Stoichiometry

1. Stoichiometry Mr Field

2. Using this slide show • The slide show is here to provide structure to the lessons, but not to limit them….go off-piste when you need to! • Slide shows should be shared with students (preferable electronic to save paper) and they should add their own notes as they go along. • A good tip for students to improve understanding of the calculations is to get them to highlight numbers in the question and through the maths in different colours so they can see where numbers are coming from and going to. • The slide show is designed for my teaching style, and contains only the bare minimum of explanation, which I will elaborate on as I present it. Please adapt it to your teaching style, and add any notes that you feel necessary.

3. Menu: • Lesson 1 – The Mole • Lesson 2 – Moles and Molar Mass • Lesson 3 - Solutions • Lesson 4 - Formulas • Lesson 5 - Equations • Lesson 6 – Theoretical Yields • Lesson 7 – Molar Volume of Gases • Lesson 8 – Ideal Gases • Lesson 9 - TEST • Lesson 10 – Test Debrief

4. Lesson 1 The Mole and the Avogadro Constant

5. Intro to Unit • Brainstorm everything you know about quantitative chemistry (equations, moles and all that stuff!)

6. Overview • Copy this onto an A4 page. You should add to it as a regular review throughout the unit.

7. Assessment • Test at the end of the unit (100%) • Approx lesson 9

8. We Are Here

9. Lesson 1: Meet the Mole • Objectives: • Understand the phrase ‘a mole’ as just a large number • Define ‘a mole’ according to Avogadro’s number • Complete an experiment to determine Avogadro’s number • Calculate numbers of particles using Avogadro’s number

10. What do the following have in common? A dozen A grand A score A pair A trio

11. A Mole • A mole is a large number: • A mole is: • 6.02 x 1023 • 602,000,000,000,000,000,000,000 • Six hundred and two thousand quadrillion • Given the symbol, L • This number is called Avogadro’s number after the Italian scientist Amedeo Avogadro who first proposed it • Picture – all the grains of sand, on all the beaches and in all the deserts in the world; that is about a tenth of a mole!

12. Some Calculations • We can use this equation to calculate a number of moles from a number of particles • Where: • n = quantity in moles • N = number of particles • L = Avogadro’s Constant (6.02x1023) • Example 1: You have 3.01x1022 atoms of carbon. How many moles is this? • n = N / L • n(C) = 3.01x1022 / 6.02x1023 • n(C) = 0.0500 mol • Example 2: You have 6.02x1024 molecules of water. How many moles of hydrogen atoms are present? • n = N / L • n(H) = 2 x n(H2O) • n(H) = 2 x 6.02x1024 / 6.02x1023 • n(H) = 20.0 • Example 3: How many atoms of hydrogen are there in 2.5 moles of methane (CH4)? • N(H) = 4 x N(CH4) • N(H) = 4 x n(CH4) x L • N(H) = 4 x 2.5 x 6.02x1023 • N(H) = 6.02x1024

13. Some Calculations • How many moles of ethene (C2H4) are there in 1.5x1022 molecules? • How many moles of oxygen atoms in 1.20x1024 molecules of sulphuric acid (H2SO4)? How many moles of oxygen molecules could they make? • How many molecules in 3.0 moles of nitrogen gas? • How many chloride ions are released on dissolving 0.050 moles of calcium chloride (CaCl2)? • BONUS QUESTION:How many moles of people are there on the planet?

14. How do you determine Avogadro’s number? • There are a number of ways including: • Electrolysis • Measurement of electron mass • X-ray crystal density • Molecular monolayers • You will be determining L by measuring the area of a molecular monolayer • Not the most accurate but feasible in our lab!

15. Key Points

16. Lesson 2 Moles and Molar Mass

17. Refresh • How many molecules are present in a drop of ethanol, C2H5OH, of mass 2.3 × 10–3 g? (L = 6.0 × 1023 mol–1) • A. 3.0 × 10 19 • B. 3.0 × 1020 • C. 6.0 × 1020 • D. 6.0 × 1026 • How many oxygen atoms are there in 0.20 mol of ethanoic acid, CH3COOH? • A. 1.2 × 1023 • B. 2.4 × 1023 • C. 3.0 × 1024 • D. 6.0 × 1024

18. We Are Here

19. Lesson 2: Moles and Molar Mass • Objectives: • Calculate the relative mass and molar mass of substances • Relate the mass of a substance to a quantity in moles • Conduct an experiment to determine the number of moles of water of crystallisation

20. Atomic and Formula Mass • Ar • is the relative atomic mass of an element • found in the periodic table • Mr • is the relative molecular mass of a compound • To calculate Mr, you add the Arof all the atoms in a compound • What does the term ‘relative’ mean? • What determines the mass of an atom and how can Ar be a decimal number?

