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Chapter 2: Amount of substance

A-Level Chemistry Teacher Pack | Page 1 | ©HarperCollinsPublishers Limited 2015

Scheme of work


CHAPTER PLANNING

16 hours

This chapter includes a substantial amount of practical work including one of the Required practicals. Some of the ideas dealt with resonate with Chapter 1, for example, Relative masses. Some of the lessons for Chapter 2 may be taught in parallel or interspersed with lessons for Chapter 1.

You may wish to use some of the later lessons that are predominately focussed on practical skills and experiments with lessons that are part of other chapters. For example, 2.11 Thermal decomposition could be part of lessons about Group 2 metals, and 2.15 Relative formula mass of succinic acid could be part of lessons about carboxylic acids.

Although Required practical 1: Make up a volumetric solution and carry out a simple acid–base titration is positioned here, there is no reason why suitable practical activities such as 2.15 Relative formula mass of succinic acid should not come later in the course and practical skills to be assessed for endorsement (required for A-level, but not AS) demonstrated then.

ONE HOUR LESSONS SPECIFICATION CONTENT

2.1 Relative masses 3.1.2.1 Relative atomic mass and relative molecular mass Relative atomic mass and relative molecular mass in terms of 12C. The term relative formula mass will be used for ionic compounds. Students should be able to define relative atomic mass (Ar) and relative molecular mass (Mr).

2.2 The mole

2.3 Solutions

2.4 Mole calculations

3.1.2.2 The mole and the Avogadro constant The Avogadro constant as the number of particles in a mole. The mole as applied to electrons, atoms, molecules, ions, formulae and equations. The concentration of a substance in solution measured in mol dm-3. Students should be able to carry out calculations using: • the Avogadro constant • mass of substance, relative molecular mass and amount in moles • concentration, volume and amount of substance in a solution.

Students will not be expected to recall the value of the Avogadro constant. MS 0.1: Students carry out calculations using numbers in standard and ordinary form, e.g. using the Avogadro constant. MS 0.4: Students carry out calculations using the Avogadro constant. MS 1.1: Students report calculations to an appropriate number of significant figures, given raw data quoted to varying numbers of significant figures. Students understand that calculated results can only be reported to the limits of the least accurate measurement.

2.5 Ideal gas equation

2.6 Relative molecular mass of a volatile liquid

3.1.2.3 The ideal gas equation The ideal gas equation pV = nRT with the variables in SI units. Students should be able to use the equation in calculations. They will not be expected to recall the value of the gas constant, R. AT a, b and k PS 3.2: Students could be asked to find the Mr of a volatile liquid. MS 0.0: Students understand that the correct units need to be in pV = nRT MS 2.2, 2.3 and 2.4: Students carry out calculations with the ideal gas equation, including rearranging the ideal gas equation to find unknown quantities.

2.7 Calculating empirical formulae

2.8 Determine the

3.1.2.4 Empirical and molecular formula Empirical formula is the simplest whole number ratio of atoms of each element in a compound. Molecular formula is the actual number of atoms of each element in a

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formulae of compounds compound. The relationship between empirical formula and molecular formula. Students should be able to calculate: • empirical formula from data giving composition by mass or data giving percentage

by mass • molecular formula from the empirical formula and relative molecular mass.

AT a and k PS 2.3 and 3.3: Students could be asked to find the empirical formula of a metal oxide.

2.9 Balanced chemical equations

2.10 Yields from reactions

2.11 Thermal decomposition

2.12 Titrations

2.13 Ethanoic acid in vinegar

2.14 Relative formula mass of a hydrated salt

2.15 Relative formula mass of succinic acid

2.16 Analysing medicines

3.1.2.5 Balanced equations and associated calculations Equations (full and ionic). Percentage atom economy is:

molecular mass of desired product

sum of molecular masses of all reactants x 100

Students should be able to • write balanced equations for reactions studied • balance equations for unfamiliar reactions when reactants and products are

specified.

Students should be able to use balanced equations to calculate masses; volumes of gases; percentage yields; percentage atom economies; concentrations and volumes for reactions in solutions. AT a, d, e, f and k PS 4.1: Students could be asked to find: the concentration of ethanoic acid in vinegar; the mass of calcium carbonate in an indigestion tablet; the Mf of MHCO3; the Mr of succinic acid; the mass of aspirin in an aspirin tablet; the yield for the conversion of magnesium to magnesium oxide; the Mr of a hydrated salt (e.g. magnesium sulfate) by heating to constant mass. AT a and k: Students could be asked to find the percentage conversion of a Group 2 carbonate to its oxide by heat. AT d, e, f and k: Students could be asked to determine the number of moles of water of crystallisation in a hydrated salt by titration. MS 0.2: Students construct and/or balance equations using ratios. Students calculate percentage yields and atom economies of reactions. MS 1.2 and 1.3: Students (a) select appropriate titration data (i.e. identify outliers) in order to calculate mean titres (b) determine uncertainty when two burette readings are used to calculate a titre value. Required practical 1: Make up a volumetric solution and carry out a simple acid–base titration.

