Practical notes Paper 6 notes

Cambridge IGCSE chemistry paper 6_practical notes Page | 223

T. Suleiman/ 0509441242 Topic 1. Practical notes Page | 223

Practical notes Paper 6 notes

Safety signs, safety precautions and lab apparatus on pages 1-4 Purification methods on pages ………. Graphs Unless instructed otherwise, the independent variable should be plotted on the x-axis (horizontal axis) and the dependent variable plotted on the y-axis (vertical axis). •• Each axis should be labelled with the physical quantity and the appropriate unit, e.g. time / s.

•• Unless instructed otherwise, the scales for the axes should allow more than half of the graph grid to be used in both directions, and be based on sensible ratios, e.g. 2 cm on the graph grid representing 1, 2 or 5 units of the variable. •• The graph is the whole diagrammatic presentation, including the best-fit line when appropriate. It may have one or more sets of data plotted on it. •• Points on the graph should be clearly marked as crosses (×) or encircled dots ( ).

•• Large ‘dots’ are penalised. Each data point should be plotted to an accuracy of better than one half of each of the smallest squares on the grid. •• A best-fit line (trend line) should be a single, thin, smooth straight-line or curve. The line does not need to coincide exactly with any of the points; where there is scatter evident in the data, Examiners would expect a roughly even distribution of points either side of the line over its entire length. Points that are clearly anomalous should be ignored when drawing the best-fit line. •• The gradient of a straight line should be taken using a triangle whose hypotenuse extends over at least half of the length of the best-fit line, and this triangle should be marked on the graph. Bets fit line Smooth straight –line or curve The x-axis and the y-axis should be labeled with the quantities and the units

Read temperature too early

Used larger amount of the limiting anomalous point: not on the smooth line / not in order (trend) with other points

(d)Pie charts •• These should be drawn with the sectors in rank order, largest first, beginning at ‘noon’ and proceeding clockwise. Pie charts should preferably contain no more than six sectors.

78% N2

21% O2

0.04% CO2

1.0% Ar+ other noble gases

anomalous point

length of Mg ribbon/ cm

chan

ge in

te

mpe

ratu

re /o C

anomalous point

time/ s

Volu

me

of

gas /

cm

3



(c) Numerical results Data should be recorded so as to reflect the precision of the measuring instrument. •• The number of significant figures given for calculated quantities should be appropriate to the least number of significant figures in the raw data used. (e) Bar charts •• These should be drawn when one of the variables is not numerical. They should be made up of narrow blocks of equal width that do not touch. (f) Histograms •• These are drawn when plotting frequency graphs with continuous data. The blocks should be drawn in order of increasing or decreasing magnitude and they should touch. Tables •• Each column of a table should be headed with the physical quantity and the appropriate unit, e.g. time / s.

•• The column headings of the table can then be directly transferred to the axes of a constructed graph. Anomalous result when a table of data is given

(a) The point not in order with other results = anomalous

Example: Time / s 1 2 3 4 5

Volume of gas/ cm3 20 30 44 50 60

Ans: point at 3 seconds clearly not in order with other results.

(b) Inconsistent is used when comparing between two results that must match. Inconsistent is used in titration to identify a titration run that doesn’t match with accurate (precise) ones

Example: Titration 1 2 3 4

Volume added / cm3 12.1 12.2 12.9 12.3

Ans: run 3 inconsistent (not concordant) with other results by more than 0.2

Methods of collecting gases - Diagrams 1. Inside gas syringe

Gas syringe for collecting gases and measuring their volumes especially gases that are soluble in water [ NH3, SO2, HCl . NO2 and Cl2]

2. Upward delivery collecting a gas which is less dense than air [ 2H , 3NH ]

3. Above water for collecting a gas which is insoluble or slightly soluble in water [ 2H , 2N , 2O 2CO , alkanes and alkenes]

4. Downward delivery collecting a gas which is denser than air [ 2CO , 2SO ,HCl ,Cl2]

Inverted measuring cylinder

water

gas

trough

gas jar

gas jar

gas



Table showing methods of collecting gases Property/ method H2 O2 CO2 NH3(g) Cl2(g) SO2 , NO2 & HCl(g) N2 Ethene

Less dense than air? slightly

Upward delivery

Denser than air? slightly

Downward delivery

Solubility in water?

very low low low Very

high moderate very high Insoluble

above water Note: all gases can be collected inside a gas syringe Labelled diagrams showing different methods of collecting and measuring the volume of gas - Method 1 for collecting a gas over water in a

measuring cylinder: The apparatus below was used to prepare hydrogen and measure the volume of gas produced (insoluble)

- Method 2 for collecting a gas in a gas syringe and measuring its volume.

Method 3 for collecting the gas by downward delivery ( below air)

Example: Sulfur dioxide is a poisonous gas which is denser

than air and soluble in water. Sulfur dioxide can be prepared by adding dilute hydrochloric acid to sodium sulfite and warming the mixture.

Sodium sulfite

dilute hydrochloric acid

heat

drying agent CaCl2

dry SO2(g)

aqueous hydrogen peroxide

Manganese (IV) oxide)

gas syringe

water metal

dilute hydrochloric acid

Measuring cylinder

Conical flask

trough



1. Laboratory preparation of H2 From reaction of dilute acids (HCl and H2SO4) with metals (Mg, Al, Zn, Fe)

(aq) (aq) 2(aq) 2(aq)Zn +2HCl H + ZnCl

First gas – jar should not be used because it contains displaced air of apparatus in addition to H2

Precaution: Do not remove heat before disconnecting the delivery tube to prevent back – suction

Lab preparation of oxygen ( on small scale)

1. By decomposition of hydrogen peroxide, H2O2 ,using manganese (IV) oxide as a catalyst

22O2H 2MnO O2H2 2O

2. By thermal decomposition of potassium chlorate, KClO3 , using MnO2(s) as a catalyst

32KClO 2MnO /heat 2KCl 23O

Big Notes on the processes - First gas jar should not be used because it does not contain pure O2, it contains O2 and displaced air

(O2 and N2) from the conical flask. - Some O2 remains in apparatus - O2

is collected above water being very slightly soluble in water.

- O2 can be collected inside a gas syringe Precaution: Do not remove heat before disconnecting the delivery tube to avoid suction – back

trough

delivery tube

clamp

Gas jar

water

O2 gas

KClO3,MnO2(s)

heat ceramic wool trough

delivery tube

clamp

Gas jar

water

O2 gas H2O2(l)

MnO2(s)

O

trough

delivery tube

clamp

Gas jar

water

H2 gas dilute HCl(aq)

Zinc powder

H2

Heat



Laboratory preparation of CO2: by reaction of CaCO3(s) and dilute HCl(aq)

3(s)CaCO + (aq)2HCl 2(g)CO + (l)2OH + 2(aq)CaCl

CO2(g) is denser than air It is collected below air by downward delivery

CO2(g) is slightly soluble in Hot water It is collected above water by water displacement

Laboratory preparation of collecting ammonia gas, NH3, Ammonia can be prepared by heating a mixture of ammonium salt such as NH4Cl with a metal oxide, such as CaO, or metal hydroxide such as NaOH. Strong bases such as NaOH, KOH, Ca(OH)2 and CaO can displace ammonia from ammonium salts. Equations of the reaction:

Sodium hydroxide + sodium hydroxide sodium sulphate + water + Ammonia (s)2NaOH + (s)424 SO(NH ) (s)42SONa + (g)O2H2 + (g)32NH

Ionic equation of the reaction: (s)-2OH + +4(s)2NH 2 (g)2H O + (g)32NH

By reaction of aqueous NaOH and NH4Cl By reaction of solid Ca(OH)2 and NH4Cl Diagram 1 Diagram1

Heat is applied under the mixture of the (NH4)2SO4 and Ca(OH)2 or CaO in diagram 2 [because they reactants are solids; they need heat enough to overcome activation energy] and under ammonium salt in diagrams 1 to speed up the reaction.

A ceramic wool is used to stop solids (reactants and products) from moving with the gases CaO or Na2SO4 can be used as drying agents to absorb the produced water. CaCl2 and H2SO4 cannot be used as drying agents because they react with ammonia. We cannot use CaCl2 and concentrated H2SO4 as drying agent because it reacts with ammonia Physical Test for ammonia gas: it turns damp red litmus paper blue Chemical test: HCl(g), they react producing white fumes

H2O(g) NH3(g)

Ceramic wool)(Filter

gas jar

dry NH3 gas

U-tube

CaO drying agent “absorbs H2O”

(NH4)2SO4 (s) and NaOH(s)

heat

heat

CaO “Drying agent

NaOH(aq) ammonia

Ammonium salt

2CO

1

dilute HCl(aq)

CO2(g

CO2(g)

CaCO3(s) downward delivery

2(g)CO

2

clamp

inverted 100 cm3 measuring cylinder

trough

hot water

CO2 gas



Ammonia cannot be collected above water because it is very soluble in water. Ammonia is collected by upward delivery (above air) because it is less dense than air. it is dried with CaO, because it is basic oxide

How to prepare ammonia solution - because NH3

gas is highly soluble in water,

- its solutions needs funnel arrangement to avoid suction back. (Rim of funnel just touches the level of water)

Reduction of copper (II) oxide with ammonia and preparation of N2 gas NH3 acts as a reducing agent. It reduces hot CuO to Cu metal. the reaction produces N2 gas CuO must be heated because cold CuO doesn’t react with NH3. The reaction produces nitrogen gas, water and copper. So it can be used to prepare N2 gas. Equations of the reaction:

ammonia + Copper(II)oxide copper + water + nitrogen (g)32NH + (s)3CuO (s)3Cu + 2 (g)3H O + (g)2N

Observations: - Black CuO turns red-brown solid because it is reduced into Cu - Condensed clear liquid found to be water - Bubbles of colourless gas N2, displace water in the gas-jar

Nitrogen is collected above water because it is insoluble in water.

Precaution: remove the delivery tube before you stop heating to avoid suck-back

“Suck-back happens when hot confined gases cool and cause water to enter the apparatus”. Hydrogen gas can also be used as a reducing agent; but no nitrogen is produced

NH3(g)

beaker

water

trough

delivery tube

clamp

Gas jar

water

N2

Black

condensed water

Copper (II) oxide

ammonia

cold water

Heat



Apparatus When a mixture of an ammonium salt and an alkali is heated, ammonia is formed. Dry ammonia gas was passed over heated copper (II) oxide using the apparatus below. Copper, nitrogen and ammonia were produced.

Ceramic wool stops solids moving with the gases [ NH3 nd H2O] gas K is nitrogen CaO drying agent to remove / absorb water moving with NH3 gas

Reduction of metal oxides such as CuO, PbO, ZnO, in school by H2 and carbon The labeled apparatus; where heat should be applied, reactions, Observations with reasons; precautions

In metals and alloys

Past papers question_0620_s16_qp_62/Q1 The diagram shows the apparatus used to reduce copper(II) oxide with hydrogen

(a) Complete the boxes to name the apparatus. [2] (b) Use an arrow to indicate where heat is applied. [1] (c) The colour of the copper(II) oxide changes from ………………………. to …………………….[2] (d) Suggest a reason why the U-tube is surrounded by ice. Ans: to cool the vapor and condense it (e) (i) Identify the colourless liquid formed.

(ii) Give a chemical test for this liquid. Test ………………………………………………………… Result……………………………………………………….

(iii) How could you show that this liquid is pure?

NH3

water

A B C

ammonium sulpahte/ calcium oxide

gas K

Ceramic wool

CaO Copper (II) oxide

heat heat



investigating the reaction of methane, CH4, and copper(II) oxide.

Sample answered question A student investigated the reaction of methane, CH4, and copper(II) oxide. She passed methane gas over hot copper(II) oxide using the apparatus shown.

The solid changed colour to red-brown and drops of liquid condensed in the cold part of the tube. (a) What was the original colour of the solid? Ans: black [ the colour of CuO] (b) Suggest the identity of

(i) the red-brown solid, Ans: copper

(ii) the drops of liquid. Ans: water [2] (c) Suggest a physical test to identify the liquid.

test Ans: measure its boiling point result Ans: 100oC sharp [2]

Equation of the reaction

Finding the % of O2 in air

1 The percentage of oxygen in air can be determined by the following experiment.

The gas syringe contains 50 cm3 of air. The large pile of copper is heated and the air is passed from one gas syringe to the other over the hot copper. The large pile of copper turns black. The gas is allowed to cool and its volume measured. The small pile of copper is heated and the remaining gas passed over the hot copper. The copper does not turn black. The final volume of gas left in the apparatus is less than 50 cm3

(i) Explain why the copper in the large pile turns black. Ans: Cu reacted with O2 and formed CuO which is black [1]

red-b4 2 2b rownlack4CuO + CH 4Cu + CO + 2H O

gas syringe containing 50 cm3 of air

large pile of copper

small pile of copper

gas syringe heat

heat

copper(II) oxide excess methane burning

methane



(ii) Why must the gas be allowed to cool before its volume is measured? Ans: hot gas has higher volume / heat causes expansion of gases/ ORA [1]

(iii) *Explain why the copper in the small pile did not turn black. Ans: all O2 has reacted with the large pile of copper

.......................................................................................................................................[1]

(iv) *What is the approximate volume of the gas left in the apparatus? Ans: O2= 20 % of air O2 = 20 % of 50 Remaining O2 = 0.8 x50=40 [1]

620_s13_qp_32_Q1(c) Investigating Effect of reacting hot metal such as Iron, magnesium, with steam Reaction of heated magnesium ribbon with steam

Word equation: magnesium + Steam magnesium oxide + hydrogen chemical equation Mg(s) + H2O(g) MgO + H2

- Before the reaction magnesium ribbon is cleaned. Suggest how and why

Ans: cleaned with sand paper to remove oxide layer on its surface. [ with acids give water in addition to H2]

- Boiling tube is used not a test tube. Why? Ans: boiling tube withstands heat, while normal test tube doesn’t - the PURPOSE of the ceramic wool is to hold water

- heat is applied under magnesium; radiated heat produces steam - before the reaction magnesium is grey

- after the reaction it turned into white powder; this powder is MgO

- Hydrogen is collected over water BECAUSE it is insoluble in water.

