PART A - John Wiley & Sonscatalogimages.johnwiley.com.au/Attachment/07314/0731405366/61_053… · EXPERIMENT 5.2 : The effect of heat on metals 15 EXPERIMENT 6.1 : Building molecular

PART A

Practical work

Experiments 3Solution preparation quantities 4

Experiments for each chapter 6

Results, answers and activities 67Each chapter contains:

• expected results for each experiment

• suggested answers to experiment questions

• additional learning activities

• internet resources.

Experiments

Solution preparation quantities 4

EXPERIMENT 1.1 : Mixing liquids 6

EXPERIMENT 1.2 : Flame colours 7

EXPERIMENT 2.1 : Making a spiral periodic table 8

EXPERIMENT 3.1 : The empirical formula of magnesium oxide 9

EXPERIMENT 3.2 : The empirical formula of a hydrated salt 10

EXPERIMENT 4.1 : Investigating calcite crystals 11

EXPERIMENT 4.2 : Ionic models 12

EXPERIMENT 4.3 : The underwater garden 13

EXPERIMENT 5.1 : Metallic trees 14

EXPERIMENT 5.2 : The effect of heat on metals 15

EXPERIMENT 6.1 : Building molecular models 16

EXPERIMENT 6.2 : Testing properties (student design) 17

EXPERIMENT 7.1 : Developing a periodic table 18

EXPERIMENT 7.2 : Trends down a group: the alkaline earth metals 22

EXPERIMENT 8.1 : Formation of esters 25

EXPERIMENT 8.2 : Breaking double bonds 26

EXPERIMENT 9.1 : Cross-linking an addition polymer to make slime 27

EXPERIMENT 9.2 : Heating plastics 28

EXPERIMENT 9.3 : Making rubber 29

EXPERIMENT 9.4 : Making nylon 30

EXPERIMENT 10.1 : Making temporary and permanent emulsions 31

EXPERIMENT 10.2 : Comparing water-in-oil and oil-in-water emulsions 32

EXPERIMENT 11.1 : Testing for solubility 33

EXPERIMENT 11.2 : Classifi cation of chemical reactions 34

EXPERIMENT 12.1 : The production of smog 36

EXPERIMENT 13.1 : Determining a solubility curve 37

EXPERIMENT 13.2 : Specifi c heat capacity 39

EXPERIMENT 13.3 : A water modelling exercise 41

EXPERIMENT 14.1 : Indicators and pH 42

EXPERIMENT 14.2 : The reactions of acids 43

EXPERIMENT 14.3 : Finding the pH of common household substances 45

EXPERIMENT 14.4 : Carbon dioxide content of fi zzy drinks 46

EXPERIMENT 15.1 : Preparation of a solution of known concentration 47

EXPERIMENT 15.2 : Stoichiometry of a reaction 48

EXPERIMENT 16.1 : Galvanic cells 49

EXPERIMENT 16.2 : Corrosion 51

EXPERIMENT 16.3 : Minimising corrosion (student design) 52

EXPERIMENT 16.4 : The reactivity of metals and their salts 53

EXPERIMENT 16.5 : Simple redox equations 54

EXPERIMENT 16.6 : Complex redox equations 55

EXPERIMENT 17.1 : Making oxygen(student design) 56

EXPERIMENT 17.2 : Preparation and properties of carbon dioxide 57

EXPERIMENT 18.1 : Measuring gas volume 59

EXPERIMENT 18.2 : The egg in the fl ask 60

EXPERIMENT 18.3 : Expanding balloon and marshmallow 61

EXPERIMENT 18.4 : Diffusion of gases 62

EXPERIMENT 19.1 : The atom economy 64

Chemistry 1: Practical work

studyon Chemistry 1 TEACHER SUPPORT KIT © John Wiley & Sons Australia, Ltd 2007

This sheet may be photocopied for non-commercial classroom use.

4

Solution preparation quantities

Chemical Formula Make up to 1 L with distilled water

agar medium galactose polymer 30 g added slowly to 1 L of H2O

warmed to 80°C, stirring constantly

aluminium nitrate — 0.1 mol L–1 Al(NO3)3.9H2O 37.5 g

aluminium nitrate — 1.0 mol L–1 375 g

barium nitrate — 0.1 mol L–1 Ba(NO3)2 26.1 g

borax — 4% NaB4O7.xH2O 40.0 g

calcium hydroxide — 0.1 mol L–1 Ca(OH)2 7.4 g

limewater 25 g, allow to settle before use

calcium nitrate — 0.1 mol L–1 Ca(NO3)2.4H2O 23.6 g

copper nitrate — 0.1 mol L–1 Cu(NO3)2.3H2O 24.2 g

copper sulfate — 0.1 mol L–1 CuSO4.5H2O 25.0 g

copper sulfate — 1.0 mol L–1 250 g

ethanoic acid — 0.1 mol L–1 CH3COOH 6 mL

iron(III) nitrate — 0.1 mol L–1 Fe(NO3)3.9H2O 40.4 g + 50 mL 2M HCl

iron(II) sulfate — 0.1 mol L–1 FeSO4.7H2O 27.8 g + 100 mL 1M H2SO4

hydrochloric acid — 0.1 mol L–1 HCl 8.75 mL

hydrochloric acid (12 mol L–1, 36%)

— 1.0 mol L–1

87.5 mL

hydrochloric acid (12 mol L–1, 36%)

— 2.0 mol L–1

175 mL

lead acetate — 2.0 mol L–1 Pb(CH3COO)2.3H2O 759 g

lead nitrate — 0.1 mol L–1 Pb(NO3)2 33.0 g

lead nitrate — 2.0 mol L–1 662 g

magnesium nitrate — 0.1 mol L–1 Mg(NO3)2.6H2O 25.6 g

magnesium nitrate — 1.0 mol L–1 256 g

magnesium sulfate — 0.1 mol L–1 MgSO4.7H2O 24.7 g

nickel nitrate — 0.1 mol L–1 Ni(NO3)2.6H2O 29.0 g

polyvinyl alcohol — 6% (C2H3OH)n

60 g added slowly to 1 L of H2O warmed to

80°C, stirring constantly

potassium dichromate — 0.1 mol L–1 K2Cr2O7 29.4 g

potassium iodide — 0.1 mol L–1 KI 16.6 g

potassium nitrate — 1.0 mol L–1 KNO3 101 g

potassium permanganate — 0.1 mol L–1 KMnO4 15.8 g

potassium thiocyanate — 0.1 mol L–1 KSCN 9.7 g

Chemistry 1: Practical work

© John Wiley & Sons Australia, Ltd 2007 studyon Chemistry 1 TEACHER SUPPORT KITThis sheet may be photocopied for non-commercial classroom use.

5

Chemical Formula Make up to 1 L with distilled water

silver nitrate — 0.1 mol L–1 AgNO3 17.0 g

silver nitrate — 2.0 mol L–1 340 g

sodium carbonate — 0.1 mol L–1 Na2CO3 11.0 g

sodium chloride — 0.1 mol L–1 NaCl 5.8 g

sodium hydroxide — 0.1 mol L–1 NaOH 4.0 g

sodium hydroxide — 1.0 mol L–1 40 g

sodium nitrate — 0.1 mol L–1 NaNO3 8.5 g

sodium dihydrogen phosphate —

0.1 mol L–1

NaH2PO4.2H2O 15.6 g

sodium sulfate — 0.1 mol L–1 Na2SO4 14.2 g

sodium sulfide — 0.1 mol L–1 Na2S.9H2O 24.0 g

starch (C6H10O5)x 10 g made into a paste with water.

Add slowly to 1 L of almost boiling water,

stirring constantly.

sulfuric acid (18 mol L–1, 98%) —

0.1 mol L–1

H2SO4 5.5 mL


0.5 mol L–1

27.5 mL


1.0 mol L–1

55 mL

zinc nitrate — 0.1 mol L–1 Zn(NO3)2.6H2O 29.7 g

zinc sulfate — 0.1 mol L–1 ZnSO4.7H2O 28.8 g

zinc sulfate — 1.0 mol L–1 288 g



6

Chemistry 1: Experiments EXPERIMENT 1.1

Mixing liquids

The particles in a liquid are generally further apart than the particles in a solid.

When two liquids are mixed their particles are able to slip into the spaces

available in each other.

AIMTo find the final volume when equal volumes of two miscible liquids are

mixed.

APPARATUS2 s 10 mL measuring cylinders

5 mL ethanol

5 mL distilled water

paper towel

2 dropping pipettes

METHOD

1. Measure 5 mL of distilled water into one of the measuring cylinders.

2. Measure 5 mL of ethanol into the second measuring cylinder.

3. Pour the ethanol into the measuring cylinder containing distilled water.

4. Observe and record the volume of the mixture.

QUESTIONS

1. What is the total volume of the two liquids?

2. Explain your observations in terms of the arrangement of particles in a

liquid.

3. If you substituted petrol for the ethanol in this experiment, would you

expect the results to be the same? Explain.


7

Chemistry 1: Experiments

Flame colours

When vaporised in the fl ame of a Bunsen burner, some metallic elements

produce a characteristic colour. The fl ame colour is the result of electrons

moving from a higher energy level or electron shell to their normal shell.

Everyday applications of fl ame colours can be seen in ‘neon’ advertising

signs (neon itself is red) and the mercury vapour and sodium vapour lamps

used in street lighting.

AIMTo observe the characteristic fl ame colours of the following metal ions: Ca2+,

Sr2+, Ba2+, Cu2+, K+, Na+, and to identify an unknown metal ion.

APPARATUS7 watch-glasses

7 nichrome or platinum wires, each attached to a holder (the wires must be

clean)

7 Bunsen burners

7 heatproof mats

2 mol L–1 solution of hydrochloric acid

small sample jars containing chlorides of the following metals: Ca2+, Sr2+, Ba2+,

Cu2+, K+, Na+, and an unknown metal ion (to be determined by the teacher)

METHOD

1. Set up seven stations, each with the pieces of equipment listed, including a

different metal salt at each station. This avoids contamination of the wire.

2. Very carefully, dip the wire into a watch-glass containing the hydrochloric

acid solution and then dip the moistened wire into the sample of the metal

salt.

3. Place the wire in the edge of the fl ame and observe the colour. Record the

result.

4. Move to a new station and repeat the procedure with a new salt and wire.

QUESTIONS

1. How can you tell that it is the metal ion and not the non-metal ion that

produces the fl ame colour?

2. Would you expect all metal ions to produce a fl ame colour? Explain.

3. Identify the unknown metal ion by its characteristic fl ame colour.

HCl is corrosive and must be

handled with extreme care.

HCl is corrosive and must be

handled with extreme care.

EXPERIMENT 1.2



8


Making a spiral periodic table

METHOD

1. Obtain a 110 cm length of 1 cm wide ticker tape (may be obtained from

your physics teacher) and mark off at 1 cm intervals.

2. Number the squares 1 to 105 and write the symbol and atomic number

of each element in each square in increasing order of atomic number, as

shown in the figure below.

1 2 3 4 5 6 7 8 9

LiH He Be B C N O

Continue up to element 105

3. Glue back-to-back squares numbered 21 to 30, 39 to

48, 57 to 70, 71 to 80 and 89 to 103.

4. Glue square 2 to one end of a pencil and wind the

ticker tape around the pencil as shown in the figure at

right.

5. Glue squares 10, 18, 36, 54 and 86 in a similar

fashion.

6. Construct a base from plasticine or another material

so that your spiral periodic table remains upright.

7. The squares of elements belonging to a particular

group in the periodic table can be given a specific

colour so that the group arrangement in your spiral

table is more evident.

QUESTIONS

1. How are elements in the same group in the periodic table arranged in your

spiral table?

2. How are the elements in the same period in the periodic table arranged in

your spiral table?

3. Discuss the advantages and disadvantages of this form of periodic table as

compared with the two-dimensional table used by chemists.

3

1

910

2

18

11

19

17

3635

37

HeH

Li

Ne Na

F

ArK

Cl

Kr Rb

Br

3

1

910

2

18

11

19

17

3635

37

HeH

Li

Ne Na

F

ArK

Cl

Kr Rb

Br

EXPERIMENT 2.1


9


The empirical formula of magnesium oxide

AIMTo deduce the empirical formula of magnesium oxide.

