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Worksheet 1.1 Precision, accuracy and significant figures INTRODUCTION For a quantity to have an exact value, it must either be defined or obtained by counting. All measured quantities have an inherent uncertainty because all instruments used to make measurements have limitations, and the people operating the instruments have varying skills. The accuracy of a measurement is an expression of how close the measured value is to the ‘correct’ or ‘true’ value. The precision of a set of measurements refers to how closely the individual measurements agree with one another. Thus, precision is a measure of the reproducibility or consistency of a result. The precision of a measurement is sometimes expressed as an uncertainty using a plus/minus (±) notation to indicate the possible range of the last digit. An alternative method is to indicate the certainty of the measurement by the use of significant figures. To clarify the number of significant figures in a measurement, the value may be written in standard form. A number written in standard form is expressed as a number greater than 1 but less than 10 multiplied by 10x, where x is an integer. When a calculation involves multiplication and division, the result should have the same number of significant figures as the factor with the least number of significant figures. For addition and subtraction calculations, the result should have the same number of decimal places as the number used with the fewest decimal places. In most calculations you will need to round off numbers to obtain the correct number of significant figures. No. Question Answer 1 Which of the following quantities would have an inherent uncertainty? A The number of pages in this book B Your measured height (in cm) C The number of mL in 6.0 L D A volume of liquid measured using a pipette 2 Shooting at targets may be used as an analogy to show the ideas of precision and accuracy in measurements. Label each of the shooting targets shown as representing one of the following situations. N for neither accuracy nor precision B for both precision and accuracy P for precision, but inaccuracy Page 1 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 1.1 Precision, accuracy and significant figures No. Question Answer 3 Why are measurements in experiments often repeated several times and the results averaged? 4 State the number of significant figures in each of the following measured quantities. a A temperature reported as 26.1°C b A burette reading of 32.34 mL c A mass reading of 0.0471 g d A time recorded as 6.000 s 5 Express each of the following numbers in standard form, ensuring you use the correct number of significant figures. a 140.7 b 5005 c 980.0 d 0.0075 6 Round each of the following numbers to three significant figures, and express in standard form. a 7.8001 b 600.5 c 98.345 d 0.000600 7 Express the number 6000 in standard form to show that it contains: a 1 significant figure b 4 significant figures. 8 Calculate each of the following and express the answers to the correct number of significant figures. a 5.6 × 120 b 0.0045 × 67.1 c 0.046 ÷ 0.023 d 63 × 7.06 9 Perform the following calculations and round off the answers to the correct number of significant figures. a 3.256 + 45.2 – 3.815 b 12.13 + 342.0 + 4.108 10 A dozen eggs have a mass of 722 g. What is the average mass of the eggs? Page 2 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 1.2 Calculations involving gases and solutions NAME: CLASS: INTRODUCTION The following formulas are essential! You must know how and when to use them. Formula What it means m mass of sample (g) n= Amount (mol) = M molar mass (g mol−1 ) N number of particles n= Amount (mol) = NA Avogadros number (mol−1 ) For solutions n=c×V Amount (mol) = concentration (mol L–1) × volume (L) For gases at standard temperatures and pressures V volume of gas (L) n= Amount (mol) = VM molar volume (L mol−1 ) where molar volume, VM, is 22.4 L mol–1 at STP (0°C and 1.0 atm) 24.5 L mol–1 at SLC (25°C and 1.0 atm) For gases pV = nRT Pressure (kPa) × volume (L) and = amount (mol) × gas constant (8.31 J K–1 mol–1) × temperature (K) m pV = RT M For a fixed amount of a gas pV = k or pressure of gas × volume of gas T = a constant temperature of gas (K) p 2V2 p1V1 = T1 T2 Unit of concentration % m/m % m/v % v/v ppm ppb What it means Mass of solute (in g) in 100 g of solution Mass of solute (in g) in 100 mL of solution Volume of solute (in mL) in 100 mL of solution Mass of solute (in g) in 106 g of solution (equivalent to mg L–1 for dilute solutions) Mass of solute (in g) in 109 g of solution (equivalent to μg L–1 for dilute solutions) For dilution of a solution: c1 × V1 = c2 × V2 where c1 = initial concentration, V1 = initial volume, c2 = final concentration, V2 = final volume. Page 1 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 1.2 Calculations involving gases and solutions No. Question Answer 1 a What amount (in mol) of oxygen is present in 5.00 g of glucose, C6H12O6? b How many atoms are present in 8.36 g of hydrogen peroxide, H2O2? 2 What volume of water should be added to 35.0 mL of 0.30 M H2SO4 in order to produce a 0.090 M solution? 3 What pressure does 0.49 mol of SO3 exert in a sealed 3.0 L vessel at 54ºC? 4 The heaviest known atom has a mass of about 4 × 10–22 g. What would be the mass of one mole of these atoms? 5 2.5 g of a gas initially occupying a volume of 600 mL, at 260 K, is heated to 325 K at constant pressure. What would its new volume be? 6 What volume will 62.0 g of carbon dioxide gas occupy at a temperature of 124°C and 210 kPa? Page 2 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 1.2 Calculations involving gases and solutions No. Question Answer 7 What volume of water must be added to 1.0 L of a solution containing 70.2 g of NaCl to produce a solution of 0.67 M NaCl? 8 2.06 g of a hydrocarbon occupies 16 L at 27ºC and 20 kPa. Find the molar mass of this compound, and so identify the hydrocarbon. 9 a What volume of 5.00% m/v cloudy ammonia cleaning solution is needed to make 250 mL of a 1.50% m/v solution? b What mass of ammonia is present in 150 mL of the 1.50% solution? 10 3.0 g of carbon dioxide occupies 687 mL at 143 800 Pa. What volume does it occupy at a pressure of 199 kPa, assuming temperature is constant? 11 0.778 g of one of the halogens (Group 17) was found to occupy a volume of 122 mL at a pressure of 99.8 kPa and a temperature of 26°C. Which halogen was it? 12 A sample of water from a waterway is found to contain 600 ppm mercury. What is this concentration in ppb? Page 3 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 1.2 Calculations involving gases and solutions No. Question Answer 13 3.5 g of Pb(NO3)2 is added to 60 mL of distilled water in a beaker and stirred to dissolve the solid. 10 mL of this solution is then transferred to another beaker and mixed with 20 mL of distilled water. What are the concentrations (in mol L–1) of these two solutions? 