PAPER 2

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4.6.1.1 Calculating rates of reactions

  • The rate of a chemical reaction can be found by measuring the quantity of a reactant used or the quantity of product formed over time:
  • mean rate of reaction = quantity of reactant used / time taken
  • mean rate of reaction = quantity of product formed /  time taken
  • The quantity of reactant or product can be measured by the mass in grams or by a volume in cm3.
  • The units of rate of reaction may be given as g/s or cm3/s.
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4.6.1.2 Factors which affect the rates of chemica

  • Factors which affect the rates of chemical reactions include:
  • the concentrations of reactants in solution, the pressure of reacting gases, the surface area of solid reactants, the temperature and the presence of catalysts. 
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4.6.1.3 Collision theory and activation energy

  • Collision theory explains how various factors affect rates of reactions.
  • According to this theory, chemical reactions can occur only when reacting particles collide with each other and with sufficient energy.
  • The minimum amount of energy that particles must have to react is called the activation energy.
  • Increasing the concentration of reactants in solution, the pressure of reacting gases, and the surface area of solid reactants increases the frequency of collisions and so increases the rate of reaction.
  • Increasing the temperature increases the frequency of collisions and makes the collisions more energetic, and so increases the rate of reaction. 
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4.6.1.4 Catalysts

  • Catalysts change the rate of chemical reactions but are not used up during the reaction.
  • Different reactions need different catalysts.
  • Enzymes act as catalysts in biological systems.
  • Catalysts increase the rate of reaction by providing a different pathway for the reaction that has a lower activation energy.
  • A reaction profile for a catalysed reaction can be drawn in the following form:
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4.6.2.1 Reversible reactions

  • In some chemical reactions, the products of the reaction can react to produce the original reactants.
  • Such reactions are called reversible reactions and are represented:
  • The direction of reversible reactions can be changed by changing the conditions.
  • For example:
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4.6.2.2 Energy changes and reversible reactions

  • If a reversible reaction is exothermic in one direction, it is endothermic in the opposite direction. The same amount of energy is transferred in each case.
  • For example:
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4.6.2.3 Equilibrium

  • When a reversible reaction occurs in apparatus which prevents the escape of reactants and products, equilibrium is reached when the forward and reverse reactions occur at exactly the same rate.
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4.6.2.4 The effect of changing conditions on equi

  • The relative amounts of all the reactants and products at equilibrium depend on the conditions of the reaction.
  • If a system is at equilibrium and a change is made to any of the conditions, then the system responds to counteract the change.
  • The effects of changing conditions on a system at equilibrium can be predicted using Le Chatelier’s Principle.
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4.6.2.5 The effect of changing concentration

  • If the concentration of one of the reactants or products is changed, the system is no longer at equilibrium and the concentrations of all the substances will change until equilibrium is reached again.
  • If the concentration of a reactant is increased, more products will be formed until equilibrium is reached again.
  • If the concentration of a product is decreased, more reactants will react until equilibrium is reached again. 
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4.6.2.6 The effect of temperature changes on equi

  • If the temperature of a system at equilibrium is increased:
  • - the relative amount of products at equilibrium increases for an endothermic reaction
  • - the relative amount of products at equilibrium decreases for an exothermic reaction.
  • If the temperature of a system at equilibrium is decreased:
  • - the relative amount of products at equilibrium decreases for an endothermic reaction
  • - the relative amount of products at equilibrium increases for an exothermic reaction. 
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4.6.2.7 The effect of pressure changes on equilib

  • For gaseous reactions at equilibrium:
  • - an increase in pressure causes the equilibrium position to shift towards the side with the smaller number of molecules as shown by the symbol equation for that reaction
  • - a decrease in pressure causes the equilibrium position to shift towards the side with the larger number of molecules as shown by the symbol equation for that reaction. 
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4.7.1.1 Crude oil, hydrocarbons and alkanes

  • Crude oil is a finite resource found in rocks.
  • Crude oil is the remains of an ancient biomass consisting mainly of plankton that was buried in mud.
  • Crude oil is a mixture of a very large number of compounds.
  • Most of the compounds in crude oil are hydrocarbons, which are molecules made up of hydrogen and carbon atoms only.
  • Most of the hydrocarbons in crude oil are hydrocarbons called alkanes.
  • The general formula for the homologous series of alkanes is
  • CnH
  • The first four members of the alkanes are methane, ethane, propane and butane.
  • Alkane molecules can be represented in the following forms:
  • C2H6 or
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4.7.1.2 Fractional distillation and petrochemical