PRIOR KNOWLEDGE

Students are likely to have been taught, at KS4, that chemical reactions can be represented by word equations or by symbol equations.

They will probably know how to calculate the percentage of an element in a compound from Ar of the element and Mf of the compound, the empirical formula of a compound from the masses or percentages of the elements in a compound, and the masses of reactants and products from balanced symbol equations.

They are also likely to know the difference between actual yield, theoretical yield and percentage yield, and the reasons why percentage yield is not 100%.

Some students may have carried out practical work such as working out the empirical formulae of CuO and MgO, and calculating yields (for example Mg burning to produce MgO or wire wool burning to produce iron oxide).

Students who studied triple science are also likely to have been taught about acid-alkali titrations and their use in analysis. They may be familiar with units for concentration: moles per dm3 (mol dm-3) and grams per dm3 (g dm-3). Some may have carried out strong acid/strong alkali titrations, such as hydrochloric acid against sodium hydroxide solution, to find unknown concentration.

WHERE IT LEADS

Later in their A-level course, students will find out how the arrangements of electrons in atoms affects the chemical and physical properties of elements and their compounds.

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Lesson plan 2.1

Lesson 2.1 Relative masses

LEARNING OUTCOMES

Students should be able to

• define relative atomic mass (Ar) , relative molecular mass (Mr) and relative formula mass (Mf).

THE JOURNEY SO FAR

Students will have learned about relative atomic mass (Ar) and relative molecular mass (Mr) in Chapter 1. In this lesson they will also use relative formula mass (Mf) in recognition that many elements and compounds do not exist as molecules.

POSSIBLE BARRIERS TO PROGRESS

Some students might still be a little confused about calculating relative atomic masses. It may appear to them that using whole numbers for relative isotopic masses means that they are using mass numbers. So, to clarify:

mass number = number of protons + number of neutrons. There are no units.

Relative masses of protons and neutrons are usually approximated to 1.

Masses of atoms are measured in atomic mass units (amu). One amu is defined currently as 1.66053873 x 10-27 kg, but it is not an SI unit. As a basis for comparison, the mass of a carbon-12 atom is taken to be 12.00000 amu. Other atoms are measured relative to this.

Low resolution mass spectrometry gives relative masses that match the mass numbers. High resolution mass spectrometry gives more precise measurements. The relative masses of atoms can be measured to several significant figures, but they are never whole numbers. The reason is that the mass of a proton is 1.00728 amu and the mass of a neutron is 1.00867 amu. So the mass of an atom is always slightly greater than its mass number, even given that, when atoms form, a very small amount of mass is converted into energy.

Taking uranium-238 as an example: mass number = 238; relative mass = 238.0508; mass of protons and neutrons to make 238U atom = 239.918.

LESSON OUTLINE

Engage and remind

This lesson is largely a recap of work from Chapter 1, with the added concept of relative formula mass (Mf).

Recap the use of mass spectrometry to measure relative isotopic masses and the relative abundances of isotopes of an element. They could be shown some mass spectra and asked to interpret them, for example, those in Activity sheet 2.1.1 Examples of low resolution mass spectra. It may be useful to remind students of the difference between low resolution and high resolution mass spectrometry.

Core activities

Ask students to read Chemistry Student Book 1, Section 2.1 and answer questions 1 and 2.

Some further questions are provided in Activity sheet 2.1.2 Relative masses.

Consolidate and look ahead

The lesson has been largely about consolidating prior learning. Briefly look ahead by asking the question, how can numbers of particles (atoms, molecules and ions) be measured out?

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Activity sheet 2.1.1

Examples of low resolution mass spectra

TASK 1: CARBON DIOXIDE

This is the mass spectrum of carbon dioxide, CO2.

01. What are the m/z values for the lines at A, B and C?

02. To which ions do the lines at A, B and C correspond?

03. What is the relative molecular mass of carbon dioxide?

TASK 2: BROMINE

This is the mass spectrum of bromine, Br2.