Note: transition metals like Fe and Zn, react with steam similar to magnesium; But magnesium reacts faster and more heat is produced being more reactive Observations with iron: Grey solid turns black / orange Observations with zinc: Grey solid turns white powder The heated boiling tube often cracks after the experiment due to one of the three reasons 1. back suction ( when the gas is passed through water) as in the diagram below

Wet ceramic wool absorbs water

magnesium ribbon

heat

Boiling tube

Inverted measuring cylinder

water

H2(g)

trough



2. overheated / or /high temperature is reached ( very important)

(a) (i) Complete the diagram to show how the hydrogen produced could be collected and its volume measured. Label your diagram.[2] Ans: add gas syring or above water inide measuring cylinder

(ii) State the effect of a lighted splint on the gas produced. Ans: burns with pop sound

(b) What liquid is absorbed on the mineral wool? Ans: Water (c) (i)Use two arrows to show two places where heat is applied.[1] Ans: shown on the diagram

heat is applied under magnesium and the ceramic wool soaked with water. Reason: under the ceramic wool to produce steam ; under magnesium because hot magnesium reacts faster with steam.

In this case heat is applied under water because it is far away from the source of heat under Mg.

3. In Dehydration of crystal (hydrated salt) The boiling tube cracks because water vapour condensed on cooler part at top of tube (1) and runs back to very hot tube

(a) If hydrated CoCl2.2H2O is heated

It turns from pink (CoCl2.2H2O )to blue (CoCl2)

If water is added to the remaining solid; it turns back to pink. Reason: the reaction is reversible

(b) If hydrated CuSO4.5H2O is heated

It turns from blue (CuSO4.5H2O)to white (CuSO4)

If water is added to the remaining solid; it turns back to blue. Reason: the reaction is reversible

Questions to answer (i) draw an arrow to show where heat is applied. Ans: Under the hydrated sal

(ii) What is the purpose of the water bath or ice?

Ans: to condense the water vapour / To change water vapour into liquid (iii) Calculate the maximum mass of water that could be collected when a sample of

hydrated copper(II) sulfate of mass 2.50 g is heated. [Mr of CuSO4.5H2O is 250]

heat

water

Hydrated salt

Boiling tube

water

beaker

Wet ceramic wool absorbs water

magnesium ribbon

heat

Boiling tube

heat



During finding the number of water of crystallisation

Sources of error improvement the solid absorbs water from the air store in a sealed container / airtight container / desiccator

Thermal decomposition nitrates and observations 1. Decomposition of all nitrates and including LiNO3 except all other group I nitrates

Example: When magnesium nitrate is heated strongly, it decomposes,

Mg(NO3)2 MgO + NO2 + O2

- Products: MgO (white solid) NO2 (brown acidic gas) and O2 ( colourless gas) - Observations: brown gas appears,

salt stays white but its size (mass) decreases. Solid melt with excessive heating

Notes:

- if we bubble the produced gases through a universal indicator it turns red as nitric acid (strong acid) is produced

- pH of the solution about 1 - we cannot collect NO2 gas above water because it is soluble in water. ( above water is a mistake) - NO2 should be collected below air by downward delivery because it is denser than air

( above air is a mistake)

- If we pass the produced gases over heated metal such as copper, it turn black because it reacts with O2 producing CuO which is black.

- We heat under copper because cold copper doesn’t react with oxygen. - The glass wool stops/ traps salts moving with the gases

2. if KNO3, NaNO3, or RbNO3 is heated strongly, it decomposes partially. - Products: KNO2 (white solid) and O2 ( colourless gas) - Observations: (mass) of white salt decreases

Colourless gas

heat

copper

heat

magnesium nitrate

glass wool

Water and Universal indicator



Thermal decomposition of carbonates and observations [ from Group I only Li2CO3 decomposes]

Example: When calcium carbonate is heated strongly, it decomposes, - Products: CaO (white solid) CO2 (colourless acidic gas) - Observations: if we bubble the produced gas through lime water, it turns milky because

insoluble CaCO3 forms.

2(aq) 2(g) 3(s) 2 ( )limewater ppt, (insoluble)

Ca(OH) +CO CaCO H O l

If we continue bubbling the gases, the solution turns clear again because the excess CO2 reacts with CaCO3 and water producing a soluble compound Ca(HCO3)2.

3(s) 2 ( ) 2( ) 3ppt

2(aq)CaCO Ca(HCO )gH O CO l

Laboratory preparation of hydrogen chloride A sample of hydrogen chloride gas can be prepared by adding concentrated sulphuric acid to sodium chloride or hydrochloric acid and solid KMnO4. Because hydrogen chloride gas is strong-smelling, denser than air and soluble in water, it can be collected by downward deliver below air. Never pass it through water as it dissolves in water.

Cracking - in organic topic - The labeled apparatus - Where heat should be applied - reactions - Observations with reasons - Precautions

Extraction of metals above H2 by electrolysis of its molten salt – in electrolysis

- The labeled apparatus - Where heat should be applied - Half – reactions - Overall redox reaction - Observations with reasons - Precautions

heat

limewater CaCO3



Preparing hydrogen gas and mistakes in the apparatus

1 The diagram below shows some apparatus was used to prepare a gas in the laboratory. (a) Name the four pieces of apparatus labeled.

A. _________________________ [1]

B. _________________________ [1]

C. _________________________ [1]

D. __________________________ [1]

(b) What two properties of the gas can you deduce from the diagram. 1. _____________________________________________________________________________________ 2. _____________________________________________________________________________________ [2]

(c) Name a gas that could be prepared using this apparatus. ___________________________________ [1] Mistakes in the apparatus/ collecting gases /O/N/96/Q2

Answer (a) A=tap funnel; B=conical flask; C=gas Jar; D; delivery tube; (b) Less dense than air being collected by upward delivery

Acidic or neutral because it is dried by sulfuric acid Insoluble in water because it passed through water

(c) Hydrogen,

Mistakes in the apparatus and collecting gases 1 Ammonia is produced when aqueous sodium hydroxide is warmed with an ammonium salt.

Ammonia is less dense than air and very soluble in water. The following apparatus was used to prepare a sample of dry ammonia gas.

(a) Name substance C. [1]

(b) Name substance D. [1]

(c) What laboratory equipment, necessary for the preparation of ammonia is missing in the diagram? _______________________________________[2]

(d) Suggest why concentarted sulphuric acid should not be used to dry ammonia. _____________________________________________________________________ [1]

(e) There are two mistakes in the apparatus shown in the diagram. Identify these mistakes.

Mistake 1 _______________________________________________________________ [1]

Mistake 2 _______________________________________________________________ [1] O/N/1993/Q4

concentrated sulphuric acid

A

B

C

D

Concentrated sulphuric acid

C

D

ammonia



Answer (a) C = Aqueous sodium hydroxide (b) D = Ammonium salt such as ammonium sulfate (c) Bunsen burner for heating (d) Because ammonia is alkali will react with sulphuric acid (e) Mistake 1: The delivery tube in and out of the sulphuric acid flask.

The inlet is short. The outlet is under the liquid surface. Mistake 2: Ammonia is collected by downward delivery despite being less dense than air

2 Nitrogen dioxide is a strong-smelling, denser than air and soluble in water. A sample of

nitrogen dioxide can be prepared by adding concentrated nitric acid to copper and warming the mixture. Study the diagram and answer the questions. (a) (i) Fill in the boxes in the diagram above to show the chemicals used. [2]

(ii) Indicate were heat is applied using an arrow [1]

(b) Identify and explain two mistakes in the diagram.

Mistake 1 ______________________________________________________________________________ [1] Mistake 2 ______________________________________________________________________________ [1]

(c) Sate and explain two precautions that should be taken when carrying out this experiment.

precaution 1 ____________________________________________________________________________ [2] precaution 2 ___________________________________________________________________________ [2]

Labelling and Mistake in the apparatus/ MAY 1995/Q4 Answer

(a) Labels on the diagram (b) Mistake 1: The gas is soluble in water and it is bubbled through water

Mistake 2: The gas is denser than air and they collected it by upward delivery. (c) Precaution 1: Add concentrated HNO3 dropwise to avoid violent reaction

Precaution 1: Heat gently to maintain controllable rate of formation of nitrogen dioxide

water

Ans: copper

Ans: conc.HNO3



Thermal decomposition of metals

7- Nitrogen dioxide is an acidic oxide that is soluble in water. Nitrogen dioxide and oxygen are produced when magnesium nitrate crystals are heated.

2Mg(NO3)2 2MgO + 4NO2 + O2 2Cu + O22CuO Brown black

May/June/1993/Q4

(a) What would be observed in parts A, B and C of the apparatus? Part A ______________________ B ______________________C____________________ [3]

(b) Explain your observations in parts A, B and C of the apparatus? (i) Part A ________________________________________________________________

(ii) Part B ________________________________________________________________ (iii) Part C _____________________________________________________________ [3]

Answer (a) A: Brown gas; B: Solution changes from green to orange; C: Brown solid turns black (b) Part A: NO2 which is brown is produced

Part B: NO2 is soluble acidic gas; it make the solution acidic Part C: Cu combines with O2 forming CuO which is black

2 A sample of natural gas was passed through aqueous bromine. The colour of the solution of bromine changed from orange to colourless. Comment on this observation. [2]

Answer The gas contains C=C/ is alkene [1] Because Bromine reacts with C=C [1]

heat heat

A B C magnesium nitrate copper

Water + universal indicator



Electrolysis of molten lead (II) bromide using graphite electrodes Because solid PbBr2 salt doesn’t conduct electricity as the ions cannot move. We heat it to melt for ions to become free to move and conduct electricity

Ions in the solution: 2+ -Pb , Br When the circuit is switched on: - 2+Pb ions move to the negative electrode (cathode): - -Br ions move to the positive electrode (anode):

Reactions at electrodes:

- At cathode: 2+ -Pb + 2e Pb reduction: gain of electrons

- At anode: 2- -2Br 2e + Br oxidation: loss of electrons

- Overall redox reaction: ( ) + 2 2PbBr Pb Brl All Observations when the circuit is switched on. - Brown fumes at the anode [ due to formation of Br2] - Grey deposit at the cathode [ due to formation of Pb]

Precaution: Must be carried out in a fume cupboard because the produced bromine vapour is dangerous to health. Same observations and precautions for all metal bromides; only the metal obtained at cathode will be different With metal iodide purple fumes at the positive electrode

With metal chloride green gas at the positive electrode

Electrolysis of water using dilute sulphuric acid The role of H2SO4 is catalyst to speed up electrolysis by its free ions Ions in the solution: + - 2-

4(aq) (aq) (aq)&H , OH SO When the circuit is switched on: - + (aq)H ions move to negative electrode (cathode): - - 2-

4(aq) (aq)&OH SO ions move to positive electrode (anode): Reactions at electrodes:

- At cathode: +2

-2H + 2e H reduction

- At anode: 2 2- -4OH 4e + 2H O + O oxidation

- Overall redox reaction: ( ) (g) + (g)2 2 22H O 2H Ol Volume ratio 2V 1V

All Observations when the circuit is switched on. - bubbles of colourless gas at each electrode

- twice the volume of gas is formed at the cathode than at the anode because mole (volume)

ratio of H2:O2 is 2:1 - bulb lights up

metalgrey Silvery

bromide (II) LeadMolten

fumesBrown

heat

Graphite anode (+)

Graphite cathode (-)

dilute H2SO4(aq)

2(g)H2(g)O



Note: - H+ and OH- ions used in electrolysis come from water not from H2SO4 - Solution becomes more acidic; pH decreases as the concentration of H+ ions increases

because volume of water decreases by electrolysis. Electrolysis of concentrated hydrochloric acid solution Ions in the solution: + - -(aq) (aq) (aq)&H , OH Cl When the circuit is switched on: - + (aq)H ions move to negative electrode (cathode): - - -(aq) (aq)&OH Cl ions move to positive electrode (anode):

Reactions at electrodes:

- At cathode: +2

-2H + 2e H reduction: gain of electrons

- At anode: 2- -2Cl 2e + Cl oxidation: loss of electrons

- Overall redox reaction: + 2 2

-2H + 2Cl H + Cl

All Observations when the circuit is switched on. - bubbles of gas at each electrode - colorless gas at the cathode - green gas at the anode - the universal indicator around the anode bleaches by the formed chlorine gas - level of liquid falls in both tubes as the gases collected over water pushes water down - bulb lights up

Important Note: Level of water must be equal in both tubes because the gases are 1:1 mole ( volume) ratio, However, water level in the positive electrode will be higher because some Cl2 dissolves in water being slightly soluble in water Electrolysis of copper (II) sulphate using (inert) graphite electrodes Ions in the solution: + - 2+ 2-

(aq), 4(aq) (aq) (aq)&H , OH Cu SO

When the circuit is switched on: - + 2+H , Cu ions move to negative electrode (cathode): - - 2-

4OH , SO ions move to positive electrode (anode): -

Reactions at electrodes: - At cathode: 2+ -Cu + 2e Cu reduction: gain of electrons

- At anode: 2 2- -4OH 4e + 2H O + O oxidation: loss of electrons

All Observations when the circuit is switched on.

