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## Atomic Structure

All material things are made from atoms. There are just over one hundred different types of atom, called elements. Atoms can join together in millions of different combinations to make all the substances on Earth and beyond.

Structure of the atom

Every atom is made of a nucleus consisting of protons and neutrons. The nucleus is surrounded by electrons. Protons and electrons are oppositely charged. Neutrons have no charge. This means the nucleus of an atom is always positively charged. It is very small compared to the size of the atom.

An atom has a neutral overall charge because it has the same number of electrons as protons.

Protons and neutrons have the same mass. Electrons have such a small mass that this can usually be taken as zero.

The atomic number (also called the proton number) is the number of protons in an atom. All the atoms of the same element have the same number of protons – a number that is unique to that element.

The mass number (also called the nucleon number) is the total number of protons and neutrons in an atom.

The elements are arranged in the Periodic Table in ascending order of atomic number so it's easy to find the name or symbol for an atom if you know the atomic number.

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## Mass number

The relative formula mass of a compound is found by adding together the relative atomic masses of all the atoms in the formula of the compound.

Structure of the atom

Each atom consists of a nucleus containing protons and neutrons, with electrons arranged around it.

Protons and neutrons both have a relative mass of 1 unit.

Electrons have a very small mass compared to protons and neutrons. Generally when working out the mass of atoms and molecules we can ignore the mass of the electrons.

The mass number of an atom is never smaller than the atomic number. It can be the same, but is usually bigger.

The full chemical symbol for an element shows its mass number at the top, and atomic number at the bottom. Here is the full symbol for carbon.

It tells us that a carbon atom has six protons. It will also have six electrons, because the number of protons and electrons in an atom is the same.

The symbol also tells us that the total number of protons and neutrons in a carbon atom is 12. Note that you can work out the number of neutrons from the mass number and atomic number. In this example, it is 12 - 6 = 6 neutrons.

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## Isotopes

Isotopes are atoms of an element with the normal number of protons and electrons, but different numbers of neutrons. Isotopes have the same atomic number, but different mass numbers.

The different isotopes of an element have identical chemical properties. However, some isotopes are radioactive.

## Isotopes of hydrogen

Most hydrogen atoms consist of just one proton and one electron, but some also have one or two neutrons.

## Isotopes of chlorine

Chlorine atoms contain 17 protons and 17 electrons. About 75 per cent of chlorine atoms have 18 neutrons, while about 25 per cent have 20 neutrons.

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## Relative atomic mass

Different atoms have different masses. Atoms have such a small mass it is more convenient to know their masses compared to each other. Carbon is taken as the standard atom and has a relative atomic mass (Ar) of 12.  Atoms with an Ar of less than this have a smaller mass than a carbon atom. Atoms with an Ar which is more than this have a larger mass than a carbon atom.

Calculating relative atomic mass

Chlorine's Ar of 35.5 is an average of the masses of the different isotopes of chlorine. This is calculated by working out the relative abundance of each isotope. For example, in any sample of Chlorine 25% will be 37Cl and 75% 35Cl. The relative atomic mass is therefore calculated using the equation:

(% of isotope 1 × mass of isotope 1) + (% of isotope 2 × mass of isotope 2) ÷ 100

So in the case of chlorine:

(75 × 35) + (25 × 37)100

= 2625 + 925100

= 35.3

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## Dot-and-cross diagrams

The electronic structure of each element can be shown simply as the number of electrons in each shell. For example, lithium is 2.1, neon is 2.8.8, and calcium is 2.8.8.2.

The arrangement of electrons can also be shown using a 'Dot-and-cross' diagram. Electron shells are drawn as circles, with the electrons on each shown as dots or crosses.

Lithium

A black dot represents the nucleus. Lithium atoms have three electrons. A small circle around the nucleus has two red dots on, representing the first shell with two electrons. A larger outer circle has one red dot on, representing the second shell with one electron.

Lithium is in Group 1 of the periodic table.

Fluorine

Fluorine atoms have nine electrons. Two of these fit into the first shell, and the remaining seven electrons fit into the second shell.

Fluorine is in group 7 of the periodic table.

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## Electronic Structure

The way electrons are arranged in an atom is called the 'electronic structure'. There is a link between an element's electronic structure and its place in the periodic table. You can work out an element's electronic structure from its place in the periodic table.

Moving across each period, you can see that the number of shells is the same as the period number.

As you go across each period from left to right the outer shell gradually becomes filled with electrons. The outer shell contains just one electron on the left hand side of the table, but is filled by the time you get to the right hand side.

Moving down each group, you can see that the number of electrons in the outermost shell is the same as the group number.