Bromine has two naturally occurring isotopes: 79 Br relative abundance = 50.5% 81Br relative abundance = 49.5%

01. What are the m/z values for the lines at A, B, C, D and E?

02. To which ions do the lines at A B, C, D and E correspond?

03. What is the relative molecular mass of bromine? Give your answer to three significant figures.

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Relative masses

TASK 1

01. Define

a. relative atomic mass

b. relative molecular mass

02. Most non-metallic elements exist as molecules, for example: H2, N2, O2, F2, Cl2, Br2, I2, P4, S8. Calculate the relative molecular masses of

a. oxygen

b. chlorine

c. phosphorus

d. sulfur

using these Ar values: H = 1, O = 16, Cl = 35.5, P = 31, S = 32

03. Now calculate the relative molecular masses of

a. oxygen

b. chlorine

c. phosphorus

d. sulfur

using these more precise Ar values: H = 1.008, O = 15.999, Cl = 35.453, P = 30.974, S = 32.065

04. Calculate the percentage error (to three significant figures) in the relative molecular mass of

a. oxygen

b. chlorine

c. phosphorus

d. sulfur

when rounded-up relative atomic masses are used rather than the precise values.

TASK 2

01. Name the types and numbers of atoms needed to make one molecule of

a. water

b. sulfur trioxide

c. nitrogen trichloride

d. hydrogen peroxide

e. sulfuric acid

f. nitric acid.

Note: The sulfuric acid and nitric acid you use on the laboratory are aqueous solutions. The pure acids are both oily liquids that exist as molecules.

02. Using the rounded-up Ar values, calculate the relative molecular masses of

a. water

b. sulfur trioxide

c. nitrogen trichloride

d. hydrogen peroxide

e. sulfuric acid

f. nitric acid

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03. Explain why we calculate the relative formula mass rather than relative molecular mass of silicon dioxide.

04. Name the types and numbers of atoms needed to make one formula unit of

a. magnesium chloride

b. calcium carbonate

c. magnesium hydroxide

d. calcium nitrate.

05. Calculate the relative formula masses of

a. magnesium chloride

b. calcium carbonate

c. magnesium hydroxide

d. calcium nitrate

using these Ar values: H, 1; C, 12; N, 14; O, 16; Mg, 24; Cl, 35.5; Ca, 40

TASK 3

Many metal salts have water of crystallisation. This mean that a fixed number of water molecules are held in the giant structure of metals ions. They are called hydrated salts.

Calculate the relative formula masses of these hydrated salts:

01. Sodium carbonate decahydrate, Na2CO3.10H2O

02. Copper(II) sulfate pentahydrate, CuSO4.5H2O

03. Aluminium sulfate octahydrate, Al2(SO4)3.8H2O

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Lesson plan 2.2

Lesson 2.2 The mole

LEARNING OUTCOMES

Students should be able to carry out calculations using

• the Avogadro constant • mass of substance, relative molecular mass (or relative formula mass) and amount in moles

THE JOURNEY SO FAR

For many students this may be their first encounter with the concept of the mole, though some that studied triple science might be familiar with units for concentration: moles per dm3 (mol dm-3).


Many students may think of the mole simply as a number: 1 mole, 2 moles and so on. It is important that they know and use it as a unit, in the same sense that (a) a couple is two, (b) a dozen is 12, (c) a gross is 144.

One mole is the Avogadro constant (6.023 x 1023) of items, and the ‘item’ can be anything. In science the item might be atoms, molecules, formula units, ions or sub-atomic particles. The mole (abbreviation mol) is the SI Unit for an amount of a substance.

When stating the number of moles it is essential to write the formula of substance. Also, on occasions, the particle being counted needs to be specified: atom, molecule or ion, for example. One mole of chlorine molecules, Cl2, is made from 2 mol chlorine atoms. A useful analogy might be a four-legged table. One table has four legs. A couple of tables have 2 x 4 = 8 legs and a dozen have 12 x 4 = 48 legs. One mole of four-legged tables would have 4 moles of legs.

LESSON OUTLINE

Engage and remind

Ask students how quantities of elements and compounds are measured out in the laboratory. Since chemists are interested in the numbers of particles (atoms, ions or molecules) that take part in a chemical reaction, pose the question: What do we need to know to get from measured quantities to numbers of particles?

Core activities

Guide students through Chemistry Student Book 1, Section 2.2. It might be helpful to display (a) measured masses or volumes of 1 mole each of a selection of elements and compounds, (b) 0.01, 0.1 and 1 mol of powdered carbon (0.012 g, 0.12 g, 1.2 g and 12 g) and air (24 cm3, 240 cm3, 2.4 dm3, 24 dm3). Alternatively, you could show photographs. Ask students to complete Activity sheet 2.2.1 Moles.