21. Calculating Mr • HCl • Ar(H) = 1.01 • Ar (Cl) = 35.45 • Mr = 1.01 + 35.45 = 36.46 • C2H4 • Ar(C) = 12.01 • Ar (H) = 1.01 • Mr = 2x12.01 + 4x1.01 = 16.06 • H2SO4 • Ar(H) = 1.01 • Ar (S) = 32.06 • Ar(O) = 16.00 • Mr = 2x1.01 + 32.06 + 4x16.00 • = 98.08 • Mg(OH)2 • Ar(Mg) = 24.31 • Ar(O) = 16.00 • Ar (H) = 1.01 • Mr = 24.31 + 2x16.00 + 2x1.01 = 58.33

22. Calculate Mr for: • Br2 • C3H8 • (NH4)2SO4 • C6H12O6

23. Molar Mass, Mm • This is the mass of one mole of something. • To calulate Mm, simply stick ‘g’ for grams on the end of Mr. • For example: • Mr(H2O) = 18.02 • Mm(H2O) = 18.02 g • Note: This is why the value of L was chosen to be what it was.

24. Relating ‘mass’ and ‘molar mass’ • You need to be able to solve problems like: • How many moles of Y is mass X? • What is the mass of X moles of Y? • X moles of Y has a mass of Z, what is it’s molar mass? • Use this equation:

25. For example: • How many moles of water are present in 27.03g? • Calc Mm(H2O): • Mm(H2O)= 2x1.01 + 16.00 • = 18.02 • Find n(H2O): • n(H2O) = M / Mm • = 27.03 / 18.02 • = 1.50 mol • What is the mass of 4.40 mol of iron (III) oxide (Fe2O3)? • Calc Mm(Fe2O3): • Mm(H2O)= 2x55.85 + 3x16.00 • = 159.70 g • Find M(Fe2O3): • M(Fe2O3) = n x Mm • = 4.40 / 159.70 • = 703 g • 1.30 moles of an unknown compound has a mass of 20.9g, what is it’s molar mass? • Mm(unknown)= M / n • = 20.9 / 1.30 • = 16.1 g/mol

26. Questions • 8.8 moles of a compound has a mass of 1.41 kg. Calculate its molar mass. • 0.010 mol of an oxide of hydrogen has a mass 0.340 g. Deduce it’s formula. • Calculate the mass of 0.10 mol of benzene (C6H6) • Calculate the mass of 0.75 mol of ammonium nitrate (NH4NO3) • What quantity of iron (III) oxide is present in a 1.0 kg sample? • What quantity of cobalt (II) chloride (CoCl2) is present in a 2.40 g sample?

27. Moles of Water of Crystallisation • Many compounds can incorporate water into their crystal structure, this is called the water of crystallisation. • CoCl2 is blue CoCl2.6H2O is pink • The ‘.’ means the water is ‘associated’ with the CoCl2 • It is loosely bonded, but exactly how is unimportant • In this experiment you will calculate the moles of water of crystallisation of a compound

28. Key Points

29. Lesson 3 Solutions

30. Refresh • A student reacted some salicylic acid with excess ethanoic anhydride. Impure solid aspirin was obtained by filtering the reaction mixture. Pure aspirin was obtained by recrystallisation. The following data was recorded by the student. • Mass of salicylic acid used 3.15 ± 0.02 g • Mass of pure aspirin obtained 2.50 ± 0.02 g • Determine the amount, in mol, of salicylic acid, C6H4(OH)COOH, used.

31. We Are Here

32. Lesson 3: Solutions • Objectives: • Understand the relationship between concentration, volume and moles • Prepare a standard solution of silver nitrate. • Pose and solve problems involving solutions (of the chemical kind not the answers kind)

33. Solutions Basics • Aqueous copper sulfate solution: + SOLUTE SOLVENT SOLUTION

34. Concentration • This is the strength of a solution. Most Concentrated Least Concentrated

35. Molarity • The number of moles of a substance dissolved in one litre of a solution. • Units: mol dm-3 • Pronounced: moles per decimetre cubed • Units often abbreviated to ‘M’ (do not do this in an exam!) • Volume must be in litres (dm3) not ml or cm3 • This is the most useful measure of concentration but there are others such as % by weight, % by volume and molality.

