the chemicals you use in the laboratory are nitrates or chlorides. If you want to make an insoluble salt, you can react together two soluble salts in a precipitation reaction.

Making an insoluble salt

Silver chloride is insoluble - you can see this from the table. You need a soluble silver salt and a soluble chloride salt to make it. Silver nitrate and sodium chloride are both soluble. When you mix their solutions together, you make soluble sodium nitrate and insoluble silver chloride:

• silver nitrate + sodium chloride → sodium nitrate + silver chloride
• AgNO₃(aq) + NaCl(aq) → NaNO₃(aq) +AgCl(s)

The silver chloride appears as tiny particles suspended in the reaction mixture - it forms a precipitate. The precipitate can be filtered, washed with water on the filter paper, and then dried in an oven.

Remember: if you want to make an insoluble salt XY, mixing X nitrate with sodium Y will always work. In the example above, X is silver and Y is chloride.

Using precipitation reactions

Precipitation reactions can be used to remove unwanted ions in solution. This is useful for treating drinking water and waste water.

Electrolysis


This is the process where ionic substances are broken down using electricity into smaller simpler substances.

Ionic substances are compounds formed when a metal reacts with a non metal.

They contain charged particles called ions.

Ions are free to move when dissolved in water so it allows electrolysis to happen.

Ions are free when the substances are molten or dissolved in water.

Cations ( positively charged ions ) move to the cathode ( negatively charged electrode) They lose electrons so they are oxidised

Anions (negatively charged ions ) move to the anode (positively charged electrode) They gain electrons so they are reduced

OILRIG

Oxidation is loss of electrons

Reduction is gain of electrons

Electrodes are conductors which are in contact with a circuit to a battery.

Electrolyte is a substance which in solution will conduct electricity.

Electroplating


This is the process where we cover metals with another metal often more expensive.

The cathode is electroplated

the anode is the metal that is used to coat the cathode.

the electrolyte should be a solution of the same metal of the anode.

E.g

Silver can be used to coat cheap jewellry,

Silver is the anode and a cheaper metal like copper is placed at the cathode

The electrolyte is silver nitrate.

This can also be used to purify copper of both the anode and cathode are copper and the solution is copper sulphate

Aluminium extraction

Aluminium is the most abundant metal on earth, however it can only be extracted using electrolysis.

Bauxite is an aluminium ore. An ore is a rock that contains enough concentration of the metal you want to be economically viable to extract.

Bauxite is purified to get aluminium oxide.

Electrolysis is done.

The aluminium oxide is dissolved in molten  cryolite, this is used to reduce the melting point so it would be cheaper than melting aluminium oxide as the melting point is over 2000 degrees.

Graphite is used as a cathode and the anode.

Aluminium is formed at the cathode and sinks to the bottom of the tank where it can be tapped off.

Oxygen is formed at the positive electrode, this reacts with the carbon to produce carbon dioxide.

As a result, the anode has to be replaced frequently but carbon is cheap.

Hydrogen and metals form positive ions (cations).

If we use electrolysis on metals more reactive than hydrogen in a solution of water, hydrogen will bubble out as the metal would displace the hydrogen ions in water and form the metal oxide.

We can only use electrolysis on solution with metals less reactive than water like copper as they would be the ones that move to the cathode.

The negative ion in solution will move to the anode, normally these are nonmetals.

Since most nonmetals are gases at room temperature, they will bubble off at the anode.

Chloride ions will form chlorine Cl2 when they are oxidised.

This also follow the rules previously

the more reactive element will be the one that leaves, oxygen will leave if the solution has sulfur ions.

Chlorine

We can make chlorine and hydrogen by  electrolysis of dissolved sodium chloride (salt)

Since hydrogen is more reactive then sodium, the H+ ions leave at the cathode

Since chloride ions are more reactive than oxygen ions, Cl- leave at the anode.

The sodium ions and oxygen ions are left behind and form a sodium hydroxide solution NaOH which is an alkaline.

Hydrogen can be used as fuel and for making ammonia.

Chlorine can be used to kill bacterial and to make bleach + plastics

Sodium hydroxide can be used to make soap and bleach.

Half equations.

Reactants are ions and electrons

Products are molecules.

Electrons are shown as e- we add or take away electrons.

2Cl- - 2e- = Cl2

ons are called ionic bonds.

Ions

How ions form

Ions are electrically charged particles formed when atoms lose or gain electrons. This loss or gain leaves a complete highest energy level, so the electronic structure of an ion is the same as that of a noble gas - such as a helium, neon or argon.

Metal atoms and non-metal atoms go in opposite directions when they ionise:

• Metal atoms lose the electron, or electrons, in their highest energy level and become positively charged ions
• Non-metal atoms gain an electron, or electrons, from another atom tobecome negatively charged ions

Positively charged sodium and aluminium ions

Negatively charged oxide and chloride ions

How many charges?

There is a quick way to work out what the charge on an ion should be:

• The number of charges on an ion formed by a metal is equal to the group number of the metal
• The number of charges on an ion formed by a non-metal is equal to the group number minus eight
• Hydrogen forms H⁺ ions.