Students may well question how the number 6.023 x 1023 can be worked out. Either ask students to carry out the experiment described in Activity sheet 2.2.2 Estimating the Avogadro constant – stressing the need for care when doing this experiment – or perform it as a demonstration. While results may sometimes disappoint students, with care, a reasonable estimate of the Avogadro constant can be made. The main objective, however, is for students to grasp the principle of the determination. Discuss the reasonableness of assumptions made in the calculation, for example, oleic acid is shaped like a cube.

Here and in subsequent lessons in this chapter, students will gain experience in these mathematical areas:

• MS 0.1: Recognise and use expressions in decimal and ordinary form. • MS 0.4: Use calculators to find and use power, exponential and logarithmic functions. • MS 1.1: Use an appropriate number of significant figures.


Ask students to answer questions 3 and 4 from Chemistry Student Book 1, Chapter 1, either at the end of the lesson or as homework. Explain that, during the next lesson, students will carry out a quantitative analysis of a solution, using the mole in calculations.

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Moles The mole (abbreviation mol) is the SI Unit for an amount of a substance. One mole is the Avogadro constant (6.023 x 1023) number of atoms, molecules, formula units, ions or sub-atomic particles.

Relative atomic masses are measured relative to an atom of the isotope carbon-12 which is assigned the value 12.0000. One mole of carbon-12 atoms is 6.023 × 1023 atoms and has a mass of 12.000 g

TASK 1: HOW MANY ATOMS?

01. There are many words used to name a specific number of items. What specific number of items do the following words indicate?

a. couple

b. dozen

c. baker’s dozen

d. score

e. gross

f. ream

g. mole

02. How many atoms are there in

a. 3 mol carbon

b. 0.01 mol argon

c. 500 mol gold?

TASK 2: MAKING SUBSTANCES FROM ATOMS

01. How many atoms are needed to make

a. 1 mol of water

b. 10 mol of ammonia

c. 2 mol methane

d. 0.001 mol hexane, C6H14?

02. How many atoms are needed to make

a. 100 mol copper(II) oxide, CuO

b. 0.1 mol sliver nitrate, AgNO3

c. 1 x 104 mol silicon dioxide, SiO2

d. 2 x 10-4 mol calcium carbonate, CaCO3?

03. Explain the difference between compounds listed in question 01 and compounds listed in question 02.

TASK 3: STRUCTURE OF MOLECULES

Look at these diagrams. They show the structures of three molecules.

In each case, (a), (b) and (c), name the atoms needed to make the molecules and the number needed to make one mole of each compound.

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Estimating the Avogadro constant

THE INVESTIGATION

In this investigation, a relatively simple experiment is used to estimate the value of the Avogadro constant.

EQUIPMENT AND MATERIALS

• dropping pipette • clean bowl with non-stick surface inside, such as Teflon® • glass plate, large enough to cover the bowl, with four corks to support it above the bowl • wash bottle containing tap water • lycopodium powder in dispenser • scissors • tracing paper, pencil and scissors • solution containing 0.0068 cm3 of oleic acid dissolved in 10 cm3 pentane • 1 mol dm-3 hydrochloric acid

PROCEDURE

Wear eye protection. Pentane is very volatile and flammable. Make sure you work away from any naked flames.

01. Fill the bowl with tap water until it is almost full. Add a drop of 1 mol dm-3 hydrochloric acid. Use a wash bottle to top up the water level until its surface bulges up above the rim of the bowl.

02. Lightly dust lycopodium powder over the surface of the water.

03. Carefully place two drops (about 0.1 cm3) of the oleic acid solution on the water’s surface. (If you have a graduated pipette, use it to measure the 0.1 cm3 more accurately.) The pentane will evaporate and the oleic acid will spread out into a monolayer, pushing back the lycopodium powder.

04. Place a glass plate on supporting stoppers over the bowl making sure it is not touching the water. Support it at the corners with the four corks. Put the tracing paper on the glass plate and trace the outline of the oleic acid film.

05. Carefully cut out the outline from the tracing paper and place it on a piece of graph paper. By counting the squares, estimate the area of the oleic acid film.

DATA ANALYSIS

The oleic acid solution contained 0.006 g (0.0068 cm3) of oleic acid dissolved in 10 cm3 of pentane.

1. Calculate the number of moles of oleic acid in the 0.10 cm3 of oleic acid solution. Mr[C18H34O2] = 282.

2. Assume that an oleic acid molecule has the shape of a cube with sides of length d and volume d3.

Vfilm = A x d, where Vfilm = the volume of the film, A = area and d = thickness.

Rearranging,

d = Vfilm

A

Calculate d and estimate the volume of an oleic acid molecule, Vmolecule.

Remember the volume of oleic acid.

