Chemistry Paper 3.pdf

Paper 3 experiment

spm chemistry

Form 4 (2.1) Aim: To study the diffusion of particles Materials: Liquid bromine, copper (II) sulphate crystals, potassium manganite(VII) crystals, hot liquid agar, water Apparatus: Gas jar and cover, dropper, beaker, test tube, rubber stopper

Procedure: 1. Two drops of liquid bromine are dropped into a gas jar and covered with glass cover. 2. Water is poured into a beaker and then put several small potassium manganite (VII) crystals into the beaker. 3. A potassium manganite (VII) crystal is placed on the surface of the solidified agar in a test tube which is then clamped inverted and stoppered with a rubber stopper. Observation 1. Brown liquid (reddish brown) vaporises to form brown vapour that completely fills the gas jar 2. Purple potassium manganite (VII) spreads slowly throughout the water 3. Purple potassium manganite(VII) spreads slowly throughout the agar. Conclusion Diffusion takes place because bromine and potassium manganite(VII) are made up of tiny and discrete particles that move continuously and randomly from a region of high concentration to region of low concentration through particles of air (A), particles of water (B) or particles of agar (C)

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Paper 3 experiment

spm chemistry

Form 4 (2.2) Aim: To deter mine the melting and freezing points of naphthalene Materials: powdered naphthalene, water Apparatus: Test tube, beaker, conical flask, thermometer, retort stand and clamp. Bunsen burner, tripod stand, wire gauze, stopwatch

Procedure 1. Naphthalene is heated in a test tube and stirred with a thermometer as in apparatus set-up P. 2. The temperature of the naphthalene is recorded every half minute until its temperature reaches 90 °C. 3. Liquid naphthalene is allowed to cool down while stirred with a thermometer in a conical flask as in apparatus set-up Q. 4. The temperature of liquid naphthalene is recorded every half minute until its temperature goes down to 60 °C. 5. Graphs of temperature against time of heating and cooling of naphthalene are plotted Result Table for heating and cooling of naphthalene Time/s

0

30

60

90

120

150

180

210

240

270

300

330

360

390

Tempe rature/ °C

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Paper 3 experiment

spm chemistry

Analysis

Conclusion The melting point and freezing point of naphthalene are 80°C Discussion 1. Temperature of naphthalene does not change during melting. This is because the heat supplied is absorbed to waken (or overcome) the attractive force between naphthalene molecules. 2. Temperature of naphthalene does not change during freezing. This is because the heat lost to surrounding is balanced by heat released during the formation of attractive force between naphthalene molecules 3. During heating the naphthalene needs to be stirred continuously to make sure the temperature of the naphthalene is uniform 4. A hot water bath is used when heating the solid naphthalene to make sure heating temperature is uniform 5. Liquid naphthalene must be stirred and placed in a conical flask for slow and constant cooling. Otherwise, supercooling may take place. 6. The boiling point of naphthalene is 120°C. If solid naphthalene is heated directly without using hot water bath, naphthalene may sublime and produce vapours that are flammable and poisonous. A hot water bath reaches a maximum temperature of 100°C and melts the naphthalene but prevents it from subliming 7. If the naphthalene used is not pure, its melting and freezing points will be below 80°C

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Paper 3 experiment

spm chemistry

Form 4 (3.1) Aim: To determine the empirical formula of copper oxide Materials: copper oxide powder, zinc granules, hydrochloric acid 2 mol𝑑𝑚−3 ,anhydrous calcium chloride. Apparatus: Round bottom flask, combustion tube with a small hole, stopper and delivery tube, chemical balance, retort stand and clamp, thistle funnel, U-tube, asbestos paper.

Procedure 1. Combustion tube and asbestos paper is weighed. 2. Dry copper oxide is placed onto the asbestos paper and apparatus is weighed again 3. Dry hydrogen gas is passed through the apparatus for a few minutes. 4. Gas coming out from the small hole X collected into a test tube and tested with a burning wooden splinter. When the ‘pop’ sound is o longer heard, hydrogen gas at X is ignited. 5. The copper oxide in the combustion tube is heated. 6. Heating is stopped when no more changes takes place in the combustion tube. 7. Hydrogen gas is passed through the combustion tube continuously until the apparatus cools down to room temperature. 8. After it has cooled down, the combustion tube and asbestos paper containing copper is weighed again. 9. The heating, cooling, and weighing processes are repeated several times until a constant mass is obtained.

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Paper 3 experiment

spm chemistry

Results Mass of combustion tube + asbestos paper = x g Mass of combustion tube + asbestos paper + copper oxide = y g Mass of combustion tube + asbestos paper + copper = z g Calculation Element

Mass (g) Number of moles Simplest ratio

Cu

O

z–x

y-z

z – x / 64

y – z / 16

p

q

Conclusion The empirical formula of copper oxide is CupOq Discussion 1. Hydrochloric acid reacts with zinc granules to produce hydrogen gas. 2. 3.

Zn + 2HCl → ZnCl2 + H2

Hydrogen is dried by passing through anhydrous calcium chloride. In the combustion tube, copper oxide is reduced by hydrogen gas to copper and water is formed. If copper (II) oxide, CuO

CuO + H2 → Cu + H2O

If copper (I) oxide, Cu2O

Cu2O + H2 → 2Cu + H2O

4.

The combustion of hydrogen and oxygen gas forms water droplets at the end of the combustion tube. 2H2 + O2 → H2O

5.

The empirical formula of the other metals like lead (II) oxide, PbO, and iron(II) oxide, FeO, can also be determined by this method.

6. Precautions a) All connections in the apparatus must be air-tight. b) Hydrogen gas is flowed through the apparatus for a few minutes to displace all air present in the apparatus. A mixture of hydrogen and oxygen gas is explosive if it is ignited.

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Paper 3 experiment

spm chemistry

c) Hydrogen gas is allowed to flow continuously to prevent outside air from diffusing into the apparatus and mixing with hydrogen gas. This may cause an explosion. d) Hydrogen gas is allowed to continue flowing when cooling the copper to room temperature so that the hot copper is not oxidized again by oxygen in air to copper oxide. e) The processes of heating, cooling and weighing the product are repeated until a constant mass is obtained to make sure all the copper oxide is reduced to copper. 7.

Oxides of metals that are more reactive than hydrogen such as magnesium oxide cannot be reduced by hydrogen to its metal in a combustion tube. The empirical formula of metal oxides such as magnesium oxide need to be determined by carrying out experiment 3.2

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Paper 3 experiment

spm chemistry

Form 4 (3.2) Aim: To determine the empirical formula of magnesium oxide Material: Magnesium ribbon, sand paper Apparatus: Crucible with lid, Bunsen burner, tongs, pipe clay triangle, balance, tripod stand

Procedure 1. An empty crucible with lid is weighed. 2. 20cm magnesium ribbon is cleaned with sandpaper, rolled and put into the crucible. 3. The lid and its contents are weighed again. 4. The crucible is heated strongly until the magnesium starts to burn. 5. The moment the magnesium starts burning, the crucible is closed with its lid. 6. Heating is continued and the lid of the crucible is opened and closed quickly every once in a while to ensure the white smoke is not lost to the surroundings during heating. 7. When the magnesium ribbon is no longer glowing, the lid is opened and the crucible is heated strongly to ensure combustion of magnesium is complete. 8. The crucible and its contents is cooled down to room temperature and weighed again. 9. The crucible and its contents is repeatedly heated, cooled down and weighed again until a constant mass is obtained.

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Paper 3 experiment

spm chemistry

Result Mass of crucible + lid = a g Mass of crucible + lid + magnesium = b g Mass of crucible + lid + magnesium oxide = c g

Calculation Element

Mg

O

Mass 9g)

b-a

c–b

Number if moles

b – a / 24

c – b / 16

m

n

Simplest ratio

Empirical formula of magnesium oxide = MgmOn Conclusion The empirical formula of magnesium oxide is MgmOn Discussion 1. Magensium reacts with oxygen to produce magnesium oxide 2. Precautions a. Magnesium ribbon is cleaned with sandpaper to remove the oxide layer on its surface b. The crucible needs to be closed with lid when heating the magnesium to prevent some of the magnesium oxide from being lost to the surroundings as white smoke c. The lid of the crucible is open once a while during heating to allow oxygen (air)from outside to diffuse inside for complete combustion of the magnesium to take place d. The process of heating, cooling and weighing is repeated until a constant mass of product is obtained. This is to make sure the magnesium reacts completely with oxygen. 3. Inaccurate results may be obtained if a. Some of the smoke of magnesium oxide is lost to the surrounding b. Combustion of magnesium maybe incomplete.

