AQA GCSE Chemistry C2.3 How Much?

revision notes :)


C2.3.1 Mass of Atoms

  • An atom has a nucleus containing postively charged protons and neutrons which have no charge. Electrons, negatively charged, are arranged in shells around the nucleus.
  • Each atom has the same number of protons and electrons. The number of protons an atom has is called the atomic number.
  • The mass of a proton and a neutron is equal, and electrons are a lot lighter which means that when working out the relative mass of an atom, the electrons can be ignored.
  • Therefore, the relative mass of an atom is all in the nucleus, and is the mass of the protons added to the mass of neutrons. Instead of calculating the actual mass, we call the total number of protons and neutrons the mass number.
  • number of neutrons=mass number - atomic number.
  • Atoms of the same element but with different numbers of neutrons are called isotopes-the atomic number is still the same. Sometimes, the extra neutrons make the nucleus unstable, so it is radioactive. This is only for some isotopes.
  • Isotopes of hydrogen are deuterium (mass number of 2) and tritium (mass number of 3). They have identical chemical properties but different physical properties. Tritium is radioactive but all react with oxygen to make water.
  • The electronic structure is the same in isotopes.
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C2.3.2 Masses of Atoms and Moles

  • The actual mass of an atom is so small it wouldn't be useful to use in experiments/calculations, so instead the relative mass is used; (Ar)
  • The relative atomic mass of an element is usually the same/similar to the mass number of its most common isotopes of the element which is found naturally-an average mass.
  • Relative formula masses (Mr) of compounds:
  • E.g.: CO2
  • Ar of C=12
  • Ar of O=6 there are two oxygen atoms in the compound so 6x2
  • Mr=12+6x2=24
  • Moles:
  • This is sort of a shorthand of 'relative atomic mass in grams'.
  • A mole is equal to 6.02 x 10^23.
  • For example, in 6g of oxygen atoms, there are 6.02 x 10^23 oxygen atoms.
  • The same way, the relative atomic mass for carbon in grams is 12g, which means there is a mole of carbon atoms in 12g of it.
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C2.3.3 Percentages and Formulae

  • The formula mass of a compund can be used to calculate the percentage mass of each element in it.
  • Example- in order to work out the percentage mass of magnesium in magnesium oxide:
    • The formula of magnesium oxide is MgO.
    • The Ar of magnesium is 24, and the Ar of oxygen is 16. Adding these together, the Ar is 24 + 16 = 40.
    • In 40g magnesium oxide, 24g is actually magnesium, so the percentage of Mg in MgO is 24 ÷ 40 = 60%.
  • Empirical formula:
  • The percentage of each element in a compound can be found by experiments, then the simplest ratio to each type of atom in the compund can be found-the empirical formula.
  • Sometimes this is the same as the actual number of atoms in one molecule-the molecular formula, but not always. for example, hydrogen peroxide has the empirical formula HO, but molecular formula of H2O2.
  • Example: - A hydrocarbon has 75% carbon and 25% hydrogen by mass. What's the empirical formula?
    • If there was 100g of the compound, there would be 75g carbon and 25g hydrogen.
    • Then, work out the number of moles by dividing the mass of each by the relative atomic mass:
      • Carbon- 75/12=6.25 moles of carbon atoms. Hydrogen-25/1= 25 moles of hydrogen atoms.
    • This shows that there are 6.25 C moles with every 25 H moles; the ratio is 6.25:25. Then you divide both these numbers by the smallest number in the ratio, getting 1:4. So, there are 4 hydrogen atoms for every carbon atom.
    • The empirical formula is CH4.
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C2.3.4 Equations and Calculations

  • Masses can be calculated from balanced chemical equations.
  • If the equation is: H2 + Cl2 --> 2HCl:
  • The Ar of hydrogen=1 and the Ar of chlorine is 35.5.
  • The mass of 1 mole of H2 = 2x1=2g. The mass of one mole of Cl2=  2x35.5 =76g
  • Mass of HCl= 1+35.5=36.5. The mass of one mole of HCl=36.5g
  • The equation states that 1 mole of hydrogen reacting with one mole of chlorine gives two moles of HCl, so if these are turned into masses:
  • 1 mole of hydrogen =1x2g=2g; 1 mole of chlorine=1x71=71g; 2 moles of HCl=2.36.5=73g.
  • Calculations
  • Sodium hyroxide+chlorine gas-->bleach. 2NaOH+ Cl2--> NaOCl+NaCl+H2
  • If there is 100g of sodium hydroxide, how much Cl2 is needed to make bleach?
  • Ar of H=1; of O2=16, Ar of Na=23, Ar of Cl2=35.5.
  • The mass of 1 mole of NaOH would be 23+16+1=40. Mass of 1 mole of Cl2=35.5x2=71.
  • Therefore, 100g of NaOH=100/40=2.5 moles.
  • The equation shows that for every 2 moles of sodium hydroxide, 1 mole of chlorine is needed, so 2.5/2=1.25 moles of chlorine.
  • 1 mole of chlorine has a mass of 71g, so 1.25x71=88.75g. 88.75g of chlorine is needed to react with 100g of sodium hydroxide.
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C2.3.5 Yield part 1