Each element in a group therefore has the same number of electrons in its outer shell.

Group 0 is a partial exception to this rule, since although it comes after group 7 it is not called 'group 8'; and it contains helium, which has only two electrons in its outer shell.

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## Ions

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 to become negatively charged ions

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

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.

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## Forming an ionic bond

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.

When a non-metal forms a bond the name ending changes. In these reactions the ending is –ide showing only that element is present. If the ending was –ate it means that oxygen is also present as well as the element.

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.

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## 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.

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## Formulae of ionic compounds

Sea water contains a number of dissolved salts. When these salts dissolve in water, the ions separate. Sea water therefore contains a mixture of ions. The most common ions in sea water are shown in the table:

Name of ion                    Formula and charge Chloride Cl Sodium Na+ Sulfate SO42− Magnesium Mg2+ Calcium Ca2+ Potassium K+ Carbonate CO32− Bromide

Br

You can use the charge on the ions shown in the table to work out the formulae of the ionic compounds.

Sodium ions each have a single positive charge. Chloride ions each have a single negative charge. For the charges to cancel out in the neutral salt sodium chloride, they must be in a ration of 1:1. So the formula of sodium chloride is NaCl. Magnesium ions each have two positive charges. Chloride ions each have a single negative charge. For the charges to cancel out in the neutral salt magnesium chloride, they must be in a ration of 1:2. So the formula of magnesium chloride is MgCl2.

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## Ionic structure

The ions in a compound such as sodium chloride are arranged in a lattice structure. This regular arrangement results in the formation of a crystal.

This pattern is repeated in all directions, giving a giant three-dimensional lattice structure in sodium chloride crystals.

Because of the strong electrostatic forces between them, it takes a great deal of energy to separate the positive and negative ions in a crystal lattice. This means that ionic compounds have high melting points and boiling points.

Solid ionic compounds do not conduct electricity, because the ions are held firmly in place. They cannot move to conduct the electric current. But when an ionic compound melts, the charged ions are free to move. Molten ionic compounds do conduct electricity.

When a crystal of an ionic compound dissolves in water, the ions separate. Again, the ions are free to move, so a solution of an ionic compound in water also conducts electricity.

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## Properties of ionic compounds

• High melting and boiling points - Ionic bonds are very strong - a lot of energy is needed to break them. So ionic compounds have high melting and boiling points.
• Conductive when liquid - Ions are charged particles, but ionic compounds can only conduct electricity if their ions are free to move. Ionic compounds do not conduct electricity when they are solid - only when dissolved in water or melted.

Sodium chloride,  NaCl -  High melting point: 800ºC

Non-conductive in its solid state, but when dissolved in water or molten NaCl will conduct electricity.

Magnesium oxide, MgO - Higher melting point than sodium chloride: around 2,800ºC. This is because its Mg2+ and O2- ions have a greater number of charges, so they form stronger ionic bonds than the Na+ and Cl- ions in sodium chloride.

Because magnesium oxide stays solid at such high temperatures, it remains non-conductive. It is used for high-temperature electrical insulation.

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## Flame Tests

Metals change the colour of a flame when they are heated in it. Different metals give different colours to the flame, so flame tests can be used to identify the presence of a particular metal in a sample.

Barium - Pale green

Calcium - Yellow-red

Copper - Green-blue

Lithium -  Red

Sodium - Orange

Potassium - Lilac

## Example

Flame tests are useful for confirming the results of a precipitate test. For example, an unknown solution that produced a pale blue precipitate with sodium hydroxide solution, and a green-blue flame test, must contain a copper compound.

To identify an alkali metal, a flame test must be used instead of a sodium hydroxide precipitate test. This is because the alkali metals do not form precipitates with sodium hydroxide

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## Anion tests

Testing for the Sulfate Ion - You can test to see if a solution contains sulfate ions by using barium chloride. If barium chloride solution is added to a sample of water containing sulfate ions, barium sulfate is formed. Barium sulfate is insoluble in water, and will be seen as a white precipitate.

Barium chloride solution + sodium sulfate solution → sodium chloride solution + solid barium sulfate

Testing for the Halide Ions - You can test to see if a solution contains chloride, bromide or iodide ions by using silver nitrate. If silver nitrate solution is added to a sample of water containing halide ions the silver halide is precipitated. This is because the silver halides are all insoluble in water. The results look like this:

• Silver chloride is a white precipitate
• Silver bromide is a cream precipitate
• Silver iodide is a pale yellow precipitate

Silver nitrate solution + sodium bromide solution → sodium nitrate solution + solid silver bromide

Testing for Carbonate Ions - Limewater is used to test for the presence of carbonate ions (CO32-). Acid is added to the test compound. If carbonate ions are present then carbon dioxide gas bubbles off. If this is passed through limewater is turns the limewater from clear to cloudy.