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Paper 3 experiment

spm chemistry

Form 4 (4.1) Aim: To study the reactivity of Group 1 elements (alkali metals) with oxygen. Problem statement: Do elements of group 1 show different reactivities in their reaction with oxygen? Hypothesis: The further down Group 1, the higher the reactivity of the metal towards oxygen Variables: Manipulated: Type of alkali metal/ Type of Group 1 element Responding: reactivity of alkali metals Constant: Size of alkali metal Materials : Lithium, sodium, potassium, oxygen, filter paper

Apparatus: Gas jar with cover, gas jar spoon, Bunsen burner knife

Procedure 1. A small piece of lithium is cut using the knife. 2. The layer of oil on the lithium is dried using filter paper. 3. The lithium is placed on the gas jar spoon and heated until it starts to burn. 4. The gar jar spoo is quickly placed into the gas jar filled with oxygen. 5. The colour of the flame, how vigorous the reaction is, and the properties of the residue is observed and recorded. 6. The experiment is repeated using sodium and potassium

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Paper 3 experiment

spm chemistry

Observation Lithium burns slowly with reddish flame. Sodium burns vigorously with a yellowish flame Potassium burns very vigorously with a reddish-purple flame. White fumes and a white solid are formed in all cases. Analysis The alkali metals burns in oxygen gas with a bright flame to form white fumes of metal oxide that becomes a while solid when cooled. 4Li + O2 → 2Li2O 4Na + O2 → 2Na2O 4K + O2 → 2K2O Conclusion The reactivity of group 1 elements with oxygen increases when going down the group. The hypothesis is accepeted. Discussion If the metal oxide formed (combustion 4N product between alkali metal and oxygen) is mixed with water, an alkali solution is formed. If the phenolphthalein indicator is dropped into the solution formed, the colourless solution turns pink. Li2O + H2O → 2LiOH (lithium hydroxide) Na2O + H2O → 2NaOH (sodium hydroxide) K2O +H2O → 2KOH (potassium hydroxide)

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Paper 3 experiment

spm chemistry

Form 4 (4.2) Aim: To study the reactivity of Group 1 elements (alkali metals) with water. Problem statement: Do element of Group 1 show different reactivities in their reaction with water? Hypothesis: The further down group 1, the higher the reactivity of the metal with water Variables Manipulated: Type of alkali metal/type of group 1 element Responding: Reactivity of the reaction Constant: Size of metal Materials Lithium, sodium, potassium, water, filter paper

Apparatus Glass trough, forceps, knife

Procedure 1. A glass trough is half filled with water. 2. A small piece of lithium is cut with a knife 3. The layer of oil on the lithium is dried using filter paper. 4. The piece of lithium is carefully placed on the surface of the water. The reactivity of the lithium with water is observed and recorded. 5. The experiment is repeated using sodium and potassium

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Paper 3 experiment

spm chemistry

Observation Lithium moves slowly and randomly on the surface of water. Sodium moves quickly and randomly with a “hssing” sound. Potassium moves randomly and very quickly with a reddish purple flame and produce “hiss” ad “pop” sounds. In each case, a colourless solution if formed. Analysis The alkali metals react vigorously with cold water to form colourless alkali solutions and release hydrogen gas. 2Li + 2H2O → 2LiOH + H2 2Na + 2H2O → 2NaO + H2 2K + 2H2O → 2KOH+ H2 Conclusion The reactivity of Group 1 elements with cold water increases down the group. The hypothesis accepted. Discussion Group 1 element are very reactive 4Nand react with oxygen and water vapour in the air. Therefore, they need to be kept in paraffin oil.

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Paper 3 experiment

spm chemistry

Form 4 (4.3) Aim: To study the reaction of halogen with sodium hydroxide solution Problem statement: How do chlorine, bromine and iodine react with aqueous sodium hydroxide solution? Hypothesis: The halogens show similar chemical properties when they react with sodium hydroxide solution but the reactivity decreases down the group from chlorine to iodine. Variables (a) Manipulated variable: Types of halogen used (b) Responding variable: The products of reactions Constant variable: Concentration of sodium hydroxide solution Materials: Chlorine gas, liquid bromine, iodine crystals, sodium hydroxide solution Apparatus: Test tube, stopper, test tube holder and teat pipette. Procedure (A) Reaction of chlorine gas with aqueous sodium hydroxide solution 1. Chlorine gas is bubbled into aqueous sodium hydroxide solution. 2. The colour change of chlorine is recorded. (B) Reaction of bromine with aqueous sodium hydroxide solution 1. Two drops of liquid bromine are added to aqueous sodium hydroxide solution using a teat pipette. 2. The test tube is tightly closed with a rubber stopper and the mixture is shaken. 3. The colour change of bromine is recorded.

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Paper 3 experiment

spm chemistry

(C) Reaction of iodine with aqueous sodium hydroxide solution 1. Some iodine crystals are added to aqueous sodium hydroxide solution. 2. The test tube is tightly closed with a rubber and the mixture is shaken. 3. The colour change of iodine crystal is recorded. Results Halogen

Observation

Chlorine

The greenish chlorine gas dissolves quickly in NaOH solution to form a colourless solution.

Bromine

The brownish liquid bromine dissolves steadily in NaOH solution to form a colourless solution.

Iodine

The dark iodine crystal dissolves slowly in NaOH solution to form a colourless solution

Conclusion 4N 1. Chlorine, bromine and iodine react with sodium hydroxide solution to form two types of salts and water. X2 (g) + 2NaOH (aq) NaX + NaOX + H2O, where X=Cl, Br, I 2.

The reactivity of halogens with sodium hydroxide solution decreases down the group from chlorine to iodine

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Paper 3 experiment

spm chemistry

Form 4 (4.4) Aim: To study the reaction between chlorine(halogen) and iron Problem statement: How do chlorine, bromine and iodine react with iron? Hypothesis: The halogens show similar chemical properties when they react with iron but the reactivity decreases down the group from chlorine to iodine. Variables (a) Manipulated variable: Types of halogen used (b) Responding variable: Products of reactions and rate of the reactions (c) Constant variable: Iron wool Materials: Chlorine gas, liquid bromine, iodine crystals, soda lime, potassium manganate( VII), concentrated hydro-chloric acid and iron wool Apparatus: Combustion tubes, Bunsen burner, retort stand and clamp, conical flask and thistle funnel Procedure Reaction of chlorine gas with iron wool

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Paper 3 experiment 1. 2. 3. 4.

spm chemistry

A small roll of iron wool is placed in the middle of a combustion tube. The iron wool is then heated strongly. Chlorine gas is prepared in the laboratory by adding concentrated hydrochloric acid to potassium manganate VII). The chlorine gas produced is allowed to pass through the heated iron wool. The excess chlorine gas is absorbed by the soda lime.

(B) Reaction of bromine gas with iron wool

1. 2. 3. 4.

A small roll of iron wool is placed in the middle of a combustion tube and is heated strongly. The liquid bromine is warmed4N up by using a Bunsen burner. The bromine is vaporised and bromine gas passed through the heated iron wool. The excess bromine gas is absorbed by the soda lime.

(C) Reaction of iodine with iron wool

1. 2. 3.

4.

A few crystals of iodine are placed in a boiling tube. A small roll of iron wool is then placed in the middle of a combustion tube. The iron wool is heated strongly first, followed by the iodine crystals (sublimation will take place). The iodine vapour produced is allowed to pass through the hot iron wool.

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Paper 3 experiment

spm chemistry

Results Halogen

Observation

Chlorine

Hot iron wool glows brightly when chlorine gas is passed over it. A brown solid is formed.

Bromine

Hot iron wool glows moderately bright when bromine gas is passed over it. A brown solid is formed.

Iodine

Hot iron wool glows dimly when iodine vapour is passed over it. A brown solid is formed.

Conclusion 1. Chlorine, bromine and iodine show the same chemical properties when they react with iron wool, producing brown iron (II) halides. 2. The reactivity of the halogen decreases down the group from chlorine to iodine.

4N

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Paper 3 experiment

spm chemistry

Form 4 (5.1) Aim: To prepare ionic compounds Materials: Magnesium ribbon, sodium, chlorine gas, iron wool and sodium hydroxide solution Apparatus: Tripod stand, clay pipe triangle, Bunsen burner, crucible and lid, sandpaper, gas jar, gas jar spoon, combustion tube, filter funnel, retort stand , clamp and beaker (A) Preparation of magnesium oxide Procedure 1. A 5cm length of magnesium ribbon is cleaned with a piece of sandpaper. 2. The magnesium ribbon is placed in the crucible. 3. The magnesium ribbon is heated strongly. Any changes that occur are recorded.

(B) Preparation of sodium chloride 1. A small piece of sodium metal is placed in a gas jar spoon and is heated carefully until it begins to ignite. 2. The ignited sodium is placed in a gas jar filled with chlorine gas. Any changes that occur are recorded.

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Paper 3 experiment

spm chemistry

(C) Preparation of iron( III) chloride 1. A little iron wool is placed inside a combustion tube. 2. The end of the combustion tube is connected to a filter funnel inverted into a beaker with some sodium hydroxide solution. 3. The iron wool is heated strongly until it glows. 4. Chlorine gas is passed through the iron wool while being heated. Any changes that occur are recorded.

Method

Observation

Heating of magnesium in air

•

Burning of sodium in chorine gas

•

•

• •

Heating of iron in chlorine gas

• •

Inference

The magnesium ribbon burns with a bright flame. White powder is formed. 4N

The white powder formed is magnesium oxide.

Sodium burns with a bright yellow flame. The yellowish-green colour of chlorine gas is decolourised. White fumes are produced and deposited as white powder.

The white powder formed in sodium chloride.

The iron wool continues to glow brightly in the chlorine gas A brown powder is formed

The brown powder formed is iron(III) chloride.

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Paper 3 experiment

spm chemistry

Conclusion 1. Generally, the reaction between metals and non-metals produces ionic compounds. 2. Ionic compounds such as magnesium oxide, sodium chloride and iron(III) chloride can be prepared by direct combination of the metal and non-metal elements. Metal Non-metal Ionic compound Magnesium + oxygen → magnesium oxide Sodium + chlorine → sodium chloride Iron + chlorine → iron(III) chloride

4N

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Paper 3 experiment

spm chemistry

Form 4 (6.1) Aim: To identify electrolytes and non-electrolytes

Problem statement: How to identify electrolytes and nonelectrolytes? Hypothesis: Substances that, in the molten state or in aqueous solution, conduct electricity and then undergo chemical reactions are electrolytes. Substances that do not conduct electricity in any state are non-electrolytes.

