[Edu.joshuatly.com] Pahang JUJ SPM 2011 Chemistry

8/10/2019 [Edu.joshuatly.com] Pahang JUJ SPM 2011 Chemistry

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CHAPTER 6: ELECTROCHEMISTRY

A. Electrolytes and Non Electrolytes

1) Electrolytes

Electrolytes are substancesthat can conduct electricity when they are in molten states or aqueoussolutionand undergo chemical changes.

Examples of electrolytes are:

a) dilute acids and alkalisb) molten zinc bromidec) sodium chloride solution

Electrolytes are able to conduct electricity because there are free ions presentin the moltenor aqueous

state.

2) Non-electrolytes

Non electrolytes are substances thatcannotconduct electricity either in molten state or aqueous solution.

Examples of non-electrolytes are:

a) naphthalene,b) benzenec) alcohold) tetrachloromethane

In other words, all non-electrolytesare covalent compounds which do not contain ions and thus, they arenot able to conduct electricity.

Battery bulb is lighted

Diagram 6.1:An electrolyte can conduct electricity because of the presence of free moving ions.

1. What is an electrolyte?Substances that can conduct electricity when they are in molten state oraqueous solution and undergo chemical changes.

2. Explain why a sugar solution does not conduct electricity while a common

salt solution conducts electricity.Sugar solution does not have freely moving ions while salt solution do have freely movingions.

MoltenPbBr2

NaCl


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Experiment: To Investigate the electrolysis of molten lead (II) bromide, PbBr2.

1. Lead (II) bromide, PbBr2is an ionic compound.2. Hence, it consists of the positive lead (II) ions, Pb

2+and

the negative bromide ions,Br-.

3. In solid lead (II) bromide, PbBr2, these ions do not move freely but are held in afixed positions in a lattice. When it melts, the ions are free to move.

Battery mentol switch

Diagram 6.2: Electrolysis of molten lead (II) bromide, PbBr2

4. During themolten lead(II) bromide, PbBr2,

i) the negative bromide ions,Br-are attracted to the anode,

ii) the positive lead(II) ions,Pb2+

are attracted to the cathode.

In the molten state, heat energy that is provided breaks down theelectrostatic forces so that ions are able to move freely. Thus we have:

PbBr2(s) Pb2+ +2Br-

What happens to the cathode and anode?

5. At the anode, two bromine atoms combine to form a bromine gas.The half equation that we have are as below:

2Br-(l) Br2 (g) + 2e-

At the cathode, lead(II) ions, Pb2+

undergo discharge whereby each of the ionsaccepts two electrons to form a lead atom. The half equation are as follows:

Pb

2+

(l) + 2e

-

Pb(s)

7. Combining the two half equations, we get the overall equation that represents the electrolysis ofmolten lead (II) bromide.

Pb2+

(l) + 2Br-(l) Pb (s) + Br2 (g)

The process of discharge can either be donating electrons or accepting electronsBromine gas, Br2is red-dish brown in color and is poisonous.

Anode

Lead(II)ion, Pb

2+

Cathode

molten lead(II) bromide,PbBr2


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Test :

1. State the meaning of the following terms.

a) anode

An electrode which is connected to the positive terminal of an electric source.

b) cathode

An electrode which is connected to the negative terminal of an electric source.

c) electrolysis

Aprocess whereby compounds in either molten or aqueous state are brokendown into their constituent by passing electricity through them.

2. A molten oxide, Q2O3 is electrolysed using carbon electrodes.

a) Draw a labeled diagram to show the set-up of apparatus for the electrolysis.

Battery mentol switch

b) Write the formulae for the ions present in the electrolyte.

Formulae ions are Q3+

and O2-

c) State the ions that move to each of the electrodes during the electrolysis.

i) the negative ions, O2-

are attracted to the anode,ii) the positive ions, Q

3+are attracted to the cathode.

d) Write the half equation of the reaction at each of the electrodes.

Half equations: anode terminal: O2-

O2 + 2e-

Half equations: cathode terminal: Q3+

+ 3e- 3Q

d) Named the substances formed at each of the electrodes.

Positive electrodes: an atom of substance Q is performed.Negative electrodes: an oxygen gas is released.

Anode

Carbonelectrodes

Cathode

molten oxide,Q2O3


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CHAPTER 6 :ELECTROCHEMISTRY

Voltaic Cell

Example of simple voltaic cell

1. Voltaic cell also known as galvanic cell.2. Denial cell is another example of a voltaiccell.3. In Denial cell, the two solutions are connected through a salt bridgeor a

porous pot.4. The solution used in the salt bridge is usually a diluted acid or a solution of

sodiumor potassiumsalts.5. The function of the salt bridge or porous pot is to stop the two electrolytes from mixing.

But allow the movementof the ions in order to complete the electric circuit.

Electrolytic cellElectrical energy is supplied to produced

chemical reaction

Chemical cell Chemical reaction produced the electricalenergy


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How does a voltaic cell work?

The electrochemical series

Li

Ca

MgAlZn

SnPb

CuHgAg

Increasing

tendency

of metals

to lose

electron

K

Na

Fe

H

Reaction in the voltaic cell using magnesium ribbon and copper plate as electrodes


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The apparatus set-up of a Daniell Cell

using a salt bridgeThe apparatus set-up of a Daniell Cell

using a porous pot

Zinc is higher than copper in the electrochemical series. Thus it releases electronsmoreeasily than copper. Each zinc atom donatestwo electrons to form a zinc ion.The zinc ions

are released into the electrolyte. The zinc platedissolves gradually.Since the Zinc is higher than copper in the electrochemical series. It acts as the cathode ofthe cell(negative terminal).The ionic equation of cathode:Zn(s)Zn2+ (aq) + 2eCopper as the anodeof the cell (positive terminal).The electrons are acceptedby thecopper(II) ions to form copper metal. The copper plate becomes thickergradually.The ionic equation of anode: :Cu2+ (aq) + 2eCu(s)The intensity of the blue colour of the copper(II) sulphate solution decreasesas theconcentration of copper(II) ions decreases gradually.

Decreases donates ion electrons zinc plate anode cathode thicker accepted


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Cells Used in Advantages Disadvantages

Torches

Radio

Electrical toys

Cassetteplayers

Nospillage

Small in size

Easily carried about

Produces regular current andvoltage

Obtained in different sizes

Cannot be recharged

Not long lasting

Electrolyte tends to leak

Low voltage is produced

Motor vehicles Can be recharged

Produces a high voltage and usefor a long period of time

Produces high current(up to 75A)for heavy duty purposes

Big in size

Expensive

Leak charges when not in used

Spillage of acid can occurr

Heavy and difficult to carriedabout

Cassette players

Electrical toys

Appliances which are inuse for long periods

Last longer than dry cell

Produces constant current

No liquid electrolyte

Expensive

Cannot be recharged

Electrolyte leakage can occur

http://edu.joshuatly.com


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Cells Used in Advantages Disadvantages

Cameras

Watches

Calculators

Small in size

Last a longer time

Very mobile and easily used

Very expensive

Cannot be recharged

Mercury is not environmentfriendly

Electrical toys

Cassette players

Radios

Can be recharged

Smaller size than accumulator

No spillage

Can be recharged up to 500times

Expensive

Transformer is needed forrecharging



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Differences between electrolytic and voltaic cell

Characteristic Electrolytic cells Chemical cells

EnergyElectricalenergy produces chemical

reactions

Chemicalreactions produce electricalenergy.