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## Spectroscopy

All atoms give off light when heated, although sometimes this light is not visible to the human eye. A prism can be used to split this light to form a spectrum, and each element has its own distinctive line spectrum. This technique is known as spectroscopy.

Chemists use spectroscopy to detect very small amounts of an element. Some examples of what line spectra look like are shown here:

A prism can be used to split light

Scientists have used line spectra to discover new elements. In fact, the discovery of some elements, such as rubidium and caesium, was not possible until the development of spectroscopy. The element helium was discovered by studying line spectra emitted by the Sun.

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## Covalent bonds

A covalent bond is formed between non metal atoms, which combine together by sharing electrons. Covalent compounds have no free electrons and no ions so they don't conduct electricity.

Non metals combine together by sharing electrons. The shared pair of electrons holds the two atoms together. It's called a covalent bond. The group of atoms bonded together in this way is called a molecule.

Hydrogen (H2) Water (H2O)

Covalent compounds are usually gases or liquids with low melting points or boiling points and they don't conduct electricity.

Low melting and boiling points - The covalent bonds binding the atoms together are very strong but there are only very weak forces holding the molecules to each other (the intermolecular forces). Therefore, only a low temperature is needed to separate the molecules when they are melted or boiled.

Non-Conductors - Covalent compounds have no free electrons and no ions so they do not conduct electricity.

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## Dot and cross models

Hydrogen (H2) Chlorine (Cl2) Water (H2O)

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## Simple molecular compounds

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

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 not conduct electricity. This is because they do not have any free electrons or an overall electric charge.

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.

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## Giant molecular structures

Giant covalent structures contain a lot of non-metal atoms, each joined to adjacent atoms by covalent bonds. The atoms are usually arranged into giant regular lattices - extremely strong structures because of the many bonds involved.

## Properties of giant covalent structures

• Very high melting points - 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 3,600ºC.
• Variable conductivity - Diamond does not conduct electricity. Graphite contains free electrons, so it does conduct electricity. Silicon is semi-conductive - that is, midway between non-conductive and conductive.

## Graphite

Graphite is a form of carbon in which the carbon atoms form layers. These layers can slide over each other, so graphite is much softer than diamond. It is used in pencils, and as a lubricant. Each carbon atom in a layer is joined to only three other carbon atoms. Graphite conducts electricity.

## 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.

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## Uses of carbon compounds - Diamonds

A diamond is one giant molecule of carbon atoms.

Diamond is extremely hard and has a high melting point. For this reason it is very useful in cutting tools. The cutting edges of discs used to cut bricks and concrete are tipped with diamonds. Heavy-duty drill bits, such as those used in the oil exploration industry to drill through rocks, are made with diamonds so that they stay sharp for longer.

Diamond is insoluble in water. It does not conduct electricity.

Every atom in a diamond is bonded to its neighbours by four strong covalent bonds leaving no free electrons and no ions. This explains why diamond does not conduct electricity. The bonding also explains the hardness of diamond and its high melting point. Significant quantities of energy would be needed to separate atoms so strongly bonded together.

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## Uses of carbon compounds - Graphite

Graphite is formed from carbon atoms in layers.

Graphite is black, shiny and opaque. It is not transparent. It's also a very slippery material. It's used in pencil leads because it slips easily off the pencil onto the paper and leaves a black mark. It is a component of many lubricants, for example bicycle chain oil. Graphite is insoluble in water. It has a high melting point and is a good conductor of electricity, which makes it a suitable material for the electrodes needed in electrolysis.

Each carbon atom is bonded into its layer with three strong covalent bonds. This leaves each atom with a spare electron, which together form a delocalised 'sea' of electrons loosely bonding the layers together. These delocalised electrons can all move along together – making graphite a good electrical conductor. Because the layers are only weakly held together they can easily slip over one another. This explains why graphite is so slippery. Melting graphite is not easy, though. It takes considerable energy to break the strong covalent bonds and separate the carbon atoms.

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## Chromatography

Chromatography can be used to separate mixtures of coloured compounds. Mixtures that are suitable for separation by chromatography include inks, dyes and colouring agents in food.

Simple chromatography is carried out on paper. A spot of the mixture is placed near the bottom of a piece of chromatography paper and the paper is then placed upright in a suitable solvent, eg water. As the solvent soaks up the paper, it carries the mixtures with it. Different components of the mixture will move at different rates. This separates the mixture out.