  • In a reaction, not all of the reactants are used up, and less product than the mass of the reactants added together is made. The mass of the product is the yield, and the percentage yield is the percentage of the amount made compared to the maximum amount that could have been made.
  • Percentage yield=amount of product produced ÷ maximum amount of product possible x100%
  • Limestone is mostly made of CaCO3, which decomposes in a kiln to make CaO and CO2.
  • If 200 tonnes of CaCO3 is processes per day, and 98 tonnes of CaO is collected. What's the percentage yield? 
  • Ar : Ca= 40, C=12, O=16. CaCO3 --> CaO + CO2.
  • M of CaCO3=40+12+16x3=100. Mr of CaO=40+16=56.
  • Therefore, 100 tonnes of CaCO could make 56 tonnes of CaO if the yield was 100%, and 200 tonnes of CaCO3 could make 112 (56x2) tonnes of CaO. Therefore, 98/112=87.5%.
  • As an explanation, some limestone could be lost as dust in the crushing process.
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C2.3.5 Yield part 2

  • Few reactions do have 100% yield because:
  • Some reactions are reversible
  • Some react to give unexpected products
  • Some of the product may be lost in handling or left in the apparatus
  • The reactants might not be fully pure and some chemical reactants produce more than one product and it may be hard to seperate the product that is wanted from the mixture.
  • Reactions with high yields in industry help conserve energy resources and reduce waste.
  • Chemical factories/plants are designed by chemical engineers so it works as safely and economically as possible, so as little energy and raw material possible is wasted, helping the company to make a profit.
  • It is also better for the environment; conserving our limited resources and reducing pollution.
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C2.3.6 Reversible Reactions

  • In some reactions, the products react to produce the reactants again.
  • Examples:
  • It is difficult to make a solution that's exactly neutral with an alkaline solution and an acid.
  • An indicator has to be used to tell when they right amount of acid has been added-they react in different pH conditions to formed differently coloured compounds.
  • Litmus is a complex molecule. If represented as HLit (H=hydrogen), HLit is red, but when alkali is added, it turns into Lit ion by losing an H+ ion. If more acid is added, blue Lit becomes red HLit etc: HLit  D H+ + Lit -                                      
  • Ammonium chloride breaks down on heating and forms ammonia gas and hydrogen chloride gas-example of thermal decomposition. When the two gases cool down they react and re-form ammonium choride:
  • ammonium chloride (NH3)   ammonia (NH4Cl) +  Hydrogen chloride (HCL)
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C2.3.7 Analysing Substances

  • Substances added to food to extend shelf life, taste or appearance-a food additive. Additives approved for use in Europe are given E numbers, which are like codes to identify the additives. E.g.E102-yellow food colouring called tartrazine.
  • Many instruments can be used to identify unknown compounds, including food additives.
  • One technique-chromatography-it works as some compounds in a mixture dissolve better than others in particular solvents and their solubility determines how far they travel across the paper.
  • Once compounds in a food have been seperated with chromatography, they can be identified-we can compare the chromatogram with others obtained from known substances, and the same solvent and temperature must be used.
  • Instrumental methods-rapid and accurate methods, quicker and enable small samples to be analysed.
  • Also important in fighting pollution; careful monitoring of the environment using sensitive instruments is common, and this type of analysis is used all the time in healthcare.
  • Disadvantages-usually very expensive, takes special training to use and they often give results that can only be interpreted by comparison with data from known substances.
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C2.3.8 Instrumental Analysis-Gas Chromatography

  • Samples being analysed-often mixtures of different compounds, so the first step-seperate the compounds. Gas chromatography-mass spectrometry (GC-MS) is used.
  • Gas chromatography-seperates compounds , which are then passed on to the mass spectrometer.
  • Gas chromatography-technique similar to paper. Instead of a solvent moving over paper, it has gas moving through a column packed with a solid. The mixture is first vapourised, then a 'carrier' gas moves it through the coiled column, and different compounds have different attractions to the material in the column-the compounds with stronger attractions take longer to get through the column, with a longer retention time.
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C2.3.8 Instrumental Analysis-Mass Spectrometer

  • The unknown substances in the sample can be identified by comparing the chromatograph with results for known substances and the analysis for the known substances had to have taken place in identical conditions to compare retention times.
  • The gas chromatography apparatus can be attached directly to a mass spectrometer, which identifies the substances very quickly and accurately and can detect very small quantities in the sample.
  • Measuring relative molecular masses:
    • A mass spectrometer also provides an accurate way of measuring the Mr of a compound-the peak with the largest mass corresponds to an ion with just one electron removed. The mass of the electron is so tiny it is ignored when looking at the mass of atoms.
    • The peak is called the molecular ion peak-always found as the last peak on the right when you see a mass spectrum.
    • The pattern of peaks, the mass spectrum, acts as a 'fingerprint' for unknown compounds, and the pattern is quickly matched against a database of known compounds stored on the computer.
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