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- 1. C5a Moles and Empirical Formulae • The relative atomic mass of an element is the average mass of an atom of the element compared to the mass of _______ of an atom of _____________. • It can be found on the periodic table. From the relative atomic mass and atomic number we can deduce the number of ______ and ______ in the nucleus of an atom. Relative Formula Mass Calculations – revision Calculate the relative formula mass of the following compounds (show calculations). 1. NaOH = 2. CaSO4 = 3. Fe2O3 = 4. Na2O = 5. Ca(NO3)2 = Note that the relative formula mass has no units Molar Mass Calculations • The unit for the amount of a substance is the mole. You don’t need to know anything about moles at the moment. It is just an easy way for chemists to handle very large numbers of atoms. A pair of shoes = __ shoes A trio of apples = __ apples A dozen apples = __ A mole of apples = _________ i.e. 602 000 000 000 000 000 000 000 apples • The molar mass of a substance is its relative formula mass in grams. (it is the same number as the relative formula mass but it has units) Relative Formula Mass Molar Mass H2O 18 18 g/mol NaOH CaSO4 Fe2O3 Na2O
- 2. Higher It is important for chemists to know exactly how much of a chemical to use in a reaction for two reasons: Cost - wasting chemicals costs money Safety – some chemicals are more hazardous than others and we must make sure that they react completely Consider the following reaction: We want to react 10 grams of lithium with iodine to make lithium iodide lithium + iodine lithium iodide 2Li I2 2LiI 10g Two Problems: Lithium reacts vigorously with water and oxygen so we need to make sure that there is enough iodine to completely react all of the lithium. Iodine is hazardous so we don’t want to have too much left over. 1. Work out how many moles (atoms) there are in 10g of Li The relative atomic mass (from the periodic table) tells us that each lithium atom has a molar mass of ___ g/mole moles = mass Therefore in 10g there must be molar mass 2. Look at the symbol equation above 2 atoms of lithium react with ___ molecule of iodine 2 moles Li __ mole I2 2.5 moles 3. Mass of Iodine needed to react with lithium The relative atomic mass of iodine is ____ So the molar mass = _____ g/mol mass = moles x molar mass So 12.5 moles = = g of iodine So, to make sure that all 10 grams of lithium reacts, we must have a minimum of 159g of iodine 4. Supposing we wanted to add a 2% excess of iodine to make certain the lithium has reacted. How much iodine in total would we need?
- 3. Practice Questions 1. Determine the number of moles of an element from the mass of that element moles = mass / molar mass a) 2 g of hydrogen, H2 f) 184g of sodium b) 1 g of hydrogen g) 4.6g of sodium c) 10g of hydrogen h) 32g of oxygen d) 56g of iron i) 4g of oxygen e) 14g of iron j) 10.8g of Al 2. determine the number of moles of a compound from the mass of that compound a) 10g of NaOH b) 272g of CaSO4 c) 32g Fe2O3 d) 31g Na2O e) 16.4g Ca(NO3)2 f) 58.5g of NaCl g) 75g CaCO3 h) 5.1g Al2O3 3. Determine the masses of the different elements present in a given number of moles of a compound
- 4. The Law of Conseravtion of Mass • Matter cannot be created or destroyed. • Atoms of elements can rearrange to form new substances but atoms cannot disappear completely. • In a chemical reaction mass is conserved • This means that the total mass of the reactants must be equal to the total mass of the products. mass of reactants = mass of products e.g. lithium + iodine lithium iodide 2Li I2 2LiI 10g 159g ____ g 1. Calcium carbonate, CaCO3, decomposes on heating to produce calcium oxide, CaO and carbon dioxide, CO2. a) Write a balanced symbol equation for the reaction: b) If 100g of CaCO3 produces 56g of CaO, how many grams of CO2 are produced? c) If 100 tonnes of CaCO3 produces 56 tonnes of CaO, how many tonnes of CO2 are produced? d) If 300g of CaCO3 produces 168g of CaO, how many grams of CO2 are produced? e) If 25g of CaCO3 produces 14g of CaO, how many grams of CO2 are produced? Higher f) On a separate sheet of paper, calculate out how many moles of CaCO3, CaO and CO2 are produced in b) to e). Do you notice any patterns in terms of the number of moles? 2. 64 g of copper reacts with 16 g of oxygen. What mass of copper oxide forms? 3. 160g of anhydrous copper sulfate, CuSO4, is dissolved in water. When it crystallises it forms hydrated copper sulfate crystals, CuSO4.5H2O, with a mass of 250g. What mass of water has been used to make the crystals?
- 5. Extension: for question 2 a) calculate how many moles of each substance there are and b) work out the molecular formula for copper oxide.
- 6. HIGHER – Practice calculations on reacting quantities 1. If you add 5.3g of sodium carbonate, Na2CO3 to sulphuric acid what mass of sodium sulfate Na2SO4 would be made? Na2CO3 + H2SO4 Na2SO4 + CO2 + H2O 2. 200 tonnes of calcium carbonate is thermally decomposed to produce calcium oxide, CaO, and carbon dioxide, CO2. Write a balanced symbol equation and then calculate how many tonnes of calcium oxide can be made. CaCO3 CaO + CO2 3. Ammonium nitrate fertiliser is made in the following reaction: NH3 + HNO3 NH4NO3 How many tonnes of ammonia are needed to make 2400 tonnes of ammonium nitrate? NH4NO3 is made from NH3 4. How much carbon is needed to reduce 223g of lead oxide? 2PbO + C 2Pb + CO2 5. Calculate the mass of water produced when 1g of hydrogen burns completely in air. 2H2 + O2 2H2O 6. What mass of magnesium is needed to produce 20g of magnesium oxide? 2Mg + O2 2MgO
- 7. 7. Zinc oxide is heated with carbon to form zinc and carbon monoxide: zinc oxide + carbon zinc + carbon monoxide ZnO + C Zn + CO a) How much zinc oxide do you need to make 130g of zinc? b) How many moles of ZnO is used? c) What mass of carbon is needed to react with the zinc oxide? d) What mass of carbon monoxide will be released? e) The carbon monoxide reacts with oxygen in the air to make carbon dioxide: CO + O2 CO2 Using the information above, calculate what mass of carbon dioxide will be produced. There are 2 ways you could do this, see if you can find both methods.