Current Current supplyinto the cell. Current producedby the cell

Anode/CathodeCathode: terminal negative

Anode : terminal positive

Cathode: terminal positive

Anode : terminalnegative

Positiveterminal

Anion gives out electrons to anode Electronsreceive by positive terminal.

NegativeTerminal

Cationaccept electron from negativeterminal.

Electronsgiven out from negative terminal.

Type ofelectrode

Pairs of graphite, platinumor suitablemetals.

Pair of different metals.



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CHAPTER 8 : SALTS

8.1 Definitions of salts

A salts is a compound formed when the hydrogen ion, H+from an acid is replaced by a metal ion

or an ammonium ion, NH4+

Example;

8.2 Types of salts

Soluble saltsalts that can be dissolve in water at room temperature

Insoluble saltsalts cannot be dissolve in water at room temperature

8.3 Summarised of the solubility of salts

All sodium , potassium and ammonium salts are soluble in water

All nitrate salte are soluble in water

All chloride salts are soluble in water except :

- silver chloride (AgCl)

- mercury (I) chloride (HgCl)

- Lead (II) chloride (PbCl2) - soluble in hot water but insoluble in cool water



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All sulphate salts are soluble in water except

- Lead (II) sulphate ( PbSO4)

- Barium Sulphate (BaSO4)

- Calcium Sulphate (CaSO4)

All carbonate salts are insoluble except

- Sodium carbonate (Na2CO3)

- Potassium Carbonate (K2CO3)

Ammonium Carbonate (NH4)2CO3



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Exercise 1

a) Complete the table with the chemical formulae of saltsb) Marked the insoluble salts from the table.

Metal ionSulphate salts

H2SO4Nitrate salts

HNO3

Carbonate

SaltsH2CO3

Chloride saltsHCl

Iodide saltsHI

K+ K2SO4 KNO3 K2CO3 KCl KI

Na+ Na2 SO4 NaNO3 Na2CO3 NaCl NaI

Ca2+

CaSO4 Ca(NO3)2 CaCO3 CaCl2 CaI2

Mg2+

MgSO4 Mg(NO3)2 MgCO3 MgCl2 MgI2

Zn2+

ZnSO4 Zn(NO3)2 ZnCO3 ZnCl2 ZnI2

Pb2+

PbSO4 Pb(NO3)2 PbCO3 PbCl2 PbI2

NH4+ (NH4)2SO4 NH4NO3 (NH4)2CO3 NH4Cl NH4I

Ba2+

BaSO4 Ba(NO3)2 BaCO3 BaCl2 BaI2

Hg+ Hg2SO4 HgNO3 Hg2CO3 HgCl HgI

Ag+ Ag2SO4 AgNO3 Ag2CO3 AgCl AgI

8.4 Preparation and purification of soluble salts

- Sodium salts

- Potassium salts Acid + alkali salts + water

- Ammonium salts

Soluble Salts

Acid + metal oxide salts + water

- Others salts Acid + metal salts + hydrogen gas

Acid + metal carbonate salt + water + carbon dioxide

The salts formed during preparation of soluble salts contain impurities. Therefore, these salts need to be

purified through a process known as recrytallisation.



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Exercise 2

Fill in the blank with the suitable reactants

1. Nitic acid + sodium hydroxide sodium nitrate + water

2. Hydrochloric acid + magnesium oxide magnesium chloride + water

3. sulphuric acid + zinc zinc sulphate + hydrogen gas

4. hydrochloric acid + calcium carbonate calcium chloride + water + carbon dioxide

8.5 Preparation insoluble salts

Prepared by precipitation method / double decomposition reaction.

8.6 Chemical and ionic equations

In the formation of the precipitate of barium sulphate, BaSO4, the chemical equation can be written:

Example : BaCl2(aq) + Na2SO4 (aq) BaSO4(s) + 2NaCl (aq)

Ions Ba2+

+ Cl- + Na

+ + SO4

2-BaSO4 + Na

+ + Cl

-

Ionic equation : Ba2+

+ SO42-

BaSO4

(shows the ions that take part in the reaction)

Exercise 3

Complete the table

Insoluble Salt Ions Ionic equation

ZnCO3 Zn2+

, CO32-

Zn2+

+ CO32-

ZnCO3

AgCl Ag+ , Cl

- Ag

++ Cl

- AgCl

BaSO4 Ba2+

, SO42-

Ba2+

+ SO42 BaSO4

PbCl2 Pb2+

, Cl- Pb

2+ + Cl

- PbCl2

PbSO4 Pb2+

, SO42-

Pb2+

+ SO42-

PbSO4

CaCO3 Ca2+

, CO32-

Ca2+

+ CO32-

CaCO3



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8.6 Numerical problem involving stoichiometric reaction in the preparation of salt

A balanced chemical equation for a reaction in preparation of a salt can be used to calculate the

stoichiometric quantities of the following

Masses of reactants

Volumes and concentrations of reactants

Masses of products

Volumes of products

Example;

Ammonium phosphate, (NH4)3PO4 is use as a fertilizer. 29.8g of this salt is prepared by neutralizing

phosphoric acid, H3PO4with ammonium gas, NH3. Calculate the volume of ammonium gas, NH3reacted at

room conditions.

[Relative atomic mass; H=1: N =14: P=31; O= 16; Molar volume; 24 dm3mol

-1at room conditions]

Solutions;

a. Calculate the number of moles

2.88 g

[3(14) + 12(1) + 31 + 4(16)

= 0.2 mol

b. Write a balanced chemical equation

Compare the mole ratio of NH3and(NH4)3PO4

H3PO4(aq) + 3NH3(aq) (NH4)3PO4(aq)

c. Calculate the number of moles of NH3base onthe mole ratio

= 3 X 0.2 mol

= 0.6 mol

d. Calculate the volume of NH3

Volume = number of mole X volume

= 0.6 mol X 24 dm3mol

-1

= 14.4 dm3

=

3 mol 1 mol



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B. Qualitative Analysis of Salts

Inthe qualitative analysis of salts, we need to identify the ions that are present in salts. This canbe done by analysing their physical and chemical properties.

Observations on the physical properties of salts

1. Colour and solubility in water

Certain physical properties of salts such colour and solubitity in water are observed to help us infercertain cations and anions that are present in salts.