3. Calculate the number of molecules in the oleic acid film using

Vfilm

Vmolecule

4. The Avogadro constant is the number of molecules in 1 mol of substance (oleic acid in this experiment). Calculate it from the moles of oleic acid in the 0.10 cm3 of oleic acid solution used to make the film and the number of molecules in the oleic acid film.

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Technician notes for Activity sheet 2.2.2

Estimating the Avogadro constant This experiment may be done as a class exercise or a demonstration.


• dropping pipette • clean bowl with non-stick surface inside, such as Teflon® • glass plate, large enough to cover the bowl, with four corks to support it above the bowl • wash bottle containing tap water • lycopodium powder in dispenser • scissors • tracing paper, pencil and scissors • solution containing 0.0068 cm3 (0.006 g) of oleic acid dissolved in 10 cm3 pentane (pentane is HIGHLY

FLAMMABLE and an IRRITANT) • 1 mol dm-3 hydrochloric acid LOW HAZARD

NOTES

Pentane is used because it is very volatile and evaporates quickly, allowing the film of oleic acid to form. However, solutions made up in pentane must be kept in a tightly-stoppered container to prevent evaporation. Also, pentane is highly flammable and students should work well away from naked flames.

Given the precision of balances available, it is likely that a solution of 0.06 g oleic acid in 100 cm3 of pentane will need to be made.

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Lesson plan 2.3

Lesson 2.3 Mole calculations

LEARNING OUTCOMES

Students should be able to carry out calculations using

• the Avogadro constant, mass of substance and amount in moles

THE JOURNEY SO FAR

This lesson reinforces learning in earlier lessons, in particular 2.2 The mole and 2.3 Solutions.

LESSON OUTLINE

Engage and remind

To remind students about moles and the Avogadro constant ask them to complete Assignment 1: How many atoms are there in the world? in Chemistry Student Book 1, Chapter 2.

Core activities

Guide students through Chemistry Student Book 1, Section 2.2, Converting mass to moles and Converting moles to masses. To check their progress, ask them to answer questions 5 and 6. As in earlier lessons, it might be helpful to display measured masses or volumes of different number of moles of a selection of elements and compounds.

Then run through Chemistry Student Book 1, Section 2.2 and, again to check progress, answer question 7.

At an appropriate time, tell students about the experiments carried out by Lord Rayleigh to estimate the size of a molecule. Ask students to complete task 1 on Activity sheet 2.3.1 Lord Rayleigh and olive oil. To reinforce the work they did in the previous lesson on the estimation of the Avogadro constant, ask students to complete task 2. Students might be asked to research the work of Lord Rayleigh and his contribution, and that of others, to working out the approximate size of molecules.


Ask students to complete Activity sheet 2.3.2 Mole calculations. Explain to students that, in the next lesson, they will use the concept of the mole to describe amounts of substances in solution.

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Lord Rayleigh and olive oil Lord Rayleigh was an English scientist and Nobel prize winner. He lived from 1842 to 1919. One of his scientific experiments was aimed at estimating the size of an oil molecule. Rayleigh did this by putting a single drop of olive oil (of known mass and density) onto the surface of clean water and measuring the area it covered when it had spread out completely on the water’s surface.

Rayleigh made a key assumption, which has been proven to be valid. It was that the oil slick was one molecule thick. He repeated his experiment several times.

TASK 1

Here are some data from a typical experiment using olive oil.

• Diameter of oil drop = 0.1 cm • Density of olive oil = 0.90 g cm-3 • Area of oil film = 0.9 m2

From data similar to these, Rayleigh was able to estimate the size of the olive oil molecule. Can you?

01. Calculate the volume in cm3 of the oil drop using V = (4/3)πr3

02. Assuming the molecule is cubic (which it isn’t, but it makes the calculation easier) estimate the length of the molecule in centimetres. Remember that the volume of the oil drop equals the volume of the oil film.

TASK 2

Olive oil is a mixture of triglycerides (made from fatty acid and glycerol) and small quantities of fatty acids and glycerol. Here is the molecular formula for a typical triglyceride:

01. Calculate the relative molecular mass, Mr[triglyceride], of this typical triglyceride. (Ar: H, 1; C, 12; O, 16)

02. Now, using the data from task 1, calculate the

a. number of typical triglyceride molecules in the oil film

b. mass of each typical triglyceride molecule

c. number of molecules in one mole of typical triglyceride molecules

03. What is the significance of this number?

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Mole calculations For these tasks you will need these Ar values:

H, 1; C, 12; N, 14; O, 16; Ne, 20; Na, 23; Mg, 24; S, 32; Cl, 35.5; Ca, 40

TASK 1: MOLES TO MASSES

Calculate the masses of:

01. 4 mol water molecules

02. 0.25 mol sulfur trioxide gas

03. 2 x 103 calcium atoms

04. 5 x 10-3 nitrogen trichloride molecules

05. 0.1 mole sodium ions

06. 0.02 mol sulfate ions

07. 4 x 10-5 neon atoms

TASK 2: MASSES TO MOLES

Calculate the number of moles of:

01. molecules in 142 g chlorine gas

02. formula units in 285 g magnesium chloride

03. magnesium ions in 285 g magnesium chloride

04. nitrate ions in 82 g calcium nitrate

05. atoms of neon in 500 g of neon gas

TASK 3: SOLUTIONS

01. Calculate the concentration in mol dm-3 of 2.85 g of magnesium chloride in 250 cm3 of solution.

02. Calculate mass of sodium sulfate needed to make 100 cm3 of 0.5 mol dm3 sodium sulfate solution.

03. If 10 cm3 of 2 mol dm-3 hydrochloric acid is diluted with water to make 250 cm3, what is the concentration of the diluted solution?

04. Calculate the concentration in mol dm-3 of nitrate ions in an 8.2 g dm-3 solution of calcium nitrate.

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Lesson plan 2.4

Lesson 2.4 Solutions

LEARNING OUTCOMES


• carry out calculations using concentration, volume and amount of substance in a solution

THE JOURNEY SO FAR

Students should be familiar with using the Avogadro constant to calculate numbers of particles (atoms, molecules and ions) and formula units. They will also be able to convert moles of substances into masses and vice versa. This lesson introduces the application of these to solutions.


In chemistry, we normally talk about mass of solute in solution rather than mass of solute in solvent, though the latter is often used when making solutions of approximate concentration. Until they are familiar with and confident using units, it may help to read them out in full. For example, ‘5 grams in 1 decimetre cubed of solution’ or ‘0.2 gram in 100 centimetre cubed of water (or other solvent)’.

LESSON OUTLINE

Engage and remind

Show students a solution of potassium manganate(VII) and begin to dilute it progressively with water. Now show them a solution of potassium manganate(VII) of unknown concentration. Ask them to suggest how the concentration might be determined and what units could be used for the concentration.

Core activities

Explain that analytical work in chemistry requires the safe and correct use of various apparatus and techniques. Ask students to complete Activity sheet 2.3.1 Analysing potassium manganate(VII) solution.

In this activity, students make a solution of known concentration and, from it, prepare a series of other solutions of known concentration. This illustrates the principle of ATe Use volumetric flask, including accurate technique for making up a standard solution. However, instead of a volumetric flask they use a measuring cylinder in this activity. The diluted solutions are made up using measuring cylinders, but two burettes (one for the standard potassium manganate(VII) solution and the other for deionised water) could be used for greater accuracy.

If a colorimeter is available, the absorbance of the potassium manganate(VII) solution could be measured and a graph plotted of absorbance against concentration. The absorbance of the unknown could be measured and the graph used to determine its concentration.

Now ask them to complete task 1 on Activity sheet 2.3.2 Calculating concentrations.


Suggest that students read the introductory part of Chemistry Student Book 1, Section 2.7, answer questions 16 and 17, and then complete task 2 on Activity sheet 2.3.2 Calculating concentrations.

As an out-of-class activity – and in preparation for the lessons to come - ask students to read:

• Making a volumetric solution from Chemistry Student Book 1, Section 2.7 and answer question 18 • Required practical: Apparatus and techniques from Chemistry Student Book 1, Section 2.7 and complete

task 3 on Activity sheet 2.3.2 Calculating concentrations.

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Analysing potassium manganate(VII) solution

THE INVESTIGATION

Potassium manganate(VII) is an intense purplish black coloured crystalline solid. It dissolves readily in water and the intensity of the colour depends on the concentration of potassium manganate(VII).

You will be given a solution of potassium manganate(VII) of unknown concentration. Your task is to make a series of solutions of potassium manganate(VII) of known concentration and to determine the concentration of the unknown by colour matching.


• 8 x 100 cm3 beakers • 1 x 250 cm3 measuring cylinder • 2 x 25 cm3 measuring cylinders • electronic balance (capable of weighing to the nearest 0.01 g or better) • piece of white card or paper • potassium manganate(VII), KMnO4(s) • solution of potassium manganate(VII) of unknown concentration

PROCEDURE

Wear eye protection. Potassium manganate(VII) stains very easily, including skin (the purple stain slowly becomes brown), so wear protective gloves when working with solid potassium manganate(VII) and its solutions.

Prepare a standard solution of potassium manganate(VII)

01. Place a 100 cm3 beaker on the balance pan and zero the balance. Measure 0.2 ±0.1 g potassium manganate(VII) into the beaker and record the mass to the nearest 0.01 g.