APPARATUSelectronic balance

Bunsen burner

crucible with lid

heatproof mat

pipeclay triangle

tripod

matches

metal tongs

10 cm strip of clean magnesium ribbon

METHOD

1. Accurately weigh the crucible and lid.

2. Roll the magnesium ribbon into a loose coil, then place into the crucible

and reweigh.

3. Place the crucible on the pipeclay triangle and tripod. Use the Bunsen

burner to heat the crucible, gently at first and then progressively more

strongly, with the lid covering about three-quarters of the crucible top.

4. Continue heating for 5 minutes with the lid off. When the magnesium has

burned, allow the crucible and contents to cool completely. (If the contents

are grey, add 4 or 5 drops of water and reheat.)

5. Cool and reweigh the crucible, lid and contents.

6. Record your results.

RESULTS

1. Calculate and record the mass of oxygen in the magnesium oxide.

2. Calculate the mole ratio of magnesium to oxygen in the compound and

hence determine the empirical formula of the compound.

QUESTIONS

1. How does your calculated empirical formula of magnesium oxide compare

with the known formula of magnesium oxide?

2. Write a balanced equation, including states, for the formation of magnesium

oxide.

3. Suggest sources of error (not your own mistakes!) in this experiment,

explaining how they may affect your empirical formula calculations.

4. Nitrogen in the air may react with magnesium. Suggest how this could

affect your empirical formula calculations.

EXPERIMENT 3.1



10


The empirical formula of a hydrated salt

Water molecules form a part of the crystal lattices of many salts. The number

of water molecules per formula unit is usually fixed. The physical properties of

the crystal may be changed by the presence of water molecules. For example,

CuSO4.5H2O is blue, whereas the anhydrous (dehydrated) form, CuSO4, is

white.

AIMTo determine the empirical formula of a hydrated salt.

APPARATUSBunsen burner

crucible with lid

electronic balance

heatproof mat

pipeclay triangle

tripod

matches

spatula

2 g hydrated magnesium sulfate (or 2 g hydrated barium chloride)

METHOD

1. Accurately weigh the crucible and lid.

2. Add the hydrated salt to the crucible and reweigh.

3. Heat the contents strongly, with the lid ajar, for 10 to 15 minutes.

4. Allow to cool and then reweigh.


RESULTS

1. Use your experimental results to calculate:

(a) number of moles of anhydrous salt

(b) number of moles of water.

2. Determine the empirical formula of the hydrated salt.

QUESTIONS

1. Write a balanced equation, including states, for the dehydration of the

salt.

2. List the sources of error (not your mistakes!) in the experiment.

EXTENSIONA student accurately weighed 2.00 g of hydrated barium sulfate into a crucible

but lost some of the hydrated salt from the crucible when she sneezed. How

would this loss in mass affect her results if she based her calculations on an

initial 2.00 g sample?

EXPERIMENT 3.2


11


Investigating calcite crystals

AIMTo observe the nature of crystal structure using calcite crystals.

APPARATUSlarge calcite crystal

reinforced blade

light hammer or block of wood

METHOD

1. Study the calcite crystal, observing the way its grain runs across its surface

in lines.

2. Put the crystal on a bench and carefully place the reinforced blade on the

crystal so that its cutting edge is running with the grain rather than across

it.

3. Give the blade a light tap with a hammer or small block of wood.

4. Observe what happens to the crystal.

QUESTIONS

1. What are the lines that can be observed on the surface?

2. Describe the appearance of the pieces that form after cleavage.

3. What does this experiment tell you about the crystal structure of calcite?

4. What generalisations can be made about ionic crystals?

Use caution when handling the

blade.

Use caution when handling the

blade.

EXPERIMENT 4.1



12


Ionic models

Chloride ions are approximately twice the size of sodium ions. When sodium

ions and chloride ions combine, an ionic lattice is formed.

AIMTo construct a model of sodium chloride and to use it to simulate some

properties of sodium chloride.

APPARATUS3 overhead projector (acetate) sheets, 267 mm s 267 mm

2 overhead projector pens (1 green, 1 orange)

plasticine

METHOD

1. Use a green circle of radius 3 cm and an orange circle of radius 1.5 cm to

represent the chloride ion and the sodium ion respectively.

2. Draw the arrangements of ions, as in the figure below, on the three overhead

projector sheets, with one arrangement per sheet.

3. Overlay the three sheets with the second one in the middle. Use four balls

of plasticine, one at each corner, to separate sheets 1 and 2, and another

four balls of plasticine to separate sheets 2 and 3. This simulates a three-

dimensional model of the crystal.

– –+ +

–– ++

– –+ +

–– ++

sheet 1

–– ++

– –+ +

–– ++

– –+ +

sheet 2

plasticine

– –+ +

–– ++

– –+ +

–– ++

sheet 3

4. Viewing from above, describe what you see in terms of the coordination

number of sodium chloride.

5. Move the second sheet so that all the positive ions are superimposed when

you look downward through your ‘crystal’. This movement is simulating

an applied force. Explain what will happen to the crystal. What property of

ionic crystals are you modelling when you do this?

6. Another property of ionic crystals is that they conduct electricity in the

molten and aqueous states. How could you simulate this property using

your model?

QUESTIONS

1. Why is the radius of the chloride ion larger than the radius of the sodium ion?

2. This model is useful but has limitations. How adequately does the model

illustrate the structure and properties of ionic compounds?

EXTENSION

1. Using existing ionic models in the laboratory, discuss the strengths and

limitations of these models.

2. Compare these models with the model you have just constructed. Which

model best describes ionic structure? Justify your choice.

EXPERIMENT 4.2


13


The underwater garden

AIMTo grow crystals.

APPARATUS500 mL water glass containing sodium silicate solution of density 1.1 g L 1

large beaker or rectangular trough

large crystals or walnut-sized lumps of any of the following solids:

CuSO4

CoCl2CaCl2MgSO4

Al2(SO4)3

KAl(SO4)2

FeSO4

FeCl3NiSO4

Sr(NO3)2

METHOD

1. Fill the beaker or trough with the water glass, and place it on a stable

surface safe from shocks or vibrations.

2. Carefully drop the crystals onto the bottom of the container.

3. Leave the container for a couple of days. (Some crystals will start to grow

immediately.)

QUESTIONS

1. Describe the colours obtained from each crystal.

2. Which crystal grew the fastest?

3. Write the correct chemical name for each of the formulas listed in the

‘Apparatus’ section of this experiment.

EXPERIMENT 4.3



14


Metallic trees

Metal crystals may be made by displacement reactions. In this activity, you may

make metallic trees from either silver crystals (more expensive) or lead crystals

(less spectacular). The crystals may also be grown in agar medium in a fl at

petri dish to enable you to examine them more closely under the microscope.

AIMTo produce metal crystals of silver or lead.

APPARATUScotton

glass rod

beaker

metal cutters

For silver tree:

square of heavy copper foil, 4 cm s 4 cm

2 mol L–1 solution of silver nitrate

For lead tree:

heavy zinc foil

2 mol L–1 solution of lead nitrate or lead acetate with a few drops of nitric acid

added to hasten the process

For preparation of agar medium:

beaker

petri dish with lid

agar

METHOD

1. Cut the shape of a pine tree out of the copper (for the silver tree) or the zinc

(for the lead tree).

2. Fill the beaker with the silver nitrate solution (for the silver tree) or with

the solution of lead acetate and nitric acid (for the lead tree).

3. Pass the cotton through a hole made in the tip of the tree, and suspend the

tree from a glass rod resting across the top of the beaker.

4. Results may take several hours or days, depending on the thickness of the

metallic tree.

5. If you wish to examine the metal crystals in agar, prepare a mixture of agar

and warm solution in a beaker (make sure the agar has dissolved). Pour the

mixture into the petri dish. When nearly set, immerse the tree in the agar

medium and cover with the petri dish lid.

RESULTSDraw the crystals as they appear under a stereomicroscope.

QUESTIONS

1. What metal displaced silver to make the silver trees?

2. What metal displaced lead to make the lead trees?

Lead salts are harmful if taken

internally. Do not inhale or

swallow. Silver nitrate is caustic

and will stain your hands. Avoid

contact with eyes and skin. Use

gloves. This experiment may be

performed in a fume cupboard.

Lead salts are harmful if taken

internally. Do not inhale or

swallow. Silver nitrate is caustic

and will stain your hands. Avoid

contact with eyes and skin. Use

gloves. This experiment may be

performed in a fume cupboard.

EXPERIMENT 5.1


15


The effect of heat on metalsAIMTo investigate the malleability of quenched and annealed metals.

APPARATUS3 steel needles (e.g. sewing needles) or 3 equal lengths of steel wire

Bunsen burner and heatproof mat

2 pairs of tongs

500 mL container of iced water

safety glasses

METHOD

1. Bend the fi rst steel needle as far as you can. Note its malleability and

strength.

2. Using tongs, heat a second steel needle until it is red hot.

3. Immediately immerse the needle in iced water.

4. Bend this second needle as far as you can. Note its malleability and

strength.

5. Heat the third steel needle until it is red hot.

6. Allow the needle to cool slowly.

7. Now bend this needle as far as you can. Note its malleability and

strength.

QUESTIONS

1. What is the effect of quenching a metal?

2. What is the effect of tempering a metal?

3. Why have the properties of the metal changed?

4. Use the ball bearing model to explain why the properties of the metal have

altered.

5. In what ways does the ball bearing model fail to explain the change in

properties of the modifi ed metal?

Safety glasses should be worn in

case needles break.

Safety glasses should be worn in

case needles break.

EXPERIMENT 5.2



16


Building molecular models

AIMTo construct molecular models of H2, Cl2, HCl, H2O, NH3, CH4, CH2F2, H2S,

C2H6, C2H5Cl and CH3CH2OH.

APPARATUSplastic bag containing the following items from a molecular building kit:

35 hydrogen atoms

4 chlorine atoms

2 oxygen atoms

1 nitrogen atom

8 carbon atoms

2 fluorine atoms

1 sulfur atom

42 green straws, each 2.5 to 3 cm long

16 white straws, each 2.5 to 3 cm long

METHOD

1. Draw the structural formula of each of the molecules listed in the

experimental aim.

2. Predict the shape of each molecule.

3. Construct the molecules using green straws to represent bonding pairs of

electrons and white straws to represent non-bonding pairs of electrons.

4. Verify, from your three-dimensional model, that your earlier predicted

shape is correct.

QUESTIONS

1. Why do the bonding pairs in CH4 arrange themselves in a tetrahedral

shape?

2. Explain the effect of lone pairs on the shape of molecules such as H2S and

NH3.

EXPERIMENT 6.1


17


Testing properties(student design)AIMTo design and perform an experiment to test the properties of a number of

substances.

APPARATUSStudents to select own apparatus. Substances to be tested may include

aluminium, copper, zinc, magnesium, sodium chloride, copper sulfate, water,

solid iodine, sulfur and sugar. (Your teacher may wish to include others.)

METHODDesign your own method for testing the properties of the substances provided.

Properties to be tested may include:

(a) electrical conductivity

(b) heat conductivity

(c) malleability

(d) lustre

(e) fl ammability

(f) solubility.

Check with your teacher before proceeding.

QUESTIONClassify the substances you have tested as ionic, metallic or covalent, based

on their observed properties. Justify your choices for the classifi cations you

make.

Solid iodine is toxic; handle it in

a fume cupboard.

Solid iodine is toxic; handle it in

a fume cupboard.

EXPERIMENT 6.2



18


Developing a periodic table

Just over a century ago only two-thirds of the elements known today had been

identified and their atomic masses were not known with certainty. Nevertheless,

chemists were attempting to find patterns among the elements in order to

arrange the elements into a periodic table.

AIMTo understand how the theory of the periodicity of elements was developed by

looking for patterns in the properties of the elements known to chemists such

as Meyer and Mendeleev in 1869.

APPARATUSgraph paper

data sheet for elements as known in 1869 (table 7A)

METHODPart A: Physical properties of the elements

Although the masses of the atoms of the elements were not known in 1869,

the relative atomic mass of each atom, based on a scale where hydrogen = 1,

was known. This mass was called the atomic weight of the atom. (Hydrogen

atoms at that time were assigned a mass of 1 unit. An atom that weighed,

for example, seven-and-a-half times as much as a hydrogen atom would be

assigned an atomic weight of 7.5.)

