Chemistry Unit 1

All units from AQA Chemistry Unit 1

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Electron Arrangements and Ionisation energy

Ionisation Energy - the energy required to remove a mole of electrons from a mole of atoms in a gaseous state, which subsequently become positive ions

The first electron needs the least energy to remove, this is because it is being removed from a neutral atom. This is the first IE.

As you remove more electrons, they require more and more energy. This is because they are being removed from positive ions.

Electrons become increasingly difficult to remove as they are closer to the nucleus and so receive a stronger force from the nuclear charge (across a period). Less protection from shielding also mean more energy is required.

When electrons are forced to pair in an orbital, for example in Sulphur, this causes repulsion which makes an electron easier to remove, defying the trend of increasing IEs.

When a new sub level is started, IE also falls due to the increased shielding from the nuclear charge provided by the previous sub level.

These cases, which go against the trend, are evidence for the existence of sub levels.

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Relative atomic and molecular masses

The relative atomic mass, Ar, is the weighted average mass of an atom of an element, taking into account its naturally occurring isotopes, relative to 1/12th the relative atomic mass of an atom of Carbon-12.

The relative molecular mass, Mr, of a molecule is the mass of that molecule compared to 1/12th the relative atomic mass of an atom of Carbon-12

Add the Ar's of all the atoms present in the molecule to find the Mr

Relative formula masses are calculated in the same way

The Avogadro Constant is the number of atoms in 12g of Carbon-12

The amount of substance that contains 6.022 x 10(23) particles is called a mole.

Mol = Mass / Mr

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The Ideal Gas Equation

PV = nRT

R = 8.31,

Volume is measured in metres cubed (m³)

Temperature is measured in Kelvins

Pressue is measured in Pa

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Empirical and Molecular Formulae

The empirical formula represents the simplest ratio of atoms of each element present in a compound.

Work out the number of moles in the amounts of each of the elements in the molecule, then divide through by the smallest to get the EF

The molecular formula gives the actual number of atoms of each element in one molecule of the compound (it only applies to substances that exist as molecules)

To find the number of units of the EF in the molecular formula, divide the Mr by the Mr of the empirical formula.

For example, ethene is C2H4, but its empirical formula is CH3. 28/14 = 2

Therefore multiply CH3 by 2 to get C2H6.

Combustion analysis is used to find the empirical formula of new compounds.

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Moles in Solutions

The concentration of a solution tells us how much solute is present in a known volume of solution

· Concentration is measured in moles per cubic decimetre

· 1molddm-3 means there is 1 mole of solute per cubic decimetre of solution, 2moldm-3 means there are 2 moles of solute etc.

· 1 cubic decimetre = 1000 cubic centimetre

Concentration = number of moles / volume in dm3

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Balanced Equations and Related Calculations

Balanced symbol equations use the formulae of reactants and products. There are the same number of atoms of each element on both sides of the arrow. This is because atoms are never created or destroyed in chemical reactions.

Ionic equations consider the ions present in a reactions, this is usually when the substances are in solution (aq)

In an ionic equation, we split each of the ionic substances into their constiuent ions, then we cross out any ions that appear on both sides of the reaction.

We call these spectator ions, as they do no take part in the reaction.

Whenever an acid reacts with an alkali, the ionic equation will be H(+) + OH(-) --> H2O

The charges on each side of the ionic equation must be the same.

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Balanced equations, atom economies and percentage

Percentage atm economy measures the amount of useful product formed during a reaction in relation to all of the products formed

% Atom economy = mass of desired product / total mass of reactants X 100

Some reactions, in theory, have no wasted atoms and so therefore have a 100% atom economy. E.g. C2H4 + Br2 --> C2H4Br2

The percentage yield of a reaction tells us how much of a product was obtained in reality compared to the maximum amount that coul be obtained

% Yield = grams of product obtained / theoretical maximum grams of product X 100%

The percentage yield of a reaction is often, if not always, less than 100% due to a number of reasons:

  • Reaction does not reach completion
  • Some product is left on equipment, cannot be collected
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Ionic Bonding

When atoms bond together they share or transfer electrons to achieve a more stable electron arrangement, often a full outer main level of electrons, like the noble gases.

Ionic bonding occurs between metals and non-metals. The metals lose electrons, and the non-metals gain electrons, forming ionic bonds. Positive and negative ions are formed.

The ionic bonds form as the oppositely charged ions are attracted by electrostatic forces

Ionic compounds always exist in a lattice with balanced attractive and repulsive forces, so ions of the same charge are never next to each other.