Graphite anode (+)


CuSO4(aq)

Graphite anode (+)


Conc. HCl(aq) + UI

2(g)H2(g)Cl



- Blue colour of the solution fades as the concentration of Cu2+ is decreasing, until it becomes colourless at the end

- colorless gas at the anode - brown deposit forms at cathode due to formation of Cu atoms from reduction of Cu2+ ions - remaining solution is sulphuric acid ( + 2-

42H , SO ) - solution changes from CuSO4(aq) to H2SO4(aq)

Electrolysis of copper (II) sulphate using copper electrodes Ions in the solution: + - 2+ 2-

(aq), 4(aq) (aq) (aq)&H , OH Cu SO When the circuit is switched on: - + 2+H , Cu ions move to negative electrode (cathode): - - 2-

4OH , SO ions move to positive electrode (anode): Reactions at electrodes:

- At cathode: 2+ -Cu + 2e Cu reduction: gain of electrons

- At anode: 2+ -Cu Cu + 2e oxidation: loss of electrons All Observations when the circuit is switched on.

- Solution stays blue because Cu2+ ions removed from the solution at cathode are replaced by Cu2+ formed from oxidation of anode;[i.e, concentration of Cu2+ remains the same for same reason]

- brown deposit forms at cathode; so cathode becomes thicker ( its mass increase) - Anode becomes thinner because Cu atoms are oxidised to Cu2+ that dissolved in the solution.

- The solution stays CuSO4(aq) During reactions of Aluminum with acids, the reaction is slow at start because it has protective Al2O3 layer on its surface.

After some time the reaction becomes very fast because 1. the oxide layer is removed and Al becomes exposed to the acid. Also the 2. reaction is exothermic, the produced heat speeds up the reaction.

Sample past papers question Concentrated hydrochloric acid can be electrolysed using the apparatus shown. [Total: 6] (a) Label the position of the electrodes on the diagram. [1] (b) Give two observations when the circuit is switched on.

1 …………………………………………………………………….

2 …………………………………………………………………… [2]

(c) (i) Name the product at the positive electrode. ………………………. [1] (ii) State a test for this product and the result of the test. Test ……………………………………………………………………. result ……………………………………………………………………. [2]

copper anode (+)

copper cathode (-)

CuSO4(aq)



Question Answer

(b) bullb lights up bubbles green gas at the anode level of liquid inside the tubes falls

(c) chlorine (gas), Cl2, (d) damp blue litmus paper, bleaches

How to extract the coloured pigments from solid such as leaves – in chromatography topic Details of chromatography – in chromatography topic Solved Model past papers questions 1 A student extracted the colours present in some leaves using the apparatus below.

(a) Complete the boxes to identify the pieces of apparatus used. [2] (b) Use labelled arrows to indicate

(i) the solvent, (ii) the solution of colours. [2]

(c) Chromatography was used to separate the colours. The chromatogram obtained is shown.

(i) On the diagram, label the solvent front. (ii) How many colours were present?

Ans: three colours [three spots]

(c)(i)solvent front

(a) electrodes

Funnel

Mortar

Pestle

(b)(i)solvent

(b)(ii)Solution of colours



Acid – Base Titration Definition: a lab techniques that measures the exact volume of acid needed to neutralize a measured volume of an alkali and vice versa. Uses of titration

1. to find the unknown concentration of an acid, an alkali, oxidising or reducing agent…] [ main use] 2. to prepare soluble alkali metals salts 3. to compare between the concentrations of strong acids 4. to compare between the concentrations of strong alkalis

Apparatus used in titration: - Pipette: transfers exact volume of the analyte to the conical flask - Burette: measures the exact volume needed to neutralise the analyte in the conical - Conical flask: where the reaction takes place. Its volume must be larger than the total volume of

the titrant and the analyte.

before titration - first rinse the burette with distilled water to remove trace of chemicals previously used - then rinse with the titrant to remove traces of water

* If traces of water remain in the burette, the added volume will be larger than the actual and so the calculated number of moles / concertation of the analyte will be larger than its actual amount.

End point: when all number of moles of the acid or base in the conical flask is used up /neutralised. Therefore, at the end point pH=7 ( in all strong acid – strong base titration)

Methods for determining the end point 1. Using a suitable acid - base indicator [phenolphthalein, methyl orange or litmus paper] End point =

when the indicator changes colour permanently 2. Using pH – meter [ we plot a graph – the end point at the vertical line of the sudden jump 3. Measuring temperature change during titration [plot a graph – end point where the two lines meet]

Titration using indicator In acid - base titration we need an indicator because almost all acids and alkalis are colorless.

Indicator: A substance that gives different colors in acids and bases. When the indicator changes color permanently, this is the end-point

Indicator Colour in acid Colour in alkali In neutral Litmus paper / solution Red Blue purple phenolphthalein colourless pink colourless Methyl orange Red Yellow *Its color at end point is orange

in strong acid strong alkali titration, a U.I can be used where the color at the end point is green.



How to find the unknown concentration of NaOH(aq) by titration[ Practical details outline]

Example: titrating NaOH(aq) with H2SO4(aq) • Fill the burette with 0.5 mol/dm3 H2SO4(aq) • Using pipette, transfer 25cm3 of NaOH(aq) to the conical flask, then add 2 to 3 drops of

phenolphthalein indicator. • Start adding H2SO4 (aq) from the burette, 2 cm3 at a time, and swirl until the indicator just changes

colour ( from colourless to pink) • Record the volume • Repeat the titration many times until concordant results – result that have max. 0.2 cm3 difference • take average of the concordant results and carry out calculation • Carry out calculation if needed – use the calculated moles and the stoichiometry of the balanced equation

Interpretation diagram for titration steps

- The end point is recognized by colour change of the indicator - at the end point all the alkali in the conical flask is neutralized by the added acid - pH at end point is 7 - Equation of the reaction: H2SO4 + 2NaOH Na2SO4 + 2H2O

Notes Put a white tile under the flask. REASON: To see colour change more clearly/ easily it is not necessary to dry the conical flask before repeating the titration: REASON: Water doesn’t

affect the number of moles of the acid/alkali in the conical flask.

Adding more acid / alkali beyond the end point is called overshot the endpoint when the conical flask is filled with strong acid, overshot the endpoint makes the solution strong basic pH 13 when the conical flask is filled with strong alkali, overshot the endpoint makes the solution strong acidic pH 1

the calculated number of moles/ concentration of the acid will be more than the actual

First titration is rough (anomalous); it only gives information about the volume needed for titration. It is not included in calculating the average volume because it is not consistent [concordant] with other data We do two or three precise titrations to take average and increase the reliability of the results *when we repeat titration, smaller volume will be used as the repeated experiment will be more accurate Precautions (risk assessment): when working with acids and strong alkalis such as HCl(aq) and NaOH, wear gloves and safety goggles to protect your eyes and hands because strong acids and strong alkalis are corrosive. Note: In titration we use dilute solutions because

End point

4

turns colourless

3cm

50

40

30

20

10

0

3

dil.H2SO4(aq) 0.1mol/dm3

3cm

50

40

30

20

10

0

burette

25cm3 NaOH(aq)

phenolphthalein indicator [ pink]

2

dropper

1

25cm3 NaOH(aq)

conical flask

pipette



1. concentrated acids and alkalis are corrosive, to make them safer. 2. with more dilute titrant larger volume will be added which decreases the error/ increases the accuracy

We titrate with standard solution to get accurate results. Standard solution: the solution whose concentration is accurately measured. Anomalous result is - the one not consistent [concordant] with other values (results) [ when volumes are given] - the point not on the smooth line [ when graph is given] Worked Example A student titrated 25cm3 of KOH(aq) using dilute nitric acid. A student did the titration four times and recorded the following results.

titration number 1 2 3 4

volume of dilute nitric acid / cm3 18.1 18.9 18.3 18.2

(a) Which one of the results is anomalous? Ans: number 2 / 18.9 cm3 (b) Suggest what might have caused this result to be anomalous.

Ans: overshot end point / or/ more than 25 cm3 KOH in flask (c) Use the other results to calculate the average volume of dilute nitric acid that reacted with

the aqueous potassium hydroxide. Ans: (18.1+18.2+18.3)/3=18.2 cm3

Sources of error and improvements Sources of error Improvement

“To Increase precision and reliability of the results” 1. Doing titration only once Do titration many times and take average;

It increases reliability of the results and allows us spot the anomalous point

2. measuring or recording error repeating the experiment; increases reliability and spots the anomalous point 3. overshot end-point

4. using a measuring cylinder to transfer a solution to the conical flask.

Use pipette as it measures accurate volume [ increases the accuracy]

5. using relatively concentrated solutions. Use more dilute titrant; larger volume would be added decreases the % error / increases the accuracy

In comparison using the terms Equal , double, twice, half, more or less

Other uses of titration: 2 To compare between the concentrations of different samples of same acid

Titrate identical samples of each acid (same volume and same concentration) using same alkali. When the indicator changes colour permanently, stop and Record the volume

The acid that needs largest volume of alkali to neutralize it is most concentrated. Example When titrating 20cm3 of Q and R of same acid, HCl, with same alkali, 0.05 mol/ dm3 NaOH, the added volume on Q is double the added volume on R, then the concentration of Q is twice the concentration of R Initial reading of the burette is 0.0cm3 in all experiments

20cm3 acid Q

0.5 mol/dm3 NaOH

20cm3 acid R

0.5 mol/dm3 NaOH 10

20



3 To compare between the strength of different acids Titrate identical samples of each acid (same volume and same concentration) using same alkali. When the indicator changes colour stop and Record the volume

The acid that needs largest volume of alkali to neutralize it is the strongest. Example Three solutions of different acids: Q, R and W are titrated with the same alkali, 0.05 mol/ dm3 NaOH, the results are shown below

(a) What is the volume required to neutralize each of the acids? Q ……….…………., R …………………….., W …………………………

(b) What is your conclusion about the relative strength of the three acids …………………………………………………………………………………………………………

4 To compare between the concentrations of monoprotic acid strong acid (HA) and monoprotic strong alkali (MOH). Titrate 20cm3 of strong acid such as HCl(aq) with a strong alkali such as NaOH(aq) if larger volume of alkali is added ( such as 22cm3) then the alkali is less concentrated if smaller volume of alkali is added( such as 19cm3), then the alkali is more concentrated

Practice Example: (a) without any calculation, compare between the concentrations of

HCl(aq) and NaOH(aq) ……………………………………………………………

(b) Write the balanced equation of the reaction. NaOH + HCl NaCl + H2O …………………………………………………………

(c) Calculate the concentration of NaOH, Moles of HCl = cxv= 0.5 x (20/1000)= 0.01mole Moles of NaOH = 0.01 moles Concetration of NaOH = n/v=0.01/ (10/1000) = 1 mol//dm3

(d) Does your answer in (c) confirm your answer in (a) …………………………………………………………

5 Using pH – meter we can draw the graph which shows how the pH changes during titration.

Example: The graph shows the change in the pH when aqueous potassium hydroxide is added to 25.0 cm3 of dilute nitric acid. A pH meter was used to track the change in pH with addition of alkali. Point X is the end point ( pH=7, volume=25 cm3). Measure change in pH during titration. Record your results and plot the graph

20 cm3 0.5mol/dm3

HCl(aq) in conical flask

?? mol/dm3 NaOH

3cm

50

40

30

20

10

0

10 cm3

Q

20

10 cm3

W

30

10 cm3

R

10



Titrating strong acid with strong alkali The burette is filled with alkali KOH(aq)

The conical flask contains acid [HCl(aq)]

Titrating strong acid with strong alkali The burette is filled with acid HCl(aq)

The conical flask contains alkali [NaOH(aq)]

pH increases slightly then Sudden increase in the pH at the end point Volume at end point = 25cm3. [ look at the arrow]

pH decreases slightly then Sudden decrease in the pH at the end point Volume at end point = 25cm3. [ look at the arrow]

Neutralization – temperature graph Using temperature change graph to determine the exact volume of acid needed for to neutralize a measured volume of alkali and vice versa. Extrapolate/ extend the lines until they meet, this is the end point (neutralisation point)

1. *In this type of graph only temperature change is important, so the scale can be taken from any temperature.

2. Why does the temperature increase? Ans: Reaction is exothermic

3. Why does the temperature decrease? Ans All the acid is used up and the added alkali is colder.

4. Reason of anomalies of points A,B and C A: larger volume is added which caused

higher temperature rise. B: added smaller volume of acid or base /

read temperature before highest temp rise is reached [took shorter time to read highest temperature rise]

C: took longer time to read the temperature] 5. The solutions are left to stand for about 30

minutes before neutralization starts. Reason: to reach room temperature./ to ensure all solutions are at same temperature as surrounding

6. Polystyrene beaker is used instead of a glass beaker because it is good insulator to minimise heat loss. 7. Keep stirring the mixture. Reason: to ensure uniform distribution of heat. 8. To increase the accuracy add smaller volumes to take more points ; more accurate plot ( graph) 9. Use data logger – records temperatures more accurately- more data – more accurate plot

10 20 30 40 50 0 0 Volume of added acid/ cm3

pH X

2

4

6

8

10

12

14

10 20 30 40 50 0 0 Volume of added alkali/ cm3

pH X

2

4

6

8

10

12

14

Tem

pera

ture

/ o C

neutralizationV

C

B

A

20

30

50

40

60

70

10 20 30 40 50 60010

endpoint



Redox titration Procedure: same as in acid base titration; except no indicator. In general, no need for an indicator because “there is already colour change at endpoint” as each substance has its own colour, (coloured), after all the substance in the conical flask ( analyte) has been used up / reacted, only one excess drop changes the colour indicating the end point. This is why redox reactions are called self-indicating Example: if we are titrating colourless Fe2+ solution by adding purple KMnO4 when all Fe2+ are used up only one (excess)drop of KMnO4 makes the solution pink https://www.youtube.com/watch?v=UCReOPY-YDU Important exception when iodine is being titrated (iodometric titration) When iodine is in the conical flask and being titrated, starch solution is added as indicator because just before the end point the colour is very light yellow ; so it is difficult to recognize/identify the change from faint yellow to colourless at the end point. Big note: add starch solution just before the end point, it makes iodine solution blue black; at the end point (when all iodine is used up) colour changes from dark blue to colourless.

https://www.youtube.com/watch?v=ahBZzj-ETCk

note: titration reaction should be rapid, if it is slow then either heat , change the titrant or add a catalyst

Acid – Base Titration [comparison between the different concentrations of same alkali]

1 A student investigated the reaction between hydrochloric aicd and aqueous solution of potassium hydroxide of different concentrations, labeled P, Q and R.