14 30 mL of a sodium carbonate solution is made up to a total volume of 300 mL with distilled water. The resultant solution has a sodium carbonate concentration of 0.108 M. What mass of sodium carbonate was present in the original solution? 15 50 mL of a 2.0% m/v glucose solution is mixed with 50 mL of a 6.0% m/v glucose solution. The solution is made up to a total of 300 mL with distilled water. What is the concentration of the final solution? 16 A sample of N2 gas collected at 25°C and 750 mmHg pressure occupies 190 mL. What volume will it occupy at STP? 17 50 mL of a 16% m/v silver nitrate solution is added to an equal volume of distilled water. What is the concentration of the dilute solution in ppm? 18 A solution of silver nitrate (AgNO3) is made by dissolving 2.33 g of solid in 398 mL of distilled water. What is the concentration of this solution in a % m/v? b M? Page 4 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 1.3 Stoichiometry NAME: CLASS: INTRODUCTION Stoichiometric problems can be tackled by remembering five basic steps. Step 1: Write a balanced chemical equation for the reaction. Step 2: List all data given, including relevant units. Remember to also write down the symbol of the unknown quantity. Step 3: Convert the data given to moles, using the relevant formulas: N pV V m n= n=c×V n= n= n= M NA RT VM Step 4: Use the chemical equation to determine the mole ratio of the unknown quantity to the known quantity. This ratio enables calculation of the number of moles of the unknown quantity. Step 5: Finally, convert this number of mole back into the relevant units of the unknown. In all calculation questions, take care to show all steps in your working, including reacting ratios where relevant. Also, take care to add the correct unit to your answer (e.g. mol, L) and give your answer to the correct number of significant figures. No. Question Answer 1 SO2 in the atmosphere contributes to acid rain. The equation for formation of the acid is represented by the equation: 2SO2(g) + O2(g) + 2H2O(l) → H2SO4(aq) What mass of sulfuric acid will form from 50.0 L of sulfur dioxide at SLC? 2 Phosphoric acid can be generated by the oxidation of phosphorus with nitric acid according to the following equation: P(s) + 5HNO3(aq) → H3PO4(aq) + H2O(l) + 5NO2(g) If sufficient reactants are available to produce 1.00 kg of phosphoric acid (H3PO4), what mass of nitrogen dioxide will also be generated in the reaction? Page 1 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 1.3 Stoichiometry No. Question Answer 3 23.8 mL of sulfuric acid is just neutralised by 29.9 mL of 2.86 M sodium hydrogen carbonate solution, according to the equation: 2NaHCO3(aq) + H2SO4(aq) → Na2CO3(aq) + CO2(g) + H2O(l) Determine the concentration of the sulfuric acid solution used. 4 In a car accident, the impact triggers ignition of a detonator cap in the air bag, which causes sodium azide (NaN3) to decompose explosively. a Write the equation to show the decomposition of sodium azide into solid sodium and nitrogen gas. b If the bag contained 75 g of sodium azide, what volume of gas would form at 30ºC and 101.3 kPa? 5 Sodium thiosulfate reacts with bromine in alkaline solution according to the equation: Na2S2O3(aq) + 4Br2(l) + 10NaOH(aq) → 2Na2SO4(aq) + 8NaBr(aq) + 5H2O(l) In order to completely react with 10.0 mL of bromine of density 3.12 g mL–1, 239 mL of NaOH was added along with excess Na2S2O3. Determine the concentration of the sodium hydroxide solution required. 6 Ammonium sulfate, an important fertiliser, can be prepared by the reaction of ammonia with sulfuric acid according to the equation: 2NH3(g) + H2SO4(l) → (NH4)2SO4(aq) Calculate the volume of NH3 needed to react with 19.56 g of H2SO4 at 87°C and 2.99 atm. Page 2 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 1.3 Stoichiometry No. Question Answer 7 When magnesium carbonate is heated strongly, it decomposes to magnesium oxide and carbon dioxide gas. What volume of carbon dioxide would be produced at STP when 100 g of magnesium oxide was generated in this reaction? 8 When solid calcium carbonate reacts with nitric acid solution, neutralisation takes place. a Write the equation for this reaction. b If 10.0 g of calcium carbonate reacts with 100 mL of 0.500 M nitric acid, what volume of carbon dioxide is formed at SLC? 9 Arsenic undergoes oxidation by a hot, concentrated solution of sodium hydroxide to produce sodium arsenate and hydrogen gas according to the equation: 2As(s) + 6NaOH(aq) → 2Na3AsO3(s) + 3H2(g) 6.57 g of arsenic is reacted with 250 mL of 0.779 M sodium hydroxide solution. Calculate the mass of hydrogen gas evolved in the process. 10 8.00 g of barium hydroxide is dissolved in 120 mL of 1.886 M hydrochloric acid to produce barium chloride solution and water. a Write an equation to represent this neutralisation reaction b Determine the concentration of the barium chloride solution that results. Page 3 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 2.1 Determining the molecular formula of a gas NAME: CLASS: INTRODUCTION Empirical formula: The formula showing the smallest whole-number ratio of atoms in a substance. Molecular formula: The formula showing the actual number of atoms of each element present in a molecule of a compound. For example, the molecular formula of glucose is C6H12O6 and its empirical formula is CH2O. No. Question 1 A compound used in ceramics contains, by mass, 22.8% Na, 21.5% B and 55.7% O. What is its empirical formula? 2 A sweet-smelling organic compound contains 0.0556 mol of carbon, 0.112 g of hydrogen atoms and 9.57 × 1021 oxygen atoms. Calculate its empirical formula. 3 If a compound contains 75.7% arsenic and 24.3% oxygen by mass, and has a molar mass of 395.6 g mol–1, what is its molecular formula? Answer Page 1 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 2.1 Determining the molecular formula of a gas No. Question 4 A compound has the following composition by mass: carbon 40.0%, hydrogen 6.67% and oxygen 53.3%. If the compound has a molecular mass of approximately 60, determine both its empirical and molecular formulas. Answer The apparatus shown below may be used to determine the empirical and molecular formulas of an unknown compound of general formula CxHyOz. The questions that follow refer to this apparatus. No. Question 5 What is the purpose of the CaCl2(s)? 