## Rf values

Different chromatograms and the separated components of the mixtures can be identified by calculating the Rf value using the equation:

Rf = distance moved by the compound ÷ distance moved by the solvent

The Rf value of a particular compound is always the same - if the chromatography has been carried out in the same way. This allows industry to use chromatography to identify compounds in mixtures.

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## Separation techniques

Mixtures of liquids can be separated according to their properties. The technique used depends on whether the liquids dissolve in each other, and so are miscible, or if they are immiscible. Fractional distillation is a technique used to separate liquids according to their boiling points. Chromatography is used to separate mixtures of coloured compounds.

# Separation of liquids

Liquids can be described in two ways – immiscible and miscible. The separation technique used for each liquid depends on the properties of the liquids.

## Immiscible liquids

Immiscible means that the liquids don't dissolve in each other – oil and water are an example. It is possible to shake up the liquids and get them to mix but they soon separate. Separating immiscible liquids is done simply using a separating funnel. The two liquids are put into the funnel and are left for a short time to settle out and form two layers. The tap of the funnel is opened and the bottom liquid is allowed to run. The two liquids are now separate.

## Miscible liquids

Miscible liquids are harder to separate as they dissolve in each other. Miscible liquids are often separated using fractional distillation. This is possible as miscible liquids have different boiling points.

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## Fractional distillation of liquid air

You need to be able to explain how nitrogen and oxygen are obtained from the air.

About 78 per cent of the air is nitrogen and 21 per cent is oxygen. These two gases can be separated by fractional distillation of liquid air.

## Liquefying the air

Air is filtered to remove dust, and then cooled in stages until it reaches –200°C. At this temperature it is a liquid. We say that the air has been liquefied.

Here's what happens as the air liquefies:

1. Water vapour condenses, and is removed using absorbent filters
2. Carbon dioxide freezes at –79ºC, and is removed
3. Oxygen liquefies at –183ºC
4. Nitrogen liquefies at –196ºC

The liquid nitrogen and oxygen are then separated by fractional distillation.

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## Metallic structure

The particles in a metal are held together by strong metallic bonds. It takes a lot of energy to separate the particles. That is why they have high melting points and boiling points. Solid metals are crystalline. This means that the particles are close together and in a regular arrangement. There are strong electrostatic forces holding the particles together.

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## Transition metals

The elements in the centre of the periodic table, between groups 2 and 3, are called the transition metals. Most of the commonly used metals are there, including iron, copper, silver and gold.

## Common properties

The transition metals have the following properties in common:

• They form coloured compounds
• They are good conductors of heat and electricity
• They can be hammered or bent into shape easily
• They are less reactive than alkali metals such as sodium
• They have high melting points - but mercury is a liquid at room temperature
• They are usually hard and tough
• They have high densities
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## Strong alkalis

The hydroxides formed in all of these reactions dissolve in water to form alkaline solutions. These solutions turn universal indicator purple, showing they are strongly alkaline. Strong alkalis are corrosive, so care must be taken when they are used, for example, by using goggles and gloves.

## Why does the reactivity increase down the group?

All alkali metals have one electron in the outer shell. In a reaction, this electron is lost and the alkali metal forms a +1 ion. As you go down group 1, the number of electron shells increases – lithium has two, sodium has three etc. Therefore, the outermost electron gets further from the nucleus. The attraction from the positive nucleus to the negative electron is less. This makes it easier to remove the electron and makes the atom more reactive.

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## Halogens

Melting and Boiling points - The halogens have low melting points and boiling points. This is a typical property of non-metals. Fluorine has the lowest melting point and boiling point. The melting points and boiling points then increase as you go down the group.

State at room temperature - Room temperature is usually taken as being 25°C. At this temperature, fluorine and chlorine are gases, bromine is a liquid, and iodine and astatine are solids. There is therefore a trend in state from gas to liquid to solid down the group.

Colour - The halogens become darker as you go down the group. Fluorine is very pale yellow, chlorine is yellow-green, and bromine is red-brown. Iodine crystals are shiny purple - but easily turn into a dark purple vapour when they are warmed up.

Predictions -  When we can see a trend in the properties of some of the elements in a group, it is possible to predict the properties of other elements in that group. Astatine is below iodine in Group 7. The colour of these elements gets darker as you go down the group. Iodine is purple, and, as we would expect, astatine is black.

USES OF HALOGENS

Halogens are bleaching agents. They will remove the colour of dyes. Chlorine is used to bleach wood pulp to make white paper. Halogens kill bacteria. Chlorine is added to drinking water at very low concentrations. This kills any harmful bacteria in the water, making it safe to drink. Chlorine is also added to the water in swimming pools.