Variables (a) Manipulated variable : Types of compounds (b) Responding variable: Electrical conductivity (c) Constant variable: Numbers of batteries, type of light bulb and amount of substance used Apparatus: Crucible, spatula, carbon (graphite) electrodes, batteries, light bulb, switch, rheostat, connecting wires, tripod stand, clay pipe triangle and Bunsen burner Materials: Glucose, naphthalene, lead(I1) bromide and potassium iodide. Procedure (A) To investigate the electrical conductivity of substances in the solid state and in the molten state 1. A crucible is half-filled with lead(II) bromide solid. 2. The crucible with its contents is placed on a clay triangle on a tripod stand. 3. Two carbon electrodes are dipped in the lead (II) bromide solid and are connected to the batteries, rheostat, switch and a light bulb with connecting wires (Figure 6.1.) 4. The switch is turned on and the light bulb is checked if it lights up.

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Paper 3 experiment 5. 6.

spm chemistry

The lead(II) bromide solid in the crucible is heated up until itt melts. The switch is turned on again to check if the light bulb lights up. Steps 1 to 5 of the experiment are repeated using naphthalene in place of lead(II) bromide.

(B) To investigate the electrical conductivity of substances in the solid state and in aqueous solutions 1. Three spatulas of potassium iodide solid are put in a beaker. 2. Two carbon electrodes are dipped in the potassium iodide solid and then connected to the batteries, rheostat, switch and a 4N light bulb with connecting wires (Figure 6.2). 3. The switch is turned on and the light bulb is checked if it lights up. 4. Distilled water is added to the beaker and the mixture is stirred until all the potassium iodide has dissolved. 5. The switch is turned on again and the light bulb is checked if it lights up. 6. Steps 1 to 5 of the experiment is repeated using glucose in place of potassium iodide.

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Paper 3 experiment

spm chemistry

Results Chemical substances

Physical state

Does the light bulb lights up?

Observation Does reaction occur?

Inference

Lead(II) bromide

Solid

No

No noticeable change

Nonelectrolyte

Liquid (molten)

Yes

Brown gas is evolved

Electrolyte

Solid

No

No noticeable change

Nonelectrolyte

Liquid (molten)

No

No noticeable change

Nonelectrolyte

Solid

No

No noticeable change

Nonelectrolyte

Aqueous solution

Yes

Solution turns to a brown colour

Electrolyte

Solid

No

No noticeable change

Nonelectrolyte

Aqueous solution

No

No noticeable change

Nonelectrolyte

Naphthalene

Potassium iodide

Glucose

4N

Conclusion 1. Lead(II) bromide is an electrolyte in the liquid but not in the solid state. 2. Potassium iodide is an electrolyte in aqueous solution but not in the solid state. 3. Lead(II) bromide and potassium iodide are ionic compound. Ionic compounds are electrolytes in the molten state or aqueous solution but are non-electrolytes in the solid state. 4. Naphthalene and glucose are covalent compounds and are non-electrolytes in any state.

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Paper 3 experiment

spm chemistry

Form 4 (6.2) Apparatus: Crucible, spatula, graphite electrodes, batteries, light bulb, ammeter, switch, rheostat, connecting wires, tripod stand, clay pipe triangle and Bunsen burner. Materials: Lead(II) bromide. Procedure 1. A crucible is half-filled with lead(II) bromide solid. 2. The solid lead(II) bromide is heated strongly until it melts to a molten state. 3. Two carbon electrodes are dipped in the molten lead(II) bromide and are then connected to batteries, rheostat, switch and light bulb by the connecting wires (Figure 6.4). 4. Electric current is allowed to flow through for 15 minutes and the changes that occur at the light bulb, ammeter, cathode and anode are recorded.

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Paper 3 experiment

spm chemistry

Results Apparatus

Observation

Inferences

Light bulb

Light bulb lights up

Ammeter

Ammeter neddle is deflected.

Molten lead (II) bromide conducts electricity.

Anode

Pungent brown gas that changes damp blue litmus paper to red is evolved

Bromine gas evolved

Cathode

Shiny grey metal id disposed

Lead metal is formed

Conclusion 1. The lighting up of the bulb and the deflection of the ammeter needle shows that molten lead(II) bromide is an electrolyte and can conduct electricity. 2. Electrolysis of molten lead(II) bromide produces bromine gas at the anode and lead metal at the cathode.

4N

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Paper 3 experiment

spm chemistry

Form 4 (6.3) Aim: To investigate the effect of the concentration of ions on the selective discharge of ions and the products of electrolysis of aqueous solutions. Problem statement: How does the concentration of ions determine the types of ions discharged during electrolysis? Hypothesis: lons of higher concentration will be selectively discharged during electrolysis.

Variables (a) Manipulated variable : Concentration of ions in the solution (b) Responding variable: Types of ions to be discharged at the anode and cathode (c) Constant variable: Types of ions in the electrolyte, types of electrodes, duration of electrolysis Apparatus: Batteries, electrolytic cell, carbon electrodes, ammeter, switch, connecting wires with crocodile clips and test tubes Materials: Aqueous 0.2 mol dm-3 copper(II) chloride, CuCl2 solution, and aqueous 0.001 mol dm-3 copper(II) chloride solution. Procedure 1. Concentrated aqueous copper(II) chloride solution of 2.0 mol dm-3 is put into an electrolytic cell with carbon electrodes.

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Paper 3 experiment 2. 3.

4. 5.

spm chemistry

Two test tubes, filled with copper(ll) chloride solution are inverted over the carbon anode and cathode respectively (Figure 6.6). The switch is turned on and electric current is allowed to flow for 15 minutes. Any change in colour of the electrolyte and any other changes that occur around the carbon electrodes are recorded. Steps 1 to 4 of the experiment are repeated using the dilute copper(II) chloride solution of 0,001 mol dm-3 to replace the concentrated copper(lI) chloride solution.

Results Electrolyte

Observation

Concentrated copper(II) chloride solution of 2.0 mol dm-3

At the cathode • Brown deposit is formed

Copper metal is produced.

At the anode • Bubbles of pungent greenish-yellow gas are produced. The 4Ngas turns the damp blue litmus paper to red and then bleaches it

Chlorine gas is produced.

Colour of electrolyte • The blue colour of the solution becomes paler

Concentration of Cu2+ ion in copper(II) chloride solution decreases.

At the cathode • Brown deposit is formed

Copper metal is produced.

At the anode • Bubbles of colourless gas are produced. • The gas lights up a glowing wooden splint

Oxygen gas is produced.

Colour of electrolyte • The blue colour of the solution becomes paler

Concentration of Cu2+ ion in copper(II) chloride solution decreases.

Diluted copper(II) chloride solution of 0.001 mol dm-3

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Inference

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Paper 3 experiment

spm chemistry

Conclusion 1. In the electrolysis of concentrated aqueous copper(II) chloride solution, copper metal is produced at the cathode and chlorine gas is produced at the anode. At the anode, the Clions are selectively discharged, producing chlorine gas because the concentration of Cl- ions is higher than that of OHions. 2. In the electrolysis of dilute aqueous copper(II)chloride solution, copper metal is produced at the cathode and oxygen gas is produced at the anode. At the anode, OH- ions are selectively discharged, producing oxygen gas because the concentration of Cl- ions is low. 3. The type of ions that is selectively discharged at the electrode is determined by the concentration of the ions. The hypothesis is accepted.

4N

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Paper 3 experiment

spm chemistry

Form 4 (6.4) Aim: To investigate the effect of the types of electrodes on the selective discharge of ions and the products of electrolysis of aqueous solution. Problem statement: How do the types of electrodes determine the types of ions discharged during electrolysis? Hypothesis: The products of electrolysis of copper(II) sulphate solution with copper electrodes are different from that with carbon electrodes. Variables (a) Manipulated variable : Types of electrodes (b) Responding variable: Products of electrolysis (c) Constant variable: Types of ions in the electrolyte and the concentration of ions Apparatus: Batteries, electrolytic cell, carbon electrodes, copper electrodes, ammeter, switch, rheostat, connecting wires with crocodile clips and test tubes. Materials: Aqueous 1.0 mol dm-3 copper(II) sulphate, CuSO4 solution Procedure

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Paper 3 experiment 1. 2. 3. 4. 5.

spm chemistry

Aqueous 1.0 mol dm copper(II) sulphate solution is put into an electrolytic cell with carbon electrodes. A test tube filled with copper(II) sulphate solution is inverted over the carbon anode (Figure 6.7). The switch is turned on and the electric current is allowed to flow for 15 minutes. Any change in colour of the electrolyte and any other changes that occur around the carbon electrodes are recorded. Steps 1 to 4 of the experiment are repeated using copper electrodes to replace carbon electrodes.

Result Types of electrodes

Carbon

Copper

Observation

Inference

At the cathode: Brown deposit is formed

Copper metal is deposited.

At the anode: Bubbles of colourless gas are produced The gas lights up a4N glowing wooden splint

Oxygen gas is produced.

Colour of electrolyte: The blue colour of the solution becomes paler

Concentration of Cu2+ ion decreases.

At the cathode: Formation of brown deposit makes the cathode thicker

Copper metal is produced.

At the anode: Anode corrodes and becomes thinner

Copper anode dissolves to form Cu2+ ions

Colour of electrolyte: The blue colour of the solution remains unchanged

Concentration of Cu2ions in copper(II) sulphate remains constant

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Paper 3 experiment

spm chemistry

Conclusion 1. In the electrolysis of aqueous copper(II) sulphate solution: (a) If a carbon electrode is used as the anode, OH- ions are discharged and oxygen gas is produced. (b) If a copper electrode is used as the anode, both OH- ions and SO42- ions are not discharged. Instead the copper anode dissolves to produce Cu2+ ions. (c) Cu2+ ions are discharged at the cathode producing copper metal whether the cathode used is a carbon electrode or a copper electrode. 2. The types of electrodes used during electrolysis determine the types of ions discharged and the products of electrolysis. The hypothesis is accepted.