	Group 1	Group 2	Group 3	Group 4	Group 5	Group 6	Group 7	Group 0
Example element	Na	Mg	Al	C	N	O	Cl	He
Charge	1+	2+	3+	Note 1	3-	2-	1-	Note 2
Symbol of ion	Na+	Mg2+	Al3+	Note 1	N3-	O2-	Cl-	Note 2

Note 1: carbon and silicon in Group 4 usually form covalent bonds by sharing electrons.

Note 2: the elements in Group 0 do not react with other elements to form ions.

Representing positive ions

You need to be able to show the electronic structure of some common metal ions, using diagrams like these:

Lithium, Li

Lithium is in Group 1. It has one electron in its highest energy level. When this electron is lost, a lithium ion Li⁺ is formed.

Sodium, Na

Sodium is also in Group 1. It has one electron in its highest energy level. When this electron is lost, a sodium ion Na⁺ is formed.

Neon atom

Note that a sodium ion has the same electronic structure as a neon atom (Ne).

But be careful - a sodium ion is not a neon atom. This is because the nucleus of a sodium ion is the nucleus of a sodium atom and has 11 protons - but the nucleus of a neon atom has only 10.

Magnesium, Mg

Magnesium is in Group 2. It has two electrons in its highest energy level. When these electrons are lost, a magnesium ion Mg²⁺ is formed.

A magnesium ion has the same electronic structure as a neon atom (Ne).

Calcium, Ca

Calcium is also in Group 2. It has two electrons in its highest energy level. When these electrons are lost, a calcium ion Ca²⁺ is formed.

A calcium ion has the same electronic structure as an argon atom (Ar).

Representing negative ions

You need to be able to show the electronic structure of some common non-metal ions, using diagrams like these:

Fluorine, F

Fluorine is in Group 7. It has seven electrons in its highest energy level. It gains an electron from another atom in reactions, forming a fluoride ion, F^-.

Note that the atom is called fluorine, but the ion is called fluoride.

Neon atom

Note that a fluoride ion has the same electronic structure as a neon atom (Ne).

Once again, a fluoride ion is not a neon atom, because thenucleus of a fluoride ion is the nucleus of a fluorine atom, with 9 protons, and not of a neon atom, with 10.

Chlorine, Cl

Chlorine is in Group 7. It has seven electrons in its highest energy level. It gains an electron from another atom in reactions, forming a chloride ion, Cl^-.

Oxygen, O

Oxygen is in Group 6. It has six electrons in its highest energy level. It gains two electrons from one or two other atoms in reactions, forming an oxide ion, O^2-.

Ionic bonding

When metals react with non-metals, electrons are transferred from the metal atoms to the non-metal atoms, forming ions. The resulting compound is called an ionic compound.

Consider reactions between metals and non-metals, for example:

• sodium + chlorine → sodium chloride
• magnesium + oxygen → magnesium oxide
• calcium + chlorine → calcium chloride

In each of these reactions, the metal atoms give electrons to the non-metal atoms. The metal atoms become positive ions and the non-metal atoms become negative ions.

There is a strong electrostatic force of attraction between these oppositely charged ions, called an ionic bond. The animation shows ionic bonds being formed in sodium chloride, magnesium oxide and calcium chloride.

Ionic compounds

Group 1 and Group 7

The elements in Group 1 of the Periodic Table are called the alkali metals. They form ionic compounds when they react with non-metals. Their ions have a single positive charge. For example, sodium forms sodium ions, Na+.

The elements in Group 7 of the Periodic Table are called the halogens. They form ionic compounds when they react with metals. Their ions have a single negative charge. For example, chlorine forms chloride ions, Cl–.

Sodium chloride

Sodium chloride, NaCl, forms when sodium and chlorine react together. It contains oppositely charged ions held together by strongelectrostatic forces of attraction – the ionic bonds. The ions form a regular lattice in which the ionic bonds act in all directions.

Dot-and-cross diagrams

You need to be able to draw dot-and-cross diagrams to show the ions in some common ionic compounds.

Sodium chloride, NaCl

Sodium ions have the formula Na⁺, while chloride ions have the formula Cl^-. You need to show one sodium ion and one chloride ion. In the exam, make sure the dots and crosses are clear, but do not worry about colouring them.

Magnesium oxide, MgO

Magnesium ions have the formula Mg²⁺, while oxide ions have the formula O^2-. You need to show one magnesium ion and one oxide ion.

Calcium chloride, CaCl₂

Calcium ions have the formula Ca²⁺. Chloride ions have the formula Cl^-.

You need to show two chloride ions, because two chloride ions are needed to balance the charge on a calcium ion.

Formulae of ionic compounds

Ionic compounds are represented by formulae. Symbols and numbers show the atoms in the compound.

For example: ZnCO₃ is the formula for zinc carbonate.

One zinc atom (Zn) and one carbon atom (C) are chemically bonded with three oxygen atoms (O₃). Notice that we don't need to write a 1 next to the Zn or C.

Brackets in formulae

Sometimes brackets are used.

For example: Fe(OH)₃ is the formula for iron(III) hydroxide.

Iron(III) hydroxide consists of one iron atom joined with three oxygen and three hydrogen atoms. The formula is written like this because the oxygen and hydrogen atom often act together.