6 What is the purpose of the NaOH(aq)? Answer Page 2 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 2.1 Determining the molecular formula of a gas No. Question 7 14.34 g of a hydrocarbon was burned in excess oxygen. The mass of the NaOH(aq) increased by 43.54 g, while the mass of CaCl2(s) increased by 22.26 g. What is the empirical formula of the hydrocarbon? Answer A method used to determine the relative molecular mass of a volatile liquid is shown below. Steam is passed through the outer jacket surrounding a graduated syringe that contains a small, measured volume of air. When the temperature has stabilised, a weighed sample of a few mL of liquid is injected into the graduated syringe using a small hypodermic syringe. The liquid vaporises and the final volume of air plus sample is recorded. The questions which follow refer to this apparatus. No. Question 8 Why would this method be unsuitable for a liquid with a boiling point above 90°C? 9 Using the apparatus shown, 0.16 g of liquid was vaporised to produce a volume of 46 mL, at a temperature of 100°C and a pressure of 1.0 atm. Determine the molar mass of the liquid. Answer Page 3 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 2.1 Determining the molecular formula of a gas No. Question 10 In one experiment an 18.99 g sample of the compound was burnt in excess oxygen. When the gases evolved were passed through anhydrous CaCl2, its mass increased by 11.38 g. The remaining gases, when bubbled through a NaOH solution, increased its mass by 27.83 g. In a separate experiment, a 6.21 g sample of the compound was vaporised. The vapour occupied 2.17 L at 200°C and 1.25 × 105 Pa. Calculate the molecular formula of the compound. 11 Which instrumental method is routinely used to determine the molar mass of an unknown organic compound? Answer Page 4 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 2.2 Gravimetric analysis of a fertiliser NAME: CLASS: INTRODUCTION The purity of a sample of ammonium sulfate fertiliser was determined by gravimetric analysis in which the sulfate was precipitated as barium sulfate. The following steps were used in this procedure: 1 A sample of about 5 g of the fertiliser was dried in an oven for an hour at 110ºC. 2 The dried fertiliser was ground into fine powder using a mortar and pestle. 3 1.00 g of the dried, ground fertiliser was weighed out into a 250 mL beaker. 4 150 mL of distilled water and 2 mL of 10 M HCl were added to the fertiliser with stirring. 5 The solution was heated to nearly boiling. 6 The fertiliser solution was maintained at that temperature while the precipitating agent was prepared. If the fertiliser was pure, it was calculated that 37.9 mL of the 0.200 M BaCl2 stock solution would be required. 45 mL of the stock solution was placed in a 100 mL beaker and heated to nearly boiling. 7 The fertiliser solution was stirred vigorously as the hot barium chloride solution was added slowly. 8 Heating of the beaker continued while the precipitate was allowed to settle. A few drops of the supernatant liquid were tested for complete precipitation by adding them to a few drops of barium chloride solution on a watchglass. No cloudiness was observed. 9 The fertiliser–barium chloride mixture was covered and left to ‘digest’ just below boiling point for one hour. (This digestion process is necessary for experiments involving barium sulfate because the crystals formed are initially too small to filter. Digestion allows larger crystals to form.) 10 The hot solution was filtered under vacuum filtration into a pre-weighed sintered glass crucible. 11 The precipitate was washed several times with hot water. 12 The crucible and precipitate were dried to constant mass in an oven at 60ºC over two successive days; 1.63 g of precipitate was obtained. No. Question 1 Why was it necessary to dry and grind the fertiliser before weighing? 2 If the fertiliser was 100% ammonium sulfate, what amount (in mol) of ammonium sulfate would be present in 1.00 g? Answer Page 1 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 2.2 Gravimetric analysis of a fertiliser No. Question 3 This procedure assumes that there were no insoluble components in the fertiliser. If there had been, what extra step would have been necessary? 4 Step 6 states that 37.9 mL of barium chloride was needed. Show how this volume was calculated. 5 If 37.9 mL was the maximum volume needed, why was 45 mL used? 6 How was it shown that complete precipitation had taken place? 7 Why was the precipitate washed in hot water? 8 Define the term ‘weighed to constant mass’. 9 Use the final mass of precipitate to calculate the % m/m of ammonium sulfate in the fertiliser. 10 Where could errors have occurred in this experiment? Answer Page 2 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 2.3 Gravimetric analysis problems NAME: CLASS: INTRODUCTION Analysis by precipitation sounds very simple, but, in practice, obtaining accurate results by this method requires planning and careful experimentation. The method involves considering such variables as the degree of solubility, the size of the precipitate particles formed and the possibility that other ions present might interfere. Gravimetric analysis involves precipitating one of the ions of interest and weighing the precipitate. The amount of precipitate tells us the amount of the ion of interest that must have been present in the original solution. Sometimes the positive ion is precipitated, sometimes the negative ion. No. Question 1 Gravimetric analysis cannot be used for solutions of sodium nitrate. Why is this? 2 Silver nitrate is added to a solution containing a mixture of sodium chloride and potassium chloride. a Write a balanced equation for the reaction that occurs between: i silver nitrate and sodium chloride ii silver nitrate and potassium chloride. b Explain why this procedure can be used for an analysis of the chloride ion concentration, but not the potassium ion concentration. 3 Most precipitates are not 100% insoluble. What is the main precaution taken during a gravimetric procedure to ensure that this does not affect the accuracy too much? Answer Page 1 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 2.3 Gravimetric analysis problems No. Question 4 Use a solubility table to suggest a compound that could be added to each of the following solutions to produce a precipitate. a Na2CO3 b MgI2 c NaOH 5 The sulfate content of fertiliser can be found by adding barium nitrate, Ba(NO3)2, to precipitate the sulfate as barium sulfate, BaSO4. a Write an ionic equation for the reaction between barium nitrate and sulfate ions. b A student suggests that the sulfate content could also be found by boiling the water from a solution of the fertiliser and weighing the solid left behind. Why does this procedure not work? 6 When barium sulfate is collected as a precipitate, care must be taken to use a filter paper that has very fine pores. Explain why this is needed. 7 a Write a balanced equation for the reaction between silver nitrate and lithium chloride solutions. b A solution contains 0.02 mole of lithium chloride. A student wishing to precipitate this lithium chloride adds 25 mL of 2 M silver nitrate solution. Has the student added an excess of silver nitrate? Answer Page 2 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 2.3 Gravimetric analysis problems No. Question 8 An organic compound has a formula C8H15Cl3. The purity of a sample of this compound is to be checked by the addition of silver nitrate to precipitate the chlorine atoms as silver chloride. It is difficult to write a complete, balanced equation for this reaction but fortunately only a partial equation is required. a Write a partial equation for this process. b A 2.0 g sample leads to 0.478 g of precipitate. Calculate the % purity by mass of the organic compound in the sample. 