Ionic compounds:

  • Are always solids at room temperature,
  • Have gaint structures and therefore high melting points,
  • Conduct electricity when molten or in solution,
  • Are brittle as the ions which make up the lattice are different sizes - do not slide
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Covalent Bonding

Non-metal atoms need to receive electrons to fill the spaces in their outer shells

  • A covalent bond forms between a pair of non-metal atoms
  • The atoms share electrons so that each atom has a noble gas arrangement
  • A covalent bond is a shared pair of electrons

Covalently bonded molecules have a neutral charge because no electrons have been transferred from one atom to another

Atoms with covalent bonds are held together by the electrostatic attraction between the nuclei and the shared electrons. This takes place within the molecule.

Double covalent bonds are when four electrons are shared, for example in molecules of oxygen. This is a double bond and is represented by O=O

Molecular substances are often gases or liquids at room temperature, as the molecules are only held together by weak intermolecular forces

They are poor conductors of electricity, as they contain no charged particles

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Co-ordinate Bonding

Co-ordinate bonding or dative covalent bonding is when one atom provides both of the electrons in a covalent bond.

In a co-ordinate bond:

  • The atom that accepts the elctron pair is an atomc that does not have a filled outer main level of electrons, we say the atom is electron-deficient
  • The atom that is donating the electrons has a pair of electrons that is not being used in a bond, called a lone pair

Co-ordinate covalent bonds are represented by an arrow ( -->). The arrow points towards the atom that is accepting the electron pair.

Co-ordinate bonds have exactly the same strength and length as ordinary covalent bonds between the same pair of atoms.

The ammonium ion has covalently bonded atoms but is a charged particle. One of the H+ ions forms a co-ordinate bond with the nitrogen atom.

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Electronegativity is the power of an atom to attract the electron density in a covalent bond towards itself

The Pauling scale is used as a measure of electronegativity. It runs from 0 to 4. The greater the number, the more electronegative the atom.

Electronegativity depends on:

  • The nuclear charge
  • The distance between the nucleus and the outer shell electrons
  • The shielding of the nuclear charge by electrons in inner shells

The smaller the atom, the closer the nucleus is to the shared outer main level electrons and the greater its electronegativity.

The larger the nuclear charge, the greater the electronegativity.

The most electronegative atoms are found at the top right-hand corner of the periodic table, they are the smallest atoms with the relatively largest nuclear charge.

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Polarity of covalent bonds

Polarity is about the unequal sharing of the electrons between atoms that are bonded together covalently. It is a property of the bond.

When both atoms are the same, they have the same electronegativity and so the electrons in the bond are shared equally between the atoms. We say the bond is non-polar.

If we think of the electrons as being in a cloud of charge, then the cloud is uniformly spread between the two atoms.

In a covalent bond between two atoms of different electronegativity, the electrons in the bond will not be shared equally between the atoms. For example in HF, Fluorine is more electronegative than Hydrogen and so the electrons in the covalent bond are more attracted to the fluorine than the hydrogen.

The electron cloud is distorted towards the fluorine, adding partial charges to the formula:

Hδ+ - Fδ- Covalent bonds like this are said to be polar. The greater the difference in electronegativity, the more polar is the covalent bond.

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Metallic bonding

Metals are shiny elements made up of atoms that can easily lose up to three outer electrons, leaving positive metal ions.

In metals, all the of metal atoms lose their outer level electrons, which become a 'sea' of delocalised electrons, this means they are not tied to a particular atom.

The sea of electrons holds the positive ions together through electrostatic attraction.

  • The number of delocalised electrons depends on how many electrons have been lost by each metal atom
  • The metallic bonding spreads thoughout so metals have giant structures

Metals are good conductors of heat and electricity. The delocalised electrons are free to carry charge and can vibrate to help the metal conduct heat.

The strength of a metal depends on it's charge and the size of the ion. The larger the charge, the more delocalised electrons and so the stronger the electrostatic attraction.

The smaller the ion, the closer the electrons are to the positive nucleus and the stronger the bond. Metals are also malleable and ductile, and have high melting points.

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Intermolecular forces - Dipole-dipole forces

Atoms in molecules and in giant structures are held together by strong covalent, ionic or metallic bonds. Molecules and seperate atoms are attracted to one another by other, weaker forces called intermolecular forces.

If the intermolecular forces are strong enough, then molecules are held closely together to be liquids or even solids.

Dipole-dipole forces act only between certain types of molecules, they are weaker than hydrogen bonds, but stronger than van der Waals forces.

Dipole-dipole forces act between molecules that have permanent dipoles. For example in HCl. Two molecules that both have dipoles will attract one another.