Since the densities of many elements were known, the relative atomic volume,

and thus the relative sizes of atoms, could be calculated using the formula:

atomic weightatomic volume = density

Enter the data provided in table 7A into a spreadsheet. Generate a graph

of atomic weight (horizontal axis) vs atomic volume (vertical axis) on your

computer for the elements in the list.

Part B: Chemical properties of the elements

The formula of the simplest compound that each element forms with hydrogen

has been listed in table 7A. The valency of an element may be considered to

be equivalent to the number of atoms of hydrogen that will combine with one

atom of that element. Using this definition, calculate the valency of the first 17

elements listed in table 7A, then plot their valency (vertical axis) against atomic

mass in increasing order (horizontal axis).

QUESTIONSPart A: Physical properties of the elements

1. (a) Describe the shape of the graph produced.

(b) Can you explain why the atomic volume does not continue to rise with increasing atomic mass?

2. (a) Identify the elements forming the five main peaks on the graph.

(b) Refer to table 7A and describe any similarities between these five elements.

Part B: Chemical properties of the elements

1. Does the valency vary periodically?

2. State the relationship between the valencies of the five related elements in

part A.

EXPERIMENT 7.1

Chemistry 1: ExperimentsCONTINUED


19

TABLE 7A Data sheet for elements as known in 1869

Element Symbol

Atomic weight (1869)2

Density (g mL–1)

Atomic volume

Hydrogen compound Valency Class1

Melting point (ºC)

hydrogen H 1 — — H2 1 N ?3

lithium Li 7 0.53 13 LiH M 180

beryllium Be 9.4 1.8 5.2 BeH2 M 1280

boron B 11 2.5 4.4 B2H6 N 2030

carbon C 12 2.26 5.3 CH4 N >3500

nitrogen N 14 — — NH3 N ?3

oxygen O 16 — — H2O N ?3

fluorine F 19 — — HF N ?3

sodium Na 23 0.97 24 NaH M 98

magnesium Mg 24 1.74 14 MgH2 M 650

aluminium Al 27.4 2.70 10 AlH3 M 660

silicon Si 28 2.4 12 SiH4 N 1410

phosphorus P 31 1.82 17 PH3 N 44

sulfur S 32 2.07 15 H2S N 113

chlorine Cl 35.5 — — HCl N ?4

potassium K 39 0.86 45 KH M 64

calcium Ca 40 1.55 26 CaH2 M 838

titanium Ti 50 4.5 11 M 1670

vanadium V 51 5.96 8.6 M 1900

chromium Cr 52 7.1 7.3 M 1900

manganese Mn 55 7.2 7.6 M 1250

iron Fe 56 7.86 7.1 M 1540

nickel Ni 59 8.90 6.6 M 1450

cobalt Co 59 9.0 6.5 M 1490

copper Cu 63 8.92 7.1 CuH M 1083

zinc Zn 65.2 7.14 9.1 ZnH2 M 419

arsenic As 75 5.7 13 AsH3 N 613

selenium Se 79.4 4.7 17 H2Se N 217

bromine Br 80 3.12 26 HBr N –7

rubidium Rb 85.4 1.53 56 RbH M 770

strontium Sr 87.6 2.6 34 SrH2 M 770

yttrium Y 89 5.51 16 M 1500

zirconium Zr 90 6.4 14 M 1850

niobium Nb 91 8.4 11 M 2420

molybdenum Mo 96 10.2 9.4 M 2610

rhodium Rh 104.4 12.5 8.3 M 1970

ruthenium Ru 104.4 12.2 8.5 M 2300

palladium Pd 106.8 12 8.9 M 1550

silver Ag 108 10.5 10 M 961

cadmium Cd 112 8.6 13 CdH2 M 321

indium In 115 7.3 16 InH3 M 156

tin Sn 118 5.8 20 SnH4 M 232

antimony Sb 122 6.0 20 SbH3 N 631

iodine I 127 4.93 26 HI N 114

tellurium Te 128? 6.1 21 H2Te N 450

cesium Cs 133 1.90 70 CsH M 29

barium Ba 137 3.5 39 BaH2 M 714

Notes:

1 Classes of elements are based on their conductivities

(M = metal, N = non-metal).

2 ‘Atomic weight’ is relative atomic mass on a scale on which the mass

of the hydrogen atom is taken as 1.

3 The elements H, N, O and F were known only as gases in 1869.

4 The melting point of chlorine, Cl, was unknown, though its boiling

point (–34.7°C) was known.

EXPERIMENT 7.1




20

METHODPart C: A periodic table

The organisation of elements into a periodic table groups elements with similar

physical and chemical properties.

1. Copy the table below and label the vertical columns (groups I, II . . . VII)

and horizontal rows (periods 1, 2 . . .) as shown.

Periodic table of the elements known in 1869I

Period 1 H II III IV V VI VII

Period 2

Period 3 IIIB IVB VB VIB VIIB VIIIB IB IIB

Period 4

Period 5

Period 6

Note: This activity uses an old system of group numbering with Roman numerals. This system was replaced in the

late 20th century by numbering with Arabic numerals, as shown in the textbook.

2. Hydrogen has been placed in group I of period 1 of the periodic table.

Write the symbols for the next seven elements in order from table 7A in

period 2, starting from the left side of the table. Repeat this process for the

next seven elements and period 3.

3. Look at the valencies of the elements in each group and complete the

following table:

Group

numberI II III IV V VI VII

Valency of

elements

4. Transfer the next two elements from table 7A to the start of period 4. Does

the vertical valency relationship still hold?

5. We can extend the group I elements from the information obtained in parts

A and B.

(a) List the three elements located so far in group I.

(b) List the five related elements identified in part A.

(c) Assuming a connection, which element should begin: (i) period 5? (ii) period 6?

6. (a) Which element should be placed in group II of:

(i) period 5? (ii) period 6?

(b) Do their valencies match the others of group II?

7. (a) From table 7A, which element precedes the one at the beginning of

period 5?

(b) Assuming that this element belongs to period 4, it logically would be placed in group VII of the periodic table. Does its valency match that of the other elements in this group?

8. Place the two elements preceding this one in period 4. Do their valencies

fit the established pattern?

EXPERIMENT 7.1



21

9. (a) Continuing backward along period 4, which element is next in the list?

(b) Considering its valency, explain why this element is a member of group IIB rather than group IV.

(c) Which element would be located in period 4, group IB?

10. (a) Refer to table 7A and determine which element, based on its atomic

weight, would be expected to fit into period 5, group VII.

(b) Considering its valency, explain why this element would fit better into group VI than into VII.

(c) Which element should then be placed into group VII?

METHODPart D: Making predictions using the periodic table

The table you have constructed has blank spaces for the unknown elements in

groups III and IV in period 4.

1. Suggest a reason for these vacancies.

2. It is possible to predict the atomic weights of these unknown elements by

taking a mathematical average, as explained below.

(a) Complete the following table by inserting atomic weights for the required elements and then calculating their mean.

Atomic weights Group V Group VI Group VII

period 3 element

period 5 element

average of these two

period 4 element

(b) Summarise any relationship you notice from the above atomic weights.

(c) Apply this technique to the unknown group III and group IV elements and estimate their atomic weights.

(d) Refer to a modern periodic table and identify these unknown elements. How do their estimated atomic weights compare with their actual atomic masses? Can you explain any differences?

3. Elements in the periodic table may be classified as metals or non-metals.

Use the data table 7A to determine:

(a) where the metals are located in the periodic table

(b) where the non-metals are located in the periodic table

(c) whether the metallic character of the elements increases or decreases down a group from top to bottom

(d) whether the metallic character of the elements increases or decreases across a period from left to right.

4. In 1894 two new elements, helium (atomic weight 4) and argon (atomic

weight 40), were discovered. Both were found to be colourless gases of

low boiling point that formed no compounds with hydrogen or any other

element.

(a) On the basis of these properties, does the creating of a new group seem justified? Explain your answer.

(b) On the basis of their atomic weights, where do these elements fit into your table?

EXPERIMENT 7.1



22


Trends down a group: the alkaline earth metals

All group 2 elements (Be, Mg, Ca, Sr, Ba and Ra) have two electrons in their

outer shell. Both families of elements lose their outer shell electrons readily to

form stable ions. The compounds of group 2 elements have many important

uses such as CaSiO3 for making glass, Mg(OH)2 as an antacid for upset

stomachs, and CaCO3 for making chalk and cement. In this experiment you

will investigate the trends in some of the properties of elements and compounds

of group 2.

AIM To investigate some of the physical and chemical properties of four group 2

elements (Mg, Ca, Sr and Ba) and compounds:

the appearance, hardness and solubility of the elements

the reaction of elements with acids

the solubility of the oxides, hydroxides, chlorides, sulfates, nitrates and

carbonates of the elements

the thermal stability of the nitrates and carbonates of the elements.

APPARATUS500 mL beakers, test tubes, test-tube rack, matches, splinter

magnesium, calcium, strontium and barium

oxides of magnesium, calcium, strontium and barium

nitrates of magnesium, calcium, strontium and barium

carbonates of magnesium, calcium, strontium and barium

1 mol L 1 solutions of Na2SO4, NaOH, NaCl, NaNO3 and Na2CO3

1 mol L 1 solutions of MgCl2, CaCl2, SrCl2 and BaCl22 mol L 1 hydrochloric acid

phenolphthalein

universal indicator

METHOD

1. Set up a table as below.

Element Appearance Hardness

Reaction

with HCl

Solubility

in water

Yes/

No

Acidic/

Basic

Mg

Ca

Sr

Ba

2. Describe the appearance of each of the four elements. Scratch the surface

of each element. Record your observations.

3. Put a few drops of phenolphthalein into a beaker of water. Place a small

(rice-grain size) piece of magnesuim into the water. Record what happens.

Repeat for the other three elements.

4. Place 2 mL of HCl into a test tube. Add a small piece of magnesium. Test

the gas produced for the presence of hydrogen gas. The test for hydrogen

gas is the ‘pop’ test with a burning splinter. Record your results. Repeat for

the other three elements.

•

•

•

•

EXPERIMENT 7.2



23


Compound

Solubility

Yes/No pH Acidic/Basic

MgO

CaO

SrO

BaO

6. Put a small amount (rice-grain size) of MgO into half a test tube of water.

If the compound dissolves in water, test the pH with universal indicator.

Record your results. Repeat for the other three elements.


Na2SO4(aq) NaCl(aq) NaOH(aq) Na2CO3(aq) NaNO3(aq)

MgCl2(aq)

CaCl2(aq)

SrCl2(aq)

BaCl2(aq)

8. Set up five test tubes in a test-tube rack.

(a) Place 2 mL of Na2SO4 into the first test tube.

(b) Place 2 mL of NaCl into the second test tube.

(c) Place 2 mL of NaOH into the third test tube.

(d) Place 2 mL of Na2CO3 into the fourth test tube.

(e) Place 2 mL of NaNO3 into the last test tube.

(f) Add 5 drops of MgCl2 into each of the five test tubes of solution. Note if a precipitate forms and record your results in a table like the one above.

9. Repeat step 8 for CaCl2, SrCl2 and BaCl2.

10. Set up a table like the one below.

Mg(NO3)2 Ca(NO3)2 Sr(NO3)2 Ba(NO3)2 MgCO3 CaCO3 SrCO3 BaCO3

What happens

when heated?

Gas present

(O2 or CO2 or

none)?

11. Place a small amount of Mg(NO3)2 into a dry test tube. Heat over a Bunsen

burner flame. Test the gas produced for the presence of oxygen gas. Oxygen

gas should rekindle a glowing splinter. Record your results. Repeat for the

nitrates of the other three elements.

12. Repeat step 11 using carbonates of the four group 2 elements.

EXPERIMENT 7.2




24

QUESTIONS

1. What physical trends are evident down group 2?

2. (a) Arrange the metals in order of increasing reactivity with water.

(b) Predict the reactivity of group 1 elements with water. Explain your prediction.

3. Arrange the metals in order of increasing reactivity with acid.

4. Arrange the metal oxides in order of increasing reactivity with water.

Comment on the trend of their acidity.