Experiment 1 A burette was filled to zero mark with hydrochloric acid. A 20cm3 of aqueous potassium hydroxide P was added to conical flask, together with 4 drops of phenolphthelein indicator. The acid was added gradually to the flask. When the colour of phenolphthelein changed, the burette reasding was noted.

Experiment 2: Experiment 1 was repeated using solution Q instead of solution P. Experiment 3: Experiment 2 was repeated using solution R instead of solution Q.

colourless Fe2+(aq)

burette 10

0.

20

40

30

50

cm3

purple KMnO4(aq)

begining

one excess drop of KMnO4(aq) makes the solution pink

burette

cm3

purple KMnO4(aq)

10

0.

20

40

30

50

End point



(a) Use the diagrams to complete the table of the volumes of hydrochloric acid used. Experiment 1 2 3 solution P Q R

Diagram of burette

Volume of hydrochloric acid /cm3 Ans: 10.9 Ans: 21.6 Ans: 6.1

[3] (b) What colour change would be observed when hydrochloric acid was added to the flask?

From _____________________________ to _______________________________ [2] (c) Complete the gaps in the following statements.

(i) Aqueous potassium hydroxide labeled _____________________ needed the smallest volume of hydrochloric acid to change the colour of the indicator. [1]

(ii) Aqueous potassium hydroxide labeled _____________________ needed the largest

volume of hydrochloric acid to change the colour of the indicator. [1] (iii) Therefore, the order of concentration of the solutions of potassium hydroxide is

most concentrated least concentrated

Q P R [3]

(iv) What would be used to measure the 20cm3 portions of aqueous potassium hydroxide? ______________________________________________________________________ [1]

(v) **If experiment 1 were repeated, would the volume of hydrochloric acid used be the same as in the original experiment? Explain your answer.

_________________________________________________________________________ [3]

(vi) Suggest another method of investigating the order of concentration of these solutions of potassium hydroxide. Explain how your method works.

_________________________________________________________________________

[3] O/N/1993/Q5

Model past paper question about redox titration [0620_M_J_2012_61_Q2]

6

7

8

21

22

23

10

11

12

pink colourless

R

Q

Pipette [ because it is transferred to the conical flask]

No, it will be less because the experiment will be carried out more accurately

Use a very sensitive type of universal indicator (or a pH meter). The solution which gives a colour that corresponds to the highest pH will be more concentrated



2 A student investigated the reaction between a solution of deep purple aqueous potassium manganate(VII), and two different colourless solutions, B and C, of an acidic solution of a sodium salt. Two experiments were carried out. Experiment 1 A burette was filled with the solution of potassium manganate(VII) to the 0.0 cm3 mark. Using a measuring cylinder, 25 cm3 of solution B was poured into the conical flask. The potassium manganate(VII) solution was added slowly to the flask and shaken to mix thoroughly. Addition of the solution was continued until there was a permanent pink colour in the contents of the flask

(a) Use the burette diagram to record the volume in the table of results and complete the table. [2]

final reading / cm3 26

initial reading / cm3 0.0

difference / cm3 26

Experiment 2 Experiment 1 was repeated using solution C instead of solution B. (b) Use the burette diagrams to record the volumes in the table and complete the table.

final reading / cm3 32

initial reading / cm3 19

difference / cm3 13

[2] (c) (i) What colour change was observed in the contents of the flask when potassium

manganate(VII) solution was added to the flask in Experiment 1?

from Ans: colourless to Ans: pink / purple [1]

(ii) Why was an indicator not added to the flask? Ans: potassium manganate solution is already pink; one excess drop colours the solution purple indicating end point [ best answer][1]

(d) (i) In which experiment was the greater volume of potassium manganate(VII) solution used?

Ans: Experiment 1 [1]

(ii) Compare the volumes of potassium manganate(VII) solution used in Experiments 1 and 2. Ans: volume of solution B is twice the volume of solution C [note: use double or twice in your comparison as it is clear in the results] [1]

(iii) Suggest an explanation for the difference in volumes in (d)(ii). Ans: concentration of B (1) is double that of C (1) [2]

(e) *If Experiment 2 was repeated using 12.5 cm3 of solution C, what volume of potassium

initial reading

18

19

20 Final reading

31

32

33

Final reading

25

26

27



manganate(VII) solution would be used? Explain your answer. Ans: 6.5 (1) cm3 (1) , half the volume used for C in experiment (2) (1) [3]

(f) A redox reaction occurs when potassium manganate(VII) reacts with solutions B and C. Explain the term redox reaction.

Ans: A reraction which involves loss and gain of electrons or oxidation and reduction [2]

(g) Give one advantage and one disadvantage of using a measuring cylinder for solution C.

Advantage: Ans: quick disadvantage Ans: not accurate [2]

(b) Why was the flask rinsed with distilled water at the end of each experiment? ***Ans: to make sure no traces of the previously used reactants remain in it. [2]

(c) Suggest, with a reason, the effect on the results of using a less concentrated solution of manganate (VII). ***Ans: larger volume of manganate (VII) will be added to same number of moles;

therefore, smaller error. [2] Redox titration past papers questions 3 A student investigated the reaction between sodium thiosulphate and potassium iodate.

Two experiments were carried out. Experiment 1 A burette was filled up to the 0.0 cm3 mark with sodium thiosulphate solution. By using a measuring cylinder, 20 cm3 of solution. A of potassium iodate was placed into a conical flask. Dilute sulphuric acid and potassium iodide were also added to the flask. The flask was shaken to mix the contents and produce a red solution of iodine.

The sodium thiosulphate solution was added to the flask. When the contents of the flask were yellow, 1 cm3 of starch solution was added to the flask. Addition of sodium thiosulphate to the flask was continued until the solution turned colourless. Use the burette diagram to record the final volume in the table and complete the column in the table of results on the table below.

Experiment 2 Experiment 1 was repeated using a different solution of potassium iodate, solution B. Use the burette diagrams to record the volumes and complete the table below.

Table of results

38

39

40

Final

28

29

30

initial

8

9

10



Experiment 1 Experiment 2 final reading / cm3

initial reading / cm3

difference / cm3

(a) Suggest why the starch was used. …………………………………………………………………… [1] (b) (i) In which experiment was the greatest volume of sodium thiosulphate solution used?

………………………………………………………………………………………………… [1]

(ii) Compare the volumes of sodium thiosulphate solution used in Experiments 1 and 2. ……………………………………………………………………………………………… [1]

(iii) Suggest an explanation for the difference in the volumes. ………………………………………………………………………………………………………

………………………………………………………………………………………………… [2]

(c) Predict the volume of sodium thiosulphate solution which would be needed to react completely with 10 cm3 of solution B.

………………………………………………………………………………………………… [2]

(d) Explain one change that could be made to the experimental method to obtain more accurate results,

- change …………………………………………………………………………………………….

- explanation …………………………………………………………………………………….. [2] Answer

Experiment 1 Experiment 2 Final reading 39.2 28.7 initial reading 0.0 8.1 Difference 39.2 20.6

(a) as an indicator.

(b) (i) Experiment 1 (1) (ii) Has greater volume than experiment 2 (1) (iii) solution A is more concentrated than B (1)

(c) 320.6 =10.3 cm2

correct

value (1) unit stated (1)

(d) Changes: Use more dilute solution (1) Explanation: larger volume would be added which decreases the error [ increases the accuracy] (1)



Neutralization graph 4 A student investigated the temperature changes when dilute nitric acid neutralised aqueous

potassium hydroxide. The instructions followed are listed below. Step 1 The solutions were left at room temperature for one hour. Step 2 Using a measuring cylinder, 20 cm3 of aqueous potassium hydroxide solution was

poured into a polystyrene cup and its temperature measured. Step 3 From a burette, 5.0 cm3 of nitric acid was added to the cup. The highest temperature

reached by the mixture was measured. A further 5.0 cm3 of nitric acid was added to the mixture and the highest temperature measured. Further 5.0 cm3 additions were made until a total of 30.0 cm3 of nitric acid had been added.

(a) Use the thermometer diagrams to complete the temperatures in the table.

volume of nitric acid added / cm3 0.0 5.0 10.0 15.0 20.0 25.0 30.0

thermometer diagram

highest temperature reached / oC

23 29 35 41 39 35 31

(b) Plot the results on the grid. Draw two intersecting straight lines through the points.

20

highest temperature reached / oC

30

40

50

5 10 15 20 250 30

Total volume of nitric acid/ cm3



(b) From your graph, work out the volume of nitric acid needed to completely neutralise the 20 cm3 of aqueous potassium hydroxide. Using an arrow, show clearly on the grid this neutralisation point. Ans: 16 cm3 [2]

(c) What was the room temperature?

Ans: 23 (°C) [ the temperature before adding any acid, the reading of the first thermometer] [1] [1]

(d) Why was a polystyrene cup used instead of a glass beaker?

Ans: Polystyrene cup is good insulator; it minimises heat loss [1]

(e) Why does the temperature: Increase? Ans: the reaction is exothermic /produced heat energy. then decrease? Ans: all the amount of KOH is used up and the added acid is cooler/ neutralised/ the reaction finished [1]

(g) What type of chemical reaction is this neutralisation? Ans: exothermic [1]

Preparing salts in acid, bases and salts chapter Other notes about preparation of salts 1. Suggest the effect of heating the solution of potassium nitrate to boiling point and then heating

for a further ten minutes. Ans: Evaporation Solid crystals form or you can say crystal decomposes

2. Describe the effect of boiling the solution of iron(II) sulfate for several minutes. Ans: evaporation of water/steam (1) solid/residue/crystals formed (1)

colour changes turns brown/darker green (1) effect of heat on solid breaks down (decomposes) (1) max 3 (1) Reason for los in mass/ yield during preparatiion of a salt

Ans: spillage (1) inaccurate weighing (1) loss by spitting (1) reaction not complete (1) some solid left in the beaker (1)



1- A student investigated the reaction of magnesium with dilute hydrochloric acid. Experiment 1 A 10 cm3 sample of dilute hydrochloric acid was placed in a boiling tube. The initial temperature of the acid was measured. A 2.5 cm length of magnesium ribbon was added to the acid in the boiling tube. The maximum temperature reached was measured. The gas given off was tested and gave a pop with a lighted splint.

(a) Name the gas given off ________________________________________________________ [1]

(b) Record the temperatures in the space next to the thermometer diagrams. [2] Initial temperature of hydrochloric acid Maximum temperature reached when magnesium was added

____________________ oC ____________________ oC

Temperature rise produced by 2.5 cm of magnesium ribbon = ______________________ oC Experiment 2 A 10 cm3 sample of the same hydrochloric acid was added to boiling tube. The initial temperature of the acid was measured. A 3cm length of magnesium ribbon was added to the acid in the boiling tube and the maximum temperature reached was measured. Experiment 3: Experiment 2 was repeated using a 4 cm length of magnesium ribbon. Experiment 4: Experiment 2 was repeated using a 5 cm length of magnesium ribbon. Experiment 5: Experiment 2 was repeated using a 6 cm length of magnesium ribbon.

(c) Record the temperatures in the space next to the thermometer diagrams and calculate the temperature rise in each case.

experiment Length of ribbon

Initial temperature of acid/oC

maximum temperature reached /oC

temperature rise /oC

2 2 cm

3 4 cm

4 5 cm

5 6 cm

[5]

20

25

30

20

25

30

20

25

30

40

45

50

20

25

30

20

25

30

3

4

4

30

35

40

20

25

30

30

35

40



(d) Plot the results of the Experiments 1 to 5 on the grid and draw straight line graph.

[4]

(a) From your graph, deduce the temperature rise of the mixture when a 1cm length of magnesium ribbon reacts with 10 cm3 of hydrochloric acid of the same concentration.

Show clearly on your graph how you worked out your answers [2]

(b) What word is used to describe a chemical reaction where the temperature increases? Ans: exothermic [1]

(c) Give two observations, other than temperature increaase, expected when magnesium reacts with

hydrochloric acid.

1. Ans: bubble of gas [1]

2. Ans: Mg disappears [1]

(d) Predict the temperature of the reaction mixture in Experiment 5 after 1 hour. Explain your answer. Ans: 25oC; room temperature [2]

(e) Explain one change that could be made to the experimental procedure to obtain more accurate

results. (f) Ans: 25oC; use powdered magnesium, faster reaction; so less heat loss [2]

M/J/ 1997/Q4 repeated 2017 and others

20

25

Temperatur

Length of magnesium ribbon /cm 1 2 3 4 5 6 0

5

10

15



Investigating the solubility of a salt Investigating solubility of a salt

To find the solubility of a substance measure 100cm3 of water, using pipette, 4 times / burette, 2 times heat/boil add lid to prevent evaporation add excess salt and stir until no more dissolves then cool to the desired temperature, then filter off the excess solid evaporate the filtrate to constant mass mass of residue = solubility

Worked typical question A student investigated the solubility of salt D in water at various temperatures. Four experiments were carried out. (a) Experiment 1:

4 g of salt D was added to a boiling tube. A burette was filled with distilled water and 10.0 cm3 of water added to the boiling tube. The mixture of salt D and water was heated carefully until all of the solid had dissolved. The boiling tube was removed from the heat and the solution allowed to cool. The solution was stirred gently with a thermometer.