Whatever their starting positions, the molecules with dipoles will 'flip' to give an arrangement where the two molecules attract

Hδ+ - Clδ- Clδ- - Hδ+ ----------> Hδ+ - Clδ- -- Hδ+ - Clδ-

Molecules must have a dipole moment in order to form dipole-dipole forces

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Intermolecular forces - van der Waals forces

All atoms and molecules are made up of positive and negative charges even though they are all neutral overall. These charges produce very weak electrostatic attractions between all atoms and molecules. We call these van der Waals forces.

Imagine a helium atom, it has two electrons and two protons in its nucleus. It has an overall neutral charge, however at any moment in time the electrons could be anywhere, this means the distribution of charge is changing at every instant.

The atoms are said to have temporary dipoles, and this affects the electron distribution of nearby atoms, causing them to have induced temporary dipoles.

  • Van der Waals forces act between all atoms or molecules at all times
  • They are in addition to any other intermolecular forces
  • The dipole is caused by the changing position of the electron cloud, so the more electrons there are, the larger the instantaneous dipole will be

This means that the size of the van der Waals forces increases with the number of electrons present. This explains why the boiling points of noble gases increases at the atomic numbers increase and of hydrocarbons as the chain length increases.

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Hydrogen Bonding

Hydrogen bonds can only form when:

  • A hydrogen atom is bonded to a very electronegative atom. This will produce a stron partial charge on the hydrogen atom
  • A very electronegative atom with a lone pair of electrons (F, O or N). These will be attracted to the partially charged hydrogen atom in another molecule and form a bond.
  • For example, ammonia molecules NH3, form hydrogen bonds with water H2O.

The boiling points of hydrides shows the presence of hydrogen bonds, as they are much higher in water, hydrogen fluoride and ammonia. If only van der Waals forces were present, their boiling points would not be as high.

Ice is an example of where hydrogen bonding has great importance. Many hydrogen bonds form between molecules of H2O, holding the molecules in fixed positions. Ice is less dense than water and so therefore floats, helping to sustain life in lakes and rivers as it insulates the water.

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States of matter


  • Regular arrangement of particles - solids have a definite shape
  • Particles held closely together - solids are not easily compressed
  • Particles vibrate about a point - diffusion rate is very slow


  • Random arrangement of particles - liquid changes shape to fill the bottom of a cup
  • Particles close together - liquids are not easily compressed
  • Rapid 'jostling' of particles - diffusion is slow, liquids evaporate


  • Random arrangement of particles - a gas will fill its container
  • Particles spaced far apart - gases are easily compressed
  • Rapid movement of particles - gases diffuse rapidly - exert pressure
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Molecular crystals

Molecular crystals consist of molecules held in a regular array by one or more of the three types of intermolecular forces.

They have low melting points and low enthalpies od fusion as these forces are much weaker than covalent, ionic or metallic bonds.

Iodine is an example of a molecular crystal. A strong covalent bond holds pairs of iodine atoms together to form I2 molecules. Since iodine molecules have a large number of electrons, the van der Waals forces are strong enough to hold the molecules together as a solid. But these are much weaker than covalent bonds, giving iodine the following properties:

  • Its crystals are soft and break easily
  • It has a low melting temperature and sublimes readily to form gaseous iodine molecules
  • It does not conduct electricity because there are no charged particles to carry charge.
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The shape of molecules and ions

Molecules are three dimensional and they come in many different shapes

We can predict the shape of a simple covalent molecule by using the idea that:

  • Each pair of electrons around an atom will repel all other electron pairs
  • The pairs of electrons will therefore take up positions as far apart as possible to minimise repulsion
  • This is called the electron pair repulsion theory

Electron pairs may be: a shared pair or a lone pair

The shape of a simple molecule depends on the number of pairs of electrons that surround the central atom.

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Shapes of molecules continued...

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Molecules with lone pairs of electrons

Some molecules have unshared (lone) pairs of electrons. The lone pairs affect the shape of the molecule. Ammonia and water are good examples of molecules where lone pairs affect the shape.

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Periodicity - the periodic table

The periodic table is a list of all the elements in order of increasing atomic number. We can predict the properties of an element from its position in the table. We can use it to explain the similarities of certain elements and trends in their properties, in terms of their electronic arrangements.

Areas of the periodic table are labelled s-block, p-block, d-block and f-block.

The block an element is in depends on the sub level location of its highest level electron.

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Periodicity continued...

A group is a vertical column of elements. The elements in the same group form a chemical 'family'; they have similar properties. Elements in the same group have the same number of electrons in the outer main level.