5. Write the chemical equation for reactions where a precipitate is formed.

6. As you move down the group, what happens to the solubility of the sulfates,

chlorides, hydroxides, nitrates and carbonates?

7. Comment on the results that you have obtained for the thermal stability of

the nitrates and carbonates of group 2 elements.

EXPERIMENT 7.2


25


Formation of esters

When carboxylic acids are added to alcohols in the presence of a dehydrating

agent, such as concentrated sulfuric acid, an ester may be formed. An ester is

characterised by a fruity smell.

TABLE 8A Odours of some common esters

Carboxylic acid

component of ester

Alcohol component

of ester

Odour of ester

produced

ethanoic acid ethanol nail polish remover

ethanoic acid pentanol pears

ethanoic acid pentan-2-ol bananas

ethanoic acid butanol raspberries

ethanoic acid butan-2-ol strawberries

salicylic acid methanol oil of wintergreen

AIMTo prepare esters.

APPARATUSdropper bottles containing ethanol,

ethanoic acid (acetic acid), pentanol

(amyl alcohol) methanol and

concentrated sulfuric acid

1 g salicylic acid

10 mL measuring cylinder

test tubes

250 mL beaker

one-hole rubber stopper

50 cm length of 6 mm glass tubing

(air condenser)

retort stand and clamp

hotplate and safety glasses

METHOD

1. Set up a 250 mL beaker two-thirds full of water on a hotplate. The beaker

will serve as a water bath.

2. Pour 2 mL of ethanol into a test tube.

3. Add 2 mL of acetic acid (glacial ethanoic acid).

4. Add 10 drops of concentrated sulfuric acid to the mixture in the test tube.

5. Use a retort stand to clamp the test tube and insert the stopper and glass

tubing into the test tube.

6. Lower the test tube into the water bath and heat the water bath until the

reaction mixture starts to bubble slowly. Continue moderate boiling for

about fi ve minutes. (Heating should not be done over a naked fl ame.)

7. Allow the test tube to cool for a few minutes. Remove the condenser and

carefully smell the odour of the ester formed. (Precaution: use your hand

to waft the odour towards your nose. Do not smell directly.)

8. Repeat the procedure using amyl alcohol (pentanol) instead of ethanol.

9. Repeat the procedure using 1 g of salicylic acid instead of acetic acid and

methanol instead of ethanol.

QUESTIONS

1. Compare the odour of each ester that you produced with the expected odour

from the table.

2. What was the role of the sulfuric acid in this experiment?

3. Name each ester produced.

Concentrated sulfuric acid is

corrosive. Avoid contact with

skin and eyes. Wear gloves

and safety glasses. Wash spills

immediately with copious

quantities of water.

Concentrated sulfuric acid is

corrosive. Avoid contact with

skin and eyes. Wear gloves

and safety glasses. Wash spills

immediately with copious

quantities of water.

EXPERIMENT 8.1



26


Breaking double bonds

If a halogen is added to an alkene, it will break the double bond and form a

colourless haloalkane. If a halogen is added to an alkane, no reaction will take

place and the colour of the halogen will not change.

AIMTo investigate the relative reactivity of the alkenes.

APPARATUS3 mL cyclohexane

3 mL cyclohexene

dropper bottle containing bromine in trichloroethane

teat pipette

2 test tubes

METHOD

1. Place the cyclohexene in one test tube and the cyclohexane in the other.

2. Add fi ve drops of bromine in trichloroethane to each test tube.

3. Observe any colour changes.


RESULTSDescribe any colour changes.

QUESTIONS

1. What is the difference between an alkane and an alkene?

2. What is happening in each solution? Explain your results.

EXTENSIONDevise an experiment that would demonstrate that margarine is unsaturated.

It is recommended that this

experiment be done as a teacher

demonstration. It should be

carried out in a fume cupboard.

Bromine is an irritant.

Fumes of cyclohexane and

cyclohexene are dangerous.

Teachers may prefer to do

the extension activity rather

than performing the original

experiment.

It is recommended that this

experiment be done as a teacher

demonstration. It should be

carried out in a fume cupboard.

Bromine is an irritant.

Fumes of cyclohexane and

cyclohexene are dangerous.

Teachers may prefer to do

the extension activity rather

than performing the original

experiment.

EXPERIMENT 8.2


27


Cross-linking an addition polymer to make slime

Slime is made from a linear polymer called polyvinyl alcohol and borax, which

contains the borate ion. The monomer used is vinyl alcohol or hydroxyethene.

Vinyl alcohol has the following structure:

C — C

—

—

—

—

H OH

H H

—

Vinyl alcohol undergoes addition poly merisation to form the linear polymer

poly vinyl alcohol:

. . . C — C — C — C — C — C . . .

—OH

H

—

—H

H

—

—OH

H

—

—H

H

—

—OH

H

—

—H

H

—

Borax has the following structure:

HO

B–

—

—

—

—HO

OH

OH

The polar —OH group on the polymer and the —OH group on the borate

ion undergo hydrogen bonding. The linear chains are thus held together in a

loose network with water trapped between them. The result is a gel-like slime

in which the hydrogen bonds can break and reform easily.

AIMTo make slime.

APPARATUS100 mL beaker

paddle pop stick

50 mL s 6% solution of polyvinyl alcohol

10 mL s 4% solution of borax

food dye

METHOD

1. Pour the polyvinyl alcohol into the beaker and add a few drops of the food

dye.

2. Add the borax solution and stir with the paddle pop stick. It will take a few

minutes for the slime to appear.

QUESTIONS

1. Describe the properties of the slime produced.

2. How can the properties of the slime be explained by its structure and

bonding?

Although the slime can be

handled, wash your hands

afterward and do not touch

your mouth during and after

handling it.

Although the slime can be

handled, wash your hands

afterward and do not touch

your mouth during and after

handling it.

EXPERIMENT 9.1



28


Heating plastics

AIMTo heat polymer samples and classify them as either thermosetting or

thermoplastic polymers.

APPARATUSBunsen burner and heatproof mat

tongs

a collection of plastics

safety glasses

METHOD

1. In a fume cupboard, hold a small piece of plastic in tongs and heat it,

gently at fi rst, above the fl ame. If no effect is noticed after a minute, heat

it more strongly in the fl ame.

2. Repeat with the other plastics.

3. Tabulate your results under the headings ‘Thermosetting plastics’ and

‘Thermoplastics’, noting the following properties:

(a) ease of melting

(b) ease of burning

(c) smoke colour

(d) colour of fl ame

(e) odour

(f) nature of the residue.

tongs

Bunsen burner

piece of plastic

QUESTIONS

1. How do thermosetting and thermoplastic polymers differ?

2. Why is the incineration of ‘plastics’ considered to be dangerous?

3. What methods of plastics disposal are in use in your community?

This experiment should be done

in a fume cupboard.

This experiment should be done

in a fume cupboard.

EXPERIMENT 9.2


29


Making rubber

AIMTo make a rubber ball from latex, and to observe some properties of natural

rubber.

APPARATUS40 mL latex

20 mL s 2 mol L 1 HCl

100 mL beaker

glass stirring rod


disposable gloves

safety glasses

METHOD

1. Place latex in the beaker.

2. Slowly add 15 to 20 mL of HCl, stirring. Stop adding the acid when the

polymer forms.

3. Form the polymer into a ball. Wearing gloves, wash the rubber ball

thoroughly in plenty of cold water. Squeeze gently to ensure that all the

acid is removed.

4. Remove the ball from the beaker and test the elasticity of the rubber.

5. Using tongs, heat the glass rod over the Bunsen burner and place it on the

rubber. Observe what happens.

RESULTSWrite a report describing your observations.

QUESTIONS

1. Would you classify the polymer produced in this experiment as thermosetting

or thermoplastic?

2. (a) The monomer in natural rubber is isoprene or 2-methylbuta-1,3-diene.

Draw structural and semi-structural formulas for isoprene.

(b) Write an equation for the chemical reaction in which isoprene is polymerised to form natural rubber.

3. Is isoprene a saturated or unsaturated hydrocarbon? Explain your answer.

4. Is natural rubber formed by addition or condensation polymerisation?

Explain your answer.

5. Explain the properties of the rubber you tested in terms of chemical

bonding.

It is important to stir the solution

well or the acid may become

trapped in the rubber as it forms.

It is important to stir the solution

well or the acid may become

trapped in the rubber as it forms.

EXPERIMENT 9.3



30


Making nylonAIMTo make nylon 6:6.

APPARATUS200 mL beaker

100 mL beaker

4.4 g 1,6-diaminohexane

1 mL adipyl chloride in hexane

tweezers

1 mL s 1.0 mol L 1 sodium hydroxide solution

dye (optional)

disposable gloves

METHOD

1. In the 200 mL beaker, prepare a solution of 4.4 g of 1,6-diaminohexane

in 50 mL of distilled water. Add a few drops of 1.0 mol L 1 sodium

hydroxide.

2. Dissolve 1 mL of adipyl chloride in 50 mL of hexane in a 100 mL beaker.

(Add dye to the solution if desired.)

3. Carefully add the hexane solution to the water solution. The hexane solution

must be poured on top of the water solution so that the two do not mix.

4. A small amount of solid will form at the interface between the two solutions.

Use the tweezers to pull out some of the solid and carefully wind the thread

around the tweezers or around another small beaker. Keep doing this until

the solutions are used up.

QUESTIONS

1. Explain why nylon 6:6 is a condensation polymer.

2. Draw the structure of the monomers from which the polymer nylon 6:6 is

derived.

3. Write an equation, using structural formulae, to show the formation of

nylon 6:6.

The substitution of decan-dioxyl

chloride (sebacoyl chloride) for

the adipyl chloride dissolved in

1,2-dichloroethane will result in

the formation of nylon 6:10. The

same procedure may be used.

Some fi nd this polymer slightly

stronger than nylon 6:6.

The fumes given off by some

of the reagents used in this

experiment are unpleasant and

may cause discomfort. Skin

irritations may occur if reagents

touch the skin. Use disposable

gloves and masks. Carry out the

experiment in a fume cupboard.

Some teachers may prefer to do

this as a demonstration.

The substitution of decan-dioxyl

chloride (sebacoyl chloride) for

the adipyl chloride dissolved in

1,2-dichloroethane will result in

the formation of nylon 6:10. The

same procedure may be used.

Some fi nd this polymer slightly

stronger than nylon 6:6.

The fumes given off by some

of the reagents used in this

experiment are unpleasant and

may cause discomfort. Skin

irritations may occur if reagents

touch the skin. Use disposable

gloves and masks. Carry out the

experiment in a fume cupboard.

Some teachers may prefer to do

this as a demonstration.

EXPERIMENT 9.4


31


Making temporary and permanent emulsions

AIMTo make and compare several temporary and permanent emulsions.

APPARATUS20 mL cooking oil

20 mL water

2 s 100 mL beakers and labels

blue food colouring (or other water-soluble dye)

Sudan(III) dye (or other oil-soluble dye)

egg yolk

detergent

stirring rod

stoppered flask

METHOD

1. Pour 20 mL cooking oil into each of two 100 mL beakers, labelled ‘A’

and ‘B’.

2. Add 20 mL water to each of the two beakers.

3. Add 5 drops of blue food colouring to beaker A. Mix contents of the beaker

with a stirring rod and record your results.

4. Sprinkle a small amount of Sudan(III) dye into beaker B. Mix contents of

the beaker with a stirring rod and record your results.

5. Add the contents of beaker A to the contents of beaker B. Mix well and

record your results.

6. Pour the contents of beaker B into a stoppered flask. Shake well and

immediately divide the temporary emulsion that forms into beakers A and B.

7. Add a teaspoonful of egg yolk into beaker A and mix well. Record your

observations and determine whether a temporary or permanent emulsion

has formed.

8. Add a teaspoonful of detergent into beaker B and mix well. Record your

observations and determine whether a temporary or permanent emulsion

has formed.

QUESTIONS

1. At which points in the experiment was a temporary emulsion formed?

Explain your decisions.

2. At which points in the experiment was a permanent emulsion formed?

Explain your decisions.

3. Draw a fully labelled diagram to show how an emulsion was formed using

the egg yolk or detergent.

4. Consider the results you obtained in step 5.

(a) Explain your observations.