Constructing formulae

The formula of a compound can be worked out if the ions in it are known. For example, the compound formed from Na⁺ and SO₄^2- will consist of two Na⁺ ions to every one SO₄^2- ion so that the compound is neutral overall. The formula is therefore Na₂SO₄.

Here are the formulae of some common ions.

Positive ions (cations)

Name	Formula	Name	Formula	Name	Formula
ammonium	NH4+	magnesium	Mg2+	zinc	Zn2+
hydrogen	H+	calcium	Ca2+	lead	Pb2+
lithium	Li+	barium	Ba2+	iron(II)	Fe2+
sodium	Na+	silver	Ag+	iron(III)	Fe3+
potassium	K+	copper(II)	Cu2+	aluminium	Al3+

Negative ions (anions)

Name	Formula	Name	Formula	Name	Formula
fluoride	F-	hydrogen carbonate	HCO3-	sulfide	S2-
chloride	Cl-	hydroxide	OH-	sulfate	SO42-
bromide	Br-	nitrate	NO3-	carbonate	CO3-
iodide	I-	oxide	O2-

A covalent bond is a strong bond between two non-metal atoms. It consists of a shared pair of electrons. A covalent bond can be represented by a straight line or dot-and-cross diagram.

Hydrogen and chlorine can each form one covalent bond, oxygen two bonds, nitrogen three, while carbon can form four bonds.

A shared pair of electrons

You will need to understand what covalent bonding is, and to remember some of the properties of molecules that are formed in this way.

A covalent bond forms when two non-metal atoms share a pair of electrons. The electrons involved are in the highest occupied energy levels - or outer shells - of the atoms. An atom that shares one or more of its electrons will complete its highest occupied energy level.

Covalent bonds are strong - a lot of energy is needed to break them. Substances with covalent bonds often form molecules with low melting and boiling points, such as hydrogen and water.

The animation shows a covalent bond being formed between a hydrogen atom and a chlorine atom, to form hydrogen chloride.

After bonding, the chlorine atom is now in contact with eight electrons in its highest energy level - so it is stable. The hydrogen atom is now in contact with two electrons in its highest energy level - so the hydrogen is also stable.

How many bonds?

Atoms may form multiple covalent bonds - that is, share not just one pair of electrons but two or more pairs. Atoms of different elements will form either one, two, three or four covalent bonds with other atoms.

There is a quick way to work out how many covalent bonds an element will form. The number of covalent bonds is equal to eight minus the group number (you can brush up on group numbers by reading through the section in AQA GCSE Science on thePeriodic Table). The table below gives more detail on this rule:

	Group 4	Group 5	Group 6	Group 7
Example	Carbon	Nitrogen	Oxygen	Chlorine
Number of bonds	8 - 4 = 4	8 - 5 = 3	8 - 6 = 2	8 - 7 = 1

Hydrogen forms one covalent bond. The noble gases in Group 0 do not form any.

Representing covalent bonds

Covalent bonds can be represented in several different ways.

Straight lines and models

Straight lines are the most common way to represent covalent bonds, with each line representing a shared pair of electrons. 2D or 3D molecular models are especially useful for showing the relationship between atoms in multiple covalent bonds. Below are some examples of straight lines and images of 3D models.

Models for covalent bonds

Element	Formula	Chemical structure	Ball-and-stick model
Hydrogen	H2
Water	H2O
Ammonia	NH3
Methane	CH4

Double and triple bonds

Note that molecules can have a double covalent bond - meaning they have two shared pairs of electrons - or a triple covalent bond - three shared pairs of electrons. A double covalent bond is shown by a double line, and a triple bond by a triple line.

A molecule of oxygen (O₂) consists of two oxygen atoms held together by a double bond, like this: 

A molecule of nitrogen (N₂) has two nitrogen atoms held together by a triple bond, like this: 

Dot-and-cross diagrams - elements

Dot-and-cross diagrams

Dot-and-cross diagrams are used to represent covalent bonds. The shared electron from one atom is shown as a dot, while the shared electron from the other atom is shown as a cross.

When drawing dot-and-cross diagrams for covalent bonds, you only need to show the electrons in the highest occupied energy level, as only these are involved.

The animations show covalent bonds represented by both displayed formulae (which use straight lines to represent bonds) and dot-and-cross diagrams:

Covalent bonding between two hydrogen atoms to form a molecule of hydrogen gas, H₂.

Covalent bonding between two oxygen atoms to form a molecule of oxygen gas, O₂.

Elements

For your examination, you need to be able to draw dot-and-cross diagrams for hydrogen, chlorine and oxygen.

You do not need to use colours in your answers.

Dot-and-cross diagrams - compounds

You will also need to be able to draw dot-and-cross diagrams representing thecovalent bonds in the molecules of some common compounds:

Hydrogen chloride, HCl

Hydrogen atoms and chlorine atoms can each form one covalent bond. One pair ofelectrons is shared in a hydrogen chloride molecule (HCl).

Water, H₂O

Hydrogen atoms can each form one covalent bond, while oxygen atoms can each form two covalent bonds. Two pairs of electrons are shared in a water molecule (H₂O).