4N

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Paper 3 experiment

spm chemistry

Form 4 (6.5) Aim: To investigate the electroplating of an iron spoon with copper using electrolysis Problem statement: How is electrolysis used to electroplate an iron spoon with copper metal? Hypothesis: Electroplating of an iron spoon with copper occurs if the iron spoon is used as the cathode, copper metal is used as the anode and aqueous copper(II) sulphate solution as the electrolyte.

Variables (a) Manipulated variable : The position of the iron spoon as an electrode (b) Responding variable: The deposition of copper on the iron spoon (c) Constant variable: Type of electrolyte and arrangement of apparatus Apparatus: Batteries, electrolytic cell, beaker, connecting wires with crocodile clips, ammeter and rheostat. Materials: 0.5 mol dm-3 copper(II) sulphate solution, copper plate and iron spoon. Procedure 1. About 200 cm 3 of 0.5 mol dm-3 copper(II) sulphate solution is poured into a beaker.

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Paper 3 experiment 2. 3.

4.

5.

spm chemistry

An iron spoon is polished using sandpaper and is connected to the negative terminal of the batteries. The spoon acts as the cathode. A piece of copper metal, as the anode, is connected to the positive terminal of the batteries. The circuit is completed using the connecting wires, rheostat and ammeter. The iron spoon and the copper metal are immersed in the copper(II)sulphate solution. The solution is electrolysed for30 minutes using a small current (0.5 A). Steps I to 4 of the experiment are repeated by interchanging the positions of the iron spoon and copper metal, whereby the iron spoon is made the anode and the copper metal is made the cathode.

Results Set

Observation

Set 1: Iron spoon as the cathode, copper metal as the anode

At the cathode • A brown metal is deposited on the surface of the iron 4N spoon

The iron spoon is plated with copper metal

At the anode • The copper anode becomes thinner

The copper anode dissolves to form Cu2+ ions

Colour of electrolyte • Colour intensity of the blue solution does not change

Concentration of Cu2+ ions in the electrolyte remains constant

At the cathode • The copper plate becomes thicker

Copper metal is deposited on the copper electrode

At the anode • No noticeable change in the appearance of the iron spoon

Electroplating of copper on the iron spoon does nnot take place

Colour of electrolyte • The blue colour of the solution becomes paler

Concentration of Cu2+ ions in the electrolyte decreases.

Diluted copper(II) chloride solution of 0.001 mol dm-3

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Inference

33

Paper 3 experiment

spm chemistry

Conclusion 1. In electroplating an iron spoon with copper using electrolysis, the on spoon is made the cathode and a piece of copper metal is made the anode. 2. Copper metal is transferred from the copper anode to the iron spoon and is deposited as a thin layer of copper metal. 3. Electroplating does not take place if the iron spoon and is made the anode. The hypothesis is accepted.

4N

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34

Paper 3 experiment

spm chemistry

Form 4 (6.6) Aim: To investigate the production of electricity from chemical reactions in a simple voltaic cell Problem statement: How does a chemical reaction produce electrical energy in a simple voltaic cell? Hypothesis: Electric current is produced when two different metals connected by wires are immersed in an electrolyte. Variables (a) Manipulated variable : Pairs of different metals (b) Responding variable: Deflection of a voltmeter needle by the electric (c) Constant variable: Types of electrolyte and arrangement of apparatus Apparatus: Voltmeter, beaker, connecting wires with crocodile clips and sandpaper.

Materials: 1 mol dm-3 sodium chloride solution, copper plates and magnesium plate. Procedure

1. 2.

A piece of magnesium plate and a piece of copper plate are polished with sandpaper. Both pieces of the magnesium and copper plates are immersed in 200 cm3 of aqueous sodium chloride solution in a beaker as shown in Figure 6.12.

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35

Paper 3 experiment 3. 4.

spm chemistry

Both plates are connected by the connecting wire to a voltmeter. The experiment is repeated using two pieces of copper plates as electrodes.

Results Type of metal used as electrodes

Observation

Magnesium metal and copper metal

•

• •

Two pieces of copper metal

• •

Inference

Voltmeter needle deflects but the deflection decreases after awhile Magnesium metal corrodes Bubbles of colourless gas are evolved around the copper metal

•

4N does Voltmeter needle not show a deflection No noticeable change occurs at the copper electrode

•

•

•

•

Electric current is produced. The voltage produced is not constant and decreases rapidly Magnesium dissolves to form Mg2+ ions Hydrogen gas is produced Electric current is not produced No reaction occurs

Conclusion 1. An electric current is produced when a chemical reaction occurs in a simple voltaic cell consisting of two different metals, connected by wires externally and immersed in an electrolyte. 2. In a simple voltaic cell, chemical energy released from chemical reactions is converted into electrical energy. 3. No electric current will be produced if both electrodes are of the same material because there is no potential difference between them. The hypothesis is accepted.

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36

Paper 3 experiment

spm chemistry

Form 4 (6.7) Aim: To construct the electrochemical series from displacement reaction. Problem statement: How to construct the electrochemical series based on the ability of a metal to displace another metal from its salt solution? Hypothesis: A metal that can displace another metal from its salt solution is placed at a higher position in the electrochemical series. The greater the number of metals that can be displaced by a metal from their solutions, the higher its position in the electrochemical series. Variables (a) Manipulated variable : Different types of metal and their salt solution (b) Responding variable: Deposition of metals or colour change in the salt solution (c) Constant variable: Concentration of nitrate salt solutions Apparatus: Test tubes, test-tube rack and sandpaper. Materials: Pieces of magnesium, zinc, iron, tin, lead and copper metals, solutions of copper(I) nitrate, lead(lI)nitrate, tin(lI) nitrate, iron(II) nitrate, zinc nitrate and magnesium nitrate (concentration and volume of all salt solutions are 0.5 mol dm-3 and 10 dm-3 respectively).

Procedure 1. Pieces of magnesium, zinc, copper, tin, lead and iron metals are polished with sandpaper. 2. 10 cm of 0.5 mol dm-3 solutions of copper(II)nitrate, lead(lI) nitrate, tin(I) nitrate, iron(lII)nitrate, zinc nitrate and magnesium nitrate are placed into separate test tubes. 3. A piece of magnesium metal is placed in the solution of every test tube except that of its salt solution (Figure 6.21).

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37

Paper 3 experiment

4.

5. 6.

7.

spm chemistry

Observations are made after awhile to check if (a) there is any colour change in the solution, (b) there are any solid deposits on the magnesium metal (c) magnesium metal dissolves If any of the above occurrences (a), (b) or (c) is observed, displacement reaction has taken place: a tick symbol, (√) is marked in the table of results. If there is no noticeable observation, a cross symbol, (X) is marked at the table to indicate 4N that displacement reaction did not take place. The experiment is repeated using different metals and fresh solutions of ions. The results of the experiment are shown in the table below.

Results Solution Metal

Cu(NO3)

Pb(NO3)2

Sn(NO3)2

Fe(NO3)2

2

Zn(NO3)

Mg(NO3)2

2

Magnesium , Mg

√

√

√

√

√

-

Zinc, Zn

√

√

√

√

-

X

Iron, Fe

√

√

√

-

X

X

Tin, Sn

√

√

-

X

X

X

Lead, Pb

√

-

X

X

X

X

Copper, Cu

-

X

X

X

X

X

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38

Paper 3 experiment

spm chemistry

Conclusion 1. Metals can be arranged according to the number of tick symbols (√) recorded (or the number of metals displaced in reactions). The more (√)symbols, the more reactive the metal is and the position of the metal is placed higher in the electrochemical series. 2. Magnesium is placed at the highest position in the electrochemical series because it can displace all the other metals from their solutions. 3. Copper is placed at the lowest position in the electrochemical series because copper cannot displace any other metals in this experiment. 4. The result of the experiment shows that the order of the order of the positions of the metals in the electrochemical series is: Mg Zn Fe Sn Pb Cu Electropositivity of metal decreases 5. 6.

The electrochemical series can be constructed from displacement reactions. The hypothesis is accepted. 4N

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39

Paper 3 experiment

spm chemistry

Form 4 (7.1) Aim: To investigate the role of water in showing the properties of alkali. Problem statement: Is water essential for an alkali to show its alkali properties? Hypothesis: An alkali will only show its acidic properties when dissolved in water. Variables (a) Manipulated variable : Types of solvents-water and propanone (b) Responding variable: Change in the colour of red litmus paper (c) Constant variable: Type of acid and red litmus paper Apparatus: Test tube and droppers. Materials: Dry ammonia gas stopped in a test tube, ammonia gas dissolved in propanone, aqueous ammonia solution and red litmus paper Procedure 1. A piece of dry red litmus paper is placed in a stoppered test tube of dry ammonia gas and the test tube is then stoppered back immediately (Figure 7.6). 2. The effect of the dry ammonia gas on the red litmus paper is recorded. 3. Another piece of dry red litmus paper is put in 5 cm3 of aqueous ammonia solution in a separate test tube. 4. Step 3 of the experiment is repeated using ammonia dissolved in propanone to replace aqueous ammonia solution.