It is very important to take care when using the halogens because they are very reactive and poisonous. Chlorine is used only in a fume cupboard. Iodine should not be handled unless gloves and goggles are worn.

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## Displacement reactions

The reactivity of the halogens decreases as we move down the group. This can be shown by looking at displacement reactions.

When chlorine (as a gas or dissolved in water) is added to sodium bromide solution the chlorine takes the place of the bromine. Because chlorine is more reactive than bromine, it displaces bromine from sodium bromide. The solution turns brown. This brown colour is the displaced bromine. The chlorine has gone to form sodium chloride.

If you look at the equation, you can see that the Cl and Br have swapped places.

chlorine + sodium bromide → sodium chloride + bromine

Cl2(aq) + 2NaBr(aq) → 2NaCl(aq) + Br2(aq)

## Reactivity series

This type of reaction happens with all of the halogens. A more reactive halogen displaces a less reactive halogen from a solution of one of its salts.

If you test different combinations of the halogens and their salts, you can work out a reactivity series for Group 7. The most reactive halogen displaces all of the other halogens from solutions of their salts, and is itself displaced by none of the others. The least reactive halogen displaces none of the others, and is itself displaced by all of the others. It works just the same whether you use sodium salts or potassium salts.

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## Noble Gases

The elements in group 0 are called the noble gases. They belong to the right-hand column in the periodic table. The noble gases are all chemically unreactive which means they are inert.

## Common properties

The noble gases have the following properties in common:

• They are non-metals
• They are very unreactive gases
• They are colourless
• They exist as single atoms (they are monatomic)

Helium - Used in balloons and airships. It is much less dense than air, so balloons filled with it float upwards.

Neon - Used in advertising signs. It glows when electricity is passed through it, and different coloured 'neon lights' can be made by coating the inside of the glass tubing with other chemicals.

Argon - Used in light bulbs. The very thin metal filament inside the bulb would react with oxygen and burn away if the bulb were filled with air instead of argon. As argon is unreactive, it stops the filament burning away.

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## Endo and exothermic reactions

Exothermic reactions

These are reactions that 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:

• Burning (combustion)
• Neutralisation reactions between acids and alkalis
• The reaction between water and calcium oxide
• Explosions

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
• Photosynthesis
• The reaction between ammonium nitrate and water
• The thermal decomposition of calcium carbonate in a blast furnace
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## Bonds and energy changes

During a chemical reaction:

• Bonds in the reactants are broken and New bonds are made in the products

Energy is absorbed to break bonds. Bond-breaking is an endothermic process. Energy is released when new bonds form. Bond-making is an exothermic process.

Whether a reaction is endothermic or exothermic depends on the difference between the energy needed to break bonds and the energy released when new bonds form. If more heat energy is released when making the bonds than was taken in when they broke, the reaction is exothermic.

What is the energy transferred to 100 cm3 of water to raise its temperature by 20ºC?

It is useful to remember that 1cm3 of water has a mass of 1g. So 100 cm3 of water has a mass of 100 g.

Energy transferred = mass of water heated × specific heat capacity of water × temperature rise

= 100 × 4.2 × 20 = 8,400 J

It is also useful to remember that 1 kilojoule, 1 kJ, equals 1,000 J. So the energy transferred is 8.4 kJ.

If 0.5 g of fuel was used, the energy output of the fuel would be: 8.4 ÷ 0.5 = 16.8 kJ/g

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## Rates of reaction

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

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

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## 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

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.

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## Catalysts

A catalyst is a substance that can increase the rate of a reaction. The catalyst itself remains unchanged at the end of the reaction it catalyses. Only a very small amount of catalyst is needed to increase the rate of reaction between large amounts of reactants.

Different catalysts catalyse different reactions. The table summarises some common catalysts used in industry and the reactions they catalyse:

## Some common catalysts used in industry and the reactions they catalyse

Iron      -       Making amonia from Nitrogen and Hydrogen

Platimun    -     Making amonia from Nitrogen and Hydrogen

Vanadium (V) Oxide   -   Making Sulfuric acid

## Catalytic converters

Modern cars have a catalytic converter to help reduce the production of toxic gases. Catalytic converters use a platinum and rhodium catalyst with a high surface area. This increases the rate of reaction of carbon monoxide and unburnt fuel from exhaust gases with oxygen from the air. The product from this is carbon dioxide and water, which is less harmful to the environment. The catalysts are designed to work best at the high temperatures found in the engine.

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