Ammonia, NH₃

Hydrogen atoms can each form one covalent bond, while nitrogen atoms can each form three covalent bonds. Three pairs of electrons are shared in an ammonia molecule (NH₃).

Methane, CH₄

Hydrogen atoms can each form one covalent bond, while carbon atoms can each form four covalent bonds. Four pairs of electrons are shared in a methane molecule (CH₄).

Simple molecular substances consist of molecules in which the atoms are joined by strong covalent bonds. However, the molecules are held together by weak forces so these substances have low melting and boiling points. They do not conduct electricity.

Giant covalent structures contain many atoms joined together by covalent bonds to form a giant lattice. They have high melting and boiling points. Graphite and diamond have different properties because they have different structures. Graphite conducts heat and electricity well because it also has free electrons.

Simple molecules

Covalent bonds form between non-metal atoms. Each bond consists of a shared pair of electrons, and is very strong. Covalently bonded substances fall into two main types:

1. Simple molecules
2. Giant covalent structures

Simple molecules

A molecule of carbon dioxide

These contain only a few atoms held together by strong covalent bonds. An example is carbon dioxide (CO₂), the molecules of which contain one atom of carbon bonded with two atoms of oxygen.

Properties of simple molecular substances

• Low melting and boiling points - This is because the weak intermolecular forces break down easily.
• Non-conductive - Substances with a simple molecular structure do notconduct electricity. This is because they do not have any free electrons or an overall electric charge.

Higher tier only

Hydrogen, ammonia, methane and water are also simple molecules with covalent bonds. All have very strong bonds between the atoms, but much weaker forces holding the molecules together. When one of these substances melts or boils, it is these weak 'intermolecular forces' that break, not the strong covalent bonds. Simple molecular substances are gases, liquids or solids with low melting and boiling points.

Macromolecules

Macromolecules have giant covalent structures. They contain a lot of non-metal atoms, each joined to adjacent atoms by covalent bonds. Their atoms are arranged into giant lattices, which are strong structures because of the many bonds involved. Substances with giant covalent structures have very high melting points, because a lot of strong covalent bonds must be broken. Graphite, for example, has a melting point of more than 3,600°C.

Diamond

Diamond

Diamond is a form of carbon in which each carbon atom is joined to four other carbon atoms, forming a giant covalent structure. As a result, diamond is very hard and has a high melting point. It does not conduct electricity.

Graphite

Graphite

Graphite is a form of carbon in which the carbon atoms form layers. Each carbon atom in a layer is joined to only three other carbon atoms.The layers can slide over each other because there are no covalent bonds between them. This makes graphite much softer than diamond. It is used in pencils and as a lubricant. Graphite conducts electricity.

Silica

Silica

Silica, which is found in sand, has a similar structure to diamond. It is also hard and has a high melting point. However, it contains silicon and oxygen atoms instead of carbon atoms.

Polymers

Polymers have properties which depend on the chemicals they are made from, and the conditions in which they are made. For example, poly(ethene) can be low-density or high-density depending upon the catalyst and reaction condition used to make it. The table summarises some differences in their properties:

	LDPE low-density poly(ethene)	HDPE high-density poly(ethene)
Branches on polymer molecules	Many	Few
Relative strength	Weak	Strong
Maximum useable temperature	85°C	120°C

Thermosoftening polymers

Polymer with no cross-links

Thermosoftening polymers soften when heated and can be shaped when hot. The shape will harden when it is cooled, but can be reshaped when heated up again. Poly(ethene) is a thermosoftening polymer. Its tangled polymer chains can uncoil and slide past each other, making it a flexible material.

Thermosetting polymers

Polymer with cross-links

Thermosetting polymers have different properties to thermosoftening polymers. Once moulded, they do not soften when heated and they cannot be reshaped. Vulcanised rubber is athermoset used to make tyres. Its polymer chains are joined together by cross-links, so they cannot slide past each other easily.

Ionic compounds

Ionic bonds form when a metal reacts with a non-metal. Metals form positive ions, while non-metals form negative ions. Ionic bonds are the electrostatic forces of attraction between oppositely charged ions.

Melting points and boiling points

Ionic bonds are very strong so a lot of energy is needed to break them. Ionic compounds contain many of these strong bonds so they have high melting and boiling points.

Conduction of electricity

Ionic compounds conduct electricity when they are dissolved in water or when they are melted. This is because their ions are free to move and carry the current. However, ionic compounds do not conduct electricity when they are solid. This is because their ions cannot move around in their lattice structure.

Metals

Layers of atoms slide over each other when metals are bent or stretched

Metals are malleable - they can be bent and shaped. This is because they consist of layers ofatoms. These layers can slide over one another when the metal is bent, hammered or pressed.

Metals - Higher tier

Metals form giant structures in which electrons in the outer shells of the metal atoms are free to move. The metallic bond is the force of attraction between these free electrons and metal ions. Metallic bonds are strong, so metals can maintain a regular structure and usually have high melting and boiling points.

Metals are good conductors of electricity and heat. This is because the free electrons can move throughout the metal.

Alloys

An alloy is a mixture of two or more elements, where at least one element is a metal. Many alloys are mixtures of two or more metals.