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40

Paper 3 experiment

spm chemistry

Results Condition of ammonia

Observation

Inference

Dry

No noticeable colour change in the red litmus paper

Does not show any alkaline properties

Aqueous (dissolves in water)

Red litmus paper has changed to blue

Show alkaline properties

Dissolve in propanone

No noticeable colour change in the red litmus paper

Does not show any alkaline properties

Conclusion 1. Aqueous ammonia solution turns the red litmus paper to blue, indicating its alkaline property. 2. Dry ammonia gas or ammonia gas dissolved in organic solvents does not show any alkaline property. 3. An alkali shows its alkaline properties only in the presence of water. When water is present,4N ammonia ionises to produce OHions that are responsible for its alkaline properties. 4. Water is essential for the formation of hydroxide ions that cause alkalinity in an alkali. The hypothesis is accepted.

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41

Paper 3 experiment

spm chemistry

Form 4 (7.2) Aim: To prepare 100 cm3 of 2.0 mol dm-3 aqueous sodium hydroxide solution Apparatus: Electronic balance, 100 cm3 volumetric flask, filter funnel, dropper and washing bottle. Materials: Sodium hydroxide solid and distilled water. Procedure 1. The mass of sodium hydroxide (NaOH) required to prepare 100 cm3 of 2.0 mol dm-3 aqueous sodium hydroxide is calculated as follows: Mass of NaOH required MV =( ) x (23 +16+1) 1000

2.0 x 100

=( ) x 40 1000 = 8.0 g

2. 3. 4. 5.

8.0 g of sodium hydroxide, NaOH solid is weighed accurately in a weighing bottle using an electronic balance. Sodium hydroxide solid is transferred to a small beaker. Sufficient distilled water is added to dissolve all the solid sodium hydroxide. Using a filter funnel and glass rod, the dissolved sodium hydroxide is transferred to a 100 cm3 volumetric flask. The small beaker, the weighing bottle and the filter funnel are all rinsed with distilled water and the contents are transferred into the volumetric flask.

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42

Paper 3 experiment 6.

7.

spm chemistry

Distilled water is then distilled added slowly until the water level is near the level mark of the volumetric flask. A dropper is then used to add water drop by drop to finally bring the volume of solution to the 100 cm3 graduation The volumetric flask is then closed with a stopper. The volumetric flask is then shaken several times to mix the solution completely. The solution prepared is 100 cm3 of 2.0 mol dm-3 aqueous sodium hydroxide.

4N

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43

Paper 3 experiment

spm chemistry

Form 4 (7.3) Aim: To find the end point of an acid-base titration during neutralisation using an acid-base indicator cm3 Apparatus: 25 cm3 pipette, pipette filler, 50 cm3 burette, retort stand, retort clamp, conical flask, filter funnel and white tile. Materials: Sulphuric acid of unknown concentration, 1.0 mol dm-3 potassium hydroxide and methyl orange. Procedure 1. A clean 25 cm3 pipette is rinsed with distilled water and then rinsed with a little of the potassium hydroxide solution. 2. 25 cm3 of 1.0 mol dm-3 potassium hydroxide is transferred using the pipette to a clean conical flask. Three drops of methyl orange indicator are added to the alkali and the colour of the solution is noted. 3. A 50 cm3 burette is rinsed with distilled water and then rinsed with a little of the sulphuric acid. 4. The burette is then filled with sulphuric acid and is clamped to a retort stand. The initial burette reading is recorded. 5. The conical flask containing 25 cm3 of potassium hydroxide is placed below the burette. A piece of white tile is placed below the conical flask for clearer observation of the colour change(Figure 7.15). 6. Sulphuric acid is added slowly from the burette to the potassium hydroxide solution in the conical flask while swirling the flask gently. 7. Titration is stopped when the methyl orange changes colour from yellow to orange. The final burette reading is recorded.

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44

Paper 3 experiment 8.

spm chemistry

Steps 1 to 7 are repeated until accurate titration values are obtained, that is, until the difference in the volumes of sulphuric acid used in two consecutive experiments is less than 0.10 cm3 .

Results Volume of sulphuric acid

Rough

Accurate

Final burette reading (cm3)

21.00

40.95

20.15

Initial burette reading (cm3)

0.00

21.00

0.10

Volume of sulphuric acid used (cm3)

21.00

19.95

20.05

Conclusion 1. The volume of sulphuric acid used is calculated as follows: Volume of sulphuric acid used = Final burette reading – Initial burette reading 2. Average volume of sulphuric acid used 19.95+20.05 4N =( ) 2 = 20.00 cm3

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45

Paper 3 experiment

spm chemistry

Form 4 (8.1) Aim: To prepare potassium chloride by the reaction between an acid and alkali Apparatus: 25 cm3 pipette, pipette filler, 50 cm3 burette, retort stand, retort clamp, conical flask, filter funnel, filter paper, beaker, tripod stand, wire gauze and Bunsen burner Materials: 2 mol dm-3 hydrochloric acid and 2 mol dm-3 potassium hydroxide and phenolphthalein indicator.

Procedure 1. 25 cm3 of potassium hydroxide is pipetted into a clean conical flask. 2. Three drops of phenolphthalein indicator are added to the alkali and the colour of the solution is noted. 3. A 50 cm3 burette is then filled with hydrochloric acid and is then clamped to a retort stand. The initial burette reading is recorded. 4. Hydrochloric acid is added gradually from the burette to the potassium hydroxide solution in the conical flask while swirling the flask gently. 5. Titration is stopped when phenolphthalein changes from a light pink colour to colourless. The final burette reading is recorded. 6. The volume of hydrochloric acid used is calculated as follows: V cm3 = Final burette reading –Initial burette reading

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46

Paper 3 experiment

spm chemistry

7.

The experiment is repeated by adding V cm3 of hydrochloric acid to 25 cm3 of potassium hydroxide in a beaker without using phenolphthalein as an indicator. 8. The colourless solution in the beaker is evaporated to form a saturated solution (to about 1/3 of the original volume). This can be tested by dropping a drop of the solution on a piece of glass plate. If crystals are formed, then the solution is saturated. 9. The saturated solution is then cooled to allow crystallisation to occur. 10. The white crystals formed are then filtered, rinsed with a little distilled water and dried by pressing between filter paper.

4N

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47

Paper 3 experiment

spm chemistry

Form 4 (8.2) Aim: To construct a balanced ionic equation for the precipitation of lead(II) chromate(VI) using the continuous variation method Problem statement: How to determine the ionic equation for the precipitation of lead(lI) chromate(VI)? Hypothesis: The height of precipitate will increase with the increase in volume of lead(II) nitrate solution until all the potassium chromate(VI) has reacted.

Variables (a) Manipulated variable: Volumes of lead(II) nitrate solution (b) Responding variable : Height of yellow precipitate (c) Constant variable: Volume of potassium chromate(VI) solution and the size of test tubes Apparatus: Test tubes of the same size, test tube rack, 50 cm3 burette, retort stand with clamp and ruler.

Materials: 0.5 mol dm-3 lead(II) nitrate solution and 0.5 mol dm-3 potassium chromate(VI) solution. Procedure 1. A burette is filled with 0.5 mol dm-3 lead(II) nitrate solution and another burette is filled with 0.5 mol dm-3 potassium chromate(VI) solution. 2. Eight test tubes are labelled 1 to 8 and placed in a test tube rack. 3. 5.00 cm3 of potassium chromate(VI) solution from the burette is placed in every test tube. Potassium chromate( VI) solution is yellow in colour. 4. Using another burette, 1 cm3 of 0.5 mol dm-3 of lead(II) nitrate solution is added to the first test tube. Progressively increase the volume of the lead(II) nitrate solution by 1 cm3 to the rest of the test tubes until 8 cm3 of lead(II) nitrate solution is added to the eighth test tube (Figure 8.6(a)).

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48

Paper 3 experiment 5.

6. 7.

spm chemistry

Every test tube is well shaken in order to mix the solutions completely. The test tubes are then allowed to stand for 20 minutes for the yellow precipitate, lead(II) chromate(VI) to settle (Figure 8.6(b)) The height of the precipitate formed in every test tube is measured accurately using a ruler. The colour of the solution above the precipitation is noted. The result obtained is recorded in Table 8.8.

Test tube number

1

2

3

4

5

6

7

8

Volume of potassium chromate (IV) solution (cm3 )

5.0

5.0

5.0

5.0

5.0

5.0

5.0

5.0

Volume of lead (II) nitrate solution (cm3 )

1.0

2.0

3.0

4.0

5.0

6.0

7.0

8.0

Height of precipitation (cm)

0.6

0.9

1.8

2.2

2.8

2.8

2.8

2.8

Colour of solution

yellow

yellow

yellow

yellow

4N

colourless

Conclusion 1. Since the diameter of the test tubes are the same, the height of the precipitate is directly proportional to the mass of precipitate formed. 2. The ionic equation for the precipitate of lead(lI)chromate(VI) is Pb2+ + CrO4 2- → PbCrO 4 . The hypothesis is accepted.

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49

Paper 3 experiment

spm chemistry

Form 4 (9.1) Aim: To compare the hardness of a pure metal and its alloy. Problem statement: Are alloys harder than pure metals? Hypothesis: Bronze is harder than copper. When a weight is dropped onto a ball bearing placed on a metal block made of copper or bronze, a larger dent will be produced on the softer copper metal block than on the bronze block Variables (a) Manipulated variable: Types of materials (copper or bronze) to make the metal block (b) Responding variable : Diameter of the dent made by a steel ball bearing (c) Constant variable: Size of steel ball bearing, mass of weight used, height from which it is dropped Materials : Copper block, bronze block, ball bearing, 1 kg weight, metre ruler, retort stand with clamp, cellophane tape and thread. Procedure 1. A metre ruler is clamped to a retort stand, and a piece of copper block is placed on the base of the retort stand. 2. A steel ball bearing is placed on the copper block and a piece of cellophane tape is used to hold the ball bearing in place. 3. A l kg weight is hung at a height of 50 cm above the copper block. 4. The weight is dropped onto the ball bearing placed on the copper block. 5. The diameter of the dent made by the ball is measured. 6. The experiment is repeated three times using different areas on the surface of the copper block. 7. The average diameter of the dent is calculated. 8. Steps 1 to 7 are repeated using a piece of bronze block.