  • The many hydrocarbons in crude oil may be separated into fractions, each of which contains molecules with a similar number of carbon atoms, by fractional distillation.
  • The fractions can be processed to produce fuels and feedstock for the petrochemical industry.
  • Many of the fuels on which we depend for our modern lifestyle, such as petrol, diesel oil, kerosene, heavy fuel oil and liquefied petroleum gases, are produced from crude oil.
  • Many useful materials on which modern life depends are produced by the petrochemical industry, such as solvents, lubricants, polymers, detergents.
  • The vast array of natural and synthetic carbon compounds occur due to the ability of carbon atoms to form families of similar compounds.
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4.7.1.3 Properties of hydrocarbons

  • Some properties of hydrocarbons depend on the size of their molecules, including boiling point, viscosity and flammability.
  • These properties influence how hydrocarbons are used as fuels. 
  • The combustion of hydrocarbon fuels releases energy.
  • During combustion, the carbon and hydrogen in the fuels are oxidised.
  • The complete combustion of a hydrocarbon produces carbon dioxide and water. 
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4.7.1.4 Cracking and alkenes

  • Hydrocarbons can be broken down (cracked) to produce smaller, more useful molecules. Cracking can be done by various methods including catalytic cracking and steam cracking. 
  • The products of cracking include alkanes and another type of hydrocarbon called alkenes.
  • Alkenes are more reactive than alkanes and react with bromine water, which is used as a test for alkenes. 
  • There is a high demand for fuels with small molecules and so some of the products of cracking are useful as fuels.
  • Alkenes are used to produce polymers and as starting materials for the production of many other chemicals. 
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4.7.2.1 Structure and formulae of alkenes

  • Alkenes are hydrocarbons with a double carbon-carbon bond.
  • The general formula for the homologous series of alkenes is
  • Cn
  • Alkene molecules are unsaturated because they contain two fewer hydrogen atoms than the alkane with the same number of carbon atoms.
  • The first four members of the homologous series of alkenes are ethene, propene, butene and pentene.
  • Alkene molecules can be represented in the following forms:
  • C3H6
  • or
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4.7.2.2 Reactions of alkenes

  • Alkenes are hydrocarbons with the functional group C=C.
  • It is the generality of reactions of functional groups that determine the reactions of organic compounds.
  • Alkenes react with oxygen in combustion reactions in the same way as other hydrocarbons, but they tend to burn in air with smoky flames because of incomplete combustion.
  • Alkenes react with hydrogen, water and the halogens, by the addition of atoms across the carbon-carbon double bond so that the double bond becomes a single carbon-carbon bond.
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4.7.2.3 Alcohols

  • Alcohols contain the functional group –OH.
  • Methanol, ethanol, propanol and butanol are the first four members of a homologous series of alcohols.
  • Alcohols can be represented in the following forms:
  • CH3CH2OH or
  • Aqueous solutions of ethanol are produced when sugar solutions are fermented using yeast.
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4.7.2.4 Carboxylic acids

  • Carboxylic acids have the functional group –COOH.
  • The first four members of a homologous series of carboxylic acids are methanoic acid, ethanoic acid, propanoic acid and butanoic acid.
  • The structures of carboxylic acids can be represented in the following forms:
  • CH3COOH 
  • or
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4.7.3.1 Addition polymerisation

  • Alkenes can be used to make polymers such as poly(ethene) and poly(propene) by addition polymerisation. I
  • n addition polymerisation reactions, many small molecules (monomers) join together to form very large molecules (polymers).
  • For example:
  • In addition polymers the repeating unit has the same atoms as the monomer because no other molecule is formed in the reaction.
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4.7.3.2 Condensation polymerisation

  • Condensation polymerisation involves monomers with two functional groups.
  • When these types of monomers react they join together, usually losing small molecules such as water, and so the reactions are called condensation reactions.
  • The simplest polymers are produced from two different monomers with two of the same functional groups on each monomer.
  • For example: ethane diol

and hexanedioic acid


polymerise to produce a polyester:

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4.7.3.3 Amino acids

  • Amino acids have two different functional groups in a molecule.
  • Amino acids react by condensation polymerisation to produce polypeptides.
  • For example: glycine is H2NCH2COOH and polymerises to produce the polypeptide
  • (-HNCH2COO-)n and n H2O
  • Different amino acids can be combined in the same chain to produce proteins.
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4.7.3.4 DNA (deoxyribonucleic acid) and other nat