The table shows the colour of salts in solid , in aqueous solution and the solubility ofsalts in water

SaltColour in

solidSolubility in

waterColour

in Aqueous solution

Ammonium chloride, NH4Cl white soluble colourless

Ammonium nitrate,NH4(NO3)3 white soluble colourless

Calcium carbonate, CaCO3 white insoluble -

Calcium nitrate ,Ca(NO3)2 white soluble colourless

Magnesium sulphate, MgSO4 white soluble colourless

Magnesium carbonate, MgCO3 white insoluble -Zinc sulphate, Zn SO4 white soluble colourless

Zinc nitrate ,Zn(NO3)2 white soluble colourless

Lead(II) chloride , PbCl2 white insoluble -

Lead(II) sulphate , PbSO4 white insoluble -

Lead(II) carbonate , PbCO3 white insoluble -

Copper(II) chloride , CuCl2 Blue soluble Blue

Copper(II) sulphate , PbSO4 Blue soluble Blue

Copper(II) carbonate , PbCO3 Green insoluble -

Iron(II) sulphate , FeSO4 Green soluble Pale green

Iron(III) chloride , FeCl3 Brown / Yellow soluble Brown/Yellow/ Yellowish brown

Sodium nitrate , NaNO3 white soluble colourless

Sodium carbonate , Na2CO3 white soluble colourless

Potassium nitrate , KNO3 white soluble colourlessPotassium carbonate , K2CO3 white soluble colourless

What is Qualitative analysis?

Qualitative analysis is a chemical technique used to determine whatsubstances are present in a mixture but not their quantities.



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The table shows the colour of different cations in the solid form or in aqueous solution

Observation Inference

Blue solution Ion copper (Cu+) present

green solution Ion Iron(II) Fe

+

presentYellow/Yellowish-brown/brown solution

Ion Iron (III) Fe3+

present

Green solid Hydrated Fe+, CuCO3

Brown solid Hydrated Fe+salt

White solidSalts of Na

+, K

+,NH4

+, Mg

+, Ca

+Al

+, Zn

+, Pb

+

(If the anions are colourless)

Colourless solution Na+, K

+,NH4

+, Mg

+, Ca

+, Al

+, Zn

+, Pb

+

The table shows the solubility of different types of salts in water

Compounds Solubility in water

Sodium salts

Potassium saltsAmmonium salts All are soluble

Nitrate salts All are soluble

Chloride salts All are soluble except AgCl, HgCl and PbCl2(soluble inhot water)

Sulphate salts All are soluble except BaSO4, PbSO4and CaSO4

Carbonate salts All are insoluble except sodium carbonate, potassiumcarbonate and ammonium carbonate

2. Tests for gases

Gases are often produced from reactions carried out during laboratory tests on salts. By identifying the

gasesEvolved , it is possible to infer the types of cations and anions that are present in a salt.

The table shows the test and the result of different gases

Gas Method Observation

Oxygen gas,O2

Put in glowing wooden splinterInto the test tube

Wooden splinter is rekindled /lighted

Hydrogen gas ,H2

Put in lighted wooden splinterInto the test tube

popsound produced

Carbon dioxide gas ,CO2

Bubble up the gas throughlime water

Lime water turns milky/cloudy/chalky

Ammonia gas, NH3 Put in moist red litmus paperInto the test tube

Moist red litmus paper turns blue

Chlorine gas, Cl2 Put in moist blue litmus paperInto the test tube

Moist blue litmus paper turns red andthen turns white/bleaches

Hydrogen chlorinegas , HCl

drops of concentrated ammoniaNH3solution into the test tube

Dense white fumes produced

Sulphur dioxide gas ,SO2

Bubble up the gasinto purple acidified potassiummanganate (VII), KMnO4solution

Purple acidified potassiummanganate (VII),KMnO4solutiondecolourises

Nitrogen dioxide gas ,NO2

Test with moist blue litmus paper moist blue litmus paper turns red



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3. Action of heat on salts

Effect of heat on carbonate salts

Carbonate saltColour of saltbefore heating

Colour of residueEffect on lime waterHot cold

Copper (II) carbonate,CuCO3

Green powder Black powderBlack

powderThe gas liberated turns

lime water milky

Zinc carbonate ,ZnCO3

White solid Yelow solid White solidThe gas liberated turns

lime water milky

Lead(II) carbonate,PbCO3

White solid Brown sold Yelow solidThe gas liberated turns

lime water milky

Sodium carbonate,Na2CO3

White solid White solid White solid No change

Calcium carbonate,CaCO3

White solid White solid White solidThe gas liberated turns

lime water milky

Potassium carbonate,K

2CO

3

White solid White solid White solid No change

Magnesiumcarbonate, MgCO3

White solid White solid White solidThe gas liberated turns

lime water milky

Effect of heat on nitrate salts

Nitrate SaltColourof salt

Colour of residueTest on gases liberated

Hot cold

Copper (II) nitrate,Cu(NO3)2

Bluesolid

Blackpowder

Blackpowder

A brown gas that turns blue litmus paper tored is liberated. The gas liberated alsoignites a glowing splinter

Zinc nitrate,Zn(NO3)2

Whitesolid

Yellowsolid

Whitesolid


Lead(II) nitrate,Pb(NO3)2

Whitesolid

Brownsolid

Yellowsolid


Sodium nitrate,NaNO3

Whitesolid

Whitesolid

Whitesolid

A colorless gas that rekindles a glowingsplinter is liberated

Calcium nitrate,Ca(NO3)2

Whitesolid

Whitesolid

Whitesolid


Potassium nitrate,KNO3

Whitesolid

Whitesolid

Whitesolid

A colorless gas that rekindles a glowingsplinter is liberated

Magnesium nitrate,Mg(NO3)2

Whitesolid

Whitesolid

Whitesolid


Iron(II) nitrate,Fe(NO3)2

PaleGreensolid

PaleGreensolid

PaleGreensolid


Iron(III) nitrate,Fe(NO3)3

Brownsolid

Reddish-Brownsolid

Reddish-Brownsolid

A brown gas that turns blue litmus paperred is liberated. The gas liberated alsoignites a glowing splinter



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The table shows the comparison of the effect of heat on carbonate and nitrate salts

Metal Effect of heat on carbonate salt Effect of heat on nitrate salt

PotassiumSodium Are not decomposed by heat Decompose to nitrite salt and oxygen gas.

CalciumMagnesiumAluminiumZincIronTinLeadCopper

Decompose to metal oxide and carbondioxide gas.

Decompose to metal oxide, nitrogen dioxidegas and oxygen gas.

MercurySilverGold

Decompose to metal, carbon dioxidegas and oxygen gas.

Decompose to metal , nitrogen dioxide gasand oxygen gas.

Most sulphate salts are not decomposed by heat. Only a few sulphate such as iron (II) sulphate, zincsulphate and copper sulphate decompose to sulphur dioxide or sulphur trioxide gas when heated.

All chloride salts are stable when heated except ammonium chloride. Ammonium chloride sublimes anddecomposes to produce ammonia gas and hydrogen chloride gas.