02. Add about 20 cm3 of deionised water and swirl the contents to dissolve the potassium manganate(VII).

03. Pour the solution into a 250 cm3 measuring cylinder, using a small funnel. Rinse the beaker using a wash bottle to make sure all the potassium manganate(VII) solution has been washed into the cylinder. Make the solution up to the 250 cm3 graduation mark. This is your standard solution.

Colour matching analysis

01. Label six 100 cm3 beakers A to F. Use the two measuring cylinders (one for the standard potassium manganate(VII) solution and the other for deionised water) to measure these volumes into the beakers:

Beaker A B C D E F

Volume of standard solution of potassium manganate(VII) 20 16 12 8 4 0

Volume of water 0 4 8 12 16 20

Concentration of potassium manganate(VII) / g dm-3

Measure the volumes as carefully as you can. Use dropper pipettes (one for the standard potassium manganate(VII) solution and the other for deionised water) to make the final additions to the measuring cylinders.

02. Stand the beakers in a line on a white background, such as a sheet of white card or paper.

03. Measure 20 cm3 of solution of potassium manganate(VII)) of unknown concentration into a 100 cm3 beaker.

04. Compare the unknown with the solutions you prepared by dilution of the standard potassium manganate(VII) solution. Find which solution its colour most closely matches. Estimate the concentration of potassium manganate(VII) in the unknown.

DATA ANALYSIS

Calculate the concentration of the potassium manganate(VII) standard solution in g dm-3.

Calculate the concentrations in solutions A-E, also in g dm-3.

Estimate the concentration of potassium manganate(VII) in the unknown solution.

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Technician notes for Activity sheet 2.4.1

Analysing potassium manganate(VII) solution Students should wear eye protection and protective gloves, such as nitrile gloves. Potassium manganate(VII) stains very easily, including skin (the purple stain slowly becomes brown), so wear protective gloves when working with solid potassium manganate(VII) and its solutions.

A risk assessment must be made before students start work.


• 8 x 100 cm3 beaker • 1 x 250 cm3 measuring cylinder • 2 x 25 cm3 measuring cylinders • electronic balance (capable of weighing to the nearest 0.01 g or better) • piece of white card or paper • solid potassium manganate(VII) HARMFUL and OXIDISING • solution of potassium manganate(VII) of unknown concentration

NOTES

A volumetric flask could be used instead of the 250 cm3 measuring cylinder. However, it is suggested that students understand the principle first and build up the skills of accurate measurement over the up-coming lessons. Making an accurate standard solution is also more time-consuming.

Students could use two 50 cm3 burettes to prepare the various diluted standards. This is certainly more accurate, would introduce students to the use of burettes and would not be very time-consuming, especially if the burettes are filled in advance of the lesson.

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Calculating concentrations

TASK 1

The standard solution of potassium manganate(VII) you prepared for Activity 2.3.1 Analysing potassium manganate(VII) solutions had a concentration of about 0.8 g dm-3. The intensity of the colour was related to the concentration in g dm-3, but as you have seen already, chemists are more interested in numbers of particles than they are in masses.

01. Explain why the potassium manganate(VII) solution you prepared is called a standard solution.

02. Potassium manganate(VII) is an ionic compound. Write an equation to show what happens when it dissolves in water.

03. Which ion in solution causes the characteristic purplish pink colour? Explain you answer.

04. Calculate the relative formula mass of potassium manganate(VII). (Ar values: O, 16; K, 39; Mn, 55)

05. Calculate the concentrations in mol dm-3 of

a. potassium manganate(VII)

b. potassium ions

c. manganate(VII) ions

in your standard solution.

06. Complete this table using your calculations and results from for activity 2.3.1.

Beaker A B C D E F

Volume of standard solution of KMnO4(aq) 20 16 12 8 4 0

Volume of water 0 4 8 12 16 20

Concentration of MnO4-(aq) / mol dm-3

TASK 2

01. You are given the concentration and the volume of a solution and asked to calculate the number of moles in the solution. Write a word equation to show the relationship between moles, concentration and volume. Include the units.

02. You are told the volume of a solution and the number of moles of solute in it. Write a word equation for the concentration of the solution. Include the units.

03. There are 10.6 g of Na2CO3 in 40 cm3 of aqueous solution.

a. How many moles of Na2CO3 does 10.6 g represent?

b. What is the concentration of the solution?

c. How many moles in 1.5 dm3?