9 A precipitate of Fe2O3 is obtained from an iron(III) chloride, FeCl3, solution. The mass of precipitate obtained is 0.644 g. a Write a partial equation for this reaction. b If the mass of sample used was 2.0 g, determine the percentage by mass of iron(III) chloride in the sample. 10 As households use more grey water on their gardens, the phosphorus content of the water becomes more relevant. Phosphorus can be precipitated as Mg2P2O7. During an analysis of a water sample, the Mg2P2O7 precipitate is found to weigh 0.744 g. What will be the effect on the calculated phosphorus content if: a the precipitate is not completely dry? b some of the precipitate passed through the filter paper? Answer Page 3 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 3.1 Revision of acids and bases NAME: CLASS: INTRODUCTION These reactions are typical of acids: acid + metal → salt + hydrogen gas (a redox reaction) acid + metal hydroxide → salt + water acid + metal oxide → salt + water acid + metal carbonate → salt + water + carbon dioxide Acids are proton donors; bases are proton acceptors. Acid–base reactions involve the transfer of protons. Acidic species may be described as monoprotic (donate one proton), diprotic (donate two protons), triprotic (donate three protons) or amphiprotic (both donate and accept protons). Acid–base strength is a measure of how readily protons are donated or accepted. The pH scale is used as a measure of the acidity or basicity of a solution. The scale is usually applied over the range 0 to 14 (but does extend beyond these values). pH = –log10[H3O+]. pH is a logarithmic scale, so a difference of one unit on the pH scale means a 10-fold difference in the hydrogen ion concentration. For dilute solutions at 25°C, Kw = [H3O+] × [OH–] = 10–14. No. Question 1 Write balanced chemical equations for the reactions occurring when the following chemicals are mixed. a Solid potassium hydrogen carbonate and sulfuric acid b Iron(III) oxide and nitric acid c Calcium hydroxide solution and hydrochloric acid 2 Write balanced equations to illustrate the following reactions. a Dissociation of Ca(OH)2 in aqueous solution b Successive ionisations of H2SO4 c Neutralisation of Ba(OH)2 solution with HCl solution Answer Page 1 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 3.1 Revision of acids and bases No. Question 3 Write the formulas of: a the conjugate bases of HS– and OH– b the conjugate acids of HSO4– and H2PO4–. 4 For each of the following identify the Lowry–Brønsted conjugate acid–base pairs involved in the reaction. a NH3(g) + HCl(g) → NH4Cl(s) b K2O(s) + H2O(l) → 2KOH(aq) c Na2CO3(s) + H2SO4(aq) → Na2SO4(aq) + H2O(l) + CO2(g) 5 HSO4– is an amphiprotic ion. Write chemical equations to show this ion acting as: a an acid b a base. 6 Give two reasons why two acid solutions of equal concentration could have different pH values. 7 Give an explanation for each of the following observations. a The electrical conductivity of a 1.0 M solution of methanoic acid (HCOOH) is less than that of a 1.0 M solution of hydrochloric acid (HCl). b Ethanoic acid (CH3COOH) is monoprotic, even though it contains four hydrogen atoms. 8 List the following 1.0 M solutions in order of decreasing pH. Give reasons for your order. NaOH, H2O, NH3, CH3COOH, H2SO4, HNO3 Answer Page 2 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 3.1 Revision of acids and bases No. Question 9 Calculate the pH of each solution. a 0.050 M HNO3 b 0.50 M Ba(OH)2 10 20.0 mL of a solution of pH 3.0 is diluted to produce a total volume of 200.0 mL. What is the pH of the resulting solution? 11 What volume of water must be added to 50.0 mL of a hydrochloric acid solution of pH 2.0 to increase the pH to 2.5? 12 Calculate the pH of a solution formed when 20.00 mL of 0.00100 M HCl is mixed with 20.00 mL of 0.00100 M Ba(OH)2. Answer Page 3 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 3.2 Acid–base titrations NAME: CLASS: INTRODUCTION Probably the most used analytical procedure of old is the acid–base titration. The calculations involved use the same stoichiometry procedures encountered earlier in your studies. There are certain practical steps necessary, however, to ensure accurate results. These involve the choice of indicators, choice of liquids for rinsing glassware and the correct preparation of standard solutions. No. Question 1 Explain the steps required to prepare 250 mL of a standard solution of 0.0500 M sodium carbonate using anhydrous Na2CO3. 2 20 mL of 0.10 M nitric acid, HNO3, sits in a flask. Several drops of indicator are added to the flask. When sodium hydroxide solution, NaOH, of an unknown concentration is added dropwise into the flask, there is a colour change after 5 drops of sodium hydroxide has been added. What conclusion can you draw about the concentration of the sodium hydroxide? 3 20 mL of 0.1 M nitric acid sits in a flask. Indicator is added. When sodium hydroxide is added, a colour change occurs after exactly 5 mL of sodium hydroxide has been added. Without calculating any mole quantities, what must the concentration of the sodium hydroxide be? Can you use this ratio technique to solve all titration calculations? Answer Page 1 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 3.2 Acid–base titrations No. Question 4 a Write a balanced equation for the reaction between solutions of sodium carbonate, Na2CO3, and nitric acid. b What volume of 0.20 M sodium carbonate is required to neutralise 40 mL of 0.10 M nitric acid? 5 a Write a balanced equation for the reaction between solutions of ammonia, NH3, and hydrochloric acid, HCl. b 0.1 M ammonia in a flask is to be titrated with 0.1 M hydrochloric acid. Ammonia is a weak base. The approximate pH of the ammonia solution before the titration begins is 11. i What do you think the pH of the solution in the flask will be at the equivalence point? (Look at the products of the reaction.) ii After significant extra hydrochloric acid has been added, what will the approximate pH be? iii Sketch the pH curve for this titration. 6 A 25.0 mL sample of ethanoic acid is diluted to 100.0 mL. A 20.00 mL aliquot is then titrated with 0.114 M sodium hydroxide. The titre required is 15.45 mL. Calculate the concentration of the original ethanoic acid solution. Answer Page 2 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 3.2 Acid–base titrations No. Question 7 Suggest two reasons why an indirect (back) titration might be performed, rather than a direct titration. 