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50

Paper 3 experiment 5.

6. 7.

spm chemistry

Every test tube is well shaken in order to mix the solutions completely. The test tubes are then allowed to stand for 20 minutes for the yellow precipitate, lead(II) chromate(VI) to settle (Figure 8.6(b)) The height of the precipitate formed in every test tube is measured accurately using a ruler. The colour of the solution above the precipitation is noted. The result obtained is recorded in Table 8.8.

Metal block

Diameter of the dent (mm) I

Copper

3.2

Bronze

2.4

II

III

Average

3.3

3.2

3.23

22.5

2.5

2.47

4N

Conclusion 1. The average diameter of the dents made by the steel ball bearing on the copper block is bigger than that on the bronze block. 2. Hence, bronze, a type of alloy, is harder than pure copper metal. The hypothesis is accepted.

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51

Paper 3 experiment

spm chemistry

Form 5 (1.1) Aim: To investigate the effect of the surface area of a reactant on the rate of reaction Problem statement: How does the surface area of a solid reactant affect the rate of reaction? Hypothesis: The smaller the size of the reactant particles, that is, the larger the total surface area of the reactant particles, the higher the rate of reaction. Variables (a) Manipulated variable : Size of the marble chips (b) Responding variable : Volume of gas given off at 30-second intervals (c) Constant variables : Temperature of the experiment, mass of marble chips, concentration and volume of hydrochloric acid Apparatus: Conical flask, delivery tube, retort stand and clamp, burette, measuring cylinder and stopwatch. Materials: Marble chips, powdered marble and 0.08 mol dm-3 hydrochloric acid Procedure

1.

2.

A burette is filled with water and inverted over a basin containing water. The burette is clamped to the retort stand. The water level in the burette is adjusted and the initial burette reading is recorded. 5.0 g of marble chips are placed in a small conical flask.

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52

Paper 3 experiment 3. 4. 5.

spm chemistry

50 cm3 of 0.08 mol dm-3 hydrochloric acid is added to the marble chips. The conical flask is then stoppered and the stopwatch is started immediately. The burette readings are recorded at 30-second conical flask.

Experiment I The rate of reaction using large marble chips Time (s)

0

30

60

90

120

150

180

210

240

Burette reading (cm3 )

50.0

45.5

41.5

38.0

35.0

33.0

31.0

29.0

28.0

Volume of gas (cm3 )

0.0

4.5

8,5

12.0

15.0

17.0

19.0

21.0

22.0

Experiment II The rate of reaction using powdered marble Procedure 1. Step 1 to 4 in Experiment I are repeated using 5.0 g of powdered marble. All other conditions such as temperature, 4N volume and concentration of hydrochloric acid are kept constant. 2. The results of the experiment are recorded in the following table. Time (s)

0

30

60

90

120

150

180

210

240

Burette reading (cm3 )

50.0

42.0

35.0

29.5

25.5

22.0

19.5

17.5

16.0

Volume of gas (cm3 )

0.0

8.0

15.0

20.5

24.5

28

30.5

32.5

34.0

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53

Paper 3 experiment

spm chemistry

Conclusion 1. The rate of reaction in Experiment II is higher than the rate of reaction in Experiment I as powdered marble is used in Experiment II, Thus, the rate is higher with powdered marble than with marble chips. Hence, we can conclude that the smaller the particle size, the larger the total surface area exposed for reaction and the higher the rate of reaction. 2. The hypothesis is accepted.

4N

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54

Paper 3 experiment

spm chemistry

Form 5 (1.2) Aim: To study the effect of concentration on the rate of reaction between sodium thiosulphate solution and dilute sulphuric acid. Problem statement: How does the concentration of a reactant affect the reaction between sodium thiosulphate and dilute suhphuric acid? Hypothesis: The more concentrated the sodium thiosulphate solution, the higher the rate of reaction. Variables (a) Manipulated variable : Concentration of sodium thiosulphate solution (b) Responding variable : Time taken for the cross ‘X’ to disappear (c) Constant variables : Concentration and volume of dilute sulphuric acid as well as the temperature of the solutions Apparatus: 10 cm3 and 100 cm3 measuring cylinders, 100 cm conical flask, white paper marked with a cross ‘X’ and stopwatch Materials: 0.2 mol dm-3 sodium thiosulphate solution, 1.0 mol dm-3 sulphuric acid and distilled water Procedure

1. 2.

50 cm3 of 0.2 mol dm-3 sodium thiosulphate solution is measured out using a 100 cm' measuring cylinder. The solution is then poured into a clean, dry conical flask. The conical flask is placed on a piece of paper with a cross ‘X' marked on it (Figure 1.17).

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55

Paper 3 experiment 3.

spm chemistry

5 cm3 of dilute sulphuric acid is measured out by using a 10 cm3 measuring cylinder. The acid is then quickly poured into sodium thiosulphate solution. The stopwatch is started immediately The reaction mixture is swirled once and the cross ‘X' is viewed from above. A yellow precipitate will appear slowly in the conical flask. The stopwatch is stopped as soon as the cross disappears from view and the time taken is recorded. Steps 1 to 5 are repeated with different mixtures of sodium thiosulphate solution and distilled water as shown in the following table.

4. 5. 6.

Experiment

1

2

3

4

5

Volume of Na2S2O3 (cm3)

50

40

30

20

10

Volume of water (cm3)

0

10

20

30

40

Volume of H2SO4(cm3)

5

5

5

5

5

0.20

0.16

0.12

0.08

0.04

24

30

42

62

111

0.042

0.033

0.024

0.016

0.009

Concentration of Na2S2O3 (cm3) Time taken (s) 1 𝑇𝑖𝑚𝑒

(s-1)

4N

Conclusion 1. The more concentrated the sodium thiosulphate solution, the higher the rate of reaction. 2. The hypothesis is accepted.

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56

Paper 3 experiment

spm chemistry

Form 5 (1.3) Aim: To study the effect of temperature on the rate of reaction between sodium thiosulphate solution and dilute sulphuric acid

Problem statement: How does temperature affect the rate of reaction between sodium thiosulphate solution and sulphuric acid? Hypothesis: The higher the temperature of the reactant, the higher the rate of reaction. Variables (a) Manipulated variable: The temperature of sodium thiosulphate solution (b) Responding variable: The time taken for the cross ‘X’ to disappear (c) Constant variables: The concentrations and volumes of both sodium thiosulphate solution and dilute sulphuric acid Apparatus: Conical flask, 10 cm3 measuring cylinder, thermometer, stopwatch, white paper marked with a cross ‘X', wire gauze, tripod stand and Bunsen burner Materials: 0.1 mol dm-3 sodium thiosulphate solution and 1.0 mol dm-3 sulphuric acid. Procedure: 1. 50 cm3 of 0.1 mol dm-3 sodium thiosulphate solution is poured into a clean, dry conical flask. 2. The temperature of the sodium thiosulphate solution is measured with a thermometer. 3. The conical flask is placed on a white paper marked with a cross 'X' (Figure 1.20). 4. 5 cm3 of 1 mol dm-3 sulphuric acid is quickly poured into the sodium thiosulphate solution.

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57

Paper 3 experiment 5. 6. 7. 8.

9. 10. 11. 12. 13. 14.

spm chemistry

The stopwatch is started immediately and the conical flask is swirled gently. The cross ‘X' is viewed from above. The stopwatch is stopped as soon as the cross disappears from view and the time taken is recorded. The solution in the conical flask is poured out. The conical flask is washed thoroughly and dried. 50 cm3 of 0.1 mol dm-3 sodium thiosulphate solution is poured into the conical flask. The solution is heated over a wire gauze until the temperature reaches about 45 °C (Figure 1.21).

The hot conical flask is placed over a white paper with a cross ’X’. 5 cm3 of 1 mol dm-3 sulphuric acid is measured out using a 10 cm3 measuring cylinder. When the temperature of sodium thiosulphate solution falls to 40°C, the sulphuric acid is quickly poured into the thiosulphate solution . The stopwatch is started immediately and the conical flask is swirled gently. The cross ‘X, is viewed from the top and the time taken for the cross to disappear from view is recorded. Steps 7 to13 are repeated at higher temperatures as shown in the following table.

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Paper 3 experiment

spm chemistry

Results

Experiment

1

2

3

4

5

Temperature (°C)

30

40

50

55

60

Time (s)

52

27

16

12.0

15.0

1 (s-1 ) 𝑇𝑖𝑚𝑒

0.019

0.037

0.063

0.077

0.100

Based on the results of the experiment, a graph of temperature of 1 sodium thiosulphate solution against is plotted (Figure 1.22). 𝑇𝑖𝑚𝑒

Conclusion The higher the temperature of the experiment, the higher the rate of reaction.

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59

Paper 3 experiment

spm chemistry

Form 5 (1.4) Aim: To study the effect of a catalyst on the rate of decomposition of hydrogen peroxide

Problem statement: How do catalysts affect the rate of decomposition of hydrogen peroxide? Hypothesis: Manganese(IV) oxide increases the rate of decomposition of hydrogen peroxide. Variables (a) Manipulated variable: The temperature of sodium thiosulphate solution (b) Responding variable: The release of oxygen gas (c) Constant variables: Volume and concentration of hydrogen peroxide Apparatus: Test tube and wooden splint Materials: Hydrogen peroxide and manganese(IV) oxide

Procedure: 1. A test tube is half-filled with hydrogen peroxide. 2. A glowing splint is placed at the mouth of the test tube to test for the gas evolved (Figure 1.23).

