Nota Sains Bab 5.Pptx

8/12/2019 Nota Sains Bab 5.Pptx

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4.05.01 Physical and Chemical Changes

Physical Change

1. A physical change is a change which does not produce a new substance. Only the physical

state of the substance has changed.

2. A certain substance undergoes a physical change when there is a change in the arrangement

of the particles.

3. Energy is absorbed or released in order to rearrange the particles.

4. Examples of physical change are given in Table 4.05.01.

Physical change Examples

a. Melting Ice heat water The ice only changed its state from a solid to a liquid.

b. Evaporation Water heat water vapour The water has changed from a liquid to a gas.

c. Dissolving a

substance

Sugar + water dissolve sugar solution No new substance is formed.

d. Crystallisation Hot saturated sodium chloridesolutioncool

sodium chloridecrystals

No new substance is formed.

e. Sublimation Iodinecrystalsheat sublimed iodine

The iodine has changed from crystals to powder.

f. Freezing Water cool ice The water has changed from a liquid to a solid.

Tabl e 4.05.01 Examples of physical changehot saturatedsolution ofsodiumchloridecrystals of sodium chlorideiodine crystalsheat

Figure 4.05.01 Examples of physical change coolclampsublimediodineCrystallisationSublimation

Chemical Change


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1. A chemical change is a change which produces a new substance. The new substance has

chemical properties and a composition different from the original substance.

2. Examples of chemical change are given in the table below.

Chemical change Examples

a. Burning

magnesiu

m in air

Magnesium +oxygen burn magnesium oxide

A new substance is formed with different chemical properties. Energy isreleased.

b. Mixing

zinc with

copper

sulphatesolution

Zinc + copper sulphatesolutionreact copper + zinc sulphate

Copper is displaced from the solution.

c. Heating a

mixture of

iron and

sulphur

Iron +sulphurheat iron sulphide

A new substance is formed with different chemical properties.

d. Heating

mercury

oxide

Mercury

oxideheat mercury + oxygen

Mercury oxide is decomposed into mercury and oxygen.

e. Adding

calcium

into water

Calcium +waterreact

calcium hydroxide +hydrogen

Two new substances are formed and heat energy isreleased.

f. Mixing

copper

chloride

solution

and

sodium

hydroxide

solution

Copper chloride solution + sodium hydroxidesolution copper

hydroxide + sodium chloride Two new substances are formed.

Tabl e 4.05.02 Examples of chemical change iron filings + sulphur powder

magnesiumribbon


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brightdazzlinglightwater

Figure 4.05.02 Examples of chemical change test-tube holderhydrogenHeating a mixture of iron and sulphurheatcalciumBurning magnesium ribbonReaction of calcium with water

Consequences of Chemical Change

1. A chemical change can produce the following results.

Change in colour

Chemical change Release of gasRelease of lightFigure 4.05.03

Absorption or release of heatProduction of sound

2. Example : Heating copper carbonate produces a chemical change

heat

copper carbonate

test tube

limewater turns chalky

(carbon dioxide is released)

Figure 4.05.04 Heating copper carbonate

boiling tube

The following changes are observed:

a. New substances are formed. (Copper carbonate decomposes into copper oxide and

carbon dioxide.)

b. There is a change in colour of the original substance. (Green copper carbonate is

changed into black copper oxide.)

c. Heat energy is absorbed. (Copper carbonate has absorbed energy.)

d. A gas is released, causing a change in mass of the original substance. (Carbon

dioxide is released, causing copper carbonate to lose mass.)


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e. The process cannot be reversed to get back the original substance. (Copper

carbonate cannot be obtained back.)