The table shows the deduction of the types of ion present based on the gas produced

Type of gas produced Type of ion present(anion)

CO2 Carbonate ion (CO3-) present except Na2CO3and K2CO3

O2 Nitrate ion (NO3-) presentNO2 Nitrate ion (NO3

-) present except NaNO3and KNO3

SO2 Sulphate ion (SO4-) present

NH3 Ammonim ion (NH4+) present

Exersice:

1. State three examples ofa) soluble salts b) insoluble salts

Potassium carbonate Magnesium carbonateLead(II) nitrate Lead(II) sulphateAmmonium chloride Argentum chloride

2. Which of the following salts is soluble

Lead(II) chloride Sodium carbonate

Calcium sulphate Barium sulphate



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3. Identify the gas that turns moist red litmus paper blue

Ammonia gas

4. Gas X has the following properties

Gas X is carbon dioxide gas

5. Heat

Colour of metal oxide X is yellow when hot and white when cold. Gas Y turns lime water milky.

a) Name gas Y carbon dioxide gasb) Name metal oxide X zinc oxidec) Name salt P zinc carbonated) Write an equation to represent the action of heat on salt P

ZnCO3 (s) ZnO (s) + CO2(g)

6. A sample of copper(II) nitrate, Cu(NO3)2was heated strongly. Write down the expected

observation.

Copper(II) nitrate decompose to produce black colour of residue when hot and cold. A brown gasthat changed moist blue litmus paper to red and colourless gas that lighted up a glowing woodensplinter are produced.

Quantitative Analysis of Salts

Tests for anions

Reagent / Condition Observation AnionIonic Equation

(if any)

Add dilute hydrochloric acid /nitric acid / sulphuric acid into atest tube gas liberated isimmediately bubbled throughlime water.

Effervescence.Colourless gasturns lime water

milky.

CO32-

ionCO3

2-+ 2H

+

CO2 + H2O

Colourless

Acidic gas

Turns lime water milky

Salt P Metal oxide X Gas Y+



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Reagent / Condition Observation AnionIonic Equation

(if any)

add nitric acid and silvernitrate solution

White precipitateformed.

Cl-ion Ag

+ + Cl

- AgCl

Add dilute hydrochloric acid /nitric acid and barium chloride/ barium nitrate solution

White precipitateformed.

SO42-

ionBa

2+ + SO4

2

BaSO4

Add dilute sulphuric acid andass iron(II) sulphate solution.Then drop carefully and slowlyconcentrated sulphuric acid

Brown ringformed

NO3-ion

-

Confirmatory Test for Fe2+

, Fe3+

, Pb2+

, NH4+

IonsConfirmatory Test for Fe

2+and Fe

3+Reagent Observation Conclusion

Add Sodium hydroxide solutionuntil excess

Green precipitate formed Fe 2+ ion present

Brown precipitate formed Fe+

ion is present

Add Sodium hydroxide solutionuntil excess

Green precipitate formed Fe 2+ ion present

Brown precipitate formed Fe+

ion is present

Potassium hexacyanoferrate(II) solutionPale blue precipitate Fe

+ion present

Dark blue precipitate Fe+

ion is present

Potassium hexacyanoferrate(III) solutionDark blue precipitate Fe

+ion is present

Greenish-brown solution Fe+

ion is present

Potassium thiocyanate solutionPale red colouration Fe

+ion is present

Blood red colouration Fe+

ion is present

Confirmatory Test for Pb2+

Method Observation Ionic Equation

Using aqueous solution of chloride- 2 cm

3of any solution of Cl

-+

2 cm3of any solution of Pb

2+

dilute with 5 cm3of distilled water

heat until no further change occursallow the content to cool to roomtemperature using running water fromthe tap

A white precipitateis formed

When heateddissolve in water toform colourless solution

When cooledwhiteprecipitate reappear

Pb2+

+ 2Cl- PbCl2

Using aqueous solution of iodide- 2 cm

3of any solution of I

-+

2 cm3of any solution of Pb

2+

dilute with 5 cm3of distilled water

heat until no further change occursallow the content to cool to roomtemperature using running water fromthe tap

A yellow precipitateis formed

When heateddissolve in water toform colourless solution

When cooledyellowprecipitate reappear

Pb2+

+ 2I- PbI2



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(e) sodium chloride and sodium sulphate

Silver nitrate solution is poured into the test tube.If a white precipitate is formed, the solution is sodium chloride.If no change occurs, the solution is sodium sulphate.

Or

Barium chloride solution is poured into the test tube.If a white precipitate is formed, the solution is sodium sulphate.If no change occurs, the solution is sodium chloride.



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CATION SOLUTION

PrecipitateColourless/

Unchan ed

ColouredWhite

Green

Brown

Blue

Add / put inNaOH SOLUTION

Pb

+

Zn

+

Al

+

Mg+ Ca

+

Fe+

Fe+

Cu+

NH4+

Dissolved/soluble in excess

NaOH solution

Undissolved / not soluble

In excess NaOH solution

To add/ put in

EXCESSNaOH

Pb2+

Zn2+

Al3+

Mg2+

Ca2+

add KIsolution

Yellow

Precipitate

[PbI]

White

Precipitate

[PbSO4 ]

add ion SO42-

solution

Pb2+

Zn2+

Al3+

To add/ put

in

White

Precipitate

colourless

Mg2+

Ca2+

add / put in

NH3 solution

Colourless

Using Reagent : Sodium hydroxide,NaOH solution

colourless White precipitate

Zn2+

Al3+



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CHAPTER 8 : SALTS

Exercises:

1. A salt is

l an ionic compoundll formed from an acid and baseslll an acidic compoundlV an alkaline compound

A l and ll onlyB. ll and lVC. l, ll dan lll onlyD. l, lll dan lV

2. A salt is formed when the hydrogen ion, H+ of an acid is replaced by

l a metal ion,ll an ammonium ionlll an anionlV a hydroxide ion

A l and ll onlyB. ll and lVC. l, ll dan lll only

D. l, lll dan lV

3. Which of the following salt are soluble in water

l sodium nitratell potassium carbonatelll silver chloridelV ammonium carbonate

A l and ll onlyB. l and lll onlyC. l, ll dan lV onlyD. l, lll dan lV

4. Which of the following salts is insoluble in water?

A. lead nitrateB. copper nitrsteC. barium sulphateD. calcium chloride



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5. Insoluble salt can be prepared through

A. CrystallisationB. RecrystallisationC. Precipitation reactionD. Neutralization reaction

6. Calcium carbonate is an insoluble salt. It can be prepared by using

A. calcium oxide and sodium chlorideB. calcium chloride and sodium hydroxideC. calcium chloride and sodium carbonateD. calcium hydroxide and sodium carbonate

7. Calculate the volume of 0.2mol/dm3 sulphuric acid , H2SO4, thats is needed to reactcompletely with 10cm3of 0.5mol/dm3of sodium hydroxide solution

A. 5 cm3B. 10cm3C. 12.5cm3D. 15cm3

8. Nitric acid react with metal X to produce a salt according to the chemical equation

X + 2HNO3 X(HNO3)2+ H2

Calculate the volume of 1,5 mol/dm3nitric acid that is needed to completely react with0.54g of metal X

(Relative atomic mass: X=24)

A. 25cm3B. 30cm3C. 35cm3D. 40cm3

9. If 20cm3of a metallic hydroxide, 0.5 mol dm-3reacts completely with 20cm3of hydrochloricacid, 1 mol dm -3, what is the formula of the chloride of metal M which results?