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60

Paper 3 experiment 3. 4.

spm chemistry

The changes that take place inside the test tube and on the glowing splint are recorded. 0.5 g of manganese(IV) oxide, MnO2 is added to hydrogen peroxide and shaken. The changes that take place in the test tube and on the glowing splint are recorded.

Experiment

Observation

Inside the test tube

On the glowing splint

H2 O2 without MnO2

No effervescence.

The glowing splint does not light up.

H2 O2 with MnO2

Bubbles of oxygen gas are produced.

The glowing splint is rekindled and burns brightly.

Conclusion The rate of evolution of oxygen gas increases when manganese(IV) oxide is added to hydrogen peroxide. This proves that manganese(TV) oxide acts as a catalyst and speeds up the decomposition of hydrogen peroxide to water and oxygen. The hypothesis is accepted.

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61

Paper 3 experiment

spm chemistry

Form 5 (2.1) Aim: To compare the chemical properties of alkanes and alkenes having the same number of carbon atoms.

Apparatus: Porcelain dish, wooden splint, dropper and Bunsen burner Materials: Hexane, hexene, liquid bromine and acidified potassium manganate(IV) solution. (In this experiment, hexane (C6H12) is used to represent an alkene) (A) Combustion of alkanes and alkenes in air

Procedure: 1. About 1 cm3 of hexane and hexene are placed separately in two porcelain dishes. 2. The organic liquids are ignited with a glowing splint as shown in Figure 2.8. 3. A filter paper is placed above the flame and the sootiness of the flame is observed. (B) Reactions with bromine Procedure 1. About 1 cm3 of liquid bromine is added to a test tube. 2. About 2 cm3 of hexene is the added to the liquid bromine. 3. The mixture is shaken gently. 4. The colour change that takes place in the test tube is recorded. 5. Steps 1 to 4 are repeated using hexene.

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62

Paper 3 experiment

spm chemistry

(C) Reaction with acidified potassium manganate(VII) solution 1. A few drops of potassium manganate(VII)solution are added to a test tube. About 1 cm3 of dilute sulphuric acid is then added to the KMnO4 solution. 2. About 2 cm3 of hexane is added to the acidified potassium manganate(VII) solution prepared in step 1. 3. The mixture is shaken gently. 4. The colour change that occurs in the test tube is recorded. 5. Steps 1 to 4 are repeated using hexene. Result

Test

Observation Hexane

Hexene

(A) Combustion

Burns in air with a sooty yellow flame.

Burns in air with a yellow flame which is more sooty.

(B) Reaction with liquid bromine

The brown colour of liquid bromine remains unchanged.

The brown colour of liquid bromine is decolourised.

(C) Reaction with acidified potassium manganate (VII) solution Conclusion

The purple colour of The purple colour of potassiunm potassiun manganate(VIl) manganate(VIl) is remains decolourised. The chemical properties unchanged. of alkenes are different from those of alkanes in terms of the sootiness of flame, reactions with liquid bromine and acidified potassium manganate(VII) solution.

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63

Paper 3 experiment

spm chemistry

Form 5 (2.2) Aim: To investigate the properties of ethanol. Apparatus: Evaporating dish, wooden splint, test tubes, boiling tube, glass wool, porcelain chips, retort stand with clamp and test tube holder. Materials: Ethanol, concentrated sulphuric acid, potassium dichromate(VI), blue litmus paper and liquid bromine. (A) Combustion of alkanes and alkenes in air

Procedure: 1. About 1 cm3 of ethanol is added to an evaporating dish. 2. The ethanol is ignited using a lighted wooden splint as show in Figure 2.12. 3. The flammability of ethanol and the nature of the flame are recorded. Results

Test

Observation

Flammability

Flammable

Colour of flame

Burns with pale blue flame

Sootiness of flame

Non-sooty flame

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Paper 3 experiment

spm chemistry

Conclusion Ethanol undergoes complete combustion to form carbon dioxide and water.

(B) Oxidation of ethanol

1. 2. 3. 4. 5. 6. 7.

The boiling tube is filled with approximately 5 cm3 of potassium dichromate(VI) solution, K2C2O7(aq). About 5 cm3 of dilute sulphuric acid is added to potassium dichromate(VI) solution. About 5 cm3 of ethanol is added to the acidified potassium dichromate(VI) solution. A rubber stopper fitted with a delivery tube is inserted into the boiling tube. The delivery tube is inserted into a test tube placed in a beaker half-filled with ice-cold water (Figure 2.13). The mixture of ethanol and acidified potassium dichromate(VI) is boiled slowly. The distillate is collected in the test tube. The colour and the odour of the distillate are recorded. The distillate is tested with a piece of blue litmus paper.

Results

Test

Observation on the distillate

Colour

Colourless

Smell

Vinegar smell

Action on blue litmus paper

Turns blue litmus paper red

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Paper 3 experiment

spm chemistry

Conclusion When ethanol is boiled with acidified potassium dichromate(VI) solution, it is oxidised to ethanoic acid which has the smell of vinegar. (C) Dehydration of ethanol

1. 2. 3. 4. 5. 6.

About 2 cm3 of ethanol is put into a dry test tube. A small amount of glass wool is inserted into the test tube to absorb ethanol. Some porcelain chips are placed in the middle section of the test tube. The test tube is closed with a rubber stopper fitted with a delivery tube. Another test tube is filled with water and inverted into a beaker as shown in Figure 2.14. The porcelain chips are heated strongly. The Bunsen burner flame is then shifted to the glass wool to vaporise the ethanol absorbed in it. The gas released is collected in two test tubes. The following tests are carried out on the gas collected. (a) The flammability of the gas (b) The reaction of the gas with liquid bromine.

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Paper 3 experiment

spm chemistry

Results

Test

Observation

Flammability of the gas

Burns easily with a yellow and sooty flame.

Liquid bromine

Decolourise liquid bromine immediately.

Conclusion When ethanol vapour is passed over porcelain chips(aluminium oxide), dehydration occurs and ethene is produced.

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67

Paper 3 experiment

spm chemistry

Form 5 (2.3) Aim: To prepare ethyl ethanoate Apparatus: Round-bottomed flask, Liebig condenser, tile chips, retort stand with clamp and oil bath. Materials: Pure ethanol, glacial ethanoic acid and concentrated sulphuric acid Procedure

Procedure: 1. About 30 cm3 of pure ethanol is placed in a round-bottomed flask, followed by about 25 cm3 of glacial ethanoic acid and 23 pieces of tile chips. The tile chips are added to prevent bumping and to ensure smooth boiling. 2. About 5 cm3 of concentrated sulphuric acid is added cautiously (slowly and carefully) to the reaction mixture. The mixture is shaken gently.(Caution! Concentrated sulphuric acid is very corrosive). 3. The Liebig condenser is fitted vertically to the round-bottomed flask as shown in Figure 2.21. The mixture of ethanol, ethanoic acid and concentrated sulphuric acid is boiled under reflux for about 30 minutes.

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Paper 3 experiment

spm chemistry

4.

After boiling, the mixture produced contains ethyl ethanoate together with the impurities: ethanol, ethanoic acid and concentrated sulphuric acid (the catalyst). 5. The impurities are then removed from ethyl ethanoate by distillation. Pure ethyl ethanoate is obtained. 6. From the ethyl ethanoate obtained, the following tests are carried out: a) The physical state, colour and odour of ethyl ethanoate are determined. b) By using a dropper, a few drops of ethyl ethanoate are added to a test tube filled with 2 cm3 of distilled water. The mmixture is shaken gently.

Conclusion Ethyl ethanoate is produced when ethanoic acid and ethanol are heated in the presence of concentrated sulphuric acid as a catalyst.

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69

Paper 3 experiment

spm chemistry

Form 5 (2.4) Aim: Comparing the properties of vulcanised rubber with unvulcanised rubber.

Apparatus: Clip, retord stand with clamp, weight, ruler and thread Materials: One strip each of vulcanised rubber and unvulcanised rubber. Procedure

Procedure: 1. The original length of the vulcanised rubber strip is measured. 2. A weight of 50 g is hung on the strip of the vulcanised rubber. 3. The increase in the length of the vulcanised rubber strip is measured. 4. The weight is removed and the final length of the vulcanised rubber strip is measured. 5. Steps 1 to 5 are repeated using unvulcanised rubber strip instead of vulcanised rubber strip.

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Paper 3 experiment

spm chemistry

Results

Type of rubber

Original length

Final length when weight is hung (cm)

Increase in length (cm)

Final length when weight is removed (cm)

Unvulcanise d rubber

𝑥

𝑥1

𝑥1-𝑥 = y2

𝑥1 where 𝑥2 > 𝑥

Vulcanised rubber

𝑥

𝑥3

𝑥3 -𝑥 = y2

𝑥4 where 𝑥4 = 𝑥

Conclusion Vulcanised rubber is more elastic than unvulcanised rubber.

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Paper 3 experiment

spm chemistry

Form 5 (3.1) Aim: To investigate the effect of other metals with different electropositivity on the rusting of iron

Problem statement: What is the effect of other metals with different electropositivity on the rusting of iron? Hypothesis a) A metal more electropositivity than iron will protect iron from rusting. b) A metal less electropositivity than iron will increase the rate of rusting.