Differences between Physical Change and Chemical Change

Physical change Chemical change

No new substances are formed(Water cool ice)

New substances are formed(Coppercarbonate

heatcopper oxide + carbondioxide)

Involves very little heat change Energy (heat/light) is absorbed orreleased

The change is reversible(Water ice)

The reaction is irreversible

The substances keep their original

mass (no change in mass)

The original substance changes its mass (may be due to

the release of a gas or combination with another substance)

Tabl e 4.05.03

Examples of Physical and Chemical Changes in Daily Life

1. Physical and chemical changes occur in daily life.

2. Some examples are given in the table below.

Physical change in daily life

a. Melting of ice Freezing of water Boiling of water Condensation of steam

b. Evaporation of sweatc. Formation of sugar and common salt crystalsd. Sublimation of moth ballse. Vaporisation of petrol

Chemical change in daily life

a. Rusting of iron b. Burning of fuelsc. Decay of organic matterd. Respiration of glucose in the body cellse. Digestion of food into simple substancesf. Manufacture of ammoniag. Making soap, plastic and synthetic rubberh. Making of food by green plants in sunlight

Tabl e 4.05.04

4.05.02 Heat Change in Chemical Reactions


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1. During a chemical reaction, heat energy is

a. released to the surroundings (to form new chemical bonds); or

b. absorbed from the surroundings (to break up chemical bonds).

2. Therefore, chemical reactions may be classified into two types.

Exothermic Reaction

What is an Exothermic Reaction?

1. An exothermic reaction is one which releases heat energy to the surroundings.

2. This means that during the reaction, more heat energy is released than absorbed.

3. Therefore, an exothermic reaction has the following characteristics.

a. It sets free heat energy.

b. The temperature of the surroundings rises.

c. The reactants (reacting substances) have more energy than the products of the

reaction.

energy released

4. Energy

5. products

6. Time

7. Figure 4.05.05 Exothermic reaction

8. reactants9. thermometer

temperature of water risesheat energy releasedsodium hydroxide pellets

Examples of Exothermic Reaction

The following are examples of exothermic reaction.

1. The burning of a candle

a. When a candle burns, heat and light are released.b. Hydrocarbon + oxygen water + carbon dioxide

2. Dissolving calcium oxide (quicklime) in water

a. Calcium oxide dissolves in water and releases a lot of heat energy.

b. Calcium oxide + water calcium hydroxide

3. Reaction between magnesium and dilute sulphuric acid

a. Magnesium reacts with dilute sulphuric acid and produces a lot of heat.

b. Magnesium + dilute sulphuric acid magnesium sulphate + hydrogen

4. Neutralisation of an acid with an alkali

a. This process produces heat energy.b. Dilute hydrochloric acid + sodium hydroxide solution sodium chloride + water


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5. Decay of organic matter

a. When organic matter decays due to bacteria or fungi action, heat energy is released.

b. Organic matter + oxygen carbon dioxide + water

6. Respiration in the body cells

a. The process of respiration occurring in the body cells produces energy. Some of thisenergy is in the form of heat to keep the body warm.

b. Glucose + oxygen carbon dioxide + water + energy

glucose

INPUT

carbon dioxide

energy

Figure 4.05.06 Cell respiration

OUTPUT

oxygen

water

cell

Endothermic Reaction

What is an Endothermic Reaction?

1. An endothermic reaction is one which absorbs heat energy from the surroundings.

2. This means that during the reaction, more heat energy is absorbed than released.

3. Therefore, an endothermic reaction has the following characteristics.

a. It absorbs heat energy.b. The temperature of the surroundings falls.

c. The reactants have less energy than the products of the reaction.

energy absorbed

4. Energy

5. products

6. Time

7. Figure 4.05.07 Endothermic reaction

8. reactants


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9. thermometer

temperature of water fallsheat energy is absorbedammonium nitrate

Examples of Endothermic Reaction

1. Dissolving ammonium chloride in water a. Heat is absorbed when ammonium chloride dissolves in water.

b. Ammonium chloride + waterammonium chloride solution

2. Boiling water a. Heat is absorbed when water is boiled.

b. Water heat steam

3. Melting wax a. Wax absorbs heat in order to melt.

b. Wax(solid)

heat liquid wax

4. Heating copper sulphate crystals a. When copper sulphate crystals are heated, they absorb heat and set free water of

crystallization.

b. Copper sulphate crystals(blue)

heatanhydrous copper sulphate (white) +water

5. Decomposition of mercury oxide by heat a. On heating, mercury oxide absorbs heat and decomposes into mercury andoxygen.

b. heatMercury oxidemercury + oxygen

Heat Changes in Industrial Chemical Reactions

Production of Ammonia by the Haber Process

1. Ammonia is produced by the Haber process.

2. During this process, one part of nitrogen by volume is combined with three parts of hydrogen

by volume to form ammonia.

