Chemistry: Developing fuels (DF)

    a Gas volumes:

    • 1 is 10 × 10 × 10 cm, or 1 000
    • At any given temperature and pressure, one mole of any gas always has the same volume
    • At room temperature and pressure (298 K [25 °C] and 100 ), the molar volume (volume taken by one mole) of a gas is about 24.0
    • This can be used to calculate volumes from equations, e.g. 3H2(g) represents 3 moles of hydrogen gas, so it will occupy 3 × 24.0 = 72.0
    • The ratios of moles of each substance can be determined from the balanced symbol equation:
      For an equation 2A + 3B 4C + D, for every two moles of A, four moles of C are formed - due to the balancing numbers in the equation
    • If a substance is in excess, there is not enough of another reactant for the substance to fully react (some will be left over at the end)

    The ideal gas equation (pV = nRT):

    • pressure p () × volume V () = amount of gas n () × gas constant R () × temperature T (K)
    • The gas constant is always equal to 8.314
    • 1 = 1 000

    Measuring the volume of a gas:

    • Use a gas syringe, horizontally clamped
    • Alternatively, used an inverted burette/measuring cylinder placed in a trough of water

    b Bonding in organic compounds:

    • In a single C-C bond, electrons are arranged in an area of increased electron density between the carbon atoms, a σ bond (sigma bond). This area is where the orbitals for the two bonding electrons overlap in a straight line, as shown below:
    • One π bond (pi bond) has two areas of negative charge and high electron density, above and below the σ bond. It is formed when two p orbitals overlap sideways. A double bond contains both a σ bond and a π bond
    • A π bond is weaker than a σ bond

    c 3-D representations:

    • In a three-dimensional diagram, bonds in the plane of the paper are straight lines, bonds behind are dashed wedges and bonds in front are solid wedges
    • For example, methane has one bond behind the plane of the paper/screen and one in front, as shown below

    d Enthalpy changes:

    • In a chemical reaction, some bonds are broken and some are made. The energy change from start to finish is its enthalpy, which can be positive or negative. Remember to write the sign, even if it's positive
    • An exothermic reaction gives out energy, heating the surroundings and causing the products to have less energy than the reactants, giving a negative enthalpy
    • An endothermic reaction takes in energy, cooling the surroundings, giving a positive enthalpy


    • Enthalpy measurements need to be compared, so they are all measured under standard conditions (or standard temperature and pressure, stp). For OCR, these conditions are 298 K, 101 000 (1 atmosphere) pressure and 1 concentration for solutions
    • Under this temperature and pressure, the reactants/products are in their standard states
    • To show that an enthalpy is under standard conditions and with standard states, you can use a superscript plimsoll symbol (o)

    Enthalpy change:

    • There are several measures of enthalpy change. They all have the units
    • Standard enthalpy change of reaction (ΔrHo) is the enthalpy change when molar quantities (from the equation) react under standard conditions
    • Standard enthalpy change of combustion (ΔcHo) is the enthalpy change when one mole of a substance is completely burned in oxygen under standard conditions
    • Standard enthalpy change of formation (ΔfHo) is the enthalpy change when one mole of a compound is formed from its elements in their standard states
    • Standard enthalpy change of neutralisation (ΔneutHo) is the enthalpy change when one mole of H+ ions react with one mole of OH- ions to form one mole of water under standard conditions

    e Bond enthalpy:

    • Average bond enthalpy is the average quantity of energy required to break one mole of a particular bond into separate atoms in a gaseous state, measured in
    • It is an average because it varies slightly in different compounds, but only a generic one is often listed (e.g. the bond enthalpy of a C-C bond in an ethane molecule may differ slightly from the bond enthalpy of a C-C bond in a butanol molecule)
    • Because there is a higher attraction, double bonds have higher bond enthalpies than single bonds, and triple bond enthalpies are even higher

    Bond lengths:

    • Positively charged nuclei are attracted to the electrons in the atom they're bonded to, creating an attraction. Conversely, electrons repel electrons and nuclei repel nuclei. The two nuclei sit where these forces are balanced
    • Bonds with higher enthalpies typically have shorter bond lengths
    • Bond breaking is an endothermic process and bond-making is exothermic
      - So the enthalpy change of a reaction is calculated with enthalpy change = energy absorbed breaking bonds - energy released making bonds
      - Remember to divide your answer from the above equation by the number of moles formed since enthalpy change is for one mole

    f Measuring energy transfer when heating water:

    • Energy transferred q () = mass m (g) × specific heat capacity c () × temperature change Δ T (K)
    • The specific heat capacity of water is around 4.18 (given in the data sheet)
    • To convert from energy transferred to enthalpy, divide to 1000 to convert to and then divide this by the number of moles of the substance that caused the temperature change
    • Using a copper calorimeter and spirit burner isn't very accurate due to losses to the environment. Adding insulation and reducing air circulation (e.g. by closing windows) can reduce these losses, although the flame will still need enough oxygen for complete combustion so this method is never 100% accurate
    • Temperature changes are usually measured in K rather than °C
    • We need to assume that 1 of water weighs exactly 1 g, which is only an approximation at room temperature

    Measuring energy transfer when a reaction occurs in solution:

    • Measure the first solution into an insulated vessel. Measure the temperature every minute for 3 minutes
    • Once the stopclock reads 4:00, add the second solution to the vessel
    • Record the temperature at 5:00, 6:00, etc, until it stops changing
    • Plot a graph of temperature (y) against time (x). Draw a line of best fit through the two sections you have, and interpolate to find the approximate temperature when the second reactant was added at 4:00
    • An assumption here is that the specific heat capacity of the solution is equal to the specific heat capacity of water

    g Hess' law:

    • Measuring Δ H is not always simple; so it can also be measured indirectly using enthalpy cycles, using an indirect route
    • Here is an example of an enthalpy cycle:
    • Because the bottom right arrow is facing in the wrong direction, this enthalpy must be made negative
    • Finally, to calculate the enthalpy of the top arrow (Δ H), set the two routes equal to each other
      Δ H = -2326 + 2219 = -107
    • Some examples for common enthalpy calculations are shown below

    Calculating enthalpy changes from enthalpy of formation:

    • Requirements to calculate ΔrH: ΔfH for all reactants, and products which are compounds (ΔfH is zero for single elements)
    • Write the symbol equation at the top of the cycle diagram and the elements needed to make the reactants and products at the bottom. At the bottom, use H2, O2 etc for elements which bond together in standard states and single atoms for those that don't (e.g. C, S)
    • The arrows should both point up
    • enthalpy of reaction = sum of enthalpy of formation for products - sum of enthalpy of formation for reactants
    • Example: SO2 + 2H2S 3S + 2H2O:
      Calculation: ΔrH = ( 0 (S is an element) + (ΔfH of H2O × 2) (multiply by number of moles) ) - (ΔfH of SO2 + (ΔfH of H2S × 2))

    Calculating enthalpy changes from enthalpy of combustion:

    • Requirements to calculate ΔfH: ΔcH for the fuel, and elements which make up this fuel such as H2 and C (but exclude oxygen if it's an alcohol)
    • Write the symbol equation for the formation of the fuel from individual elements (or molecules if they bond together at standard state) at the top and the combustion products at the bottom (usually XCO2 + YH2O where X and Y are balancing numbers). Add enough oxygen next to the arrows to balance the equations
    • The arrows point down
    • enthalpy of formation = sum of enthalpy of combustion for elements making up fuel - enthalpy of combustion of fuel

    Calculating enthalpy of reaction from average bond enthalpies:

    • There is a method given for this in e, but OCR also require that we can also do the calculation with a Hess cycle
    • Requirements: average bond enthalpies of all bonds
    • Example: the formation of HCl gas from H2 and Cl2:
    • Write the symbol equation for the reaction on the top and the formation of the reactants from individual atoms on the bottom (2H + 2Cl)
    • The arrows point down
    • enthalpy of reaction = sum of bond enthalpies for reactants (top left) from atoms (bottom) - bond enthalpy for product
    • For the example: ΔrH = (436 + 243) - (2 (number of moles) × 431) = -183
      - This is the enthalpy for the formation of two moles, so halve it
        H-H: 436 , Cl-Cl: 243 , H-Cl: 431

    h Catalysts:

    • A catalyst is a substance that speeds up a reaction without being involved in it - it's chemically unchanged. Catalysts provide an alternate pathway with a lower activation energy, increasing rate of reaction
    • They can undergo physical change (e.g. crumbling), although the mass will remain unchanged
    • Catalysts don't affect the amount of product, only reaction time
    • Catalysis is the process of a catalyst speeding up a reaction
    • Heterogeneous catalysis is where the catalyst is in a different state to the reactants. Usually, this is with a solid catalyst and liquid/gaseous reactants
    • Homogeneous catalysts work when the reactants and products are in the same state