8 a Draw a pH curve for the addition of 0.1 M ethanoic acid, CH3COOH, to 0.1 M sodium hydroxide, NaOH. b What will be the approximate pH at the equivalence point? c Use your curve to explain why an indicator such as methyl red would be a poor choice for this titration. 9 10.0 mL of hydrochloric acid is added to a 250 mL volumetric flask. The flask is made up to the mark and 20.0 mL aliquots of this solution are added to conical flasks for titration with 0.100 M sodium carbonate from a burette. a State which liquid should be used to rinse each of the following items. i Volumetric flask ii Pipette iii Burette iv Conical flasks b The average titre of sodium carbonate is 18.3 mL. Calculate the concentration of the original HCl used. Answer Page 3 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 3.2 Acid–base titrations No. Question 10 A 0.10 M magnesium hydroxide, Mg(OH)2, solution is used to find the concentration of a solution of nitric acid. An aliquot of 25.0 mL of magnesium hydroxide is added to a flask. The volume of nitric acid required to neutralise it is 13.5 mL. Calculate the concentration of the nitric acid solution. Answer Page 4 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 3.3 A back titration NAME: CLASS: INTRODUCTION Marble is a metamorphic rock that is almost pure calcium carbonate. The following experiment was conducted to determine the calcium carbonate content of a marble sample. An accurately weighed sample of crushed marble was added to a measured volume of recently standardised HCl solution. The solution was heated to drive off the evolved carbon dioxide. The remaining solution was titrated with a recently standardised NaOH solution, using a methyl red indicator. The results obtained are shown below. Mass of marble sample: 1.740 g Volume of HCl solution added: 40.00 mL Concentration of the standardised HCl solution: 1.020 M Concentration of the standardised NaOH solution: 0.275 M Average titre of NaOH: 25.56 mL No. Question 1 Write an equation for the reaction of HCl with: a CaCO3(s) b NaOH(aq) 2 Calculate: a the amount (in mol) of HCl added initially b the amount (in mol) of NaOH used in the titration c the amount (in mol) of unreacted HCl d the amount (in mol) of HCl reacting with the CaCO3 e the amount (in mol) of CaCO3 in the marble sample f the mass (in g) of CaCO3 in the marble sample g the percentage by mass of CaCO3 in the marble sample. Answer Page 1 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 3.3 A back titration No. Question 3 The sodium hydroxide solution used was recently standardised. Why is sodium hydroxide unsuitable as a primary standard? 4 Why was it necessary to drive off the evolved carbon dioxide before performing the titration? How would the result be affected if this step was omitted? 5 Why was it necessary to use a back titration for this analysis, rather than a direct titration of the marble with HCl solution? 6 How would each of the following errors, if made during the analysis, alter the calculated value for the percentage CaCO3? a The 40.0 mL pipette used to deliver the HCl was rinsed only with water prior to its use. b The burette was rinsed only with water prior to its use. c The volumetric flask was rinsed only with water prior to its use. 7 An alternative method of analysis involves reacting the crushed marble with excess HCl and collecting the evolved carbon dioxide. In one such experiment, 95.0 mL of gas was collected at a pressure of 765 mmHg at 23°C when 0.411 g of marble was reacted. Determine the percentage by mass of CaCO3 in the marble sample, based on this data. 8 Suggest reasons why the value obtained using this method is smaller than that obtained using the back titration. Answer Page 2 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 4.1 Oxidation numbers and redox equations NAME: CLASS: INTRODUCTION The concept of oxidation numbers (ON) or oxidation states was applied to determine whether or not electrons had moved from one species to another in a chemical reaction. An oxidation–reduction (redox) reaction is one in which one or more atoms change oxidation numbers. Oxidation occurs when an atom’s oxidation state becomes more positive, indicating that electrons have been lost. Reduction occurs when an atom’s oxidation state becomes less positive, indicating that electrons have been gained. OXIDATION …–5 –4 –3 –2 –1 0 +1 +2 +3 + 4 +5…. REDUCTION The oxidation numbers assigned to atoms in covalent compounds are hypothetical charges, and the atoms do not really have these charges within the compound, since they are only sharing electrons. For example, in the compound CO2 we say that carbon has a +4 ON and oxygen is in a –2 ON, but the atoms do not really have these charges as electrons are not transferred in covalent compounds. Recall that the oxidant is the reactant being reduced, so its ON will decrease. The reductant is the reactant being oxidised, so its ON will increase. Rules for assigning oxidation numbers 1 The oxidation number of an element is zero. 2 For a monatomic ion, the oxidation number is the charge on the ion. 3 The oxidation number of combined hydrogen is usually +1. 4 The oxidation number of combined oxygen is usually –2. 5 The sum of all oxidation numbers of atoms in a compound is zero. 6 The sum of all oxidation numbers of atoms in an ion is equal to the charge on that ion. No. Question 1 Write the formula of a substance in which nitrogen has the following oxidation numbers: a +3 b +5 c 0 d –3 Answer Page 1 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 4.1 Oxidation numbers and redox equations No. Question 2 What is the oxidation number of the atom in bold type in each of the following? a H2SO4 b K2Cr2O7 c CO2 d H2O2 3 For the following equations, determine which are redox processes. For those that are redox reactions, identify the oxidant and the reductant. a 2OH–(aq) + Cr2O72– (aq) → 2CrO42–(aq) + H2O(l) b I2O5(s) + 5CO(g) → I2(s) + 5CO2(g) c PBr3(l) + 3H2O(l) →H3PO3(aq) + 3HBr(aq) d 2Hg2+(aq) + N2H4(aq) → 2Hg(l) + N2(g) + 4H+(aq) e 3H2S(g) + 2H+(aq) + 2NO3–(aq) → 3S(s) + 2NO(g) + 4H2O(l) f 3NO2(g) + H2O(l) → 2HNO3(aq) + NO(g) 4 The highest positive oxidation number possible for any atom is equal to the number of electrons in its valence shell. For example, nitrogen has a maximum oxidation number of +5. What is the maximum oxidation number of: a oxygen? b magnesium? c chlorine? 5 Occasionally you will find that an atom in a compound has an oxidation number of zero. What is the oxidation number of each atom in glucose, C6H12O6? Answer Page 2 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 4.1 Oxidation numbers and redox equations Rules for balancing redox half-equations in acidic media 1 Write the species undergoing oxidation or reduction on the left and its conjugate on the right, leaving out any spectator ions. 2 Balance for atoms other than hydrogen and oxygen. 3 Balance for oxygen by adding water, H2O. 4 Balance for hydrogen ions, H+. 