Layers

Alloys contain atoms of different sizes. These different sizes distort the regular arrangements of atoms. This makes it more difficult for the layers to slide over each other, so alloys are harder than the pure metal.

It is more difficult for layers of atoms to slide over each other in alloys

Copper, gold and aluminium are too soft for many uses. They are mixed with other metals to make them harder for everyday use. For example:

• Brass - used in electrical fittings - is 70 per cent copper and 30 per cent zinc
• 18-carat gold - used in jewellery - is 75 per cent gold and 25 per cent copper and other metals
• Duralumin - used in aircraft manufacture - is 96 per cent aluminium and 4 per cent copper and other metals

Shape memory alloys

Shape memory alloys can return to their original shape after being bent or twisted. Nitinol is a shape memory alloy made from nickel and titanium. It is used in dental braces and spectacle frames.

Nanoscience

Working with nanoparticles is called nanotechnology

A nanometre, 1 nm, is one billionth of a metre (or a millionth of a millimetre). Nanoparticles range in size from about 100 nm down to about 1 nm. They are typically the size of smallmolecules, and far too small to see with a microscope.

Properties and uses of nanoparticles

Nanoparticles have a very large surface area compared with their volume, so they are often able to react very quickly. This makes them useful as catalysts to speed up reactions. They can, for example, be used in self-cleaning ovens and windows.

Nanoparticles also have different properties to the same substance in normal-sized pieces. For example, titanium dioxide is a white solid used in house paint and certain sweet-coated chocolates. Titanium dioxide nanoparticles are so small that they do not reflect visible light, so cannot be seen. They are used in sun screens to block harmful ultraviolet light without appearing white on the skin.

In addition to new cosmetics such as sun screens and deodorants, nanoscience may lead to the development of:

• New catalysts
• New coatings
• New computers
• Stronger and lighter building materials
• Sensors that detect individual substances in tiny amounts

Graphite and fullerenes - Higher tier

Graphite

Graphite is soft and slippery because there are only weak intermolecular forces between its layers.

From left to right - graphite, diamond, silica

Graphite is a good conductor of heat and electricity. This is because, like metals, graphite contains delocalised electrons. These electrons are free to move through the structure of the graphite.

Fullerenes

Structure of a buckminsterfullerene molecule - a large ball of 60 atoms

Carbon exists as graphite and diamond, but it can also form fullerenes. These are cages and tubes with different number of carbon atoms. Buckminsterfullerene is one type of fullerene. Its molecules are spherical and contain 60 carbon atoms.

Fullerenes may be used for drug delivery systems in the body, in lubricants and as catalysts.

The tube fullerenes are called nanotubes. These are very strong. They are useful in reinforcing structures where lightness and strength are needed - for example, in tennis racket frames.

Atomic structure, analysis and quantitative chemistry


Mass number = number of protons and neutrons.

Isotopes are atoms of an element with different numbers of neutrons.

The relative formula mass = Relative atomic mass of all atoms in the formula of the compound.

The relative formula mass in grams is one mole of the substance.

Nucleus has protons and neutrons

Electrons are arranged around it in energy levels.

Relative mass of particle

Neutron 1

Proton 1

Electron very small 1/1836

Atomic number = number of protons.

Atoms of an element will have the same number of protons. Atomic number = same.

They can have different neutrons so the mass number changes (isotopes)

Isotopes have identical chemical properties but some are radioactive.

Mr is relative formula mass

Ar is relative atomic mass

One mole of the substance is the Mr in grams.

Method of analysis


Paper chromatography

Used to analyse coloured substances e.g. coloured pigments and artificial colours.

This work as coloured substances have different dissolving rates in  liquid and bonding with paper.

Substances that are better at dissolving than bonding with paper travel further.

Substances that travel the same distance with the same colour are likely to be the same.


Instrumental method of analysis


This method uses machines they are:

Fast

Accurate (reliable)

Sensitive (can detect small amounts in a small sample)

Gas chromatography

A mixture of compounds can be separated.

It has a long glass tube called a column with a powdered solid material inside that is put in an oven.

The sample is dissolved in a solvent and then injected into one end of the column.

An unreactive gas usually nitrogen carries the sample through

The different substances travel at different speeds so they are separated.

The separated substances leave the column one after the other, they are detected by a detector.

Peaks show number of compounds

Position of peaks show the retention time for each compound



Mass spectrometry



It identifies substances very quickly and accurately in small amounts.

It provides the relative formula mass of the substances separated by gas chromatography.

The peak furthest to the right is the molecular ion peak. this is the relative formula mass of the substance.

GC is usually linked with MS so substances can be separated and identified

The rate of a reaction can be measured by the rate at which a reactant is used up, or the rate at which a product is formed.

The temperature, concentration, pressure of reacting gases, surface area of reacting solids, and the use of catalysts, are all factors which affect the rate of a reaction.

Chemical reactions can only happen if reactant particles collide with enough energy. The more frequently particles collide, and the greater the proportion of collisions with enough energy, the greater the rate of reaction.