04. Calculate the

a. number of moles of sodium hydroxide in 28 cm3 of a solution containing 1.00 mol dm-3

b. concentration of 16 g of sodium hydroxide in 0.5 dm3 of solution

c. number of moles of potassium iodide in 120 cm3 of a solution of concentration 1.5 mol dm-3

d. concentration of 3.78 g of potassium iodide in 35 cm3 of solution

TASK 3

Reflect on the method you used to determine the concentration of potassium manganate(VII).

01. What were the main sources of error that using different apparatus and techniques could have been avoided or reduced?

02. Suggest how a more accurate value for the concentration could be determined.

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Lesson 2.5 Ideal gas equation

LEARNING OUTCOMES


• carry out calculations using the ideal gas equation pV = nRT with the variables in SI units

THE JOURNEY SO FAR

The ideal gas equation is not part of GCSE Chemistry or GCSE Physics specifications so it is unlikely that students will be familiar with the gas laws (Boyle’s law and Charles’s law).


Students’ model of gases is likely to be based on the simple particle model of matter, which they encountered first some time ago and probably revisited on a number of occasions. Some students might find it difficult to imagine an ‘ideal’ gas. They have been building up pictures in their mind of atoms, ions and molecules and have spent time discussing the nature and sizes of these particles. Now they are being asked to imagine the gas particles as zero volume. They need to understand the concept of extrapolation, in other words looking at the trend in properties of something and saying what might happen if they were extrapolating to the extreme.

There is also a school of thought that suggests it is preferable to use the ideal gas equation pV = nRT rather than to use calculations based on equations such as: P1V1/T1 = P2V2/T2.

LESSON OUTLINE

Engage and remind

Ask students to describe their model (or models) for gases and the evidence upon which they are based. Now look together at Figure 2 in Chemistry Student Book 1, Chapter 2, Section 2.3 and discuss the characteristics of an idea gas and how they differ to a real gas. These are summarised in Section 2.3, The ideal gas.

Core activities

Guide students through the effect of pressure and temperature on a gas, without dwelling on it or asking students to do any calculations, to Section 2.3, Combining Boyle’s law and Charles’s law. Discuss Avogadro’s hypothesis that equal volumes of gases contain equal number of particles. Ask students to complete task 1 on Activity sheet 2.5.1 Using the ideal gas equation.

Move from there to the ideal gas equation and explain how the molar gas constant can be determined. Ask students to read, or take them through, Section 2.3, Calculating the number of moles from the gas volume, Calculating the volume of a reactant gas and the mass of a gas product and Calculating the volume of gas produced in a reaction involving a non-gas reactant.

Ask students to complete task 2 on Activity sheet 2.5.1 Using the ideal gas equation. The final question in task 2 is quite challenging and students may need a prompt about the two temperatures and their relationship to the boiling point of water.


In the next lesson, students will carry out an experiment to determine the relative molecular mass of hexane. In preparation for this, as an out-of-class activity, students might be asked to read Chemistry Student Book 1, Chapter 2, Section 2.3, Finding the Mr of a volatile liquid and complete task 3 on Activity sheet 2.5.1 Using the ideal gas equation.

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Using the ideal gas equation

TASK 1

Boyle’s law deals with the change in the volume of a fixed mass of gas as the pressure changes.

01. What is the equation for Boyle’s law?

02. What must remain constant for the equation to apply?

Charles’s law deals with the change in the volume of a fixed mass of gas as the temperature changes.

03. What is the equation for Charles’s law?

04. What must remain constant for the equation to apply?

TASK 2

01. Write the ideal gas law equation, define each term used, and state the relevant SI units used.

02. Convert the following to pressure in Pa.

a. 100 kPa

b. 5 kPa

c. 0.2 kPa

03. Convert the following to volume in m3.

a. 200 cm3

b. 20 000 cm3

c. 100 dm3

d. 0.5 dm3

04. Convert the following to temperature in K.

a. 30 oC

b. –10 oC

c. 85 oC

d. 360 oC

05. Write the ideal gas equation with number of moles as the subject.

06. To three significant figures, how many moles of carbon dioxide are there in 800 cm3 of gas at 30 oC and 110 kPa?

07. Write the equation for the combustion of methane.

08. Calculate the volume, in m3 and dm3, of gaseous products of combustion of 0.25 mol methane at:

a. 150 kPa and 85 oC

b. 96 kPa and 180 oC

TASK 3: DETERMINING THE RELATIVE MOLECULAR MASS OF A VOLATILE LIQUID

Pentane is an alkane that boils at 36 oC. This is the structure of a pentane molecule:

In an experiment, 0.10 g of pentane was vaporised at 50 oC and 100 kPa. The volume of gas produced was 34 cm3.

01. Convert the volume of pentane at 50 oC and 100 kPa into cubic metres (m3).

02. Convert the temperature into kelvin, K.

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