The temperature at which crystals first appeared was noted. The boiling tube and its contents were kept for the remaining three experiments.

(b) Experiment 2: another 2.0 cm3 of water was added to the boiling tube and contents from experiment 1 (c) Experiment 3: another 2.0 cm3 of water was added to the boiling tube and contents from experiment 2 (d) Experiment 4: another 4.0 cm3 of water was added to the boiling tube and contents from experiment 3

The experiment was repeated exactly as in experiment 1. the table shows the temperatures at which crystals first appeared in the four experiments

Experiment number

total volume of water / cm3

temperature at which crystals first appeared / C

1 10 91 2 12 76 3 14 65 4 18 54



(a) Plot the results on the grid below and draw a smooth line graph.

(f) From your graph, find the temperature at which crystals of D would first appear if the total volume of water in the solution was 20.0 cm3. Show clearly on the grid how you worked out your answer. Ans: 52 [1] oC[1] shown clearly by extrapolation (1) [3] value from graph for 20 cm3

water, 50–53 (1) ± half a small square shown clearly by extrapolation (1) unit, °C (1)

(g) How would the student know when salt D was completely dissolved in the water? Ans: clear solution forms / no solid / crystals / salt visible (1) [1]

(h) The solubility of salt D at 100 C is 57 g in 100 cm3 of water. Suggest, with a reason, the effect of using 8 g of salt D instead of 4 g in these experiments. Ans: not all salt would dissolve (1) use of figures (1) [2] e.g. only 5.7 g would dissolve in 10 cm3 water at 100 °C

[2] (i) Salt C is less soluble in water than salt D.

Sketch on the grid the graph you would expect for salt C. Label this graph. [2] Ans: since less soluble, so same amount dissolves at highertemperature sketch graph always above line (1) label (1)

8.0 10.0 12.0 14.0 16.0 18.0 20.0

temperature at which crystals first appeared / oC

total volume of water / cm3

20

40

60

80

100



(j) Describe and explain one improvement that could be made to the experimental method to obtain more reliable results in this investigation.

Improvement

Ans: do not remove thermometer from solution - use IT method / second person to note

formation of crystals - repeat - do separate experiments - use smaller volumes of water - avoid evaporation

explanation Ans: loss of solid on thermometer

- observing formation of first crystals may vary - to take average - more results to plot on graph - method of avoiding evaporation e.g. separate

experiments, lid [2]

F/M/2015/62/Q4

2- “Grow Well” is a common garden fertilizer which is partly soluble in water. A student investigated the solubility of “Grow Well” in water at room temperature using the following procedure. stage1. Water was added to 10 g of “Grow Well” in a beaker. The mixture was boiled,

allowed to cool to room temperature and the then filtered. stage2. A 100 cm3 sample of the filtrate was evaporated to dryness. The solid remaining

had a mass of 7.5 g. (a) Why, in stage 1, was the water heated to boiling and then allowed to cool to room temperature?

Ans: Heated and boiled to speed up dissolving. allowed to cool to obtain a saturated solution.

(b) What piece of apparatus could have been used to measure the filtrate in stage 2?

(b) Ans: Pipette 4 times or burette 2 times.

(c) **How could the student show that all the water had been evaporated in stage 2? (c) Ans: Heat and reweigh many times until you reach a constant mass.

P6/Q5/June 1997

3. The information in the box is about the preparation of copper sulphate crystals.

step 1: Add a small amount of black copper oxide to some hot dilute sulphuric acid, and stir. step 2: Keep adding copper oxide until it is in excess. step 3: Remove the excess copper oxide to leave blue copper sulphate solution. step 4: Evaporate the copper sulphate solution until it is saturated. step 5: Leave the saturated solution of copper sulphate to cool. Blue copper sulphate crystals form on cooling. step 6: Remove the crystals from the solution remaining. step 7: Dry the blue crystals on a piece of filter paper.



(a) (i) Suggest a reason for using excess copper oxide in Step 2. Ans: to make sure all sulphuric acid has reacted / to neutralise all the acid

(ii) Suggest how the excess copper oxide can be removed from the solution in Step 3.

Ans: By filtration

(iii) What is meant by the term saturated solution?

Ans: Dissolved the maximum amount of solid at a certain temperature/ cannot dissolve any more solid at a certain temperature

(iv) Why do crystals form when a hot saturated solution cools?

Ans: solubility decreases (as temperature falls)/ less soluble in cold water

(v) Suggest why the blue crystals are dried in Step 7 using filter paper instead of by heating.

Ans: water of crystallization might be lost; so we get anhydrous CuSO4 instead of crystals

Investigating the effect of changing the temperature on the rate of reaction between hydrochloric acid and aqueous sodium thiosulfate

Rate and temperature 1 A student investigated the effect of temperature on the speed of reaction between hydrochloric acid

and aqueous sodium thiosulfate. When these chemicals react they form a precipitate, which makes the solution go cloudy. The formation of this precipitate can be used to show how fast the reaction proceeds, using the set up shown below.

Five experiments were carried out. Experiment 1 By using a measuring cylinder 50 cm3 of aqueous sodium thiosulfate was poured into a flask. The temperature of the solution was measured. The conical flask was placed on the printed text. 10 cm3 of hydrochloric acid was added to the flask and the timer started. The time taken for the printed text to disappear from view was recorded in the table. The final temperature of the mixture was measured.

eye

100 cm3 HCl solution

aqueous sodium thiosulfate

conical flask

printed sheet

printed X



Experiment 2 50 cm3 of aqueous sodium thiosulfate was poured into a conical flask. The solution was heated until the temperature was about 30 oC. The temperature of the solution was measured. 10 cm3 of hydrochloric acid was added to the flask and Experiment 1 was repeated. The final temperature of the liquid was measured. Experiment 3 Experiment 2 was repeated but the sodium thiosulfate solution was heated to about 40 ºC before adding the hydrochloric acid. The initial and final temperatures were measured. Experiment 4 Experiment 2 was repeated but the sodium thiosulfate solution was heated to about 50 ºC before adding the hydrochloric acid. The initial and final temperatures were measured.

Experiment 5 Experiment 2 was repeated but the sodium thiosulfate solution was heated to about 60 ºC before adding the hydrochloric acid. The initial and final temperatures were measured. Use the thermometer diagrams to record all of the initial and final temperatures in the table. (a) Complete the table of results to show the average temperatures.

Experiment thermometer diagram

initial temperature

/ oC

thermometer diagram

final temperature

/ oC

average temperature

/ oC

time for printed text

to disappear / s

1

24

24 24 130

2

33

31 32 79

3

40

38 39 55

4

51

47 49 33

5

60

54 57 26

30

25

30

25

60

55

55

50

45

40

40

35

60

55

55

50

45

40

40

35



(b) Plot the results obtained on the grid and draw a smooth line graph.

(c) (i) In which experiment was the speed of reaction greatest?

Ans: experiment 5 (1) [ the cross required least time to disappear] (ii) Explain why the speed was greatest in this experiment.

Ans: more kinetic energy (1) particles move faster (1) more frequencey of collisions (1)

(d) Why was the same volume of sodium thiosulfate solution and the same volume of hydrochloric acid used in each experiment?

Ans: for fair test; to compare effect of changing the temperature only (1)

[ in this question keeping the volumes the same means keeping the concentrations the same; so study the effect of changing temperature as the variable]

(e) (i) From your graph, deduce the time for the printed text to disappear if Experiment 2 was to be repeated at 70 ºC. Show clearly on the grid how you worked out your answer.

Ans: approx 20 sec , from the graph [ value 1 mark, unit 1 mark] extrapolation is shown (1)

e(ii)

*

*

*

**

20

40

time for printed text to disappear / s

10 20 30 40 5000

60

60average temperature/oC

100

120

80

140

70

5 points correctly plotted (3), –1 for any incorrect point smooth line graph (1)



(ii) Sketch on the grid the curve you would expect if all the experiments were repeated using 50 cm3 of more concentrated sodium thiosulfate solution. curve sketched on grid below original curve

Ans: look at the graph; it is draw and labeled Steeper and below that using 25cm3.

(f) Explain one change that could be made to the experimental method to obtain more

accurate results.

Change: Ans: 1.use of data logger (1) or 2use of insulating sheath / 3use a burette or pipette

Explanation: Ans: e.g. 1timing of reaction more accurate (1) 2to reduce heat losses / 3volumes more accurate

For investigating the rate of the reaction between thiosulphate and hydrochloric acid

equation of the reaction:

2 2 3 2 2yellow ppt

(aq) (aq) (aq) (g) ( ) (s)Na S O + 2HCl 2NaCl + SO + H O + S l

Draw a cross on a paper.

Possible questions and suggested answers

In which order should water, HCl(aq) acid and Na2S2O3(aq) should be added to the beaker?

answer: 1st : thiosulphate - 2nd : water - 3rd : hydrochloric acid

Why does the cross on the paper disappear?

Answer: the formed insoluble sulfur (1) precipitates (1) or the formed insoluble sulfur (1) turns the solution cloudy/ turbid (1).

Why was the total volume of solution kept constant?

Answer: because the concentration is proportional to volume/ for fair test/ make comparison/same depth(1)

If the experiments were repeated using a 250 cm3 beaker instead of a 100 cm3 beaker.

Suggest how the results would differ. Explain your answer.

Answer: times would be longer for the cross to disappear(1) idea of smaller depth(1) because the precipitate needs to spread over larger surface area to cover the cross”

With higher concentration the reaction is faster; the cross takes shorter time to disappear

10 cm3 hydrochloric acid

eye

beaker aqueous sodium thiosulphate

paper with cross marked on it



(ii) a 100 cm3 conical flask instead of a 250 cm3 conical flask. [2] time shorter / cross disappears faster [1] depth of solution is greater [1]

w_17_61 Suggest the effect on the results of using a 100 cm3

conical flask instead of a 250 cm3 conical

flask. Explain your answer. [2] times would be shorter / cross disappears faster; depth of solution is greater S_18_61 + S_15_62

Sketch on the grid the graph you would expect if the experiments were repeated at higher temperature (50 °C). Label this graph.

Answer: line lower down (steeper) because it is faster; i.e, takes shorter time (1) doesn’t touch the other line because the concentration of the limiting reactant is higher. (1) Either the temperature is the variable or the concentration of thiosulphate / HCl

- If temperature is the variable, then we keep the volume of all solutions the same to compare the effect of temperature as the only variable; to have fair test.

- If concentration is the variable, then we keep the volume of all solutions the same because volume is proportional to concentration; to have fair test. Sketch on the grid the graph you would expect if the experiments were repeated using higher concentration of hydrochloric acid.

To have fair test Same mass/length/ size of solid Same concentration of same solution Same volume of same solution as volume is proportional to concentration Same temperature for all experiments / tests

To have fair test in precipitation reactions Same concentration of both solutions *Same total volume of solution as volume is proportional to concentration Same width of tube to compare between the height of the precipitate Same depth of tube to compare between the height of the precipitate Same temperature for all experiments / tests

At higher temperature the error is more Reason for errors in time at high temperature during studying rates of reactions is that

At higher temperature reaction goes faster so percentage error in time increases



Investigating effect of different catalyst on decomposition reaction of H2O2

1. Hydrogen peroxide breaks down to form oxygen. The volume of oxygen given off can be measured

using gas syringe.

Solids W and X both catalyse the breakdown of hydrogen peroxide. the table below shows the results with W and X at 25 °C

Volume / cm3 Time / s catalyst W catalyst X

0 0 0 20 16 29 40 32 34 60 36 36 80 37 37 100 37 37

all correct (3) -1 each incorrect

(b) Plot a graph to show each set of results. Clearly label the curves. choice of suitable scale for y-axis (1) [ choosing the largest possible scale minimizes the error in drawing]

all points correctly plotted (3) smooth curves (1) labelled (1)

[5]

aqueous hydrogen peroxide

catalyst

gas syringe

it is a single reactant reaction; no limiting reactant

W

X

*

5

1 0

1 5

2 0

2 5

3 0

3 5

4 0

*

* * *

*

** *

Volume of gas/ cm3

Time/ s 2 0 4 0 6 0 8 0 1 0 000

2 5

3 0

3 5

4 0

oX+40 C

A catalyst increases the rate of a reaction many times more than increasing temperature



(c) Which solid is the better catalyst in this reaction? Give a reason for your choice. solid … Ans: X, (1) *reason: Ans more gas given off at 20 s and 40 s, indicating faster reaction(1)

(d) *same volume of hydrogen peroxide used in both experiments (1)

(e) *Sketch a line on the grid to show the shape of the graph you would expect if the reaction with catalyst X was repeated at 40 °C.

Ans: Smooth line sketched on grid with steeper slope than for catalyst X at 25°C (1) ( above the line of X alone) levelling out at same level (1) [ look at the graph]

In all changes the new graph levels out at same level; except with changing the number of moles of the limiting reactant [ or the number of moles of the reactant in single reactant reactions

Purification of metal such as copper by electrolysis In electrolysis Electroplating - In electrolysis

Tests for ions and gases 2 **Tests to distinguish between two substances

Describe a chemical test to distinguish between each of the following pairs of substances. An example is given.