A. MClB. MCl2

C. MCl3D. M2Cl3



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10. A white salt Y is heated strongly. It was found that a brown gas, which turned damp bluelitmus paper red, and a colourless gas which rekindled a glowing splint, were producedThe residue left behind was brown when hot and yellow when cold. The salt Y is mostprobably.

A. zinc nitrateB. silver carbonateC. copper(ll) carbonate.D. lead(ll) nitrate

1. a) What is the meaning by salt?

....................................................................................................................

b) State examples of salts used in daily life.

......................................................................................................................

c) Is sodium chloride a soluble salt ?

......................................................................................................................

2. Suggest the materials which can be used to prepared soluble salt, magnesium saltrespectively in the following table using the method for preparation of soluble salts

Methods for preparation of soluble salt

1. Acid + alkali salt + water2. Acid + metal oxide (base) salt + water3. Acid + metal salt + hydrogen gas4. Acid + metal carbonate salt + carbon dioxide + water

Materials Salt

Reaction of sodium hydroxide solution and hydrochloric acid forms a type ofsalt, sodium chloride.



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barium hydroxide, lead(ll) iodide, magnesium sulphate,potassium carbonate, copper(ll) sulphate, silver chloride,

calcium nitrate, barium sulphate

3. Classified the above salts into soluble and insoluble salts

Soluble salts Insoluble salts

4. Table below shows the reaction between acid and alkali and produced salt solution.complete the table below

Acid Alkali Salt solution Chemical equation

Hydrochloric acid(HCl)

Potassium hydroxide(KOH)

HCl + KOHKCl + H2O

Nitric acid (HNO3).

Sodium nitrate(NaNO3)

HNO3+ NaOH NaNO3+ H2O

Aqueous ammonia(NH3)

Ammonium sulphate(NH4)2SO4

.. .. ...

HCl + NH3 NH4Cl + H2O



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5. Diagram below shows a flow on the preparation of insoluble salt.

BaSO4(Ba+2, SO4

-2)

Solution A with positive + Solution B with negativeionof the salt (soluble salt) ionof the salt (soluble salt)

Ba(NO3)2 ( Na2SO4)

precipitate C solution D(insoluble salt with ( soluble salt)

Positive & negative ions)BaSO4 NaNO3

a) Name the reaction of preparation of insoluble salt

..

b) List three examples of insoluble salt

c) Based on the information given in table above, complete the table below

Solution A withpositive ion

Solution B withnegative ion

Precipitate C,insoluble salt

Solution D

Silver sulphate,Ag2SO4

Sodiumchloride, NaCl

Silver chloride,AgCl

Sodiumsulphate,Na2SO4

Barium nitrate,Ba(NO3)2

Bariumsulphate,Ba2SO4

Lead(ll) nitrate,Pb(NO3)2

Lead(ll)carbonate,

PbCO3



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6. 20cm3of 0.5 mol dm-3 silver nitrate is added to a beaker containing 0.56g R.(Relative atomic mass; R = 56)

a) How many moles of silver nitrate have reacted?

b) How many mole of R have reacted?

c) Determine the number of mole of silver nitrate to number of moles R insimplest terms.



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Heat

7. Diagram 1.1 shows the method of preparing a soluble salt.

Diagram 1.1

(a) State one step that is required to speed up the reaction betwen metal oxideand acid

.....................................................................................................................

(b) Based on Diagram 1.1, state two substances that are used to prepare zincchloride salt

(i) ...............................................................................................

(ii) ................................................................................................

Salt solution

Salt crystals are obtained

The solution is allowed to cool

Metal oxide powder

50 cm of 2.0 mol dm-

acidSalt solution

Excess metal oxide



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(c) Complete Diagram 1.1 by drawing the set up of apparatus used to obtain the saltcrystals in the space provided.

(d) Name the type of reaction for preparing the salt using this method.

............................................................................................................

(e) The following solutions are used in the reaction:

Solution X : Sodium sulphate solutionSolution Y : Lead (II) nitrate solution

Complete the ionic equation below for the formation of lead(II) sulphate salt, PbSO4

.................................................... PbSO4(s)

(f) The following reaction can be used to prepare copper(II) chloride salt.

CuCO3 + 2HCl CuCl2 + H2O + CO2

Excess copper(II) carbonate is added to react with 50 cm3of 2.0 mol dm-3hydrochloric acid to form the salt.

Calculate the mass of the salt formed.

[Relative formula mass of the salt formed is 135]

[3marks



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Essay questions:

(a) Insoluble salt can be prepared by double decomposition method.Explain with an example the meaning of double decomposition.

(b) (i) Explain an experiment on how could you prepare a dry sampleof copper(II) nitrate salt in the school laboratory.

(ii) You are given a sample of copper(ll) nitrate salt.Describe how you would test for the presence of copper ionsand nitrate ions in the sample

PAPER 3

Structure Question

5.0 cm3of 1.0 moldm-3potassium sulphate solution was filled in eight test tubes of equal size and height.Then 1.0 cm3of 1.0 moldm-3 barium chloride solution is added to the first tube, followed by 2.0 cm3, 3.0cm3, 4.0 cm3, 5.0 cm3, 6.0 cm3, 7.0 cm3and 8.0 cm3 respectively in the other test tubes. The test tube isput aside. When the precipitate salt settled, its height was measured and recorded.

Test tube 1 2 3 4 5 6 7 8

Volume of potassium sulphate solution (cm3) 5.0 5.0 5.0 5.0 5.0 5.0 5.0 5.0

Volume of barium chloride solution (cm3)1.0 2.0 3.0 4.0 5.0 6.0 7.0 8.0

Height of the precipitate (cm) 0.4 0.8 1.3 1.8 2.1 2.1 2.1 2.1

a) State the hypothesis of the experiment.

....................................................................................................................

b) Name the precipitate form in the test tube

......................................................................................................................

c) State the variables for this experiment.

i) manipulated variable

...........................................

ii) responding variable

...............................................



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iii) constant variable (pembolehubah yang dimalarkan)

..................................................

d) Plot the graph of the height of the precipitate against the volume of bariumchloride solution.



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e) From the graph drawn in (d), determine the minimum volume of barium chloridesolution required to react completely with 5 cm3 of 1.0 mol/dm3 potassiumsulphate solution.

...

f) Calculate the number of moles of sulphate ions that react with 1 mole ofbarium ions.

g) Write the ionic equation to represent this reaction.



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CHAPTER 8: SALTS

Answer:

Exersices:

1. A2. A3. C4. C5. C6. C7. C8. B9. B10. D

Subjective Question

1 a) A salt is an ionic compound formed when the hydrogen ion from acid isreplaced by a metal ion or an ammonium ion from an alkali/bes.

b) Flavour to food,plaster of paris in medical feild.Barium sulphate to identify tumor in the intestines

c) soluble salt

2

3. soluble salts insoluble saltsmagnesium sulphate lead(ll) iodidecopper(ll) sulphate, silver chloridepotassium carbonate, barium hydroxidecalcium nitrate barium sulphate

4.