5 Balance for charge by adding electrons to the side with the greater positive charge. No. Question 6 For each of the following, write the two half-equations, then add them together to get the overall redox equation. Make sure you leave out any spectator ions. a Acidified potassium dichromate solution (K2Cr2O7) is used to oxidise ethanol to ethanoic acid. The dichromate ion is converted to chromium(III) ions. b The bromate ion, BrO3– may be used in acidic solution to oxidise iodide ions to iodine. The reduction product is the bromide ion. c Purple acidified potassium permanganate solution (KMnO4) is decolourised by acidified iron(II) sulfate solution, producing Mn2+ and Fe3+. d Hydrogen peroxide (H2O2) is added to acidified potassium permanganate and oxygen is evolved. e Hydrogen sulfide (H2S) gas is bubbled into a solution of acidified potassium dichromate, producing a deposit of sulfur. Answer Page 3 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 4.2 Recovering silver from solution NAME: CLASS: INTRODUCTION Given the high price of silver, it is desirable to have an efficient method for the recovery of ‘waste’ silver from aqueous solutions containing Ag+ ions. This recovery can be achieved by using a redox reaction in which Ag+(aq) is reduced to Ag(s). This worksheet concerns the analysis and extraction of silver from a 0.1 M solution, and involves both gravimetric and volumetric techniques. PART 1: VOLUMETRIC DETERMINATION OF SILVER ION CONCENTRATION A 20.00 mL aliquot of 0.100 M NaCl solution was pipetted into a 250 mL conical flask. Approximately 1 mL of 0.1 M K2CrO4 solution was also added to the flask. The flask contents were then titrated with a solution of approximately 0.1 M AgNO3 until the first permanent redbrown colour appeared. This red-brown colour indicated that the reaction between Ag+(aq) and Cl– (aq) was complete (the CrO4– ion reacts with the Ag+ ion to form a red-brown precipitate. This reaction occurs only when there is no Cl– ion present). The titration was completed to obtain concordant titres; an average titre of 19.32 mL was required. No. Question 1 Name the liquid that should be used for the final rinse of each of the following pieces of equipment prior to their use in the titration. a Burette b Pipette c Conical flask 2 Write an ionic equation for the precipitation reaction between: a Ag+ and Cl– ions b Ag+ and CrO4– ions. 3 Using the titration results, determine the concentration of the approximately 0.1 M AgNO3 solution. 4 Suggest two sources of error in this determination. Answer Page 1 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 4.2 Recovering silver from solution PART 2: RECOVERY OF SILVER FROM SOLUTION A 20.00 mL aliquot of the approximately 0.1 M AgNO3 solution was pipetted into a large test tube. A spiral of copper wire was inserted into the test tube so that most of the spiral was in the solution. The solution was stirred occasionally using the copper spiral. The test tube and spiral were allowed to stand overnight. The spiral was removed after shaking to dislodge deposited silver, and any remaining silver washed into the test tube with distilled water. The deposited silver was collected, dried and weighed. The mass of silver recovered was 0.206 g. No. Question Answer 5 How could you ensure that the recovered silver was dry? 6 a Write an ionic equation for the reaction between Cu and Ag+ ions. b Name the oxidant in this reaction. 7 Determine the percentage of silver recovered from the solution. 8 Suggest two reasons to account for the less than 100% recovery of silver from the solution. 9 Most methods for the recovery of silver from solution involve the reduction of Ag+(aq) to Ag(s). Another reductant used is the dithionite ion, S2O42–, which is oxidised to the sulfite ion, SO32–. Write a half-equation for the oxidation of the dithionite ion. 10 In an alternative process, an electric current is forced through a solution containing silver ions. Silver deposits at the negative electrode. At the positive electrode, water is oxidised to oxygen gas and hydrogen ions. Write a half-equation for this oxidation. Page 2 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 4.3 Redox titrations NAME: CLASS: INTRODUCTION Titrations are not restricted to acid–base reactions. They can also be applied to redox reactions. The principles of rinsing and careful searching for an endpoint apply to redox titrations, but instead of acid–base indicators, the colour change that occurs when elements change oxidation state is often used. No. Question Answer 1 The presence of iodine, I2, can be detected by the use of starch as an indicator. The starch forms a blue colour when iodine is present. a Write a balanced half-equation for the reduction of iodine to iodide ions. b Explain how the starch acts as an indicator in reactions involving iodine. 2 The half-equation for the reaction of vitamin C, ascorbic acid, is shown. C6H4O2(OH)4(aq) → C6H4O4(OH)2(aq) + 2e– + 2H+(aq) a Write a balanced overall equation for the reaction between ascorbic acid and iodine. b 25.0 mL of 0.10 M iodine reacts with 17.8 mL of ascorbic acid. Calculate the concentration of the ascorbic acid. 3 The now obsolete puff bag used the reaction of potassium dichromate, K2Cr2O7, and ethanol. a Write balanced half-equations for the oxidation and reduction reactions. b Write a balanced overall equation for this reaction. c A 25.0 mL sample of whiskey is diluted to 250.0 mL and 20.0 mL aliquots of the diluted whiskey are titrated against 0.200 M potassium dichromate. The average titre required is 19.45 mL. Calculate the ethanol concentration (in M) of the whiskey. Page 1 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 4.3 Redox titrations No. Question Answer 4 The concentration of an iron solution, Fe2+, can be found by titration with permanganate, MnO4–, ions. The Fe2+ is oxidised to Fe3+ and the purple permanganate is reduced to colourless Mn2+ in acid conditions. a Write balanced half-equations for the oxidation and reduction reactions. b Write an overall equation for this reaction. c How would you judge that the endpoint has been reached in this reaction? d 20.0 mL aliquots of an iron solution are titrated with 0.110 M KMnO4. The average titre required is 12.8 mL. Calculate the concentration of the iron solution. 