Variables (a) Manipulated variable: Different metals used to wrap around iron nails (b) Responding variable : Colour change in the gelatin solution (c) Constant variable: Iron nails Apparatus: Test tubes

Materials: Iron nails, magnesium, zinc, tin and copper foils, gelatin, potassium hexacyanoferrate(III),phenolphthalein indicator and sandpaper. Procedure 1. Five pieces of iron nails are cleaned using sandpaper. 2. The first clean iron nail is placed in test tube A. 3. Strips of magnesium (Mg), zinc (Zn), tin (Sn) and copper (Cu) foils are cleaned with sandpaper. 4. Each iron nail is wrapped with a different metal foil and placed in test tubes B, C, D and respectively. 5. A solution of gelatin in hot water is prepared. A few drops of potassium hexacyanoferrate(III) solution, K3Fe(CN)6, and phenolphthalein indicator are added to the hot gelatin solution.

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Paper 3 experiment

6. 7.

spm chemistry

The mixture is stirred and then poured into each of the test tubes (Figure 3.14). The test tubes are set aside for three days and then examined. The observations are recorded in the table below.

Results Test tube

A

B

C

D

E

Fe

Fe + Mg

Fe + Zn

Fe + Sn

Fe + Cu

Low

None

None

High

High

Intensity of pink colour

None

High

High

Low

Low

Gas bubbles

None

Plenty

Plenty

Few

Few

Observation

Metal Intensity of blue colour

Conclusion 1. The rusting of iron can be prevented if iron is in contact with more electropositive metals such as magnesium or zinc. 2. The rusting of iron is speeded up if iron is in contact with less electropositve metals such as tin or copper. The hypothesis is accepted.

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Paper 3 experiment

spm chemistry

Form 5 (3.2) Aim: To deduce the reactivity series of metals Problem statement: How is the reactivity series of metals deduced from the reactions of metals with oxygen. Hypothesis The more reactive a metal, the more brightly and more rapidly the metal will burn in oxygen. Variables (a) Manipulated variable: Type of metal (b) Responding variable : The intensity of the flame (c) Constant variable: The amount of metal and potassium manganate (VII) used Apparatus: Combustion tube, retort stand with clamp, spatula and Bunsen burner Materials: Potassium manganate (VII), powdered zinc, iron, lead, copper, magnesium, glass wool and asbestos paper. Procedure 1. Two spatulas of potassium manganate(VII) crystals are placed in a combustion tube. A small quantity of glass wool is then placed inside the combustion tube to prevent potassium manganate(VI)from spilling over. (Caution! If potassium manganate(VII) is mixed with metal powder, an explosion may occur during heating). 2. A spatula of zinc powder is placed on a sheet of asbestos paper and put inside the combustion tube. The combustion tube is then clamped to a retort stand. 3. The zinc powder is heated strongly (Figure 3.20).

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Paper 3 experiment 4. 5. 6.

spm chemistry

When the zinc powder has become very hot potassium manganate(VII) is heated strongly to produce oxygen gas. The intensity of the flame or glow is recorded in the following table. When the reaction has been completed, the combustion tube is set aside to cool and the contents of the combustion tube taken out. Steps 1 to 5 are repeated using (a) iron powder (b) lead powder, (c) copper powder and (d) magnesium powder.

Results

Metal

Intensity of flame / glow

Observation Colour of hot oxide

Colour of cold oxide

Yellow

White

Reddishbrown

Reddishbrown

Zinc

• •

Burns rapidly Bright glow

Iron

• •

Burns less rapidly The glow id less bright than the burning of zinc

Lead

• •

Burns slowly Faint glow

High

Yellow

Copper

•

Faint glow

Black

Black

Magnesiu m

• •

Burns very rapidly Very bright white flame produced

White

White

Conclusion 1. Based on the results obtained in this experiment, the reactivity of the five metals with oxygen is as follows: Mg > Zn > Fe > Pb > Cu (Very reactive) (Very unreactive) Reactivity decreases The hypothesis is accepted.

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Paper 3 experiment

spm chemistry

Form 5 (4.1) Aim: To determine the heat of precipitation of silver chloride Apparatus: Measuring cylinders, thermometer and plastic cup. Materials: 0.5 mol dm-3 silver nitrate solution and 0.5 mol dm-3 sodium chloride solution Procedure

1. 2. 3. 4.

5.

25 cm3 of 0.5 mol dm-3 sodium chloride solution is measured and poured into a clean and dry plastic cup using a measuring cylinder. The initial temperature of sodium chloride solution is measured and recorded. Using another measuring cylinder, 25 cm3 of 0.5mol dm-3 silver nitrate solution is measured. The initial temperature of the silver nitrate solution is measured and recorded. The silver nitrate solution is poured quickly and carefully into the sodium chloride solution(Figure 4.17). The mixture is stirred with a thermometer throughout the experiment and the highest temperature obtained is recorded.

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Paper 3 experiment

spm chemistry

Temperature

NaCl (aq)

Highest temperature obtained (°C)

32.0

Initial temperature sodium/ potassium chloride solution(°C)

29.0

Initial temperature of silver nitrate solution(°C)

28.0

Conclusion The heat of precipitation of silver chloride is -58.87 kJ mol-1.

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Paper 3 experiment

spm chemistry

Form 5 (4.2) Aim: To determine the heat of neutralisation between an acid and an alkali

Hypothesis: How are heats of neutralisation determined and compared? Hypothesis: The heat of neutralisation between hydrochloric acid and sodium hydroxide is higher than the heat of neutralisation between ethanoic acid and sodium hydroxide. Variables (a) Manipulated variable: Different types of acids (b) Responding variable: Heat of neutralisation (c) Constant variable: Concentrations and volumes of acid and alkali used Apparatus: Thermometer, plastic cup and measuring cylinder. Materials: 2.0 mol dm-3 hydrochloric acid, 2.0 mol dm-3 ethanoic acid and 2.0 mol dm-3 sodium hydroxide solution. Procedure 1. 100 cm3 of 2 mol dm-3 sodium hydroxide solution into a plastic cup by using a measuring cylinder. The initial temperature pf the acid is recorded. 2. Using another measuring cylinder, 100 cm3 of 2 mol dm-3 hydrochloric acid is measured. The initial temperature of the acid is recorded. 3. The hydrochloric acid is then poured quickly and carefully into the sodium hydroxide solution. The mixture is stirred with a thermometer and the highest temperature obtained is recorded. 4. Steps 1 to 3 are repeated using 100 cm3 of 2 mol dm-3 acid instead of hydrochloric acid.

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Paper 3 experiment

spm chemistry

Results (A) Reaction between hydrochloric acid and sodium hydroxide solution

Initial temperature of sodium hydroxide solution

30 °C

Initial temperature of hydrochloric acid

31°C

Highest temperature obtained

43.5 °C

Average initial temperature of solutions before neutralisation 30+31 = = 30.5 °C 2 Rise in temperature during neutralisation = 43.5 – 30.5 = 13 °C (B) Reaction between hydrochloric acid and sodium hydroxide solution

Initial temperature of sodium hydroxide solution

30 °C

Initial temperature of ethanoic acid

30°C

Highest temperature reached

42 °C

Average initial temperature of solutions before neutralisation = 30 °C Rise in temperature = 42 – 30 = 12 °C Conclusion The heat of neutralisation for strong acids and strong alkalis in higher than the heat of neutralisation for weak acids and strong alkalis. The hypothesis is accepted.

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Paper 3 experiment

spm chemistry

Form 5 (4.3) Aim: To determine the heat of combustion of various alcohols Hypothesis: How does the number of carbon atoms per molecule of an alcohol affect the heat of combustion? Apparatus: As the number of carbon atoms per molecule in an alcohol increases, so does the heat of combustion. Variables (a) Manipulated variable: Type of alcohol (b) Responding variable : Heat of combustion (c) Constant variable: Volumes of water and alcohol, metal container (calorimeter) and spirit lamp Apparatus: Copper container, spirit lamp, measuring cylinder, thermometer, stirrer, electronic balance, tripod stand, asbestos screen, wooden block and Bunsen burner. Materials: Methanol, ethanol, propan-1-ol and butan-1-ol

Procedure

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Paper 3 experiment

spm chemistry

1.

Using a measuring cylinder, 250 cm3 of water is measured into a copper container. 2. The copper calorimeter is placed on the tripod stand. The initial temperature of water is measured and recorded. 3. The spirit lamp is halt-filled with ethanol. The spirit lamp and ethanol are weighed and the mass is recorded. 4. The spirit lamp is placed below the copper calorimeter and the wick is lighted (Figure 4.31). The flame of the spirit lamp is shielded from the draught (blow of wind) by using an asbestos screen. 5. The water in the calorimeter is stirred throughout the experiment. 6. When the temperature of water increases by about 30 ℃ the spirit lamp is extinguished. 7. The spirit lamp and ethanol are weighed again and the mass is recorded. 8. The experiment is repeated using other alcohols as shown below to replace ethanol: (a) Methanol (b) Propan-1-ol (c) Butan-1-ol Results Alcohol

Volum e of water

Initial temperatur e of water (°C)

Highest temperatur e of water (°C)

Rise in temperatur e (°C)

Initial mass of lamp + alcoho l (g)

Final mass of lamp + alcohol (g)

Ethanol

250

30.5 (t1)

59.5 (t2)

t2 – t1

218 (m1)

216.6 (m2)

Methanol

250

t3

t4

t4 – t3

m3

m4

Propan-1ol

250

t5

t6

t6 – t5

m5

m6

Butan-1-ol

250

t7

t8

t8 – t7

m7

m8

Conclusion 1. The heat of combustion of ethanol = -1371 kJ mol-1 2. The heat of combustion increases as the number of carbon atoms per molecule in the alcohol increases.

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