Materials Salt

a). Magnesium + hydrochloric acid Magnesium chlorideb). Magnesium + sulphuric acid. Magnesium sulphate

c) magnesium + nitric acid Magnesium nitrate

Acid Alkali Salt solution Chemical equation

Potassium chloride(KCl)

Sodium hydroxide(NaOH)

Sulphuric acid(H2SO4)

H2SO4+ NH3(NH4)2SO4 + H2O

Hydrochloric acid(HCl)

Ammonium(NH4

+)Ammonium chloride

(NH4Cl)



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5. a) Double decomposition/precipitation reactionb) 1. lead(ll) iodide

2. silver chloride3. magnesium carbonate

c)

Solution A withpositive ion

Solution B withnegative ion

Precipitate C,insoluble salt

Solution D

Potassiumsulphate,

K2SO4

Ammoniumcarbonate,(NH4)2CO3

Ammoniumnitrate, NH4NO3

6. a) No. of mol of silver nitrate = MV1000

=0.5(20) = 0.01 mol1000

b) No. of mol of R = mass = 0.56 = 0.01 mol RRAM 56

c) 1: 1

7. (a) Heat the acid / Stir the mixture

(b) i. zinc powderii. hydrochloric acid

(c )

(d) Precipitation method / double decomposition reaction

(e) Pb2+ + SO42-

PbSO4

Filter paper

Salt crystals



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(f) Number of mole hydrochloric acid = 2 x 50 = 0.1 mol1000

Number of mole of copper(I) chloride = 0.1/2= 0.05 mol

Mass of copper(II) chloride = 0.05 x 135 = 6.75 g

Essay Questions

(a) Example:-

Na2SO4(aq) + Pb(NO3)2 PbSO4(s) + 2NaNO3(aq)

Double decomposition is the reaction of exchange of the ion thattake place between lead(II) ion and sulphate ion producing lead(II) sulphate

(b)(i)1. Pour nitric acid into a beaker.2. Add copper(II) oxide /copper(II) carbonate powder until excess.3. Stir the mixture4. Filter the solution mixture5. Pour the filtrate into an evaporating dish6. heat the solution until it becomes saturated.7. Allow the solution to cool

8. Filter out the crystals.9. Dry the crystals by pressing between a few pieces of filter papers10. CuO + 2HNO3 Cu(NO3)2 + H2O

(b)(ii) 1. Solid of copper(II) nitrate is dissolved in water (and producing blue solution).

Test for Cu2+

2. Add sodium hydroxide solution into a test tube3. Blue precipitate is formed

Test for NO3-

4. Add dilute sulphuric acid into a test tube5. Add iron(II) sulphate solution6. Slowly add concentrated sulphuric acid7. Brown ring is formed.



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EXERCISE 2State which of the following equation is redox reaction or not. Explain why.

a. NaOH + HCl NaCl + H2 O

b. 2 Na + Cl2 2 NaCl

c. CaCO3 + 2 HCl CaCl 2+ H2 O + CO2

d. Zn + CuO ZnO + Cu

e. Pb + 2 AgNO3 Pb(NO3)2 + 2 Ag

f. BaCl2 + Na 2 SO4 BaSO4 + 2 NaCl

g. SO 2 + 2 H 2 S 3 S + 2 H 2 O

h. Cl 2 + 2 FeCl2 2 FeCl3

i. 2 Na + 2 H2 O 2 NaOH + H2

j. Br2 + 2 KI 2 KBr + I2

k. Na2 CO3 + ZnCl2 2 NaCl + ZnCO3

Answer:

(a) Not a redox reaction. No changes in oxidation number of reactant(b) Redox reaction. Sodium undergoes oxidation , chlorine undergoes reduction(c ) Not a redox reaction. No changes in oxidation number of reactant(d) Redox. Zinc is oxidezed and copper ion/ CuO is reduced(e) Not a redox reaction. No changes in oxidation number of reactant(f) Not a redox reaction. No changes in oxidation number of reactant

Answer:

(i) Mn =+7 (ii) Mn = +4 (iii) C =+4 (iv) C = +4(v) C = +2 (vi) N = -3 (vii) N = +5 (viii) N = +4(ix) N = -3



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(g) Redox, SO2is reduced, H2S is oxidized(h) Redox, Fe

+2is oxidized, Chlorine is reduced

(i) Redox, sodium is oxidized, H2O is reduced(j) redox , KI/ Iodide ion is oxidized, bromine is reduced(k) Not a redox reaction. No changes in oxidation number of reactant

1.3 Writing Equat ionsFor Redox Reactions

Equations for redox reactions are :

(i) Chemical Equat ionfor the reaction.

(ii) Half equatio nfor oxidation (lossof electron/ increasein oxidation number).

Half equatio nfor reduction (gainin electron/ decreasein oxidation number).

(iii) Overal l Ion ic equationfor redox reaction formed by combin inghalf equation foroxidation and half equation for reduction (the number of electronsin both the half equationsmust cancel each other).

Example1 :

Reaction Aluminium andcopper(II) sulphate

Chemical Equat ionfor the reaction 2Al + 3CuSO4 Al2(SO4)3 + 3Cu

Half equatio nfor oxidation Al Al3+ + 3e

Half equatio nfor reduction Cu2+

+ 2e Cu

Changing of the coefficient of the halfequation of oxidation*

2Al 2Al3+ + 6e

Changing of the coefficient of the halfequation of reduction*

3Cu2+ + 6e 3Cu

Ionic equation 2Al + 3Cu2+ 2Al3+ + 3Cu

*Make sure that the number of electrons released in half equation for

oxidation are equalto the number of electrons received in half equationfor reduction.

Al Cu2+ SO42-

CuSO42-

Al3+



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EXERCISE 3

1 Sodium metal reacts with water

Reaction Sodium andwater

Chemical Equationfor the reaction

Half equatio nfor oxidation

Half equatio nfor reduction


Ionic equation

Answer:

Reaction Sodium andwater

Chemical Equationfor the reaction 2Na + 2H2O 2NaOH + H2

Half equatio nfor oxidation Na Na+ + e

Half equatio nfor reduction 2H+

+ 2e H2


2Na 2Na+ + 2e

2H+ + 2e H2

Ionic equation 2Na + 2H+ 2Na+ + H2

water

sodium

Na H+ OH- H2OH

-Na+



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2 Copper(II) oxide react with dry hydrogen gas (Determine the empirical formula of copper(II) oxide)

Reaction Copper(II) oxide andhydrogen gas

Chemical Equationfor the reaction

Half equatio nfor oxidation

Half equatio nfor reduction

Ionic equation

Answer:

Reaction Copper(II) oxide andhydrogen gas

Chemical Equationfor the reaction CuO + H2 Cu + H2O

Half equatio nfor oxidation H2 2H+ + 2e

Half equatio nfor reduction Cu2+ + 2e Cu

Ionic equation H2 + Cu2+ 2H+ + Cu

Copper(II) oxide

Dryhydrogen gas

Cu2+ H2 H+Cu



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EXERCISE 5

1. Compare and contrast cell P and cell Q. Include in your answer the observation and half equations forthe reactionsn of electrode in both cells.