Example: hydrogen and carbon dioxide

test lighted splint result with hydrogen: gives a pop

result with carbon dioxide splint is extinguished

(a) zinc carbonate and zinc chloride:

- Test: add dilute hydrochloric acid

- Result zinc carbonate: Ans: effervesence / fizz test for carbonate

- Result zinc chloride: Ans: no change / no fizz

or - Test: add HNO3(aq) and AgNO3(aq)

- Result zinc carbonate: Ans: no change

- Result zinc chloride: Ans: white precipitate test for chloride

(b) Ammonia and chlorine:

- Test: damp red litmus paper

- Result with Ammonia: Ans: turns blue

- Result with chlorine: Ans: bleached

(c) iron(II) sulfate and iron(III) sulfate:

- Test: add sodium hydroxide

- Result with aqueous iron(II) sulfate : Ans: green precipitate

- Result with aqueous iron(III) sulfate : Ans: brown precipitate



3 Two solids, S and V, were analysed. S was copper (II) oxide. The tests on the solids, and some of the observations are in the following table. Complete the observations in the table. Do not write any conclusions in the table

Tests observation

tests on solid S (a) Appearance of solid S black solid

(b) Hydrogen peroxide was added to solid S in a test-tube. A glowing splint was inserted into the tube.

slow effervescence

splint relit

(c) Dilute sulfuric acid was added to solid S in a test-tube. The mixture was heated to boiling point. The solution was divided into three equal portions into test-tubes.

(i) To the first portion of the solution, excess

sodium hydroxide was added.

(ii) To the second portion of the solution, about 1 cm3 of aqueous ammonia solution was added. Excess ammonia solution was then added.

blue solution formed

(i) Ans: blue precipitate (1)

(ii) Ans: blue precipitate (1)(1)

Ans: precipitate dissolves producing deep/royal blue solution (1)(1)

(iii) To the third portion of the solution, dilute

hydrochloric acid was added followed by

barium chloride solution.

(iii) Ans: white (1) precipitate (1)

tests on solid V

(d) Appearance of solid V

black solid

(e) Hydrogen peroxide was added to solid S in a test-tube. A glowing splint was inserted into the tube.

rapid effervescence

splint relit

(f) (i) Compare the reactivity of solid S and solid V with hydrogen peroxide. Ans: V is more reactive catalyst than S (1)

(ii) Identify the gas given off in test (e). Ans: oxygen (1)

(g) What conclusions can you draw about solid V? Ans: it is transition metal compound ( being coloured) (1) It is catalyst [ because it speeded up decomposition of H2O2] (1)

It is MnO2 [ you need to remember that MnO2 is the catalyst for decomposition of H2O2](1)



5 A mixture of two solids, R and S, was analysed Solid R was the water-soluble salt aluminium sulfate, Al 2(SO4)3, and solid S was an insoluble salt.

The tests on the mixture and some of the observations are in the following table. Complete the observations in the table.

tests observations Distilled water was added to the mixture in a boiling tube. The boiling tube was shaken and the contents of the boiling tube filtered, keeping the filtrate and residue for the following tests. The filtrate was divided into five test-tubes.

tests on the filtrate (a) Appearance of the first portion of the

filtrate.

.Ans: Colourless [1]

(b) Drops of aqueous sodium hydroxide were added to the second portion of the solution and the test-tube shaken. Excess aqueous sodium hydroxide was then added to the test-tube.

Ans: white precipitate ppt dissolves (in excess) [3]

(c) Aqueous ammonia was added to the third portion, dropwise and then in excess.

Ans: white precipitate ppt doesn’t dissolve (in excess) [2]

(d) Dilute nitric acid was added to the fourth portion of the solution followed by aqueous silver nitrate.

Ans: no change [ reason: because no halides] [1]

(e) Dilute nitric acid was added to the fifth portion of the solution and then aqueous barium nitrate.

Ans: white precipitate [1] [ reason: it contains sulphate]

tests observations tests on the residue (f) Dilute hydrochloric acid was added to the

residue. The gas given off was tested. Excess aqueous sodium hydroxide was added to the mixture in the test-tube.

rapid effervescence limewater turned milky

white precipitate, insoluble in excess

(g) Name the gas given off in test (f). Ans: carbon dioxide [1] (h) What conclusions can you draw about solid S? Ans: it is CaCO3 ( contains Ca2+ and CO3

2- ion [2]



Planning an experiment (practical method) Catalyst , Reactivity of metals – rank in order , Rate , Indicator, Strength of acids and bases Concentration of acids and bases, Carbonate in an acid, Chromatography, Best fuel Heat of a chemical reaction – measurements, Fertilizer Your plan should include reagents (1) masses and volumes (1) method – crush if solid - (1) result (1) filter if necessary – reweigh if necessary – compare if necessary (1) take into account the fair test when comparing between two metals, two acidic solutions, two concrete; that is, equal mass, equal particles size, equal volumes equal concentrations. How to plan an experiment to compare between two solids, use same mass and same particles size of each solid for fair test to compare between two liquids/ solutions, use same volume and same concentration of each

solution/ liquid for fair test

to compare between the concentrations of different solutions of same acid, - titrate equal volume of each acid with same alkali solution, of same concentration

note the volume of alkali needed to neutralize each acid. the acid neutralized by larger volume of alkali is the most concentrated.

or - react equal volumes of each acid with same mass and same particles size of same metal/

metal carbonate - record 1.the volume of the produced gas every 5 seconds using gas syringe or 2.highest

temperature rise or 3.observe effervescence - compare - the acid that produces larger volume of gas/ higher temperature rise/ more bubbles is most

concentrated

To compare between the reactivity of different metals, - Add 10 cm3 of 1 mol/ dm3 HCl(aq) to 50 grams of powdered metal A - measure 1.the volume of the produced gas every 5 seconds or 2.highest temperature rise or

3.observe effervescence - Repeat the same experiment using same mass and same particles size of metal B with

same volume of same acid (same concentration) - For fair test - Compare - The metal that produces largest volume of gas, /or/ causes highest temperature rise/or/

causes most effervescence is most reactive. - The metal that doesn’t show any evidence of chemical reaction is unreactive.

***To extract a metal from its ore such as copper from malachite.

- Crush using mortar and pestle to increase the surface - Heat malachite to decompose giving CuO - reduce CuO into Cu by heating with C because carbon more reactive than C so it reduce it

Coe escapes leaving copper or



- add dilute H2SO4 to CuO to change it into CuSO4 - electrolysis using Cu – electrodes - Cu2+ ions in the solution are reduced into Cu atoms that deposit at cathode.

To find the volume of gas produced by a carbonate / metal,

- add excess acid and measure the volume of the produced gas using gas syringe if soluble in water or over water if insoluble in water [H2] or O2, CO2 Cl2, slightly soluble

To show that a substance X is a catalyst for a certain substance Y, - add a small amount of X to a volume of substance Y. observe - the reaction is faster [ evidence: more bubbles / effervescence stops in shorter time/

higher temperature rise,….] - filter, wash dry and reweigh. - If the mass of the residue equals to the added mass of X, - then X is a catalyst because it speeded up the reaction and not used up

To compare between the catalytic activity of two substances Z and W for a certain

substance, - Add 1 gram powder of Z to 50cm3 of the substance and measure the volume of the

produced gas every 5 seconds until the reaction stops. - Repeat the same experiment with 1 gram powder of W and 50 cm3 of the substance. - Compare the volumes. The one that produces larger volume in same time interval is better

catalyst.

To prepare an indicator from a plant such as red cabbage - Chop up 5 leaves of the plant and crush using mortar and pestle. - Boil in 50cm3 of water for 20 minutes. - Filter; the filtrate is indicator solution. - Check whether the solution is indicator or not - Add 5 drops of the solution to HCl(aq), to NaOH and note the colours - The colours in the acid and the alkali must be different - Note the colours down. To be used for identification of acid.

To show that a given / prepared solution is an indicator,

- Add 5 drops of the liquid to hydrochloric acid and note the colour - Add same number of drops of the liquid to an alkali, NaOH and note the colour - The two colours must be different so it is an indicator. - ( the indicator give different colours in acids and bases)

To find the acidity of soil [ compare]

- Add 20g of soil A to 50cm3 of water and stir thoroughly - Filter off the soil - Take drops of the filtrate and add to pH paper - Compare the colour produced with the universal indicator paper. Each colour indicate pH number - The soil with higher pH is less acidic

mortar

pestle



To compare between the effectiveness of two fertilizers. [which is better]

- In a pot, mix 100 grams of fertilizer A with 1000 grams of soil and plant 10 seeds of beans - Repeat the same experiment with fertilizer B and put both pots in same area to receive

same amount of light and water with same amount daily for two weeks. - The pots inside which beans shoots grow faster contains better fertilizer.

To prepare a saturated solution

- add excess soluble salt to a beaker containing water. - boil the mixture for several minutes. - leave stand to cool gradually to room temperature, then - Extra soluble salt (solute) will settle down leaving a saturated solution. - filter; the filtrate is the saturated solution

To find the solubility of a substance

measure 100cm3 of water, using pipette, 4 times / burette, 2 times [100cm3 of water =100g] heat/boil add lid to prevent evaporation add excess salt and stir until no more dissolves then cool to the desired temperature, then filter off the excess solid evaporate the filtrate to constant mass mass of residue = solubility

To know which fuel produces more energy, - Place 50cm3 of water in a beaker and measure its initial temperature - burn 5 grams of fuel A to heat the 50cm3 of water and note the highest temperature reached. - Repeat the same experiment with fuel B and note the highest temperature reached. - The one that causes higher temperature rise produces more energy.

To compare between the amounts of sulphur in a fuel,

- burn one fuel and pass the produced gas through orange K2Cr2O7, and note the time it takes to turn green.

- Repeat same experiment with the other fuel using same K2Cr2O7 solution. - Compare. - The one that takes shorter time to turn the solution green contains more sulphur.

To determine the percentage of calcium carbonate in egg shell or any other solid

- Add excess HCl solution to 100 grams of crushed egg shell, [ excess to ensure all CaCO3 reacted] - when effervescence stops, filter, wash and dry. - reweigh the residue. - Percentage of CaCO3 = 100 – mass of residue.

If the available mass of egg shell is only 50 grams, then do the same experiment % of CaCO3 = (100 - mass of residue) x 2

To investigate the colours in a solid substance, 1st Extract the pigment solution 2nd carry out chromatography



Steps of extracting a liquid from a solid 1. Crush a certain mass ( or number of leaves) using mortar and pestle [ to speed up extraction] 2. Put in a solvent [ to dissolve and extract] 3. Boil for 20 minutes [ for effective extraction] 4. Filter [ to separate the liquid from the solid] 5. Heat the filtrate to evaporate ¾ of the volume of the solvent and concentrate the solution Use chromatography as follows Steps

- Apply a spot of the concentrated solution on the pencil line of a chromatography paper - Put the paper in the solvent such that base – line is above the level of the solvent. - Run chromatography until solvent front reaches near the top. - Remove the paper and let it dry. - Compare with chromatograms of pure substances or measure Rf values and match with Rf

values of pure substances. To identify rusting chemicals

- Carry out the reaction using boiled water covered with oil layer [ no O2] label tube 1 - Repeat with iron paced in a drying agent such as CaO[ no water] label tube 2 - Repeat with open tube containing water at room temperature [ both air/ O2 and water are

present. label tube 3 - Repeat with open air [ both air/ O2 and water vapour are present. label tube 4 - Observe and not any changes every 8 hours. - The nail in tubes 3 and 4 rusts - Therefore, rusting needs both water and O2.

Effect of salt on rusting

Place one nail in salt and another nail of same size in open air without salt Observe any change every day The nail in salt rust first Therefore, salt speeds up rusting



Solved past papers plans To show that a substance is a catalyst

1 Is manganese (IV) oxide a catalyst?

A catalyst is a substance that speeds up a chemical reaction and remains unchanged.

Hydrogen peroxide, H2O2 breaks down to form oxygen. This reaction is very slow without a catalyst.

Describe an experiment to show that manganese(IV) oxide is a catalyst for this reaction.

You are provided with the following items. Hydrogen peroxide solution Manganese(IV) oxide Measuring cylinder Filtration apparatus Balance Beaker Splints/Bunsen burner Distilled water Answer:

- Add 1g of manganese(IV) oxide to 20cm3 of hydrogen peroxide and note your results - Result: Faster bubbles of gas than without MnO2. - Test the gas with glowing splint; it relights, so the gas is O2 - Filter, wash, and dry the solid - Reweigh and compare - Mass of manganese (IV) oxide is 1g; it is not used up and speeded up the reaction. - Therefore, MnO2 is a catalyst. [6]

M/J/2004/Q8

2 Concrete is made from cement, sand and small stones. The ratio (by volume) of these three components varies with the use of the concrete. Over a period of years, the cement slowly reacts with carbon dioxide in the air to form calcium carbonate. Calcium carbonate can be dissolved in hydrochloric add.

Outline an experiment to find out the composition of a sample of old concrete.

Answer - Put 20g of crushed sample of old concrete in a beaker. - Add 100 ml of dilute HCl solution and stir. - When effervescence stops, filter, then - Wash the residue on the filter paper with distilled water. - Dry the residue, and then weigh it. - Weight of residue= weight of (sand +stones) - Weight of calcium carbonate= 20g – weight of residue

[5] O/N/1998/Q6

Exothermic and endothermic reactions 3 When cement powder is added to water a reaction takes place.

Describe an experiment to show that this reaction is exothermic. [4] Plan

put 100cm3 of water in a beaker and meaure its initial temperature add cement and stir using the thermometer in the beaker (1) measure the temperature (1) there is temperature rise (2) the reaction released heat therefore exothermic.

0620_s07_qp_6[_Q7(a)



Using titration 4 A student used the pieces of equipment shown below to compare the concentration of

alkali in two liquid oven cleaners. Oven cleaners contain the alkali sodium hydroxide.

(a) Name each piece of equipment.

Answer(a) A: burette B: pipette.

(b) Outline how these pieces of equipment could be used in an investigation to compare the concentration of alkalis in the two liquid oven cleaners. Answer: use titration - Fill the burette with 0.1 mol/dm3 HCl solution. - Using a pipette, transfer 10 cm3 of one of the two alkali into a conical flask, then add 3 drops

of phenolphthalein indicator. - From the burette add HCl to the conical flask, 1cm3 at a time and swirl until pink colour

disappears permanently - Record the added volume. - Repeat the same experiment with 10cm3 of the other alkali and record the added volume. - The alkali that needed larger volume of HCl to neutralize it is more concentrated.