(b) Predict what would happen if you shook the contents of this beaker in a stoppered flask.

5. Compare the results you obtained in steps 7 and 8. Determine which

emulsion is the most stable and explain why some emulsions may be less

permanent than other emulsions.

EXPERIMENT 10.1



32


Comparing water-in-oil and oil-in-water emulsionsAIMTo make, and compare the properties of, water-in-oil and oil-in-water

emulsions.

APPARATUS4 petri dishes and labels

water

cooking oil

egg yolk

dropping pipette

4 toothpicks

blue food colouring (or other water-soluble dye)

Sudan(III) dye (or other oil-soluble dye)

multimeter

METHOD

1. Rinse two petri dishes under cold water, leaving only a few drops of water

in each dish. Label the dishes ‘A’ and ‘B’, and fill their bottoms with oil.

2. Label two petri dishes ‘C’ and ‘D’. Add two drops of cooking oil to each

dish and fill their bottoms with water.

3. Use a dropping pipette to add two drops of blue food colouring to petri

dishes A and C. Mix each well with a toothpick and record your observ-

ations. (Reserve a different toothpick for each of the four petri dishes.)

4. Sprinkle a small amount of Sudan(III) dye into petri dishes B and D. Mix

each well with a toothpick and record your observations.

5. Add a small amount of egg yolk to each of the four petri dishes. Mix well

and record your observations.

6. Use the multimeter to test the resistance, in ohms, of each of the emulsions

you have formed. Set the multimeter to its highest resistance range, place

the probes about 2 cm apart into the emulsion and record the resistance.

QUESTIONS

1. Compare and explain the results you obtained in steps 3 and 4 with those

in step 5.

2. Draw diagrams which show the arrangement of the molecules in petri

dishes A and C after the egg yolk was added.

EXTENSIONCollect samples of emulsions (e.g. mayonnaise, cosmetics, yoghurt, butter or

margarine) and test their resistance and appearance following the addition of

food colouring and Sudan(III) dye. Classify each emulsion as an oil-in-water

emulsion or a water-in-oil emulsion.

Note: Resistance is inversely

proportional to the conductivity

of a substance.

Note: Resistance is inversely

proportional to the conductivity

of a substance.

EXPERIMENT 10.2


33


Testing for solubility

By adding solutions of different salts to one another, we can test the rules for

solubility. As all nitrates are soluble, we can use a variety of nitrates to produce

a positive ion (cation), and as all sodium salts are soluble, we will use a number

of sodium salts to produce a negative ion (anion).

AIMTo investigate the formation of precipitates in aqueous solution and gain skill

in writing ionic equations to represent them.

APPARATUStest tubes

dropper bottles containing 0.1 mol L–1 solutions of any of the following nitrate

salts (cations): Ca2+, Na+, Mg2+, Al3+, Ba2+, Fe3+, Pb2+, Cu2+, Zn2+, Ag+

dropper bottles containing 0.1 mol L–1 solutions of any of the following sodium

salts (anions): Cl–, CO32–, OH–, PO4

3–, S2–, SO42–

METHOD

1. Draw a table showing all possible combinations of the nitrate and sodium

solutions that you will be using.

2. Predict which of the nitrate and sodium solutions will produce a

precipitate.

3. Mix a small quantity (about 2 mL) of the first nitrate solution with a similar

amount of the first sodium solution and record your observations.

4. Repeat step 3 for all nitrate and sodium solution combinations.

RESULTSComplete your table to show your results.

QUESTIONS

1. Were your predictions about the formation of precipitates accurate?

2. Write ionic equations for those reactions in which a precipitate formed.

3. Note any discrepancies and try to find out if these are due to the concentration

of the solutions being used.

EXPERIMENT 11.1



34


Classifi cation of chemical reactionsAIMTo investigate different types of chemical reactions, classify them and write

balanced equations to represent them.

APPARATUSsmall piece of magnesium metal

tongs and tweezers


candle

2 s 250 mL beakers

Hofmann voltameter

red litmus paper

small piece of calcium

small piece of zinc

3 mL s 1 mol L 1 hydrochloric acid solution

4 cm piece of copper wire

3 mL s 0.1 mol L 1 silver nitrate solution

dropping bottle of phenolphthalein indicator

2 mL s 1 mol L 1 sodium hydroxide solution

2 mL s 0.5 mol L 1 sulfuric acid solution

5 mL s 0.1 mol L 1 nickel(II) nitrate solution

5 mL s 0.1 mol L 1 sodium carbonate solution

6 test tubes

dropper, wax taper and wooden splint

METHODThere will be eight stations around the room, one for each reaction to be

investigated. For each station, follow the instructions that follow and clean up

before moving on to the next station.

Station 1: Use tongs to burn a small piece of magnesium metal in the fl ame of

the Bunsen burner. Hold the burning magnesium over a beaker or mat.

Station 2: Light the candle. Hold a beaker full of cold water over the fl ame (not

too close or soot will form on the bottom). Observe any moisture formed on the

sides and bottom of the beaker.

Station 3: Observe the Hofmann voltameter; this shows the electrolytic

decomposition of water. Note the relative volumes of the gases observed. (The

teacher may collect the gases and test for identifi cation. For hydrogen, a lit wax

taper will go out with a ‘pop’ sound. A glowing wooden splint will burst into

fl ame if placed in a test tube of oxygen.)

Station 4: Place a small piece of red litmus paper in a beaker of water. Using

tweezers, add the calcium metal to the beaker. Observe what happens.

Station 5: Add the zinc metal to a test tube containing 3 mL of hydrochloric

acid. Observe.

Station 6: Place a small length of copper wire in silver nitrate solution in a test

tube. Note the colour of the solution before and after the reaction. (This may

take some minutes.)

Do not look directly at the fl ame

when the magnesium ignites.

Lead salts are poisonous.

Do not look directly at the fl ame

when the magnesium ignites.

Lead salts are poisonous.

EXPERIMENT 11.2



35

Station 7: Add 2 drops of phenolphthalein indicator to a test tube containing

2 mL of sodium hydroxide. Then slowly add 2 mL of sulfuric acid. Observe.

Station 8: Add a dropper-full of nickel(II) nitrate solution to a test tube

containing an equal amount of sodium carbonate solution. Observe.

RESULTSFor each reaction, write a short report detailing:

(a) reaction type (combination, decomposition, hydrocarbon combustion)

(b) the word equation for the reaction

(c) the balanced equation for the reaction

(d) any reaction evidence (colour change, precipitation, gas bubbles, energy

change).

QUESTIONS

1. Explain the colour change that occurred in the reaction at station 8.

2. What tests identify hydrogen gas and oxygen gas?

EXPERIMENT 11.2



36


The production of smog

AIMTo produce ‘smog’ in the laboratory.

APPARATUSbell jar with stopper

bicycle pump

water

delivery tube

paper

METHOD

1. Place about 3 cm of water in the bell jar as shown in the diagram.

delivery tube

bicycle pump

stopper

bell jar

lighted paper

water

2. Attach a bicycle pump to a delivery tube joined to the stopper.

3. Light a piece of paper and drop it in the jar.

4. Quickly stopper the jar and start pumping with the bicycle pump.

5. After a reasonable pressure has built up inside the jar, remove the stopper.

RESULTSWrite a clear concise report outlining your experimental method and your

observations.

QUESTIONS

1. Describe your observations.

2. How would you explain your results?

EXPERIMENT 12.1


37


Determining a solubility curve

A solubility curve is determined for a substance so that chemists know how

much solute they can add to a solvent in order to obtain a saturated solution

(one in which no more solute will dissolve).

AIMTo obtain the solubility curve for potassium chlorate, KClO3.

APPARATUS4 g pure crystalline potassium chlorate

burette

large clean test tube and glass stirring rod

thermometer


electronic balance

600 mL beaker

Bunsen burner

tripod and gauze mat

METHOD

1. Accurately weigh about 4 g of KClO3 into a large test tube. Record this

mass.

2. Add 10.0 mL of distilled water from a burette.

3. Immerse the test tube in a beaker of boiling water so that the water level

outside the tube is at least 3 cm higher than the level inside.

4. Carefully stir the mixture with the stirring rod until all the solid has

dissolved.

5. Allow the tube to cool by removing it from the water and holding it up to

the light. Stir constantly with the stirring rod.

6. Record the temperature at which crystals first appear.

7. Use the burette to add 2.5 mL of distilled water to the test tube and repeat

the above.

8. Repeat step 7 until at least five results have been recorded.

RESULTS

1. For each temperature recorded, calculate the solubility of KClO3. (Note:

100 mL of water has an approximate mass of 100 g, although this may

change with temperature.)

2. Tabulate your results in a table with the column headings shown below.

Mass of

KClO3 (g)

Volume of water

(mL)

Solubility

(g/100 g)

Temperature

(ºC)

3. Plot a graph of temperature against solubility, placing temperature on the

x-axis and solubility on the y-axis. Draw a smooth line of best fit to obtain

the solubility curve.

Note: This experiment may

be repeated with a number of

different solutes and the results

can be superimposed on the

same axes of the graph of

temperature against solubility.

Potassium chloride and sodium

sulfate are suitable solutes.

Note: This experiment may

be repeated with a number of

different solutes and the results

can be superimposed on the

same axes of the graph of

temperature against solubility.

Potassium chloride and sodium

sulfate are suitable solutes.

EXPERIMENT 13.1




38

QUESTIONS

1. Why is the stirring so important?

2. What effect would the formation of a supersaturated solution have on your

results and how could such errors be overcome?

3. Why is it important to record the temperature at which the KClO3

recrystallises rather than the temperature at which it all dissolves?

4. Suggest where any errors may have occurred in your experiment.

5. Read from your curve what the solubility of KClO3 is at 55°C.

EXPERIMENT 13.1


39


Specific heat capacity

Different liquids have different specific heat capacities. Specific heat capacity

can be calculated using the relationship between the mass of the substance being

tested, the amount of energy being added to the substance and the temperature

change that the substance experiences. In this experiment we will measure the

amount of energy being added to a liquid by passing an electric current through

a heating coil which is placed in the liquid. The amount of energy being given

off by the coil can be found by measuring the voltage (V ) across the coil, the

current (I ) flowing through the coil, and the time (t), in seconds, for which it

flowed. The amount of energy is given by the formula:

Energy � VIt

The extent to which the liquid uses this energy can be found by measuring

the mass of liquid (m), and the change in temperature (T ) that it experiences

during the time for which the current flowed.

The energy produced by the current should equal the energy absorbed by the

liquid.

VIt � mcT (c is the specific heat capacity)

c �VIt

mT

AIMTo measure the specific heat capacity of a number of different liquids.

APPARATUS100 mL water

100 mL ethanol

100 mL glycerol

polystyrene cup

power pack

ammeter

voltmeter

stopwatch

heating coil

thermometer

balance

connecting wires

METHOD

1. Find and record the mass of the empty polystyrene cup.

2. Add 100 mL of water to the cup and find and record the mass.

3. Connect the circuit as shown in the diagram below.

variable voltagepower supply

ammeter

voltmeter

heatingcoil

A

V

EXPERIMENT 13.2




40

4. Place the heating coil in the water and turn on the power.

5. Set the current to about 2 A and turn off the power. Measure the temperature

of the water.

6. Turn on the power for five minutes, stirring the water gently using the

thermometer.

7. After five minutes, turn off the power and note the highest temperature

reached by the water.

8. Pour out the water and repeat with the alcohol, and then with glycerol (or

another suitable liquid).

9. Tabulate your results and determine the specific heat capacity of each

liquid.

QUESTIONS

1. What would happen to the accuracy of this experiment if the temperature

of the liquids approached their boiling points?

2. In which areas are errors likely to occur?

3. Which liquid has the highest specific heat capacity? (The specific heat

capacity of water is 4.2 J g 1 °C 1 or 4200 J kg 1 °C 1.)

EXPERIMENT 13.2


41


A water modelling exercise

The hydrogen and oxygen in water are held together by covalent bonds. Since

there is a difference in the electronegativity of the hydrogen and oxygen atoms,

the molecule becomes polar. The hydrogen atoms have a slightly positive charge

(D+) and the oxygen atom has a slightly negative charge (D ).

AIMTo prepare a bonding model for the water molecule, and to use the model to

explain the expansion of water upon freezing.