5 A rusty nail with a mass of 2.87 g is added to a beaker containing 100.0 mL of 1.08 M hydrochloric acid. After the reaction stops, the excess acid is neutralised by the addition of 1.00 M sodium hydroxide. The volume of sodium hydroxide required is 45.9 mL. a Write a balanced equation for the reaction between hydrochloric acid and iron. b Write balanced half-equations to show that this is a redox reaction. c Calculate the % by mass of iron in the nail. Page 2 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 5.1 Analysis of iron in iron ore NAME: CLASS: INTRODUCTION Iron ores often contain a mixture of oxides and contain both Fe2+ and Fe3+ ions. The iron content of these ores may be determined by a variety of analytical techniques. This worksheet concerns three of these techniques. PART 1: SPECTROSCOPIC ANALYSIS The iron content of the ore sample was analysed using UV-visible spectrophotometry. The iron(III) ion, Fe3+, reacts with the thiocyanate ion, SCN–, to form a complex ion with an intense red colour. This complex ion may be detected in a spectrophotometer set at a wavelength of about 580 nm. A 0.100 g sample of the ore was dissolved in concentrated hydrochloric acid. The extract was filtered, treated to ensure all the iron present was converted to Fe3+, then 20.0 mL of potassium thiocyanate solution was added. The volume was then made up to 100.0 mL with deionised water. Four standard solutions of iron(III) were similarly treated and their absorbances measured to generate the calibration curve shown below. No. Question 1 How would it have been determined that 580 nm was an appropriate wavelength for this analysis? 2 Why was it necessary to construct a calibration curve for this determination? Answer Page 1 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 5.1 Analysis of iron in iron ore No. Question 3 The absorbance of the iron ore solution was found to be 0.400. Calculate the percentage by mass of iron in the ore sample. Answer PART 2: VOLUMETRIC ANALYSIS In a second experiment the iron content of the ore was determined volumetrically. A 0.268 g sample of ore was dissolved in acidic solution, filtered, and the filtrate treated to convert all the iron present to Fe2+. This solution was titrated with a standardised, acidified 0.0335 M potassium permanganate (KMnO4) solution. The titration required 19.75 mL of the permanganate solution to reach the light pink-purple endpoint. No. Question 4 Write half-equations and a balanced redox equation for the titration reaction, given that the products of the titration reaction include Fe2+(aq) and Mn2+(aq). 5 Calculate the percentage by mass of iron in the ore sample, based on this volumetric analysis. 6 If some of the iron was present in the solution as Fe3+ prior to the titration with permanganate solution, how would this have affected the value determined for percentage of iron in the ore? Answer Page 2 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 5.1 Analysis of iron in iron ore PART 3: GRAVIMETRIC ANALYSIS In a third experiment the iron content of the ore was determined gravimetrically. A 1.01 g sample of the ore was dissolved in concentrated hydrochloric acid. The extract was filtered to remove any insoluble material. Excess sodium hydroxide solution was added, and the precipitate collected and heated to convert it to solid iron(III) oxide. 1.08 g of Fe2O3 was obtained. No. Question 7 Write an ionic equation for the precipitation reaction between the following: a Iron(II) ion and hydroxide ion b Iron(III) ion and hydroxide ion 8 Suggest why the iron hydroxide precipitate was not simply collected, dried and weighed in this determination. 9 Calculate the percentage by mass of iron in the ore sample, based on this gravimetric analysis. 10 Suggest a possible reason why the value obtained using gravimetric analysis is significantly larger than the other determined values. Answer Page 3 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 5.1 Analysis of iron in iron ore CONCLUSION No. Question 11 Which method of analysis used (spectroscopic, volumetric or gravimetric): a is the most expensive? b is the most prone to error? c requires the least specialised equipment to perform? d is likely to be the most accurate? 12 Suggest another method of analysis of the iron content of the iron ore. Answer Page 4 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 5.2 Colorimetric determination of manganese in steel NAME: CLASS: INTRODUCTION A 2.890 g sample of steel was analysed for its manganese content in the following manner. The sample was dissolved in concentrated nitric acid, producing manganese(II) ions. Potassium persulfate, K2S2O8, was added to remove the carbon in the steel. This oxidised the carbon to carbon dioxide. The virtually colourless manganese(II) ions were converted to deep purple permanganate (MnO4–) ions by boiling with a solution containing periodate (IO4–) ions. In this reaction the periodate ions are converted to IO3– ions. Phosphoric acid was added in order to convert the yellow iron(III) ions to a colourless complex (so the iron ions would not interfere with the colorimetric MnO4– ion analysis). The solution was diluted with distilled water to 1.00 L in a volumetric flask. A series of standards was prepared, containing MnO4– ranging from 1.00 to 8.00 ppm. The colorimeter was set to measure absorbance at 520 nm, and calibrated using a solvent blank. The absorbances of the standard solutions and the steel sample were taken and are given in the table below. Concentration of MnO4– (ppm) 0.00 1.00 2.00 3.00 5.00 6.00 8.00 Steel sample Absorbance 0.000 0.076 0.151 0.227 0.381 0.455 0.610 0.195 No Question Answer 1 a Write a half-equation for the conversion of Mn2+ to MnO4– in acidic aqueous solution. b Write a half-equation for the conversion of IO4– to IO3– in acidic aqueous solution. c Add the two half-equations together to obtain the overall equation for step 3. 2 Plot a graph of concentration of MnO4– (x-axis) against absorbance (y-axis). Draw a straight line of best fit through the points. Page 1 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 5.2 Colorimetric determination of manganese in steel No Question 3 Determine the concentration of MnO4– in the diluted steel solution. 4 Using your answer to question 3, find the % m/m of manganese in the steel sample. 5 After step 1, the solution contained a high concentration of iron(III) ions. Where did they come from? 6 If the standards were prepared using anhydrous potassium permanganate, calculate the mass of solute needed to prepare 100.0 mL of an 8.0 ppm MnO4– solution. 7 As well as water, what else should be added to the standards and the solvent blank and why? 8 Why was 520 nm a more appropriate wavelength than 400 nm for this analysis? 9 During calibration, the absorbance was set to zero using a solvent blank. Would your calculated percentage of Mn (m/m) have been affected had this not been done? Explain. 10 Suggest an alternative method for the determination of the manganese content of the steel sample. Answer Page 2 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 5.3 Analysing mass spectra NAME: CLASS: INTRODUCTION As early as the 1920s, mass spectrometry was used to identify the isotopes of elements, and their relative isotopic masses (RIM) and abundances. Armed with this information, accurate relative atomic masses (Ar) for elements could be determined. A flowchart of the steps in the operation of the mass spectrometer is shown below. No. Question Answer 1 Name the processes occurring in stages 1, 2 and 4. 