[5] M/J/2000/Q6

5 Oven cleaners (b) Some liquid oven cleaners contain particles of an insoluble solid, bentonite, suspended in

an aqueous solution. Outline an experiment to obtain a pure sample of bentonite from the oven cleaner. Answer : filter the solution (1) wash with water (1) dry (1) [3] do not allow: evaporate to dryness.

(c) Oven cleaners contain an aqueous solution of sodium hydroxide.

Plan an investigation to show which of two different oven cleaners, C and D, contains the more concentrated solution of sodium hydroxide. You are provided with common laboratory apparatus and chemicals.

Answer: use titration method - Fill the burette with 0.1 mol/dm3 HCl solution. - Using a pipette, transfer 10 cm3 of cleaner C into a conical flask, then add 3 drops of

phenolphthalein indicator. - From the burette add HCl to the conical flask, 1cm3 at a time, and swirl until pink colour

disappears permanently - Record the added volume. - Repeat the same experiment with 10cm3 of cleaner D and record the added volume. - The cleaner that needed larger volume of HCl to neutralize it is more concentrated NaOH(aq) [5]

0620_w14_qp_62/ Q6

A B



6 Kleen Up is a colourless liquid used to clean work surfaces and glass windows. Kleen Up contains ammonia solution, which is a weak alkali.

(a) State a chemical test to show the presence of ammonia in Kleen Up. Test: Answer approach damp red litmus paper near the mouth of Kleen up bottle result Answer turns blue [2]

(b) Plan an experiment to determine the concentration of ammonia in Kleen Up. You are provided with aqueous nitric acid of known concentration and common laboratory apparatus.

Answer: Using a pipette, measure 25cm3 of Kleen Up and transfer to a conical flask (1), then add indicator such as phenolphthalein. (1) Add nitric acid from the burette (1), [ or any other acid] until neutral point ( pink colour turns colorless)(1) note volume of acid (1) Calculate concentration (1). [5]

0620/62/O/N/2011/Q6

7 Leaves from trees contain a mixture of coloured pigments which are not soluble in water. A student was given these two instructions to investigate the pigments in the leaves.

2. Crush some leaves to extract the coloured pigments. 3. Use the liquid extract to find the number of coloured pigments in the leaves.

(a) *What would the student need in order to effectively carry out instruction 1? Answer: - Crush some leaves using mortar and pestle and boil in ethanol (solvent) for several minutes - Filter; the filtrate is solution of the extracted pigments. - Heat the filtrate to evaporate ¾ of the volume the solvent and concentrate the solution. [3]

Answer:

- Crush some leaves using mortar and pestle and boil in ethanol (solvent) for several minutes

- Filter; the filtrate is solution of the extracted pigments. - Heat the filtrate to evaporate ¾ of the volume the solvent and concentrate the solution.

[3] Extracting pigments/ chromatography / 0620/06/O/N/09/Q7

(b) Describe an experiment to carry out instruction 2.

A space has been left below if you want to draw a diagram to help answer the question. Answer:

- Use chromatography technique (1) - apply drop of lemon oil to the pencil base-line on the paper (1) - put the paper in ethanol, make sure the base line is above the level of the solvent - Cover the beaker with a lid to prevent evaporation of the solvent. (1) - Run the chromatography until the solvent front reaches near top of the paper. - The number of the separated spots equals to the number of the coloured

pigment [4] Extracting pigments/ chromatography / 0620/06/O/N/09/Q7



8 The label on an aerosol can of Kleen Air air freshener is shown.

(a) What is meant by the term solvent? Answer: a substance that dissolves another substance (1)

(b) What does the hazard sign indicate? Answer: flammable (1)

(c) What method could be used to obtain ethanol (boiling point 78 °C) from a mixture of ethanol and propanone (boiling point 56 °C)? Answer: fractional distillation (1)

(d) Describe an experiment to investigate the number of coloured substances present in a sample of the lemon oil obtained from Kleen Air.

9 The label shows the substances present in a bottle of orange fruit drink.

ORANGE FRUIT DRINK Contains: orange juice, malic acid, citric acid and natural colours (carotenes)

NO ARTIFICIAL COLOURS (E NUMBERS)

(a) A piece of pH indicator paper was dipped in the drink. (i) Predict the pH value obtained.

Answer: less than 7 [ all juices are acidic] [1] (ii) Why does the pH indicator paper give a more reliable result than adding Universal Indicator

solution to the drink? **Answer: colour of orange drink obscures indicator colour.(1)

(b) Describe an experiment you could carry out to show that only natural colours were present in the drink. A space has been left if you want to draw a diagram to help you answer the question.

Answer:

- Use chromatography technique (1) - apply drop of orange drink fruit on the pencil base-line on the paper (1) - put the paper in ethanol, make sure the base line is above the level of the solvent - Cover the beaker with a lid to prevent evaporation of the solvent. (1) - Run the chromatography until the solvent front reaches near top of the paper. - Compare Rf values of the spots with those of carotenes, they are the same [4]

Chromatography / 0620/63/O/N/10/Q7



[ plan- chromatography ( extract and separate)] Some trees have purple leaves. The purple colour is a mixture of coloured pigments.

Plan an experiment to extract and separate the coloured pigments present in the purple leaves.

You are provided with some purple leaves, sand, ethanol and common laboratory apparatus. You may draw a diagram to help you answer the question. [6]

Question Answer Marks

3(a) any 6 from: cut leaves into small pieces grind / crush with sand / ethanol using pestle/mortar decant / pour-off / filter liquid chromatography apply extract to paper (in correct location) description of separating colours -

Max 6

0620_s18_qp_63_Q 4

10 E numbers identify chemicals which are added to foods. (a) E210 is benzoic acid. How could you show that a solution of benzoic acid is a weak acid?

test......................................................................................................................................

result............................................................................................................................. [2]

(b) E211 is sodium benzoate. Name a suitable substance that would react with a solution of

benzoic acid to form sodium benzoate. ...................................................................................................................................... [1]

(c) E110 is Sunset yellow. Outline a method you could use to show the presence of E110 in a food colouring. A space has been left if you want to draw a diagram to help you answer the question.

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........................................................................................................................................ [4] Chromatography / 0620/61/O/N/10/Q7



Fuel investigation Most details in energetics topic

(a) Outline a different experiment to find out which type of coal produces most sulphur dioxide on heating. You may find it useful to draw a diagram. _________________________________________________________________________ _________________________________________________________________________ [4]

11 Sulphur dioxide gas is a common pollutant formed when fossil fuels burn in air. Sulphur dioxide can be detected by using an acidic solution of potassium dichromate(VI), The dichromate solution changes colour from orange to green when a certain amount of sulphur dioxide has reacted with it. Plan an experiment to investigate which of three different samples of coal produces most sulphur dioxide Answer:

- Burn each identical amount of each coal (same mass and same particles size) separately and pass the produced gases through orange K2Cr2O7 and note the time it takes to turn the solution green.

- Compare

- The fuel that takes the shortest time to turn the solution green contains more sulphur.[6] O/N/2003/Q6/ sulphur dioxide

(b) the apparatus below and stop watch. The coal which produces most sulphur dioxide will turn the dichromate solution from orange to green faster.

acidified dichromate solution “orange”

pump

Burning coal

cotton wool



Heat plans 12 Alcohols are liquid fuels which burn in air to produce heat energy. Provided with the following

apparatus, plan an experiment to investigate which alcohol, methanol or ethanol, produces more energy. You may draw a diagram to help your answer.

Plan [5] Answer: Set up the apparatus as shown, and measure initial temperature - Burn 5 grams of methanol to heat 100 grams of water. - Measure initial and final temperature of water. - Calculate temperature rise. - Repeat the same experiment with ethanol. - Compare. - The one that cause higher temperature rise produces more

heat.

13 Petrol is a liquid fuel obtained from petroleum (crude oil).

Bioethanol is a liquid fuel made by the fermentation of carbohydrates obtained from plants such as sugar cane. Using the apparatus below, plan an experiment to investigate which of these two fuels produces more energy. You may use the space below to draw a diagram to help you answer the question.

25.26

Spirit burner

Balance

fuel

Retort stand

water

Measuring cylinder

25.26

fueSpirit burner

Balance

thermometer

Copper can



Answer: - Burn 5 grams of petrol to heat 100 grams of water. - Measure initial and final temperatures of water. - Calculate temperature rise. - Repeat the same experiment with bioethanol. - Compare. - The fuel that cause higher temperature rise produces more heat. [6]

0620/63/62/M/J/2011/Q6 Fuel/ heat

Rusting 14 Iron rusts when in contact with air and water.

You are provided with iron nails and three different samples of water: tap water, sea water, distilled water. Plan an investigation to find out which sample of water causes iron to rust the fastest. Answer: - Add 100cm3 of water (1)

- To 5 nails (1) in a beaker (1) - leave for few days until nails rust , and note the time (1) - repeat with other water samples of water(1), using same volume of water and same

number of identical nails each time. - compare the results (1) - the water in which nails took less time to rust are fastest to cause nails rust.

[4] 0620/62/O/N/10/Q7 / Rust

15 When iron nails rust, the mass of the nails Increases as the iron reacts with oxygen and

water. You are provided with the following apparatus. Plan an experiment to investigate if iron rusts more quickty in sea water or fresh water. clean iron nails - test-tubes – beakers - a balance - an oven - measuring cylinder

Answer: - Using the measuring cylinder transfer 100 cm3 of distilled water to a beaker. - Weigh 50 grams of clean iron nails and put in the beaker of fresh water. - Repeat the same experiment with same volume of seawater, using same mass of

identical nails. - After two days remove the nails, wash, dry, reweigh and calculate increase in mass in

both experiments - Compare the increase in mass - Mass of nail increases more in seawater than in fresh water - Therefore, it rusts more quickly in seawater

[5] 0620/62/M/J/2001/Q6 / Rust

Repeated O/N 2017

16 Saucepans can be made from copper, aluminum or steel. You are provided with samples of the three different metals. Outline chemical experiments that could be carried out to show which would be most suitable to make saucepans. Marks will be awarded for practical details and expected observations. Answer: - Using Bunsen burner, heat a piece of steel for long time, colour changes to red

brown; it rusts - Heat identical piece of copper for long time, black powder forms on its surface - Repeat long heating with identical piece of aluminum, no observed change - Therefore; aluminum is most suitable for saucepan [5]

O/N/96/Q6



17 Metal cooking containers, such as saucepans, can be made from copper or steel. Outline experiments that could be carried out to show which of these metals would be most suitable for a saucepan You are provided with pieces of copper and steel foil. Common laboratory chemicals and apparatus are available.

Answer: - add vinegar to a piece of steel and heat until colour changes to red brown - repeat the same experiment with copper and heat for same time - no change with copper - therefore copper id most suitable for a saucepan than steel [6]

Uses of metals/ 0620/62/M/J/10/Q7 Worked past papers question on separating mixtures of solids, where only one dissolves in water 1. *Zinc carbonate is insoluble in water but zinc nitrate is soluble.

Outline how you could obtain from a mixture of zinc carbonate and zinc nitrate.

(a) A pure, dry sample of zinc carbonate. Ans: Put the mixture in a beaker, then add enough water to it, warm and stir until all zinc nitrate dissolve

Filter Wash the residue with pure water, then dry it [3]

(b) Crystals of zinc nitrate. Ans:: Put the filtrate from (a) in an evaporating basin. Heat gently to evaporate water to crystallization point

Cool slowly for crystals to form. Collect the crystals by filtration Dry between filter papers or in an oven at low temperature [3]

Separating gases / 0620/06/Nov/94/Q2

Note: A metal is obtained from its solution by one of two methods Displacement reaction using a more reactive metal Electrolysis: it is obtained at cathode.

Finding the volume of the gas, CO2, in lemonade 1. When lemonade is heated carbon dioxide gas is given off,

(a) Draw a labelled diagram of apparatus that could be used to find the volume of carbon dioxide dissolved in 100cm3 of lemonade.

Ans

(b) How could it be shown that all of the carbon dioxide gas had been removed from the lemonade? Ans: No more gas is added to the syringe [2]

Calcium carbonate in eggshell 1. Beach sand is a mixture of sand and broken shells (calcium carbonate). Calcium carbonate reacts

with dilute hydrochloric acid to form a solution of calcium chloride.

Plan an investigation to find out the percentage of shell material in a given sample of beach sand

Ans: crush 100 g of sand beach then add to it (1) excess (1) dilute hydrochloric acid (1) Warm and stir until bubbling (fizzing) stops. Filter (1) wash (1) dry (1) residue and weigh % percentage of CaCO3 in the mixture =100 – mass of residue (1) [6]

Separation methods-Plan investigation/ 0620/06/May/Jun/2003/Q6

100cm3 of lemonade

collecting water

ice bath

Heat

CO2 gas

gas syringe



to distinguish between strong and weak acids from their reactions with reactive metals 18 You are given solutions of lactic acid and hydrochloric acid, which have the same

concentrations. Describe how you could use some magnesium ribbon to show that lactic acid is a weaker acid than the hydrochloric acid. State how you would make it a fair test.

Answer: Add same volume of each acid to identical amounts of magnesium - same mass and same particles size,

With lactic acid less bubbles of gas, lower temperature rise (solution becomes warm), and magnesium

takes longer time to dissolve (disappear)

With hydrochloric acid: more bubbles of gas, higher temperature rise (solution becomes hot), and

magnesium takes shorter time to dissolve (disappear)

Acid – base indicator plan 19 Acid base indicators

Indicators are used to identify acids and bases. Indicators can be obtained from berries and other fruits.