APPARATUSmolecular modelling kit

METHOD

1. Decide how you could best model the bonding that exists within a water

molecule and the hydrogen bonding that exists between water molecules.

2. Prepare a number of ‘water molecules’ and consider how you could best

show the expansion of water upon freezing. Remember you are trying to

show that in the liquid state the molecular arrangement takes up less space

than the molecular arrangement of the solid (ice) state. (The figures shown

on pages 294–5 of the textbook could be useful. Mapping your molecular

arrangements onto grid paper may be a good way to show differences in

area.)

QUESTIONS

1. In what ways was your modelling exercise successful in representing the

bonding in water?

2. What limitations are there to your model?

3. Describe how your model could be extended to show how water can be

poured.

EXPERIMENT 13.3



42


Indicators and pH

An acid–base indicator is a substance whose colour in aqueous solution is

influenced by the presence of dissolved acids and bases. An indicator has a

characteristic colour in acid solutions and another in basic solutions. The pH of

the solution can be estimated by comparing the colour of the solution to a pH

chart for the indicator used. A pH chart for a range of indicators is provided

below.

Indicator name pH range for colour change

methyl violet

thymol blue (acidic range)

bromophenol blue

methyl orange

bromocresol green

methyl red

bromothymol blue

thymol blue (basic range)

phenolphthalein

alizarin yellow R

0 2 4 6 8 10 12

violetyellow

yellowred

blueyellow

yellowred

blueyellow

yellowred

blueyellow

blueyellow

pinkcolourless

redyellow

AIMTo investigate the colours of various indicators in solutions of various acids and

bases, and use a pH chart to estimate the pH of the substances tested.

APPARATUS6 semi-micro test tubes

test tube rack

1–2 mL s 0.1 mol L–1 solutions of the following substances: HCl, H2SO4,

CH3COOH, NaOH, Ca(OH)2 and Na2CO3

dropper bottles containing a selection of indicators such as universal indicator

or any of those shown in the diagram above

pH charts for the universal indicator (if used)

METHOD

1. Set up the six test tubes in the test-tube rack, and place 1–2 mL of each of

the six solutions in a separate test tube.

2. To each test tube add one or two drops of the indicator that has been

assigned to your group.

3. Observe any colour change and compare the final colour of the solution to

those in the pH chart for your indicator. Determine the pH from the chart.

RESULTSRecord your results in a table, then collate them with the rest of the class.

QUESTIONS

1. What is the purpose of an indicator?

2. Why is it necessary to have several indicators?

Note: Students should perform

this experiment in groups, with

each group testing a different

indicator.

Note: Students should perform

this experiment in groups, with

each group testing a different

indicator.

EXPERIMENT 14.1


43


The reactions of acids

AIMTo investigate the reactions of a typical acid (dilute hydrochloric acid) with

metals, metal oxides, carbonates and bases.

APPARATUSdropper bottles containing 0.1 mol L 1 solutions of:

hydrochloric acid or limewater (calcium hydroxide)

sodium hydroxide

dropper bottle of bromothymol blue indicator

small samples of the following metals: zinc, copper turnings, magnesium and

iron

copper(II) oxide powder

magnesium oxide

marble chips (calcium carbonate)

sodium carbonate

12 test tubes and test-tube holder

Bunsen burner

stopper or cork

wax taper and matches

METHOD(a) Reaction with metals

1. Place a small piece of magnesium ribbon in a test tube and add about 2 mL

of dilute HCl.

2. Stopper the test tube and allow the gas to accumulate.

3. Remove stopper and test for the gas evolved by holding a lighted taper to

the mouth of the test tube.

4. Repeat using zinc sample.

5. Repeat using copper sample.

6. Repeat using iron sample.

(b) Reaction with metal oxides

1. Place a very small amount (the size of a grain of rice) of copper(II) oxide

in a test tube and add about 2 mL of hydrochloric acid. Warm the test tube

gently. If no change is observed, allow to stand for a while before making

further observations.

2. Repeat this procedure using magnesium oxide and hydrochloric acid.

(c) Reaction with carbonates

1. Add several marble chips to a test tube.

2. Add about 2 mL of limewater to a second test tube.

3. Add the hydrochloric acid to the marble chips and allow the gas evolved to

flow into the test tube containing the limewater (hold the two test tubes at

an angle to each other with the open ends touching). After a few minutes,

shake the test tube containing the limewater.

4. Repeat this procedure using sodium carbonate in place of the marble

chips.

EXPERIMENT 14.2




44

(d) Reaction with bases

1. Place a few drops (2–3) of bromothymol blue indicator into a test tube

containing 1 mL of hydrochloric acid. Observe the colour change.

2. Place a few drops (2–3) of bromothymol blue indicator into a test tube

containing 2 mL of sodium hydroxide. Observe the colour change.

3. To the second test tube, carefully add hydrochloric acid, drop by drop, until

the colour turns green. (Note: If too much acid is added, the indicator will

turn yellow.)

EXTENSIONTry this experiment using vinegar (dilute acetic acid), or sulfuric acid.

RESULTSWrite up a table to record your observations. You might, for example, use the

following headings:

Test Observations Explanation

QUESTIONS

1. Write balanced chemical equations for the four reaction types.

2. Write an ionic equation for hydrochloric acid reacting with sodium

hydroxide.

3. Write an equation for the reaction of limewater with the gas evolved in

part (c).

EXPERIMENT 14.2


45


Finding the pH of common household substances

AIMTo determine the pH of some common substances.

APPARATUSwhite tile

semi-micro test tubes

dropper bottle containing universal indicator solution

pH meter

samples of household substances collected by students, such as:

drain cleaner

cloudy ammonia

baking soda

battery acid

cleaning agents

vinegar

soft drinks

shampoos and conditioners

toothpaste

liquid fertilisers

METHOD

1. Use a test tube or white tile to test the sample with universal indicator

solution. If the sample is a solid, dissolve it first in distilled water.

2. Record the results in tabular form.

3. If a pH meter is available, use it to find the pH of a selection of

substances.

QUESTIONHow different were the results from the universal indicator and the pH meter?

Comment on the accuracy of the two methods of measuring pH.

Note: Many of these substances

are mixtures. The pH of their

solutions may be due to the

presence of more than one

compound.

Note: Many of these substances

are mixtures. The pH of their

solutions may be due to the

presence of more than one

compound.

EXPERIMENT 14.3



46


Carbon dioxide content of fizzy drinks

AIMTo determine the CO2 content of different brands of soda water and/or

lemonade.

APPARATUSsmall bottles or cans (375 mL) of soda water and lemonade which have been

refrigerated

150 mL s 0.1 mol L 1 NaOH which has previously been standardised

dropper bottle of phenolphthalein indicator

burette and stand

measuring cylinder

250 mL conical flasks

METHOD

1. Note the brand name of the drink you are testing.

2. Using a measuring cylinder, transfer 25 mL of drink into a conical flask.

Do not be concerned about the loss of gas as the solution is supersaturated

with CO2 gas.

3. Add 2 drops of phenolphthalein indicator.

4. With NaOH in the burette, titrate until the first permanent pink is obtained.

You will need to shake well during the addition of the NaOH. Record

titre.

5. Repeat three more times.

QUESTIONUsing the following equation, calculate the concentration of the CO2 in mol L 1.

CO2(aq) + NaOH(aq) NaHCO3(aq)

EXPERIMENT 14.4


47


Preparation of a solution of known concentration

A solution of known concentration is called a standard solution. In this

procedure, a standard solution will be prepared from solid sodium carbonate as

it is readily available, reasonably pure and does not react with carbon dioxide

in air. Any moisture that the solid has absorbed can be removed by simply

drying it in the oven for one hour.

AIMTo prepare a solution of known concentration.

APPARATUSdry Na2CO3

weighing bottle or beaker

balance

250 mL volumetric flask


dropping pipette and funnel

METHOD

1. Calculate the mass of dry Na2CO3 to be used to make up a 250 mL volume

of solution with a concentration of 0.050 mol L–1.

2. Put approximately this amount of the solid into a clean, dry weighing bottle

or beaker. Using the balance, weigh the bottle and its contents accurately.

Record the mass.

3. Clean and rinse the volumetric flask with distilled water. Then, using the

funnel, transfer the solid into the 250 mL volumetric flask.

4. Accurately weigh the bottle again. Record the mass. The difference of the

two masses is the accurate mass of the Na2CO3 that was transferred to the

volumetric flask.

5. Hold the funnel a little way out of the flask to allow air to escape. Then,

using distilled water, wash any solid remaining in the funnel into the

volumetric flask.

6. Add distilled water until the flask is about half full. Stir the contents by

swirling the flask until the solid is completely dissolved.

7. Add more water until the level is about 1 cm from the calibration line

on the neck of the flask. Use the dropping pipette to carefully add water,

drop by drop, until the bottom of the meniscus is level with the calibration

line.

8. Mix the contents by inserting the stopper and inverting the flask and

shaking. Repeat until thoroughly mixed.

QUESTIONS

1. Why is thorough mixing of the contents important?

2. From the mass of the solid, calculate the concentration of the sodium

carbonate solution.

3. Given the mass of solute you used, did you expect the concentration of

your solution to be 0.050 mol L–1?

EXPERIMENT 15.1



48


Stoichiometry of a reaction

This experiment will provide you with an opportunity to determine the number

of moles of copper that are obtained when a given number of moles of iron is

reacted with copper(II) sulfate pentahydrate, CuSO4.5H2O.

AIMTo investigate quantitatively the reaction between iron and copper(II) sulfate

pentahydrate.

APPARATUSbalance

filter paper and funnel

250 mL conical flask

400 mL beaker


glass stirring rod

distilled water in plastic squeeze bottle

8 g copper(II) sulfate pentahydrate

(powder)

1.5 g degreased steel wool

METHOD

1. Determine and record the mass of the filter paper. An accurate measurement

is important.

2. Fold the filter paper in half and then in half again and place it in the filter

funnel supported by the flask.

3. Determine and record the mass of the copper(II) sulfate pentahydrate. (Use

approximately 8 g; an accurate measurement is important.)

4. Determine and record the mass of the steel wool. (Use approximately 1.5 g;

an accurate measurement is important.)

5. Using the measuring cylinder, measure 100 mL of distilled water and pour

the water into the beaker.

6. Dissolve the copper(II) sulfate pentahydrate in the water.

7. Add the steel wool to the copper(II) sulfate pentahydrate solution and stir

the mixture for about 15 minutes. Record your observations.

8. Filter the mixture carefully. Wash any remaining solids into the filter paper

with distilled water.

9. When filtration is complete, allow the filter paper and solids to dry overnight

or use an oven set at a low temperature.

10. When completely dry, determine the mass of the filter paper and the residue

in it. An accurate measurement is important.

RESULTSPrepare a data table. Include entries for the following measurements:

(a) mass of copper(II) sulfate pentahydrate

(b) mass of steel wool

(c) mass of filter paper and residue

(d) mass of filter paper

(e) mass of residue (product).

QUESTIONS

1. Write a balanced equation for the reaction.

2. Write down the mole ratio of the reactants and products.

3. What type of reaction is this?

4. Calculate the number of moles of reactants that you used.

5. Calculate the number of moles of products produced.

6. How do your results compare with your predictions in terms of the mole

ratio you stated in question 2? Write a conclusion to the experiment.

EXPERIMENT 15.2


49


Galvanic cells

A galvanic cell is made using the principle of competition for electrons. The

more reactive substance will give electrons to the less reactive. An electrolyte

is necessary to allow the flow of ions from one electrode to the other. In part A

of this experiment, we will prepare simple cells using everyday objects and will

attempt to make a battery of cells, and in part B we will construct a galvanic

cell using a salt bridge, separating the process into two half-cells.

PART AAIMTo prepare a simple cell and battery.

APPARATUS2 lemons

2 potatoes

15 cm strips of copper, zinc, aluminium, magnesium, tin, iron

connecting wires

voltmeter

METHOD

1. Insert a strip of zinc into one side of a lemon.

2. Insert a strip of copper into the other side.

3. Ensure that the metal strips do not touch each other inside the lemon.

4. Connect the zinc strip to one side of a voltmeter, and connect the copper

strip to the other.