    Catalyst poisons:

    • A substance that stops a catalyst from properly functioning is called a catalyst poison
    • For example, in a heterogeneous catalyst, the poison is adsorbed more strongly than the reactants to the catalyst surface, causing the catalyst to become inactive - catalyst poisons don't necessarily destroy the catalyst

    i How a heterogenous catalyst works:

    • The reactants are adsorbed onto the catalyst surface; this weakens the bonds
    • The bonds are break, forming radicals (atoms or molecules with unpaired electrons)
    • New bonds will form as these radicals will quickly react with others
    • The products diffuse away from the catalyst surface - desorption
    • Now the catalyst surface is free to adsorb more molecules to react

    j Cracking:

    • Cracking is the process of making one larger organic molecule into two or more smaller ones, for example, reducing the size of hydrocarbon chains like below:
      C12H26 C6H14 + C2H4 + C4H8
    • Sometimes hydrogen gas is formed (e.g. C2H6 C2H4 + H2)
    • Cracking is random - different hydrogen/hydrocarbon molecules could be formed each time
    • The process requires very high temperatures and pressures without a catalyst, but a catalyst reduces this to lower temperatures and pressures

    Method for cracking a hydrocarbon vapour:

    • Clamp a boiling tube almost horizontally, pointing slightly upwards, with mineral wool soaked in a hydrocarbon (e.g. paraffin) at the end and porcelain chips in the centre (the catalyst). Use a test tube upside down in a water-filled trough and a delivery tube connecting the boiling tube and test tube
      - Heat the catalyst strongly, ensuring that the required temperatures are reached
      - Heat the alkane gently
      - Change collection tubes once one is full
      - When removing a tube, keep heating to prevent suck-back
      - Before stopping heating, remove the delivery tube from the trough for the same reason
      - Discard the first test tube because it will mainly contain displaced air rather than hydrocarbon vapour

    k Atmospheric pollutants:

    • Particulates are small particles of solids or liquids (e.g. carbon particles) and are produced by volcanoes and burning fuels. They can cause damage to animals and humans
    • Not all fuel is typically burnt by car engines - leaving some unburnt hydrocarbons. Some can be toxic to animals and humans, and contribute to photochemical smog
    • CO (carbon monoxide) is from incomplete combustion, and can cause photochemical smog. It's also a toxic gas
    • CO2 is produced by combustion and contributes to the greenhouse effect
    • Nitrogen oxides (NO) are produced by vehicles and power stations through the combustion of fuels in air (the air intake of a car will mostly contain nitrogen, some of which will react), and create photochemical smog and acid rain
    • Sulfur oxides (SO) also lead to acid rain and are toxic. They are produced by volcanoes and burnt fuels containing sulfur impurities, for example S + O2 SO2


    • Acid rain is caused by the products of combustion of some fuels reacting with water vapour in the atmosphere to form an acid. For example, SO2 + H2O H2SO3 (sulfurous acid). Nitrogen oxides from combustion/lightning can become nitric acid in the atmosphere by reacting in a similar way
    • Acid rain corrodes limestone and causes damage to life (e.g. trees, plants and animals in bodies of water)
    • Photochemical smog (containing molecules such as O3) occurs when a pollutant absorbs light and reacts due to the increase in energy. This smog can reduce visibility and cause respiratory problems

    Catalytic converters:

    • Many pollutants (e.g. CO) can be removed with a reaction to form safer ones (e.g. CO2)
    • This is often done in catalytic converters in cars. Carbon monoxide is removed (2CO + O2 2CO2), as well as nitrogen monoxide (2NO + 2CO N2 + 2CO2) and hydrocarbons (e.g. C8H18 + 12.5O2 8CO2 + 9H2O)
    • Filters (often ceramic) can be used to catch particulates, with the carbon being burnt off occasionally by raising the temperature
    • In diesel engines, nitrogen oxide compounds are formed in large quantities. These emissions are usually reduced by cooling them to lower the temperature (and increase the volume) before injecting them into the engine's intake air. This modified mixture reduces the combustion temperature, reducing further NO production as less is created at lower temperatures
    • A reagent like ammonia could also be used (4NO + 4NH3 + O2 4N2 + 6H2O). This is usually not used in consumer vehicles as it increases upkeep costs

    Other pollutant removal methods:

    • Sulfur oxides can be removed by reacting them with calcium oxide

    l Arenes:

    • Benzene is a hydrocarbon, C6H6 (a closed ring shape rather than a long chain). It has three double bonds, but the p electrons are delocalised around the whole ring so it reacts slightly differently to other hydrocarbons with double bonds
    • [YEAR ONE ONLY]: A compound containing at least one benzene ring is aromatic, an arene. Compounds containing no benzene rings are aliphatic
    • [YEAR TWO ONLY]: A compound containing at least one ring of a planar flat molecule with delocalisation (such as benzene) is aromatic, an arene. Compounds without any of these structures are aliphatic
    • The skeletal formulae of benzene are shown below. (It can be drawn in two ways due to its delocalisation, but the second is generally preferred)


    • A saturated hydrocarbon contains only single carbon-carbon bonds, because they have the maximum number of hydrogen atoms possible
    • An unsaturated hydrocarbon contains double or triple carbon-carbon bonds
    • Alkanes are saturated hydrocarbons and alkenes are unsaturated hydrocarbons

    Functional groups:

    • A functional group is a modifier such as -OH (hydroxyl). Compounds containing this group are known as alcohols. Functional groups are responsible for the specific chemical reactions of the molecules containing them
    • Examples include: alkene, hydroxyl, carboxyl

    Homologous series:

    • A homologous series is a group of compounds with the same general formulae but different molecular formulae (such as the alkanes - methane, ethane, propane, ...)

    m General formulae:

    • Alcohols: CHOH
    • Alkanes: CH
    • Alkenes: CH
    • Cycloalkanes: CH

    Naming hydrocarbons:

    • The prefixes for chains of up to 10 carbon atoms are:
      - 1 carbon: meth-
      - 2 carbons: eth-
      - 3 carbons: prop-
      - 4 carbons: but-
      - 5 carbons: pent-
      - 6 carbons: hex-
      - 7 carbons: hept-
      - 8 carbons: oct-
      - 9 carbons: non-
      - 10 carbons: dec-

    Naming alkanes:

    • Count the number of carbon atoms in the longest chain and find the prefix from the list above. Add 'ane' to the end to indicate that the compound is an alkane.
    • Now name side chains in alphabetical order with the carbon atom it is attached to. If there is more than one of a type of side chain, use 'di', 'tri', or 'tetra' to indicate the quantity
      - For example, 2-methylpentane has 5 carbon atoms in the longest chain, with a methyl side chain (CH3) attached to the second carbon atom from the closest end of the chain
      - 2,3-dimethylbutane has methyl groups on the second and third carbon atoms
    • Alkane side chains are called alkyl groups and have the general formula CH
    • Use a hyphen between each name and number (e.g. 3-ethyl-2,4-dimethylpentan-2-ol)
    • Unless a functional group has a space in its name (e.g. 'oic acid' for a carboxylic acid), the name should not contain any spaces

    Naming alkenes:

    • Alkenes are similar, but with 'ene' instead of 'ane'
    • There should also be a number between the prefix and 'ene' indicating the position of the carbon-carbon double bond if there are more than three carbon atoms. For example, 'but-1-ene' has a carbon-carbon double bond on one of the carbon atoms on the end of the chain

    Naming cycloalkanes:

    • Use the prefix 'cyclo'
    • Use the suffix 'ane'
    • For example, a cycloalkane with 5 carbon atoms is called cyclopentane

    Naming alcohols:

    • Alcohols use '-ol' after the prefix, with the number inserted showing the carbon atom which the OH side chain is on. Example: 'pentan-2-ol'

    n Complete combustion:

    • Alcohols can be used as fuels, as well as alkenes and alkanes
    • The general equation for complete combustion is fuel(l) /(g) + oxygen(g) carbon dioxide(g) + water(g)
    • These reactions are exothermic

    Incomplete combustion:

    • Incomplete combustion occurs where there isn't enough oxygen for complete combustion
    • C and/or CO will be produced instead of CO2. If both could be produced, write them as two separate equations

    p Addition polymerisation:

    • Polymers are long molecules made from smaller monomers. For example, poly(ethene) is many ethene molecules chained together
    • Copolymers consist of multiple types of monomer
    • Polymers (in Developing Fuels) are formed with addition polymerisation. Monomers will typically contain double bonds which are broken in the process (so that the carbon atoms can bond to the other monomers)
    • A repeating unit shows how monomers join, like this:

    o, q Addition:

    • An addition reaction is where two or more molecules join, breaking a double bond to form a single product
    • The below reactions show that the C=C double bond is much more reactive than a C-C single bond


    • An electrophile has a partial (or sometimes full) positive charge (usually due to a vacant orbital), meaning that it is attracted to negatively charged regions and reacts by accepting a pair of electrons to form a covalent bond
    • Electrophiles can do this by accepting an electron pair from the double C=C bond
    • A carbocation is an ion containing a positively charged carbon atom
    • Curly arrows are used to indicate the movement of a pair of electrons

    Bromine (Br2) electrophilic addition:

    • Bromine decolourises when it reacts with alkenes at room temperature, to give a dibromo compound
    • First, the double bond induces a dipole in the bromine (see OZ). This means that one of the bromine atoms has the bonding pair slightly closer to it than the other, making the other side an electrophile
    • The Br with a partial positive charge attacks the double bond, gaining two of its electrons and forming a single dative bond
    • The other Br atom from the Br2 molecule gains the two electrons from its old bond, making it negatively charged
    • Finally, a pair of electrons from the remaining negative bromine atom bonds with the carbocation and shares two electrons
    • If another anion (e.g. Cl-) is also present, it will react with the carbocation. This can be used as evidence for the mechanism

      Testing for unsaturation:

    • The above reaction is useful for testing for unsaturation (i.e. an alkene)
    • Add a few drops of low concentration Br2(aq) (bromine water) to about 1 of the sample in a test tube. Stopper the tube and shake
    • If the solution is unsaturated, the solution will turn colourless as the dibromo compound is formed

    Hydrogen bromide electrophilic addition:

    • Hydrogen bromide also reacts at room temperature with some alkenes, to give a bromo compound
    • The mechanism is very similar to the above one, except that the HBr is already a dipole, with the H atom having a partial positive charge
    • e.g. C2H4 + HBr C2H5Br

    Electrophilic addition of an alkene and water:

    • Ethene and water (steam under the required conditions) react to create ethanol at high temperature (300 °C), high pressure (60 atm) and with a H3PO4 catalyst. The yield is low (about 5%), and the reaction is reversible
    • Without these conditions, it can also be made by adding concentrated sulfuric acid (H2SO4):
    • This forms ethyl hydrogen sulfate. Cold water is added and the product heated to produce ethanol and sulfuric acid - this step is hydrolysis (water splitting a molecule into two)
    • The sulfuric acid is a homogeneous catalyst in this reaction


    • C2H4 + H2 C2H6
    • Here, a catalyst is required to break the H-H bond
    • Platinum can be used at room temperature
    • Alternatively, nickel (which is cheaper than platinum) can be used if powdered, at 150 °C and at 5 atm pressure

    r Formulae:

    • A full structural formula (sometimes called displayed formula) is a diagram, e.g.:
    • The shortened structural formula shows what each carbon atom bonds to, from left to right or right to left, e.g. CH4 for methane or CH3-CH2-CH3 (or CH3CH2CH3) for propane
    • A skeletal formula shows carbon-carbon bonds, e.g.:

    s Isomerism:

    • A structural isomer has the same molecular formula as another isomer but a different structural formula (so it has the same number of atoms, but in a different arrangement)
    • With four or more carbon atoms in an alkane, isomers are possible. These are called chain isomers
    • Isomers usually don't share the same properties

    t Stereoisomerism:

    • A C=C double bond is rigid, so the exact arrangement needs to be specified in more detail
    • For molecules with one hydrogen atom bonded to each carbon in the C=C double bond, a Z isomer has the hydrogen groups on the same side, whilst an E isomer has the hydrogen groups on opposite sides
    • For molecules with different atoms/molecules bonded to the carbons than H, use cis if the groups that are the same are on the same side, and trans if the groups that are the same are on opposite sides
    • E, Z, cis or trans, followed by a hyphen, goes just before the molecule's prefix

    u Sustainability:

    • Biofuels (e.g. ethanol and biodiesel) are produced from plant/animal materials
      - Burning them still releases CO2
      - However, if the plants/trees used to produce them are replanted, the new ones 'reabsorb' this CO2 with photosynthesis
    • Hydrogen is another potential fuel, created through electrolysis of water. Combusting hydrogen produces only water (2H2 + O2 2H2O)
    • However, it is less energy dense than petrol and nitrogen oxides are still produced with current engines at high temperatures