In the s-block (metals), elements get more reactive as we go down a group.

To the right (non-metals), elements tend to get more reactive as we go up a group.

Periods are the horizontal rows of elements in the periodic table. There are trends in physical properties and chemical behaviour as we go across a period.

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Trends in the properties of elements in Period 3

Periodicity is the patterns in the properties of elements, usually across a period.

Periodicity is explained by the electron arrangements of elements:

  • The elements in Groups 1,2 and 3, sodium, magnesium and aluminium, are metals. They have giant structures. they lose their outer electrons to form ionic compounds
  • Silicon in Group 4 has four electrons in its outer shell with which it forms four covalent bonds. The element has some metallic properties and is classed as a semi-metal
  • The elements in Groups 5, 6 and 7, phosphorous, sulfur and chlorine, are non-metals. They either accept electrons to form ionic compounds, or share their outer electrons to form covalent compounds.
  • Argon in Group 0 is a noble gas, it has a full outer shell and is unreactive88
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The melting points of Group 3 elements

  • The melting points of the Group 3 elements increase from Sodium to Aluminium due to strengthening metallic bonds.
  • Silicon has the highest melting point due to its giant covalent structure.
  • The melting points of the non-metals with molecular structures depends on the sizes of the van der Waals forces between the molecules. As a result, Sulfur (S8) has the highest melting point, followed by Phosphorous (P4), and then by Chlorine (Cl2) and Argon (Ar)

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More trends in the properties of the elements in P

Some key properties of atoms, such as size and ionisation energy, are periodic, that is, there are similar trends as we go across each period in the periodic table.

Atomic radii tell us about the sizes of atoms

  • Atomic radius is a periodic property because it decreases across each period and there is a jump when we start the next period
  • Atoms get larger as we go down any Group

The size of the atom decreases as we go across the period due to increases in the nuclear charge of the nuceus. This increases charge pulls the electrons closer to the nucleus. There are no additional electron shells to provide more shielding.

As we go down a group in the periodic table, the atoms of each element have one extra complete main level of electrons compared with the one before, so the outer electron main level is further from the nucleus and the atomic radii increase.

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A closer look at ionisation energies

  • The first ionisation energy actually drops between Group 2 and Group 3, so that Aluminium has a lower IE than Magnesium
  • The ionisation energy drops again slightly between Group 5 (Phosphorous) and Group 6 (Sulfur)
  • Similar patterns occur in other periods, we can explain this if we look at the electron arrangements of these elements
  • The first drop is due to the outermost electron being in a new sub-level (p), which makes it easier to remove
  • The second drop is due to extra replusion caused by electrons being forced to share an orbital in the p-sub level

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Organic chemistry - carbon compounds

Organic chemistry is the chemistry of carbon compounds. Life on our planet is based on carbon and organic means 'to do with living beings'.

Carbon can form rings and very long chains, which may be branched. This is because:

  • A Carbon atom has four electrons in its outer shell, so it forms four covalent bonds
  • Carbon-carbon bonds are relatively strong and non-polar

In all stable carbon compounds, carbon forms four covalent bonds and has eight electrons in its outer shell. It can do this by:

  • Forming four single bonds. E.g. methane
  • Forming two single bonds and one double bond. E.g. ethene
  • Forming one single bond and one triple bond. E.g. ethyne

The displayed formula shows every atom and every bond in the molecule.

The structural formula shows the unique arrangement of atoms in a molecule in a simplified form, without showing all the bonds. E.g. CH3CH3. Branches in the carbon chains are showin in brackets. E.g. CH3CH(CH3)CH3

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Nomenclature - naming organic compounds

A systematic name has a root that tells us the longest unbranched hydrocarbon chain or ring.

  • 1 - meth, 2 - eth, 3 - prop, 4 - but, 5 - pent, 6 - hex
  • -ane means there are no double bonds, -ene means there is a double bond
  • Prefixes are added to the beginning of the root, like methyl (CH3)
  • If there are two of the same side chains, di is added to the front - dimethyl

Functional groups are the reactive groups attached to hydrocarbon chains

A homologous series is a family of organic compounds, with the same functional group, but different carbon chain length.

  • Members of a homologous series have a general formula
  • Each member of the series differs from the nexy by CH2
  • The length of the caron chain has little effect on the chemical reactivity of the functional group
  • They have trends in physical properties and similar chemial properties
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Isomers are molecules that have the same molecular formula but whose atoms are arranged differently.