5. Observe the reading on the voltmeter.

6. Repeat the experiment using one of the potatoes.

7. Repeat steps 1 to 6 using several different combinations of the metal

strips.

8. If the voltmeter indicates the voltage in the wrong direction, connect the

leads the other way around.

9. Use two or more lemons or potatoes in series to try to make a battery.

RESULTS

1. Make a table showing the voltages you obtained for each combination you

tried with the lemon.

2. Make a similar table for the voltages obtained for each combination with

the potato.

QUESTIONS

1. Identify the best combination of metals, and explain why this was so.

2. Identify the worst combination of metals. Explain.

3. Was there any difference between the voltages obtained with the potato and

the lemon? If so, explain why.

4. What maximum voltage did you obtain with a series of cells?

5. Why is it important to ensure that the metal strips do not touch inside the

lemon or potato?

EXPERIMENT 16.1




50

PART BAIMTo construct a galvanic cell using a salt bridge.

APPARATUS15 cm metal strips such as copper, zinc, aluminium and magnesium

150 mL s 1 mol L–1 solutions containing metal ions, such as: copper sulfate,

zinc sulfate, aluminium nitrate, magnesium nitrate

150 mL s 1 mol L–1 potassium nitrate solution

glass wool and U-tube

voltmeter and connecting wires

2 s 300 mL beakers

METHOD

1. Pour zinc sulfate solution into one beaker and copper

sulfate solution into the other beaker.

2. Fill a U-tube with potassium nitrate solution and

plug the ends with glass wool. (An agar solution

containing potassium nitrate, or filter paper dipped

in potassium nitrate, may be used.)

�� Place a strip of zinc metal into the solution of zinc

sulfate and a strip of copper metal into the solution

of copper sulfate.

�� Connect the zinc electrode to the negative terminal of

the voltmeter and the copper electrode to its positive

terminal.

�� Carefully invert the U-tube containing potassium

nitrate electrolyte so that one end is in the zinc half-

cell and the other end is in the copper half-cell.

�� Note the reading on the voltmeter as the salt bridge

is inserted.

�� If time permits, repeat the experiment using

combinations of all the metals supplied. Remember

to place the metal into a solution of its ions.

QUESTIONS

1. What is the purpose of the salt bridge?

2. Which combination of metals gave the highest voltage?

3. Which electrode acts as the anode?

4. Which electrode acts as the cathode?

5. Give equations for the reaction taking place at the anode and cathode.

6. Give an equation for the overall cell reaction.

7. Which way do electrons move through the outer circuit?

KNO3

salt bridge

copperzinc

KNO3

salt bridge

copperzinc

EXPERIMENT 16.1


51


Corrosion

We shall investigate the factors that may affect the corrosion of iron. By placing

different types of iron nails in a solution containing sodium chloride, potassium

hexacyanoferrate, NaCl, phenolphthalein and agar, we will be able to identify

the ions produced around the nail. In this way we will see which type of metal

corrodes more easily, whether stress at any point on a metal surface influences

the rate of corrosion, and whether or not galvanising or the presence of any

other metal prevents corrosion.

AIMTo investigate the corrosion of metals.

APPARATUSungalvanised iron nail

iron nail wrapped in a thin strip of copper foil

galvanised nail

40 mL solution of agar/sodium chloride/potassium hexacyanoferrate/

phenolphthalein

20 mL s 2.0 mol L 1 solution of hydrochloric acid

3 petri dishes

METHOD

1. Clean half of your galvanised nail in the hydrochloric acid.

2. Bend the ungalvanised nail in a few places to cause stress points.

3. Place a different nail in each petri dish.

4. Cover each nail completely with the agar solution.

5. Observe the colours that develop around each nail.

RESULTSDraw a diagram of each of your petri dishes, clearly indicating the colours that

formed around the nails and their location.

QUESTIONS

1. What ions have formed around each part of each nail?

2. Which parts of the clean iron nail corroded first?

3. What effect did galvanising have on the rate of corrosion?

4. How effective was the copper coating in reducing the corrosion of the iron

nail?

5. Discuss and evaluate in class your experimental results. Suggest ways in

which the experiment could be improved or modified.

Notes:

1. Nails should be about 5 cm

to 7 cm long. If necessary,

obtain the ungalvanised nail

by placing a galvanised nail

in dilute hydrochloric acid to

clean off the zinc. The half-

copper nail can be obtained

by electroplating half an iron

nail with dilute copper sulfate

solution.

2. The solution may be prepared

by dissolving approximately

0.5 g of NaCl in 40 mL of water.

About 0.5 g of agar-agar is

added. Warm the solution and

stir until the agar is dispersed.

Prepare a solution of

potassium hexacyanoferrate,

K4Fe(CN)6, by dissolving 2 g

of the solid in 100 mL of water.

Add 1 mL of this solution and

0.5 mL of phenolphthalein to

the agar solution.

3. Phenolphthalein turns pink in

the presence of OH– ions.

4. The hexacyanoferrate ion

produces a white or pale blue

precipitate with Fe2+ ions, and

goes an intense blue with Fe3+

ions.

5. The NaCl acts as an electrolyte

and speeds up the corrosion

process.

Notes:

1. Nails should be about 5 cm

to 7 cm long. If necessary,

obtain the ungalvanised nail

by placing a galvanised nail

in dilute hydrochloric acid to

clean off the zinc. The half-

copper nail can be obtained

by electroplating half an iron

nail with dilute copper sulfate

solution.

2. The solution may be prepared

by dissolving approximately

0.5 g of NaCl in 40 mL of water.

About 0.5 g of agar-agar is

added. Warm the solution and

stir until the agar is dispersed.

Prepare a solution of

potassium hexacyanoferrate,

K4Fe(CN)6, by dissolving 2 g

of the solid in 100 mL of water.

Add 1 mL of this solution and

0.5 mL of phenolphthalein to

the agar solution.

3. Phenolphthalein turns pink in

the presence of OH– ions.

4. The hexacyanoferrate ion

produces a white or pale blue

precipitate with Fe2+ ions, and

goes an intense blue with Fe3+

ions.

5. The NaCl acts as an electrolyte

and speeds up the corrosion

process.

EXPERIMENT 16.2



52


Minimising corrosion (student design)

AIMTo design and perform an experiment to investigate the effectiveness of different

methods used to minimise metal corrosion. Each group should investigate a

different method.

METHODDesign your own method and prepare a list of apparatus. Check these with your

teacher before proceeding.

RESULTSRecord your results and write a short report, including a conclusion. Present

your results to the rest of the class.

QUESTIONS

1. Evaluate your experimental design.

2. Compare your design with those of other groups.

EXPERIMENT 16.3


53


The reactivity of metals and their saltsAIMTo react various metals and metal salts with one another to determine a simple

order of reactivity.

APPARATUS5 small clean strips or pieces of each of the following metals: Mg, Zn, Ag, Cu

5 iron nails (not galvanised)

5 test tubes and test-tube rack

25 mL s� 0.1 mol L 1 solutions of the following cations: Mg2+, Zn2+, Ag+,

Cu2+, Fe2+ (The anions in these salts could be any of nitrates, sulfates or

chlorides.)

METHOD

1. Copy the table below, which shows the possible combinations of metals

and salt solutions.

2. Set up the five test tubes in the test-tube holder and add approximately

5 mL of the first salt solution from the table to each.

3. Place a piece of each of the different metals in a different test tube and

observe.

4. Discard contents of all test tubes and set up again as in step 1, using the

second salt solution from the table.

5. Repeat step 2, using fresh metal samples.

6. Continue until all five solutions have been used.

RESULTSUse your table to record if there is a reaction (�) or no reaction (�).

Metal

Salt solution

Mg2+ Zn2+ Fe2+ Ag+ Cu2+

Mg

Zn

Fe

Ag

Cu

QUESTIONS

1. Write redox equations for the displacement reactions that occurred.

2. Rank the metals in order of most reactive to least reactive, and compare

your list with the electrochemical series. Note any differences.

3. Compare the reactivity of the metals you have investigated to their positions

in the periodic table. What generalisations can you make?

EXPERIMENT 16.4



54


Simple redox equations

Steel wool is placed in a solution of copper sulfate. Electrons are transferred

from the iron in the steel wool to the copper ions.

AIMTo investigate a simple redox reaction.

APPARATUS250 mL s 1 mol L 1 copper sulfate solution

250 mL s 1 mol L 1 H2SO4 solution

500 mL beaker

steel wool and tongs

METHOD

1. Place 250 mL of the CuSO4 solution in a beaker.

2. Add 2 mL H2SO4 solution.

3. Use tongs to completely submerge a wad of steel wool in the solution.

4. Leave the steel wool submerged for a few minutes and then remove.

5. Note the colour change in the steel wool and copper sulfate solution.

QUESTIONS

1. What substance has replaced the iron in the steel wool?

2. Write half-equations to show the electron transfer that has occurred.

3. Label the oxidation half-equation and reduction half-equation as

appropriate.

4. Write an overall equation to show the entire reaction.

5. Name the oxidant and reductant in this reaction.

EXPERIMENT 16.5


55


Complex redox equationsAIMTo react some common oxidants and reductants with one another and write

their redox equations.

APPARATUS4 semi-micro test tubes

test-tube rack

dropper bottles containing 0.1 mol L 1 solutions of the following substances:

potassium iodide

iron(II) sulfate

potassium dichromate

potassium permanganate

sulfuric acid

starch

potassium thiocyanide, KSCN

METHODSet up the four test tubes as shown in the table below, using 6 drops of each

reagent. Add 3 drops of sulfuric acid to each test tube. Add the appropriate

indicator and observe.

RESULTSCopy and complete the table.

Test tube Reagents Indicator Skeleton equation Observations

1 KMnO4(aq) KI(aq) starch MnO4 + I m Mn2+ + I2

2 K2Cr2O7(aq) KI(aq) starch Cr2O72 + I m Cr3+ + I2

3 KMnO4(aq) FeSO4(aq) KSCN(aq) MnO4 + Fe2+ m Mn2+ + Fe3+

4 K2Cr2O7(aq) FeSO4(aq) KSCN(aq) Cr2O72 + Fe2+ m Cr3+ + Fe3+

QUESTIONS

1. Identify the oxidant and reductant in each reaction.

2. Complete the overall redox reaction by first writing each half-equation.

3. Explain how:

(a) starch can be used to test for the presence of iodine

(b) KSCN solution can be used to test for the presence of Fe3+ ions.

Notes:

1. Starch can be used to test for

the presence of iodine, I2.

2. KSCN solution can be used to

test for the presence of Fe3+

ions.

Notes:

1. Starch can be used to test for

the presence of iodine, I2.

2. KSCN solution can be used to

test for the presence of Fe3+

ions.

EXPERIMENT 16.6



56


Making oxygen (student design)

AIMTo design and carry out a process for the laboratory preparation, collection and

testing of oxygen gas.

APPARATUSStudents to select apparatus as needed.

METHOD

1. Use the information on page 410 to devise a method of preparing a gas jar

of oxygen gas.

2. Write your method in detail, providing a diagram of the way in which you

will set up the apparatus.

3. Obtain approval for your method and apparatus design from your teacher

before you begin.

QUESTIONS

1. How did you know you were successful in producing oxygen?

2. Write a balanced chemical equation for (a) your method and (b) an

alternative laboratory preparation of oxygen gas.

Note: The test for oxygen is to

plunge a glowing splint into the

gas jar containing the oxygen. If

the splint bursts into flame, the

gas jar contains oxygen.

Note: The test for oxygen is to

plunge a glowing splint into the

gas jar containing the oxygen. If

the splint bursts into flame, the

gas jar contains oxygen.

EXPERIMENT 17.1


57


Preparation and properties of carbon dioxide

The reaction in which an acid reacts with a carbonate is used to produce carbon

dioxide in the laboratory, using Kipp’s apparatus as in the figure below.

dilute hydrochloric acid

stopcock

marble chips

water concentratedsulfuric acid

flask

Kipp’s apparatus

Carbon dioxide gas is generated in Kipp’s apparatus (left). Hydrochloric acid is

allowed to fall onto marble chips. The resultant gas is bubbled through water to

dissolve acid spray and dried by concentrated sulfuric acid. The gas is collected

in the flask (right).

AIMTo prepare and test the properties of carbon dioxide.