Measuring rates

Different reactions can happen at different rates. Reactions that happen slowly have a low rate of reaction. Reactions that happen quickly have a high rate of reaction. For example, the chemical weathering of rocks is a very slow reaction: it has a low rate of reaction. Explosions are very fast reactions: they have a high rate of reaction.

Reactants and products

There are two ways to measure the rate of a reaction:

1. Measure the rate at which a reactant is used up
2. Measure the rate at which a product is formed

The method chosen depends on the reaction being studied. Sometimes it is easier to measure the change in the amount of a reactant that has been used up; sometimes it is easier to measure the change in the amount of product that has been produced.

Things to measure

The measurement itself depends on the nature of the reactant or product:

• The mass of a substance - solid, liquid or gas - is measured with a balance
• The volume of a gas is usually measured with a gas syringe, or sometimes an upside down measuring cylinder or burette

It is usual to record the mass or total volume at regular intervals and plot a graph. The readings go on the vertical axis, and the time goes on the horizontal axis.

For example, if 24 cm³ of hydrogen gas is produced in two minutes, the mean rate of reaction = 24 ÷ 2 = 12 cm³ hydrogen / min.

Factors affecting the rate

You will be expected to remember the factors that affect the rate of reactions, and to plot or interpret graphs from rate experiments.

How to increase the rate of a reaction

The rate of a reaction increases if:

• The temperature is increased
• The concentration of a dissolved reactant is increased
• The pressure of a reacting gas is increased
• Solid reactants are broken into smaller pieces
• A catalyst is used

Rate of reaction and changing conditions

The graph above summarises the differences in the rate of reaction at different temperatures, concentrations and size of pieces. The steeper the line, the greater the rate of reaction. Reactions are usually fastest at the beginning, when the concentration of reactants is greatest. When the line becomes horizontal, the reaction has stopped.

Collisions and reactions

You will be expected to explain, in terms of particles and their collisions, why changing the conditions of a reaction changes its rate.

Collisions

For a chemical reaction to occur, the reactant particles must collide. Collisions with too little energy do not produce a reaction.

The collision must have enough energy for the particles to react. The minimum energy needed for particles to react is called the activation energy.

Changing concentration or pressure

If the concentration of a dissolved reactant is increased, or the pressure of a reacting gas is increased:

• There are more reactant particles in the same volume
• There is a greater chance of the particles colliding
• The rate of reaction increases

Changing particle size

If a solid reactant is broken into small pieces or ground into a powder:

• Its surface area is increased
• More particles are exposed to the other reactant
• There is a greater chance of the particles colliding
• The rate of reaction increases

Changing the temperature

If the temperature is increased:

• The reactant particles move more quickly
• More particles have the activation energy or greater
• The particles collide more often, and more of the collisions result in a reaction
• The rate of reaction increases

Using a catalyst

Catalysts increase the rate of reaction without being used up. They do this by lowering the activation energy needed. With a catalyst, more collisions result in a reaction, so the rate of reaction increases. Different reactions need different catalysts.

Catalysts are important in industry because they reduce costs.

Exothermic reactions transfer energy to the surroundings. Endothermic reactions take in energy from the surroundings.

Reversible reactions are where the products can react to remake the original reactants. If the forward reaction is exothermic, the reverse reaction is endothermic.

Exothermic reactions

When a chemical reaction occurs, energy is transferred to or from the surroundings - and there is often a temperature change.

Exothermic reactions transfer energy to the surroundings. The energy is usually transferred as heat energy, causing the reaction mixture and its surroundings to become hotter. The temperature increase can be detected using a thermometer. Some examples of exothermic reactions are:

• Combustion (burning)
• Many oxidation reactions, for example rusting
• Neutralisation reactions between acids and alkalis

When a flame burns it transfers heat to its surroundings.

Exothermic reactions can be used for everyday purposes. For example, hand warmers and self-heating cans for drinks (such as coffee) use exothermic reactions.

Endothermic reactions

These are reactions that take in energy from the surroundings. The energy is usually transferred as heat energy, causing the reaction mixture and its surroundings to get colder. The temperature decrease can also be detected using a thermometer.

Some examples of endothermic reactions are:

• Electrolysis
• The reaction between ethanoic acid and sodium carbonate
• The thermal decomposition of calcium carbonate in a blast furnace

Endothermic reactions can be used for everyday purposes. For example, certain sports injury cold packs use endothermic reactions.

The animation shows an exothermic reaction between sodium hydroxide and hydrochloric acid, and an endothermic reaction between sodium carbonate and ethanoic acid.

Reversible reactions

In reversible reactions, the reaction in one direction will be exothermic and the reaction in the other direction will be endothermic.

The decomposition of ammonium chloride is a reversible reaction:

ammonium chlorideammonia + hydrogen chloride

Ammonium chloride decomposes when it is heated, so the forward reaction is endothermic - energy must be transferred from the surroundings for it to happen. The backward reaction is exothermic - energy is transferred to the surroundings when it happens.

Copper sulfate

The reaction between anhydrous copper sulfate and water is reversible:

hydrated copper sulfate (blue)anhydrous copper sulfate (white) + water

Water is driven off from hydrated copper sulfate when it is heated, so the forward reaction is endothermic - energy must be transferred from the surroundings for it to happen. The backward reaction is exothermic - energy is transferred to the surroundings when it happens. This is easily observed. When water is added to anhydrous copper sulfate, enough heat is released to make the water bubble and boil.