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2. Diagram below shows the voltaic cell. Zinc act as negative terminal and copper act as

positive terminal. The voltmeter reading is 1.1 volt.

a. State the direction of electron flow in the circuit?

b. What is the function of porous pot?

c. Write the half equation for the reaction at

(i) zinc electrode:

(ii) copper electrode:

d. Write the ionic equation for the reaction in the above cell.

e. State which electrode oxidation process take place?

.

f. (i) What will happen to the intensity of blue colour of copper(II) sulphate in beaker B?

.............................................................................................................. ...



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(d) Name the substance which is

(i) oxidized : .........................................................

(ii) reduced : .................................................................

(e) Name the

(i) oxidizing agent : .............................................................................

(ii) reducing agent : ..............................................................................

(f) State the oxidation number of chlorine in chloride ion?

................................................................................................................

4 Table below shows the concentration of copper chloride in solution X and Y.

Solution X Solution Y

0.001 mol dm-3 2.0 mol dm-3

Both solutions are electrolyzed separately using carbon as electrodes.

(a) Write the half equation of the reaction that takes place at the anode for

electrolysis of

(i) solution X :

(ii) solution Y :



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(b) Name the products formed at the cathode and anode.

(b) The products collected at the anode in the electrolysis of solutions XandYare different.Explain why.

(d) Name the substance oxidized in the electrolysis of :

(i) solution X :

(ii) solution Y :

(e) Name the substance reduced in the electrolysis of :

(i) solution X :

(ii) solution Y :

Product formed at :

Anode Cathode

Solution X

Solution Y



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1.4.3 Redox Reaction inCorrosion of Metal [ Rusting of Iron]

EXERCISE 6

1 Draw a labelled diagram to show how the rusting of iron involved the ionization of iron andthe flow of electron.

2 Diagram below shows the use of zinc plates on an iron ship to prevent rusting.

(a) Explain how the zinc plates protect the iron ship from rusting.

(b) Write the half equation for the reaction in (a).

Iron ship

Zinc plate

Sea water

TAQ 8388



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3 The diagram shows the set-up of the apparatus to study the effect of other metals on the

rusting of iron nails.

(a) What is the function of :

(i) phenolphthalein?

(ii) potassium hexacyanoferate(III) solution?

(b) State the observation for each test tube P, Q, R and S after a few days.

(i) Test tube P

(ii) Test tube Q

(iii) Test tube R

(iv) Test tube S

Copper

Jelly + phenolphthalein + potassium hexacyanoferrate(III) solution

ZincMagnesium

P Q R S

Ironnail

Iron

nailIronnail

Ironnail



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(c) Based on the observations,

(i) state the metals that can prevent the rusting of iron nail

(ii) state the metal that can accelerate the rusting of iron nail.

(iii) arrange the four metals i.e. iron, zinc, magnesium and copper inascendingorder of their electropositivity.

(d) (i) State the type of reaction that takes place when iron rusts.

(ii) Write the half equation for the reaction in (d) (i).

(iii) What is the purpose of test tube R in this experiment?

4

(a) If magnesium and iron are exposed to the atmosphere . Which metal willCorrode faster?Explain your answer.

.

.

(b) Why are the products made of aluminium self-protected from corrosion?

Metals will corrode when exposed to the atmosphere over a period of time.

The rate of corrosion depends on the position of the metal in

the Electrochemical Series.



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(c ) Electroplating is one way to control the rusting of iron. Suggest twootherways to prevent iron from rusting?

1.4.4 Redox Reaction In Displacement of HalogenFrom Its Halide Solution.

EXERCISE 7

1 Predict whether the following reactions occur or not.[ If the reaction occurs, put a tick ( ) ; if no reaction occurs, put a cross ( X ) ]

Reactants ( ) / (X) Products

1 KI + Cl2

2 KI + Br2

3 KBr + Cl2

4 NaI + Br2

5 NaBr + I2

6 KCl + Br2

7 NaCl + I2

2

(a) How do you confirm the formation of iodine in the experiment?

(b) Write the half equation for the chemical change that takes place in :

(i) bromine water :

(ii) potassium iodide :

(iii) a reducing agent :

(iv) an oxidizing agent :

(c) Write ionic equation for redox reaction.

Iodine is formed when bromine water is added to potassium iodide solution.



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(d) Suggest halogen X that can replace bromine water so that iodine is alsoformed.

1.4.5 Redox Reaction In The Change Of Fe2+ Fe3+ and Fe3+ Fe2+

EXERCISE 8

1 Chlorine water is added drop by drop to 2 cm3of iron(II) sulphate solution in a testtube. The test tube is warmed gently.

(a) Write the :

Half equation of oxidation:

Half equation of reduction:

Overall ionicequation :

2 In one experiment to investigate for the changes of Fe 3+ ions to Fe 2+ ions,hydrogen sulfide gas is pass through iron(III) chloride solution. The reaction isrepresenting with chemical equation below.

2 FeCl3 + H2S 2 FeCl2+ S + 2 HCl

(a) State twoobservation in the above experiment.

i.................................................................................................

ii.................................................................................................

(b) (i) What is the type of reaction in the above experiment?

.........................................................................................................

(ii) What is the function of hydrogen sulphide in the reaction?

.................................................................................................

(iii) Suggest another substance that can act as hydrogen suphide inthe above reaction.

.......................................................................................................



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(c) What is the oxidation number of sulphur in hydrogen sulphide?

............................................................................................................

(d) (i) State the changes in the oxidation number of iron in theexperiment?

(ii) Write the half equation for the changes in iron in the aboveexperiment?

.......................................................................................................

(iii) How to confirm iron (II) ion is produced?

.......................................................................................................

1.4.6 Redox Reaction In term of Transfer of Electron at a Distance

EXERCISE 9

1 Diagram 2 shows the set- up of apparatus to investigate the reaction between iron(II)chloride solution and potassium manganate(VII) solution through the transfer of electronsat a distance.

Dilute sulphuric acid

G

Carbon electrodeQ

Potassiummanganate(VII)

Iron(II) chloride solution

Carbon electrodeP

DIAGRAM 2

- +



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(a) What is the function of dilute sulphuric acid?

(b) In Diagram 2, draw the direction of the flow of electrons.

(c) (i) What is the colour change in the solution around electrode P?

...............................................................................................................

(ii) Describe a chemical test to determine the product formed in the solutionat electrode P.

.

...............

(d) What is the substance that is being oxidized in the experiment?Explain why.

.

(e) Write the half equation for the reaction that occurs at electrode Q.

(f) Suggest another reagent that can replace potassium manganate(VII) solution.

...........

(g) What is the change in oxidation number of manganese in the reaction?



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2 The apparatus set up below is to investigate the redoxs reaction.

(a) How do you know the reaction has started?

......................................................................................................................

(b) What is the function of potassium chloride?

.......................................................................................................................

(c) What is the color changes that can be observed at M of U tube in half an hour ofexperiment.

....................................................................................................................................

(d) Give one test to confirm the product at M?

....................................................................................................................................

....................................................................................................................................

(e) What is

reduced:..........................................................................................

oxidized:

(f) Write the half equation for the reaction which occurs at N



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(g) Name one substance that can replace chlorine water to get the same product at M

(h) (i) If potassium iodide is replaced by iron(II) sulphate, what will observed at M

.