APPARATUS5 gas jars of carbon dioxide obtained from a cylinder or from Kipp’s

apparatus

5 cover slips

wax taper

4 cm strip of magnesium


beaker

dropper bottle of universal indicator

teat pipette or syringe

20 mL limewater

test tube

METHODTest the five samples of carbon dioxide as follows:

Gas jar 1

Take the cover slip off the jar. Smell the gas and observe any colour.

Gas jar 2

Light a wax taper and plunge it into the jar. Observe.

Gas jar 3

Ignite the magnesium ribbon and plunge it into the jar. Observe.

EXPERIMENT 17.2




58

Gas jar 4

1. Add a few drops of universal indicator to 100 mL of distilled water in a

beaker. Note the colour of the water and determine its pH using universal

indicator.

2. Fill a syringe or teat pipette with carbon dioxide and bubble the gas through

the water containing the indicator.

3. Observe the colour change in the water and note the pH.

Gas jar 5

Fill a syringe or teat pipette with carbon dioxide and bubble it through a test

tube of limewater. Observe the effect.

QUESTIONS

1. Does carbon dioxide have an odour or colour?

2. The carbon–oxygen bond in CO2 is strong, preventing the gas from

supporting combustion. Why did the magnesium burn in carbon dioxide?

3. What was the pH of your water after CO2 was bubbled through? Give an

equation to show the reaction that takes place between water and CO2.

4. What happens when CO2 is bubbled through limewater? The formula for

limewater is Ca(OH)2(aq). Write an equation for the reaction between

limewater and CO2 and name the precipitate causing the milky colour.

5. Carbon dioxide can be collected by the upward displacement of air. What

does this indicate about the density of carbon dioxide?

6. Liquid carbon dioxide is unheard of at normal pressures. What property

does carbon dioxide have that causes this phenomenon?

7. List as many properties of carbon dioxide as you can. Include any that you

can think of that were not covered in the experiment.

EXPERIMENT 17.2


59


Measuring gas volume

AIMTo measure molar gas volume at standard laboratory conditions (SLC).

APPARATUS0.04 g clean magnesium ribbon

5 mL s 2 mol L–1 hydrochloric acid

test tube

delivery tube and rubber stopper


beehive shelf

trough and water

METHOD

1. Set up the apparatus shown in the diagram in the following way.

(a) Pour approximately 3 mL of hydrochloric acid into the test tube.

(b) Measure the mass of the freshly cleaned magnesium ribbon (approximately 0.04 g).

(c) Half-fill the trough with water, ensuring that it covers the beehive shelf.

(d) Fill the measuring cylinder with water. Place your hand over the mouth of the cylinder and invert it under the water in the trough. Ensure that the mouth of the cylinder is over the opening in the top of the beehive shelf.

(e) Ensuring that the outlet of the delivery tube is passing under the beehive shelf so that any gas produced will bubble up into the measuring cylinder, drop the piece of magnesium into the acid. Very quickly replace the rubber stopper connected to the delivery tube so that the gas passes into the cylinder.

2. Measure the volume of gas produced in the measuring cylinder.

RESULTSRecord the mass of magnesium, volume of hydrogen gas produced, the

temperature and atmospheric pressure. Use your results and the equation for

the reaction to calculate the molar gas volume at SLC.

QUESTIONS

1. Where are errors likely to occur in this experiment?

2. Would you get a different molar gas volume if the gas being produced was

chlorine? Explain your answer.

3. What further calculations could be carried out to obtain a more accurate

result?

measuring

cylinder

delivery

tube

water

beehive shelf

stopper

Mg + hydrochloric acid

measuring

cylinder

delivery

tube

water

beehive shelf

stopper

Mg + hydrochloric acid

EXPERIMENT 18.1



60


The egg in the flask

AIMTo demonstrate the relationship between the pressure and temperature of a

fixed mass of gas at constant volume.

APPARATUS300 mL pyrex conical flask

Bunsen burner

retort stand

bosshead

clamp

heatproof mat

bath of iced water (optional)

Vaseline

shelled, hard-boiled egg, slightly larger than the neck of the flask

METHOD

1. Smear a thick layer of Vaseline around the inside neck of the conical

flask.

2. Clamp the flask above the Bunsen burner using the retort stand. Heat the

flask for two to three minutes.

3. Remove the burner from under the flask and place the shelled, hard-boiled

egg in the neck of the flask.

4. Observe the behaviour of the egg. The process may be speeded up by

immersing the flask in a bath of iced water.

QUESTIONS

1. What happens to the air in the flask as the flask is heated?

2. The egg acts as a stopper for the flask. Explain what happens to the pressure

of the gas inside the flask as it cools.

3. How does the pressure outside the flask compare with the pressure inside

the flask just before the egg pops in?

4. Why is the egg sucked into the flask?

5. State which of the three variables (pressure, temperature or volume)

changed in the course of the experiment and hence determine the gas law

it is illustrating. State the law.

6. If the flask was inverted the egg would remain in the flask. Explain.

7. What would you have to do to the inverted flask to remove the egg?

EXPERIMENT 18.2


61


Expanding balloon and marshmallow

A marshmallow contains thousands of tiny bubbles of trapped air. By changing

the external pressure on the marshmallow, we may be able to show that the

particles in the bubbles of air are exerting a force.

AIMTo demonstrate the relationship between the pressure and volume of a fixed

mass of gas at constant temperature.

APPARATUSmarshmallow

vacuum pump

balloon

filter flask

rubber stopper

METHOD(a) Balloon in a flask

1. Blow up a balloon until it will just fit into the mouth of a filter flask.

2. Tie the end of the balloon and push it into the flask.

3. Place a stopper in the mouth of the flask.

4. Connect the flask to the vacuum pump and slowly turn it on.

5. Observe what happens to the balloon.

(b) The expanding marshmallow

1. Place the marshmallow in the flask.

2. Place a stopper in the mouth of the flask and turn on the pump.

3. Note the change in size of the marshmallow.

4. Turn off the pump and remove the stopper.

5. Note the change in size of the marshmallow.

QUESTIONS

1. Why did the balloon and marshmallow change in size?

2. Would a marshmallow change size when placed in a hot drink? Explain.

3. Suggest other everyday substances that may behave similarly to the

marshmallow.

EXPERIMENT 18.3



62


Diffusion of gases

Gas molecules are very tiny when compared to the volume of the container

that they occupy. If different gases are connected, the linear motion of the gas

molecules will result in the mixing of the gases. This is called diffusion of

gases.

AIMTo demonstrate gas diffusion and compare the rate of diffusion of two gases.

APPARATUS3 glass tubes, 30 cm long and 2 cm in diameter

6 rubber stoppers to fi t tubes

30 cm ruler

cotton wool

2 tweezers

2 needles

blue and red litmus paper to fi t tubes

10 mL concentrated HCl solution

10 mL concentrated NH3 solution

stopwatch (optional)

METHOD(a) To demonstrate the diffusion of HCl gas

1. Place a long strip of moist blue litmus paper inside the glass tubing.

2. Place the stopper in one end.

3. Using tweezers, soak a wad of cotton wool in the concentrated hydrochloric

acid solution and carefully put it in the other end of the tube. Place the

other stopper fi rmly over the open end of the tube.

4. Observe the colour of the blue litmus paper.

5. Record all observations.

(b) To demonstrate the diffusion of NH3 gas

Repeat the above procedure using ammonia solution (instead of hydrochloric

acid) and red (instead of blue) litmus paper. Record your observations.

(c) To compare the different rates of diffusion for HCl and NH3 gases

1. Place the glass tube on a bench next to the ruler. The stage of an overhead

projector can be used if doing the experiment as a demonstration.

2. Using tweezers, soak a wad of cotton wool in the concentrated hydrochloric

acid solution and another in the concentrated ammonia solution.

3. Press a needle into the inside surface of each of two stoppers. Place the

soaked wads of cotton wool on the needles as shown in the diagram

below.

4. Simultaneously place the stoppers in opposite ends of the tube and push the

stoppers in place as shown.

needle glass tube

stopper cotton woolsoaked in NH

3

needle

cotton wool soaked in HCl

stopper

Note: Part (c) may be done as a

teacher demonstration.

Concentrated hydrochloric acid

and concentrated ammonia

should be handled very carefully.

HCl is highly corrosive and the

vapour will irritate the respiratory

system. NH3 vapour produced by

NH3 solution is irritating to skin,

eyes and the respiratory system.

Note: Part (c) may be done as a

teacher demonstration.

Concentrated hydrochloric acid

and concentrated ammonia

should be handled very carefully.

HCl is highly corrosive and the

vapour will irritate the respiratory

system. NH3 vapour produced by

NH3 solution is irritating to skin,

eyes and the respiratory system.

EXPERIMENT 18.4



63

5. Note the appearance of the white smoke of NH4Cl appearing and measure

its distance from either end of the tube. The stopwatch can be used to time

the changes occurring in the tube.

6. Record your observations.

RESULTSThe rates of diffusion of HCl and NH3 gases can be compared by noting the

time taken for the NH4Cl to appear and the distance of the smoke from each

end of the tube. Calculate the rates using the equation:

rate �distance

.time

QUESTIONS

1. Explain the term ‘diffusion’.

2. Why should the litmus paper be moist for parts (a) and (b) of this experiment?

Suggest why blue litmus paper was used in part (a) and red litmus paper

was used in part (b).

3. Write a balanced chemical equation for the reaction between concentrated

hydrochloric acid and ammonia that occurred in part (c) of the experiment.

What was the white ‘smoke’ that was formed in the reaction?

4. Did the ‘smoke’ form closer to the cotton wool soaked in HCl or that

soaked in NH3? Comment on the diffusion rates of NH3 and HCl gas, with

reference to your experimental results.

5. Molecules of NH3 (Mr � 17) are lighter than those of HCl (Mr � 36.5).

Does this difference in mass explain the diffusion rates of HCl and NH3

that you observed in part (c)?

EXPERIMENT 18.4



64


The atom economy

AIMTo investigate the idea of atom economy, that is, to see how much of the

reactants are converted into the final products.

Chloroethane is a gas at room temperature. It can be used as a refrigerant

or solvent, or as an anaesthetic. It can be produced by a reaction that

combines ethane with chlorine in the presence of heat or ultraviolet light. The

products vary according to the proportion of reactants and may include other

chloromethanes.

Reaction 1: CH3CH3 + Cl2 Heat or light

CH3CH2Cl + HCl

Chloroethane can also be produced by a reaction between ethanol and hydrogen

chloride.

Reaction 2: CH3CH2OH + HCl CH3CH2Cl + H2O

APPARATUSMolecular model kit

Container for ‘waste’ atoms

METHOD

1. Use the molecular model kit to construct an ethane molecule and a chlorine

molecule.

2. Count and record the number and type of each atom present in the

reactants.

3. When these two react, chloroethane, CH3CH2Cl, and hydrogen chloride,

HCl, are produced. Convert the ethane molecule into chloroethane and

place the remaining hydrogen chloride in the ‘waste’ container.

4. Record the number and type of atoms in the final product and those in the

‘waste’ container.

5. Repeat for reaction 2.

6. Pack up the kits.

RESULTS

Reaction 1 Reaction 2

Mass of atoms in the

reactants

Mass of atoms in the final

product

Mass of atoms in the

‘waste’

EXPERIMENT 19.1



65

QUESTIONS

1. Show how these reactions support the law of conservation of matter.

2. Calculate the atom economy for each reaction.

Atom economy �mass of atoms in final product

s 100%mass of atoms in reactants

3. Compare the two values. Explain which reaction would be preferable from

the point of view of green chemistry.

4. What is the significance of the ‘waste’ produced?

5. What other aspects of green chemistry would be relevant here?

6. How is atom efficiency different from yield?

7. Chloroethane is produced industrially by reacting ethene and hydrogen

chloride over an aluminium chloride catalyst at high temperature. This

is the most economical process for the production of chloroethane. Write

the equation for this reaction. What would the atom economy be for this

reaction?

EXPERIMENT 19.1

PART A - John Wiley & Sonscatalogimages.johnwiley.com.au/Attachment/07314/0731405366/61_053… · EXPERIMENT 5.2 : The effect of heat on metals 15 EXPERIMENT 6.1 : Building molecular

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