2 Describe the processes occurring in: a stage 2 b stage 4. 3 The mass spectrum data obtained for a sample containing a mixture of two monatomic gases is shown below. Label each peak on the mass spectrum using the ZA X n notation. You can refer to the periodic table. Page 1 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 5.3 Analysing mass spectra No. Question 4 Element X has two naturally occurring isotopes with relative isotopic masses of 120.90 and 122.90. If the relative atomic mass of the element is 121.75, determine the percentage abundance of the lighter isotope. Answer The modern mass spectrometer is a powerful and accurate tool that is used extensively for the detection and identification of a wide range of compounds, rather than for RIM determinations. When a molecular substance is placed in the mass spectrometer, the mass spectrum shows a peak corresponding to the molecular ion, allowing the molecular mass to be determined. The bombarding electrons in the mass spectrometer also cause molecules to fragment, giving rise to a fragmentation pattern of lower molecular mass ions. The masses of the units broken off the molecule often give clues as to the structure of the molecule. No. Question Answer 5 The molecular structure and simplified mass spectrum of 2-pentanol are shown below. Label the peaks with the probable formula for the ions producing peaks at the following mass numbers. a 88 b 87 c 73 d 45 Page 2 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 5.3 Analysing mass spectra No. Question Answer 6 A simplified mass spectrum and the molecular structures of two possible compounds forming the spectrum are shown below. Based on an inspection of the mass spectrum, which compound is most likely to have produced the spectrum shown? Explain your choice. 7 Analysis of fragments in the mass spectrum gives clues to the structure of an unknown substance. Explain how a mass spectrum may be used to clearly identify an unknown substance. Page 3 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 5.4 Spectroscopic analysis of organic compounds I NAME: CLASS: INTRODUCTION Spectroscopy is used extensively in the analysis of organic compounds. Infrared spectroscopy is used to identify particular functional groups in compounds. Mass spectroscopy is used to determine the relative molecular mass of compounds, and is widely used to identify the structure of molecules. The fragmentation pattern obtained when molecules are broken is used to help determine their molecular structure. Nuclear magnetic resonance (NMR) spectroscopy is used for determining the precise structure of organic compounds. Absorption of radiation by 1H nuclei in different chemical environments provides information about the arrangement of hydrogen atoms within the compound. The questions that follow provide practice in the interpretation of spectroscopic data used for the identification of selected organic compounds. (You will learn more about these compounds and their functional groups in chapter 7 of the coursebook.) No. Question Answer 1 An unknown compound is thought to be either pentanal or 2,2-dimethylpropanal. The structures of these two compounds are shown below. NMR spectroscopic analysis is used to help identify the compound. The 1H NMR spectrum for each compound is produced. a How many peaks would you expect to see on the 1H spectrum of: i pentanal? ii 2,2-dimethylpropanal? b What is the expected ratio of areas for the peaks on the 1H spectrum of: i pentanal? ii 2,2-dimethylpropanal? c Given that the 1H NMR spectrum shows two peaks with the integration ratio 155:17.3, what is the identity of the unknown? d How could infrared spectroscopy be used to confirm the identity of the unknown? Page 1 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 5.4 Spectroscopic analysis of organic compounds I No. Question Answer 2 Two possible structures for an unknown organic compound are shown below. Both have the molecular formula C3H6O2. The correct structure is to be determined using spectroscopic analysis. a The infrared spectrum of the unknown compound shows a distinct absorption peak at 1710 cm–1, but no distinct peak in the 2500–3000 cm–1 range. i Briefly explain why the compound absorbs radiation in the infrared section of the electromagnetic spectrum. ii On the basis of the infrared data provided, which of the two structures is likely to be correct? Explain your choice. b The 1H nuclear magnetic resonance spectrum of the compound shows two peaks, one at a chemical shift of 1.9 ppm, the other at 3.8 ppm. The peak areas are in the ratio 1:1. i Which type of electromagnetic radiation is absorbed by atoms in molecules to produce an NMR spectrum? ii Based on the NMR spectrum, which of the two structures is most likely to be correct? Explain your choice. In more detailed NMR analysis, more information can be obtained. In high resolution NMR, absorption peaks are split into closely spaced peaks, called doublets, triplets etc. This splitting is the result of the influence of neighbouring nuclei on the nucleus causing the peak. Neighbouring nuclei are also tiny bar magnets and so they influence the effect of the applied magnetic field on the nucleus under investigation. Valuable information about these neighbouring groups can be obtained by analysing the splitting patterns of each signal in the NMR spectrum. Page 2 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 5.4 Spectroscopic analysis of organic compounds I No. Question Answer 3 An unknown compound is thought to be either compound I or II (molecular structures are shown below). The 1H NMR spectrum of the unknown compound shows two peaks with integration ratio 2:3. Which compound, I or II, is the unknown? 4 The high resolution 1H NMR spectrum of the unknown compound in question 3 is shown below. The numbers 2 and 3 represent relative peak areas. What is the relationship between the number of splits in a peak and the number of 1H nuclei adjacent to the 1H nucleus causing the peak? Page 3 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use. Worksheet 5.4 Spectroscopic analysis of organic compounds I No. Question Answer 5 An unknown compound is thought to be either compound I or II, (molecular structures are shown). The high resolution 1H NMR spectrum of the unknown compound is also shown. a How many peaks would be expected for i compound I? ii compound II? b What is the expected ratio of areas for the peaks for i compound I? ii compound II? c What is the expected splitting of the peaks for i compound I? ii compound II? d What other information from the NMR spectrum could be used to distinguish between compounds I and II? Page 4 © Pearson Education Australia (a division of Pearson Australia Group Pty Ltd) 2008. This page from the Chemistry Dimensions 2, Teacher’s Resource may be reproduced for classroom use.