(iii) Explain your answer.



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1.3.7 Redox Reaction In The Reactivity Series Of Metals And Its Applications

EXERCISE 10

1 Determine whether the following reactions occur or not.If the reaction occurs, mark and if not, mark X.

Reactants / X Chemical Equation

(a) Hydrogen + zinc oxide

(b)Magnesium oxide +carbon

(c) Copper + zinc oxide

(d)Aluminium + carbondioxide

(e) Carbon + silver oxide

(f) Hydrogen oxide + copper

(g)Iron(II) oxide + hydrogengas

(h) Magnesium + steam

(i) carbon dioxide + lead

(j) Iron + lead(II) oxide



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2 Diagram 2.1 shows the set-up of the apparatus for an experiment to construct theElectrochemical Series through the ability of metals to displace other metalsfrom their salt solution.

The experiment was repeated using copper and P metals to replace zinc andP nitratesolution to replace silver nitrate solution.Table 2.2 shows the result obtained.

Experiment Metal Silver nitrate solution P nitrate solution

I ZincSilver metal is

displaced.P metal is displaced.

II CopperSilver metal is

displaced.No reaction.

III PSilver metal is

displaced.

Table 2.2

(a) What is meant by Electrochemical Series?

.....................................................................................................................

(b) Based on the results in Table 2.2, arrange the metals silver, copper, Pand zinc in ascending order of electropositivity.

(c) Name the suitable metal P.

More electropositive

Silver nitrate solution

Zinc plate

Diagram 2.1



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(d) Based on Experiment I :

(i) Zinc can displace silver metal from silver nitrate solution. Explainwhy.

(ii) Write the chemical equation for the reaction.

(iii) What is the change in the oxidation number of zinc?

(e) Diagram 2.3 shows the set up of apparatus for the reaction betweenmetal P and copper(II) nitrate solution.

Diagram 2.3

What is the colour change of the of copper (II) nitrate solution?.Explain why.

(f) State twouses of the Electrochemical Series besides the determining theability of a metal to displace another metal from its salt solution.

.........

......

Copper(II) nitratesolutionMetal P plate



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ANSWER:

EXERCISE 41. Procedure

1. Fill four test tube with salt solution of W, X, Y and Z2 Add a strip of metal W into each test tube respectively3 Repeat the steps above for metals X, Y and Z

Observation:

Salt metal solutionMetal

W X Y Z

W

X X

Y X X

Z X X

DepositX No deposit

Inference1 Deposit formed when metal W is put in the salt solution of X, Y and Z

Metal W is more electropositive than metal X, Y and Z2 Deposit formed when metal X is put in the salt solution of Y and Z

Metal X is more electropositive than metal Y and Z3 Deposit formed when metal Y is put in the salt solution Z

Y is more electropositive than ZNo deposit formed when metal Y is put in the salt solutions of W and XMetal Y is less electropositive than W and X

4 No deposit is formed when metal Z is put in the salt solutions W, X and YMetal Z is less electropisitive than W, X and Y

The descending order is W, X, Y, Z

ConclusionMetal that is more electropisitive will displace the metal which is lesselectropositive from its salt solution.



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4. (a) (i) 4OH- 2H2O + O2 + 4e(ii) 2Cl- Cl2 + 2e

(b)

(c ) The type of ion that is selected to be discharged at anode insolution Ssolution Y is different. In solution X, OH-ion is selected because theposition ofOH- ion is lower than Cl-ion in the Electrochemical Series. In solution Y,Cl-ion is selected because the concentration of Cl-ion is higher thanOH- ion.

(d) (i) X: Hydroxide ion(ii) Y: Chloride ion

(e) (i) X: copper ion(ii) Y: copper ion

Product formed at :

Anode Cathode

Solution X Oxygen copper

Solution Y Chlorine copper



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EXERCISE 6

1.

2 (a) Zinc is higher than iron in the Electrochemical Series.Zinc is more electropositive than iron.Zinc has higher tendency to release electron than iron.Zinc is ionized/oxidisedElectron flow from zinc to iron

(b) Zn Zn2+ + 2e

3 (a) (i) To detect the presence of OH-ions.(ii) To detect the presence of Fe2+/ iron(II) ions.

(b) (i) No blue colour is formed.(ii) No blue colour is formed.(iii) Intensity of blue colour is low.(iv) Intensity of blue colour is high.

(c ) (i) Zinc and magnesium(iv) Copper(v) Copper, iron, zinc and magnesium

(d) (i) Oxidation(ii) Fe Fe2+ + 2e(iii) As a control experiment

O2

Iron

Water dropletFe2O3 . xH2O (rust)

e- e

-

Fe2+

Fe2+

Fe2+

OH

-

OH-

OH-

O2

O2



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4. (a) Magnesium.Magnesium is more electropositive than iron.Magnesium atom is easier to release electron compare to iron atom

(b) the oxide layer is non porous and firmly coated the aluminium

(c ) covering the surface with oil/ paintAlloying the iron with chromiumElectroplating with metal higher than iron in the electrochemical series

EXERCISE 71

Reactants / X Products

1 KI + Cl2 KCl + I2

2 KI + Br2 KBr + I2

3 KBr + Cl2 KCl + I2

4 NaI + Br2 NaBr + I2

5 NaBr + I2 X

6 KCl + Br2 X

7 NaCl + I2 X

2. (a) Add the starch solutionThe color turns blue

(b) (i) Br2 + 2e 2Br-

(ii) 2I- I2 + 2e(iii) Iodide ion/ KI(vi) Bromine

(c ) Br2 + 2I- 2Br- + I2

(d) chlorine



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EXERCISE 8

1. (a) Fe2+ Fe3+ + eCl2 + 2e 2Cl-

2Fe2+

+ Cl2 2Fe3+

+ 2Cl-

2 (a) 1. Greenich yellow solution formed2. yellow precipitate

(b) i. redoxii. oxidizing agentiii. iron(II) chloride

(c ) -2

(d) i. +3+ 2ii. Fe 3+Fe 2++ eiii. 1. Add sodium hydroxide solution

2. green precipitate formed

EXERCISE 9

1 (a) To allow the electron flow(b) in the diagram(c ) (i) Green to brown

(ii) Add sodium hydroxide solution.Brown precipitate is formed.(d) Fe2+/ Iron(II) ion.

Released electron // Oxidation number increased from +2 to +3.

(e) MnO4- + 8H+ + 5e Mn2+ + 4H2O

(f) Acidified potassium dichromate(VI) solution, K2Cr2O7(g) +7 to +2

2 (a) the galvanometer shows reading(b) to separate the two electrolyte and allow the movement of ion between

two solution(c ) colorless to brown

(d) add starch solutionThe color changes to blue(e) chlorine

Potassium iodide/ iodide ion(f) Cl2 + 2e 2Cl

-(g) bromine water/ acidified potassium mangganate(VII)/ acidified potassium

dichromate(VI)(h) (i) M: light green to yellow solution

(ii) Iron (II) ion is oxidized to iron(III) ion



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