(a) Plan an experiment to obtain an aqueous solution of an indicator from some berries. Ans: Plan

- Crush 5 berries using mortar and pestle (1) - Add water, heat and stir (1) - Filter (1) (decant or seive is accepted) - Collect the filtrate in a bottle

(b) Plan an experiment to use the indicator solution to show that it is an effective indicator.

Ans: add 3 drops of the collected indicator solution to hydrochloric acid and note the colour (1)

- Add another 3 drops of indicator solution to NaOH solution and note the colour (1) the two colours are different.

- Conclusion: it is effective indicator because it gave different colours in acids and alkalis (1)

Reject using base instead of alkali as not all bases are soluble in water. Acid-base indicators/ 0620/06/M/J/09/Q6

20 You are provided with cans of a fizzy drink - Koola cola. (a) What is the pH of the cola?

Ans: use pH meter, pH =5 [ carbonic acid is weak acid] or [2]

Use universal indicator, it goes orange; pH 5.

(b) How many coloured pigments does the cola contain?

Ans: Apply a spot of the concentrated solution on the pencil line of a chromatography paper - Put the paper in the solvent such that base – line is above the level of the solvent. - Run chromatography until solvent front reaches near the top. - Remove the paper and let it dry. - Number of spots = number of coloured pigments. [3]



(c) What volume of gas is released when a can of cola is opened? [Note: The can will have to be opened under water.]

Ans: open the bottle under water and collect the gas in a measuring cylinerover water. Volume of the gas = volume of displaced water. [2]

(d) Is the gas released carbon dioxide? Ans:yes; because it turns limewater [1] milky. [1]

0620/06/O/N/2001/Q6

21 ELECTROPLATING A COPPER KEY Electroplating is when a metal is coaled with another metal using electricity. To electroplate a metal a very clean surface is needed. Describe an experiment to nickel plate a copper key. You are provided with the following items.

6 V bulb and holder 6 V battery and connecting wires 250 cm3 beaker solid nickel(II) sulphate, NiSO4 steel wool/sandpaper copper key distilled water nickel rod

You can use a labelled diagram to help you answer the question. ..................................................................................................................................................................

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0620/06/M/J/ 2002/Q8 Salt solution plan 22 A solution of magnesium sulphate can be made by reacting magnesium oxide with warm sulphuric acid.

(a) Describe how you could make a solution of magnesium sulphate starting with magnesium oxide powder and dilute sulphuric acid.

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Preparing salt crystals plan (b) Describe how you would obtain pure dry crystals of hydrated magnesium sulphate,

MgSO4.7H2O, from the solution of magnesium sulphate in (a). ..................................................................................................................................................................

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............................................................................................................................................................. [3] Making solutions and salts/ 0620/06/O/N/08/Q7

23 This label is from a container of ‘Bite Relief’ solution.

(a) Give a chemical test to show the presence of ammonia in Bite Relief solution.

Test.....................................................................................................................................

Result ............................................................................................................................. [2]

(b) What practical method could be used to separate the mixture of alcohol (bp 78°C) and water (bp 100°C)? [2]

........................................................................................................................................ [1] (c) Give a chemical test to show the presence of water.

Test.....................................................................................................................................

Result ............................................................................................................................. [2]

(d) What would be the effect of touching the alcohol with a lighted splint? [1]

.................................................................................................................. …….............. [1]

0620/06/M/J/08/Q7

BITE RELIEF FOR FAST RELIEF FROM INSECT BITES AND STINGS Active ingredient: Ammonia Also contains water and alcohol DIRECTIONS FOR USE: Use cotton wool to dab the solution on the affected area of the skin



****Plan to obtain a metal from its ore 24 Cerussite is an ore of lead. Cerussite is lead carbonate, PbCO3. In the reactivity series, lead is between iron and copper.

Plan an investigation to obtain a sample of lead from a large lump of cerussite.

You are provided with common laboratory apparatus and chemicals. Answer Method 1

- crush (1) with…pestle and mortar reason…to increase the surface area (1)

- heat (1) with carbon (1) because carbon is more reactive than Pb; it reduces PbO

M/J/1996/Q6

25 Acid- base (alkalis) plan

Three aqueous solutions of sodium hydroxide of different concentrations, labeled A, B and C are available. A student used titration method to find order of concentrations of these three solutions of sodium hydroxide Suggest a different method of finding the order of concentrations of the solutions of sodium hydroxide.

Answer Method 1

reagents (1) method (1) result (1) [3] - add 10cm3 of 1mol/dm3 HCl(aq) to 10cm3 of each of A, B and C solutions separately (1) - measure temperature change (1) - the NaOH solution that causes highest temeprature rise is more concentrated (1)

[Total: 3] Answer Method 2

- add named (excess) FeCl2(aq) solution to each sodium hydroxide solution separately (1) - filter the precipitate (1), wash and dry - largest mass = more concentrated solution (1) [3]

Antacids Indigestion tablets Indigestion pain is caused by too much acid in your stomach. The acid is hydrochloric acid. Indigestion tablets contain a base which neutralises the acid. You are provided with two different brands of indigestion tablets, Painremuve and Indcure. Plan an investigation to compare which of these brands of tablet is the most effective. You are provided with dilute hydrochloric acid and common laboratory apparatus.

Answer Using pipette transfer 50 cm3 of dilute hydrochloric acid to100cm3 beaker; then add 3 drops of litmus solution. Start adding Painremuve tablets to the beaker with stirring until the solution turns purple. Stop and count the number of added tablets. Repeat the same procedure with Indcure tablet; using same volume of same acid and same number of drops of litmus solution. Compare; The brand that uses fewer tablets to neutralize the acid is more effective antacid.

[Total: 7] 0620_w13_qp_63_Q6



Tonic Water Tonic water is a solution containing citric acid. The concentration of the acid can be determined by reaction with aqueous potassium hydroxide solution. Plan an investigation to show which of two different brands of colourless tonic water, Tastyton and Slimton, contains the highest concentration of citric acid. You can use common laboratory apparatus and chemicals. S_15_62_Q7

Answer By titration - Using pipette, transfer 25cm3 of Tastyton tonic water to a conical,[1] , - then add 3 drops of phenolphthalein indicator [1] - From the burette add KOH(aq) until the solution changes from colourless to permanent pink. [1] - Stop and note down the added volume. [1] - Repeat the same experiment with Slimton water, using same volume and same KOH(aq) [1] - Record the added volume - Compare the two volumes, - water that needs greater volume of KOH(aq) to neutralize it contains the highest concentration

of citric acid [1] [6]

S_15_62_Q7 There are many possible methods. The most common is titration by either adding tonic water to KOH or KOH to tonic water. However, reagents such as Mg or carbonates would also work Generic marking points can be applied to any method:

- mp1 (fair testing):known or stated volume of tonic water; - mp2 (fair testing): repeat with other sample of tonic water; - mp3 (reagent): add or react with KOH or Mg turnings etc.; - mp4 (method): use of indicator / collect gas etc.; - mp5 (endpoint):until colour changes / until no more gas evolved / for one minute etc.; - mp6 (measurement): volume of KOH added / volume of gas evolved; - mp7 (conclusion): the higher concentration is the one that needs the greater volume of KOH /

26 Indigestion tablets contain calcium carbonate.The tablets work by neutralising the excess of acid in the stomach calcium carbonate + hydrochloric acid carbon dioxide + calcium chloride + water

You are provided with 2 different brands of indigestion tablet, F and G, dilute hydrochloric acid and common laboratory apparatus.

Plan an investigation to find which brand of indigestion tablet is best at neutralising acid. Your answer should Include details of the apparatus to be used and the main practical steps in the Investigation.

Apparatus

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plan of investigation

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0620/06/O/N/99/Q7

Metals reactivity plan 27 Two metals, A and B, each react with dilute sulfuric acid to produce hydrogen.

Plan an investigation to show which metal, A or B, is the more reactive metal. You may include a diagram in your answer. You are provided with:

standard laboratory equipment powdered metals A and B dilute sulfuric acid

Answer - arrange the apparatus as shown below then - pour 50 cm3 of dilute sulphuric acid in a conical flask, using a burette - Weigh accurately 10 grams of metal A powder and add them to the conical flask - quickly put the bung in the falsk and start the timer. - Measure the volume of the gas collected in the measuring cylinder every 10 seconds

up to 100 seconds. - Record your data - Repeat the same experiment using the same volume of the same acid and the same

mass of powdered metal B. - Compare the results - The metal that produces larger volume of hydrogen gas per time is more reactive.

0620_M_J_2012_61_Q4

Metal powder water

Measuring cylinder

collected H2 gas



28 You are provided with samples of three metals, tin, zinc and silver. Plan an investigation to show the order of reactivity of these three metals.

You are provided with common laboratory apparatus and dilute acids.

Answer - add 50cm3 of dilute HCl solution to 10 grams of zinc powder and measure highest temperature rise

- repeat the same experiment with tin and silver; make sure to use the same volume of hydrochloric acid, same mass and same particles size of each metal

- measure any temperature change and calculate temperature rise - compare - no temperature change with silver and highest temperature rise with zinc - therefore, reactivity order: Zinc > tin > silver [6]

Answer - add 50cm3 of dilute HCl solution to 10 grams of zinc powder and measure the

volume of the produced gas at regular time intervals. - repeat the same experiment with tin and silver; make sure to use the same volume of

hydrochloric acid, same mass and same particles size of each metal - compare the volumes of the produced in same time intervals. - No gas is produced with silver and largest volume produce with zinc. - therefore, reactivity order: Zinc > tin > silver [6]

0620/62/O/N/15/Q6

Fertilizers plan 29 The farmer was given bags of three different fertilisers. Describe an experiment he could carry

out to find out which fertiliser would give the best crop of beans. Answer - To identical samples of each soil add same amount of each fertilizer.

- Plant 1 bean’s grain in each sample - Wet with same amount of water weekly – for fair test - Place samples in same place to receive same amount of light for -fair test

The grain which shows most growth has the best fertilizer [5]

30 If two different soils have correct pH and provided with the same fertilizers but produce different quantities of crops, they should have different nutrients (in one condition same type of crops) Plan an experiment to choose the best fertilizer for beans (crops) Answer - To identical samples of each soil add same amount of different fertilizers.

- Plant 1 bean’s grain in each sample - Wet with same amount of water weekly – for fair test - Place samples in same place to receive same amount of light for -fair test - The grain which shows most growth has the best fertilizer



31 The following table gives information about three solids: nickel, wax and magnesium

solid Solubility in ethanol

Attracted by a magnet

Nickel insoluble Yes Wax Soluble No

magnesium insoluble no

(a) Suggest how a mixture of nickel powder and magnesium powder could be separated.

Answer Use a magnet. Nickel will move away with the magnet. Magnesium will not

(b) Suggest how a sample of wax could be obtained from a mixture of wax and magnesium. Answer - Add ethanol to the mixture, warm and stir; Wax will dissolve

- Filter, magnesium remains as residue on the filter paper. - Warm the filtrate in water bath to evaporate ethanol. Wax will remains

Solubility

32 Potassium chloride is a salt that dissolves in water. The solubility of a salt is the mass in grams of the salt that dissolves in 100 cm3 of water at a particular temperature. Plan an investigation to determine the solubility of potassium chloride in water at 40 C. You are provided with potassium chloride and common laboratory apparatus.

Answer - measure 100cm3 of water, using pipette, 4 times / burette, 2 times [ 100cm3 water = 100g] - heat to >40 C - add excess KCl and stir until no more dissolves - then cool to 40 C and filter - evaporate filtrate to dryness - weigh solid, its mass = solubility of KCl at 40C . [6]

0620_s18_qp_61_Q4

Finding the mass of salt in water 33 Seawater contains sodium chloride and other salts.

Plan an experiment to find the mass of salts in 1 dm3 of seawater. You will be provided with a small bottle of seawater. You should include details of the method and any apparatus used. (1 dm3 = 1000 cm3) Answer - Weigh an evaporating dish and note its mass.(1)

- using measuring cylinder(1),transfer a measured volume of seawater 200 cm3(1), - into the evaporating dish (1) - Heat to evaporate (1) - to dryness/constant mass (1) - re-weigh and note the mass(1) - mass of salt in 200 cm3= mass of evaporating dish with residue- mass of empty dish. - mass of salt in 1dm3= mass of salt in 200 cm3 x 5 (1). [6]

0620/61/O/N/2011/Q6

Evaporating the salt to dryness, then measure the mass of the residue



percentage of water by mass in a hydrated salt 34 Washing soda crystals are crystals of hydrated sodium carbonate, Na2CO3.10H2O. When

exposed to the air, some of the water is lost from the crystals and a new substance is formed. This process occurs faster in hotter climates. Plan an experiment to determine the percentage of water by mass present in the new substance. You are provided with common laboratory apparatus.

Answer - Weigh exactly 100 g of Na2CO3.10H2O into a crucible - Heat the crucible and reweigh many times until constant mass is reached. - Note the mass of the residue. - % of water = 100 – weight of residue. [6]

/63/O/N/2017/Q4

To find the number of water of crystallisation

Sources of error improvement the solid absorbs water from the air store in a sealed container / airtight container / desiccator

Finding the mass of salt in water 35 Seawater contains sodium chloride and other salts.

Plan an experiment to find the mass of salts in 1 dm3 of seawater. You will be provided with a small bottle of seawater. You should include details of the method and any apparatus used. (1 dm3 = 1000 cm3)

Answer: - Weigh an evaporating dish and note its mass.(1) - using measuring cylinder(1),transfer a measured volume of seawater 200 cm3(1), - into the evaporating dish (1) - Heat to evaporate (1) - to dryness/constant mass (1) - re-weigh and note the mass(1) - mass of salt in 200 cm3= mass of evaporating dish with residue- mass of empty dish. - mass of salt in 1dm3= mass of salt in 200 cm3 x 5 (1). [6]

0620/61/O/N/2011/Q6 Evaporating the salt to dryness, then measure the mass of the residue

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