  • DNA (deoxyribonucleic acid) is a large molecule essential for life.
  • DNA encodes genetic instructions for the development and functioning of living organisms and viruses.
  • Most DNA molecules are two polymer chains, made from four different monomers called nucleotides, in the form of a double helix.
  • Other naturally occurring polymers important for life include proteins, starch and cellulose.
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4.8.1.1 Pure substances

  • In chemistry, a pure substance is a single element or compound, not mixed with any other substance.
  • Pure elements and compounds melt and boil at specific temperatures.
  • Melting point and boiling point data can be used to distinguish pure substances from mixtures.
  • In everyday language, a pure substance can mean a substance that has had nothing added to it, so it is unadulterated and in its natural state, eg pure milk. 
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4.8.1.2 Formulations

  • A formulation is a mixture that has been designed as a useful product.
  • Many products are complex mixtures in which each chemical has a particular purpose.
  • Formulations are made by mixing the components in carefully measured quantities to ensure that the product has the required properties.
  • Formulations include fuels, cleaning agents, paints, medicines, alloys, fertilisers and foods. 
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4.8.1.3 Chromatography

  • Chromatography can be used to separate mixtures and can give information to help identify substances.
  • Chromatography involves a stationary phase and a mobile phase.
  • Separation depends on the distribution of substances between the phases.
  • The ratio of the distance moved by a compound (centre of spot from origin) to the distance moved by the solvent can be expressed as its Rf value:
  • Rf =  distance moved by substance / distance moved by solvent 
  • Different compounds have different Rf values in different solvents, which can be used to help identify the compounds.
  • The compounds in a mixture may separate into different spots depending on the solvent but a pure compound will produce a single spot in all solvents.
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4.8.2.1 Test for hydrogen

  • The test for hydrogen uses a burning splint held at the open end of a test tube of the gas.
  • Hydrogen burns rapidly with a pop sound.
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4.8.2.2 Test for oxygen

  • The test for oxygen uses a glowing splint inserted into a test tube of the gas.
  • The splint relights in oxygen.
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4.8.2.3 Test for carbon dioxide

  • The test for carbon dioxide uses an aqueous solution of calcium hydroxide (lime water).
  • When carbon dioxide is shaken with or bubbled through limewater the limewater turns milky (cloudy).
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4.8.2.4 Test for chlorine

  • The test for chlorine uses litmus paper.
  • When damp litmus paper is put into chlorine gas the litmus paper is bleached and turns white.
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4.8.3.1 Flame tests

  • Flame tests can be used to identify some metal ions (cations).
  • Lithium, sodium, potassium, calcium and copper compounds produce distinctive colours in flame tests:
  • - lithium compounds result in a crimson flame
  • - sodium compounds result in a yellow flame
  • - potassium compounds result in a lilac flame
  • - calcium compounds result in an orange-red flame
  • - copper compounds result in a green flame.
  • If a sample containing a mixture of ions is used some flame colours can be masked
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4.8.3.2 Metal hydroxides

  • Sodium hydroxide solution can be used to identify some metal ions (cations).
  • Solutions of aluminium, calcium and magnesium ions form white precipitates when sodium hydroxide solution is added but only the aluminium hydroxide precipitate dissolves in excess sodium hydroxide solution.
  • Solutions of copper(II), iron(II) and iron(III) ions form coloured precipitates when sodium hydroxide solution is added.
  • Copper(II) forms a blue precipitate, iron(II) a green precipitate and iron(III) a brown precipitate.
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4.8.3.3 Carbonates

  • Carbonates react with dilute acids to form carbon dioxide gas.
  • Carbon dioxide can be identified with limewater.
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4.8.3.4 Halides

  • Halide ions in solution produce precipitates with silver nitrate solution in the presence of dilute nitric acid.
  • Silver chloride is white, silver bromide is cream and silver iodide is yellow.
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4.8.3.5 Sulfates

  • Sulfate ions in solution produce a white precipitate with barium chloride solution in the presence of dilute hydrochloric acid.
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4.8.3.6 Instrumental methods

  • Elements and compounds can be detected and identified using instrumental methods.
  • Instrumental methods are accurate, sensitive and rapid. 
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4.8.3.7 Flame emission spectroscopy

  • Flame emission spectroscopy is an example of an instrumental method used to analyse metal ions in solutions.
  • The sample is put into a flame and the light given out is passed through a spectroscope.
  • The output is a line spectrum that can be analysed to identify the metal ions in the solution and measure their concentrations.
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