Structural isomers are defined as having the same molecular formula but different structural formulae. There are three sub-divisions of structural isomerism, these are:

  • Positional isomerism - this is where the same functional groups are attached to the chain at different points
  • Functional group isomerism - this is where there are different functional groups
  • Chain isomerism - this is where there is a different arrangement of the hydrocarbon chain (such as branching)
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Alkanes are saturated hydrocarbons, they contain only carbon-carbon and carbon-hydrogen single bonds.

They are among the least reactive organic compounds. They are used as fuels and lubricants and as starting materials for a range of other compounds. The main sources of alkanes is crude oil.

The general formula for all chain alkanes is CnH2n+2. Hydrocarbons may be unbranched chains, branched chains or rings.

Ring alkanes have the general formula CnH2n because the 'end' hydrogens are not required.

The number of possible isomers of alkanes rises with the number of carbons. Decane (C10H22) has 75 and C30H62 has over 4 billion.

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Physical properties of alkanes

Polarity - alkanes are almost non-polar because the electronegativities of carbon and hydrogen are so similar. As a result, the only intermolecular forces between their molecules are weak van der Waals forces. The larger the molecule, the stronger the forces.

As the chain length of alkanes increases, so does the boiling point, this is due to the increasing strength of the van der Waals forces.

Alkanes are insoluble in water

Alkanes are relatively unreactive, but burn in a plentiful supply of oxygen to form carbon dioxide and water, or in a restricted supply of oxygen to form carbon monoxide or carbon.

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Fractional distillation of crude oil

Crude oil is a mixture mostly of alkanes, both unbranched and branched. It can also contain other compounds, like sulfur, which burns to form sulfur dioxide - this is one of the main causes of acid rain.

Fractional distillation is the seperation of different hydrocarbons from crude oil, it is done in a fractioning tower.

  • The crude oil is first heated in a furnace
  • A mixture of liquid and vapour passes into a tower that is cooler at the top than at the bottom
  • The vapours pass up the tower until they arrive at a tray which is sufficiently cool, then they condense to liquid
  • The mixture of liquids that condenses on each tray is piped off
  • The shorter chain hydrocarbons condense in trays nearer to the top of the tower, where it is cooler, as they have lower boiling points
  • The thick residue that collects at the base of the tower is called tar or bitumen. It can be used for road surfacing, but is often further processd to give more valuable products
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Industrial cracking

The naphtha fraction from the fractional distillation of crude oil is in huge demand, for petrol and by the chemical industry.

However, most crude oil has more of the longer chain fractions than is wanted and not enough of the naphtha fraction.

To meet demand for the shorter chain hydrocarbons, many of the longer chain fractions are broken into shorter lengths (cracked). This has two useful results:

  • Shorter, more useful, chains are produced, especially petrol
  • Some of the products are alkenes, which are more reactive than alkanes

The most valuable alkene is ethene, which is the starting material for poly(ethene) and a wide range of other everyday materials.

There are two types of cracking, thermal cracking and catalytic cracking.

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Combustion of alkanes

Fuels are substances that release heat energy when they undergo combustion. They also store a large amount of energy for a small amount of weight.

Incomplete combustion occurs when there is a limited supply of oxygen. The poisonous gas carbon monoxide is formed, and with even less oxygen, carbon (soot) is produced. Incomplete combustion often happens with longer chain hydrocarbons, which need more oxygen to burn compared with shorter chains.

There are many different types of polluting products formed from burning hydrocarbons:

  • Carbon monoxide - a poisonous gas produced by incomplete combustion
  • Nitrogen oxides - produced when there is enough energy for N2 and O2 in the air to combine. This happens in a petrol engine, at the high temperatures present. Can contribute to acid rain and photochemical smog.
  • Sulfur dioxide - produced from sulfur containing impurities present in crude oil.
  • Carbon particles - called particulates can exacerbate asthma and cause cancer
  • Carbon dioxide - a greenhouse gas which causes global warming
  • Water vapour - also a greenhouse gas
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Removing sulfur and catalytic converters

Power stations burn either coal, or more usually in the UK, natural gas, to produce electricity. Some chimneys now use calcium oxide or limestone to absorb the sulfur dioxide. This produces gypsum, which is used as plaster. This process is called flue gas desulfurisation.

Catalytic converters are honeycomb like structures made of a ceramic material coated with platinum and rhodium metals. These are the catalysts. The honeycomb shape provides an enormous surface area, so a little of these expensive metals goes a long way.

As the polluting gases pass over the catalyst, they react with each other to form less-harmful products.

They turn nitrogen oxide into nitrogen and unreacted hydrocarbons into nitrogen, carbon dioxide and water.

The reactions take place on the surface of the catalyst, on the layer of platinum and rhodium metals

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