36. Example 1: • 25.0 cm3 of a solution of hydrochloric acid contains 0.100 mol HCl. What is it’s concentration? • Answer: • Concentration = moles / volume = 0.100 / 0.0250 = 4.00 mol dm-3 • Note: the volume was first divided by 1000 to convert to dm3

37. Example 2: • Water is added to 4.00 g NaOH to produce a 2.00 mol dm-3 solution. What volume should the solution be in cm3? • Calculate quantity of NaOH: • n(NaoH) = mass / molar mass = 4.00/40.0 = 0.100 • Calculate volume of solution: • Volume = moles / concentration = 0.100 / 2.00 = 0.0500 dm3 = 50.0 cm3

38. Example 3: • It is found by titration that 25.0 cm3 of an unknown solution of sulfuric acid is just neutralised by adding 11.3 cm3 of1.00 mol dm-3 sodium hydroxide. What is the concentration of sulfuric acid in the sample. H2SO4 + 2 NaOH  Na2SO4 + 2 H2O • Use: n1C1V1 = n2C2V2 1 x C1 x 25.0 = 2 x 1.00 x 11.3 C1 = (2 x 1.00 x 11.3) / (1 x 25.0) = 0.904 mol dm-3 Where: n = coefficient C = concentration V = volume ‘1’ refers to H2SO4 ‘2’ refers to NaOH

39. Questions • You have 75.0 cm3 of a 0.150 mol dm-3 solution of zinc sulphate (ZnSO4). What mass of zinc sulphate crystals will be left behind on evaporation of the water? • What mass of copper (II) chloride (CuCl2) should be added to 240 cm3 water to form a 0.100 mol dm-3 solution? • A 10.0 cm3 sample is removed from a vessel containing 1.50 dm3 of a reaction mixture. By titration, the sample is found to contain 0.0053 mol H+. What is the concentration of H+ in the main reaction vessel? • The titration of 50.0 cm3 of an unknown solution of barium hydroxide was fully neutralised by the addition of 12 cm3 of 0.200 mol dm-3 hydrochloric acid solution. What concentration is the barium hydroxide solution? Ba(OH)2 + 2 HCl BaCl2 + 2 H2O

40. Preparing a Standard Solution • Prepare standard solutions silver nitrate with concentrations of 0.25-0.75 mol dm-3 • You will use these in a future lesson so make them well!

41. Problem Time • Write three stoichiometry problems using any of the ideas from the unit so far (make sure you know the answer). • In ten minutes time you will need to give your problems to a classmate to solve.

42. Key Points

43. Remember! Alcohol is not the answer but it is a solution*! *Of ethanol, water and various other bits and pieces

44. Lesson 4 Formulas

45. Refresh • Which statement about solutions is correct? • When vitamin D dissolves in fat, vitamin D is the solvent and fat is the solute. • In a solution of NaCl in water, NaCl is the solute and water is the solvent. • An aqueous solution consists of water dissolved in a solute. • The concentration of a solution is the amount of solvent dissolved in 1 dm3 of solution • A student added 27.20 cm3 of 0.200 mol dm–3HCl to 0.188 g of eggshell. Calculate the amount, in mol, of HCl added.

46. We Are Here

47. Lesson 4: Formulas • Objectives: • Understand the difference between empirical and molecular formulas • Experimentally determine the empirical formula of a compound. • Solve problems involving empirical and molecular formulas

48. Formulas • You have one minute to write down as many different chemical formulas as you can. • Go! • What is the job of a formula? No!!!!

49. Types of Formula • Empirical: • The ratio of the atoms of each element in a compound in its lowest terms • Molecular: • The number of atoms of each element in a molecule • You will meet other types in the organic chemistry unit including structural, displayed and skeletal.

50. Examples • Oxygen • Molecular: O2 • Empirical: O • Water • Molecular: H2O • Empirical: H2O • Ethane • Molecular: C2H6 • Empirical: CH3 • Glucose • Molecular: C6H12O6 • Empirical: CH2O • Sodium Chloride • Molecular: n/a* • Empirical: NaCl • Copper Sulphate • Molecular: n/a* • Empirical: CuSO4 • *Why do these two not have an empirical formula?