Acids have a pH of less than 7. Bases have a pH of more than 7. When bases are dissolved in water, they are known as alkalis. Salts are made when an acid reacts with a base, carbonate or metal. The name of the salt formed depends on the metal in the base and the acid used. For example, salts made using hydrochloric acid are called chlorides.

Acids and bases

Diagram of pH scale and universal indicator colours

Acids

Substances with a pH of less than 7 are acids. The more strongly acidic the solution, the lower its pH number. Acidic solutions turn blue litmus paper red. They turn universal indicator paper red if they are strongly acidic, and orange or yellow if they are weakly acidic.

Bases

Substances that can react with acids and neutralise them to make a salt and water are called bases. They are usually metal oxides or metal hydroxides. For example, copper oxide and sodium hydroxide are bases.

Alkalis

Bases that dissolve in water are called alkalis. Copper oxide is not an alkali because it does not dissolve in water. Sodium hydroxide is an alkali because it does dissolve in water.

Alkaline solutions have a pH of more than 7. The stronger the alkali, the higher the pH number. Alkalis turn red litmus paper blue. They turn universal indicator paperdark blue or purple if they are strongly alkaline, and blue-green if they areweakly alkaline.

Neutral solutions

Neutral solutions have a pH of 7. They do not change the colour of litmus paper, but they turn universal indicator paper green. Water is neutral.

Neutralisation reactions

Ions are charged particles which are formed when atoms, or groups of atoms, lose or gain electrons. For the examination, you need to know which ions are produced by acids, and which are produced by alkalis. You will also need to know the ionic equation for neutralisation.

State symbols

State symbols are used in symbol equations:

• (s) means solid
• (l) means liquid (not the same as dissolved in water - see below)
• (g) means gas
• (aq) means aqueous (dissolved in water)

Acids

When acids dissolve in water they produce aqueous hydrogen ions, H⁺(aq). For example, looking at hydrochloric acid:

HCl(aq) → H⁺(aq) + Cl^–(aq)

Alkalis

When alkalis dissolve in water they produce aqueous hydroxide ions, OH^–(aq). For example, looking at sodium hydroxide:

NaOH(aq) → Na⁺(aq) + OH^–(aq)

Ammonia is slightly different. This is the equation for ammonia in solution:

NH₃(aq) + H₂O(l) → NH₄⁺(aq) + OH^–(aq)

Be careful to write OH^– and not Oh^– or oh^–.

Neutralisation reaction

When the H⁺(aq) ions from an acid react with the OH^–(aq) ions from an alkali, a neutralisation reaction occurs to form water. This is the equation for the reaction:

H⁺(aq) + OH^–(aq) → H²O(l)

For example, hydrochloric acid and sodium hydroxide solution react together to form water and sodium chloride solution. The acid contains H⁺ ions and Cl^– ions, and the alkali contains Na⁺ ions and OH^– ions. The H⁺ ions and OH^– ions produce the water, and the Na⁺ ions and Cl^– ions produce the sodium chloride, NaCl(aq).

Making soluble salts

You need to be able to describe the reactions of acids with bases and metals. You should be able to work out the particular salt formed in the reaction.

Acids and bases

When acids react with bases, a salt and water are made:

• acid + metal oxide → salt + water
• acid + metal hydroxide → salt + water

Remember that most bases do not dissolve in water. But if a base can dissolve in water, it is also called an alkali.

Reactive metals

Acids will react with reactive metals, such as magnesium and zinc, to make a salt and hydrogen:

• acid + metal → salt + hydrogen

The hydrogen causes bubbling during the reaction, and can be detected using a lighted splint.

Naming salts

The name of the salt produced in a neutralisation reaction can be predicted. The first part of the name is 'ammonium' if the base used is ammonia. Otherwise, it is the name of the metal in the base. The second part of the name comes from the acidused:

• Chloride (if hydrochloric acid is used)
• Nitrate (if nitric acid is used)
• Sulfate (if sulfuric acid is used)

The table shows some examples:

Acid	+	Base	→	Salt + Water
Hydrochloric acid	+	Copper oxide	→	Copper chloride + water
Sulfuric acid	+	Sodium hydroxide	→	Sodium sulfate + water
Nitric acid	+	Calcium hydroxide	→	Calcium nitrate + water

Ammonium salts

Many artificial fertilisers are ammonium salts, made by the reaction of an acid with ammonia solution. For example:

Acid	Alkali	Fertiliser
Nitric acid	Ammonia solution	Ammonium nitrate
Phosphoric acid	Ammonia solution	Ammonium phosphate
Sulfuric acid	Ammonia solution	Ammonium sulfate

Crystallising salt solutions

You may be asked to describe how to make a soluble salt.

If the base dissolves in water, you need to add just enough acid to make a neutral solution. Check a small sample with universal indicator paper. If ammonia solution is used, you can add a little more than needed to get a neutral solution.

Warm the salt solution to evaporate the water. You get larger crystals if you evaporate the water slowly.