- 8. 8. A student adds 4.8g of magnesium to excess dilute hydrochloric acid. Mg + 2HCl MgCl2 + H2 a) What mass of magnesium chloride will be made? b) How many moles of magnesium are being used? c) How many moles of hydrochloric acid are needed? (look at the balanced equation) d) How many moles of hydrogen will be produced?
- 9. Teacher Copy Empirical Formulae An empirical formula gives the simplest whole number ratio of each type of atom in a compound. molecular formula empirical formula e.g. C2H6 C2H4 C6H6 CH4 C6H12O6 Calculating the empirical formula of a compound 1. Write the symbols of all the elements in a line 2. Write the mass of % of each element under their symbol 3. Divide by the molar mass (from periodic table) 4. Divide all the numbers by the smallest one to get a ratio Example 1: Find the empirical formula of a compound that contains 13.2% lithium, 26.4% nitrogen and 60.4% oxygen 1. Write the symbols of all the elements in a line Li N O 2. Write the mass of % of each element under their symbol 3. Divide by the molar mass 4. Divide all the numbers by the smallest to get a ratio Empirical formula = Example 2: Analysis of a sample of sodium oxide found it contained 2.3g of sodium and 0.8g of oxygen. Find the empirical formula of sodium oxide 1. Write the symbols of all the elements in a line Na O 2. Write the mass of % of each element under their symbol 3. Divide by the molar mass 4. Divide all the numbers by the smallest to get a ratio Empirical formula =
- 10. Example 3: Analysis of a sample of mercury oxide found it contained 2.17g of mercury and 0.16g of oxygen. Find the empirical formula of mercury oxide 1. Write the symbols of all the elements in a line Hg O 2. Write the mass of % of each element under their symbol 3. Divide by the molar mass 4. Divide all the numbers by the smallest to get a ratio Empirical formula = Example 4: Find the empirical formula of a compound that contains 31.1% iron, 15.6% nitrogen and 53.3% oxygen 1. Write the symbols of all the elements in a line Fe N O 2. Write the mass of % of each element under their symbol 3. Divide by the molar mass 4. Divide all the numbers by the smallest to get a ratio Empirical formula = Extension: Can you suggest a name and formula for this compound? Example 5: Analysis of a sample of gas found it contained 2.4g of carbon and 0.8g of hydrogen. Find the empirical formula of this gas and name it. 1. Write the symbols of all the elements in a line C H 2. Write the mass of % of each element under their symbol 3. Divide by the molar mass 4. Divide all the numbers by the smallest to get a ratio Empirical formula = Name =
- 11. 1. 7.2g of carbon was reacted 1.2g of hydrogen and 9.6g of oxygen. What is the empirical formula of glucose? 2. The actual formula is C6H12O6. Calculate the mass of a molecule. 3. What percentage of glucose is carbon? 4. Write a balanced symbol equation for the complete combustion of glucose in air. 5. a) If 9g of glucose is burnt, what mass of oxygen would be needed for complete combustion to take place? Use your balanced equation to help you. b) What mass of water would be produced?
- 12. Experiment to Measure the Increase in Mass on Complete Oxidation of Magnesium Ribbon lid magnesium crucible ribbon pipeclay triangle heat heatproof mat 1 Complete the table of results: 1 1 mass of empty crucible + lid 2 2 mass of crucible + lid + magnesium 3 3 mass of crucible + lid + contents after heating 4 4 mass of magnesium at the start (2−1) 5 mass of magnesium oxide formed (3−1) 6 mass of oxygen in the oxide (5−4) 2 Explain why the mass of the crucible and its contents increased. 3 Explain why it was necessary to raise the crucible lid from time to time. 4 Explain why the crucible lid had to be put back quickly if the magnesium flared up. 5 Calculate the empirical formula for magnesium oxide. (H) 6 Comment on how easy it was to tell if the formula for magnesium oxide is Mg2O, MgO or MgO2. (H) 7 Look up the formula of magnesium oxide in a textbook or data book. Suggest why the amount (in moles) of oxygen combining with the magnesium is lower than expected from the formula. 8 What other metals, if any, could have their empirical formula found by this method? (H)
- 13. Experiment to Measure the Decrease in Mass on Reduction of Copper Oxide burning gas clamp here gas from gas tap black copper oxide tube must be kept heat horizontal 1 Complete the table of results. 1 mass of empty tube 2 mass of tube + copper oxide before heating 3 mass of tube + copper after heating 4 mass of copper oxide before heating (2−1) 5 mass of copper formed (3−1) 6 mass of oxygen in the oxide (4−5) 2 State the name the brownish pink solid formed at the end of the experiment. 3 Why did the mass of the copper oxide decrease? 4 What happened to the oxygen from the copper oxide? 5 Suggest why the gas from the gas tap is kept flowing through the tube after the Bunsen burner has been removed. (H) 6 Calculate the empirical formula for black copper oxide. (H)
- 14. 7 What type of chemical reaction is taking place? (H)

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