Copper oxide, and other transition metal oxides or hydroxides, do not dissolve in water. If the base does not dissolve in water, you need an extra step. You add the base to the acid until no more will dissolve and you have some base left over (called an excess). You filter the mixture to remove the excess base, then evaporate the water in the filtrate to leave the salt behind.

Making insoluble salts

Insoluble salts do not dissolve in water. They can be made by mixing appropriate solutions of ions together.

Soluble and insoluble salts

Soluble	Insoluble
All nitrates	None
Most sulfates	Lead sulfate, barium sulfate
Most chlorides, bromides and iodides	Silver chloride, silver bromide, silver iodide, lead chloride, lead bromide, lead iodide
Sodium carbonate, potassium carbonate	Most other carbonates
Sodium hydroxide, potassium hydroxide	Most other hydroxides

Notice that all nitrates and most chlorides are soluble. This is why many of the chemicals you use in the laboratory are nitrates or chlorides. If you want to make an insoluble salt, you can react together two soluble salts in a precipitation reaction.

Making an insoluble salt

Silver chloride is insoluble - you can see this from the table. You need a soluble silver salt and a soluble chloride salt to make it. Silver nitrate and sodium chloride are both soluble. When you mix their solutions together, you make soluble sodium nitrate and insoluble silver chloride:

• silver nitrate + sodium chloride → sodium nitrate + silver chloride
• AgNO₃(aq) + NaCl(aq) → NaNO₃(aq) +AgCl(s)

The silver chloride appears as tiny particles suspended in the reaction mixture - it forms a precipitate. The precipitate can be filtered, washed with water on the filter paper, and then dried in an oven.

Remember: if you want to make an insoluble salt XY, mixing X nitrate with sodium Y will always work. In the example above, X is silver and Y is chloride.

Using precipitation reactions

Precipitation reactions can be used to remove unwanted ions in solution. This is useful for treating drinking water and waste water.

Electrolysis


This is the process where ionic substances are broken down using electricity into smaller simpler substances.

Ionic substances are compounds formed when a metal reacts with a non metal.

They contain charged particles called ions.

Ions are free to move when dissolved in water so it allows electrolysis to happen.

Ions are free when the substances are molten or dissolved in water.

Cations ( positively charged ions ) move to the cathode ( negatively charged electrode) They lose electrons so they are oxidised

Anions (negatively charged ions ) move to the anode (positively charged electrode) They gain electrons so they are reduced

OILRIG

Oxidation is loss of electrons

Reduction is gain of electrons

Electrodes are conductors which are in contact with a circuit to a battery.

Electrolyte is a substance which in solution will conduct electricity.

Electroplating


This is the process where we cover metals with another metal often more expensive.

The cathode is electroplated

the anode is the metal that is used to coat the cathode.

the electrolyte should be a solution of the same metal of the anode.

E.g

Silver can be used to coat cheap jewellry,

Silver is the anode and a cheaper metal like copper is placed at the cathode

The electrolyte is silver nitrate.

This can also be used to purify copper of both the anode and cathode are copper and the solution is copper sulphate

Aluminium extraction

Aluminium is the most abundant metal on earth, however it can only be extracted using electrolysis.

Bauxite is an aluminium ore. An ore is a rock that contains enough concentration of the metal you want to be economically viable to extract.

Bauxite is purified to get aluminium oxide.

Electrolysis is done.

The aluminium oxide is dissolved in molten  cryolite, this is used to reduce the melting point so it would be cheaper than melting aluminium oxide as the melting point is over 2000 degrees.

Graphite is used as a cathode and the anode.

Aluminium is formed at the cathode and sinks to the bottom of the tank where it can be tapped off.

Oxygen is formed at the positive electrode, this reacts with the carbon to produce carbon dioxide.

As a result, the anode has to be replaced frequently but carbon is cheap.

Hydrogen and metals form positive ions (cations).

If we use electrolysis on metals more reactive than hydrogen in a solution of water, hydrogen will bubble out as the metal would displace the hydrogen ions in water and form the metal oxide.

We can only use electrolysis on solution with metals less reactive than water like copper as they would be the ones that move to the cathode.

The negative ion in solution will move to the anode, normally these are nonmetals.

Since most nonmetals are gases at room temperature, they will bubble off at the anode.

Chloride ions will form chlorine Cl2 when they are oxidised.

This also follow the rules previously

the more reactive element will be the one that leaves, oxygen will leave if the solution has sulfur ions.

Chlorine

We can make chlorine and hydrogen by  electrolysis of dissolved sodium chloride (salt)

Since hydrogen is more reactive then sodium, the H+ ions leave at the cathode

Since chloride ions are more reactive than oxygen ions, Cl- leave at the anode.

The sodium ions and oxygen ions are left behind and form a sodium hydroxide solution NaOH which is an alkaline.

Hydrogen can be used as fuel and for making ammonia.

Chlorine can be used to kill bacterial and to make bleach + plastics

Sodium hydroxide can be used to make soap and bleach.

Half equations.

Reactants are ions and electrons

Products are molecules.

Electrons are shown as e- we add or take away electrons.

2Cl- - 2e- = Cl2