Nitrogen + hydrogenmercury + oxygen

3. This process is reversible.

4. During the process, heat is produced. Therefore, this is an exothermic reaction.

Production of Sulphuric Acid by the Contact Process

1. Sulphuric acid is produced by the Contact process.


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a. Sulphur + oxygen sulphur dioxide

b. Sulphur dioxide + oxygen sulphur trioxide + heat

c. Sulphur trioxide is dissolved in concentrated sulphuric acid (99.5%).

d. Concentrated sulphuric acid is diluted in water (98.0%).

2. a. Sulphur is burnt in air to obtain sulphur dioxide gas.

b. Sulphur dioxide is made to react with more oxygen to form sulphur trioxide.

c. Sulphur trioxide is dissolved in concentrated sulphuric acid to get a 99.5%concentration called oleum.

d. The highly concentrated sulphuric acid is diluted with water to produce 98.0%sulphuric acid.

3. During the process, heat is released. Therefore, this is an exothermic reaction.

4. This process uses a catalyst, a pressure of 1 2 atmospheres and a controlled temperature

of 450C 500C.5. Sulphuric acid has many uses in industry.

4.05.03 Reactivity Series of Metals

Reactivity of Metals with Water

1. Potassium, sodium, calcium and magnesium react with cold water to form alkalis and set free

hydrogen.

Reactive metal + water alkali + hydrogen

2. Potassium is too reactive and explosive, and should not be used by students.

3. The reaction of sodium, calcium and magnesium with water can be observed as follows.

Activity Observation/Explanationsodium

water

Sodium and water

1. The sodium floats on water (less dense than water) and rolls into a ball.

2. It runs all over the surface of the water and reacts very vigorously withthe water.

3. At the end the sodium bursts into a yellow flame.4. A gas (hydrogen) is set free and heat is produced.5. The water at the end of the reaction turns red litmus paper blue,

showing an alkali has been formed.


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Sodium + water hydrogen + sodium hydroxide

hydrogen (burns with a popsound) watercalcium

Calcium and water

6. The calcium sinks in water, showing that it is denser than water.7. It reacts vigorously with the water and sets free a gas which burns with

a pop sound (hydrogen).8. Heat is produced by the reaction.9. At the end the water turns red litmus paper blue, showing an alkali is

produced.10. The solution turns chalky when carbon dioxide is bubbled through it,

showing that the solution is limewater.

Calcium + water hydrogen + calcium hydroxide (limewater)

bubblesof hydrogenwater

polished magnesium ribbonmagnesium

powder

Magnesium and water

11. The polished magnesium ribbon reacts very slowly with water, settingfree a gas (hydrogen).

12. Magnesium powder reacts more actively with water because it has amuch larger surface area.

13. A very weak alkali, magnesium hydroxide, is formed at the end of thereaction.

Magnesium + water hydrogen + magnesium hydroxide

Reactivity series: Potassium > Sodium > Calcium > Magnesium

Tabl e 4.05.05

Reactivity of Metals with Dilute Acids

1. Reactive metals react with dilute acids such as dilute sulphuric acid and dilute hydrochloric

acid to produce hydrogen and a salt.

Suitable metal + dilute acid hydrogen + salt

2. This reaction also produces heat, i.e. it is an exothermic reaction.

3. Metals which are reactive react more vigorously with the acids than metals which are less

reactive.

4. Metals do not react with dilute nitric acid .

5. The hydrogen collected in a test tube burns with a pop soun d.

6. The salts obtained depend on the metals and the acids used.

7. The salts can be obtained by evaporating their solutions to dryness.

Examples


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a. Magnesium + dilute sulphuric acid hydrogen + magnesium sulphate

b. Zinc + dilute sulphuric acid hydrogen + zinc sulphate

c. Iron + dilute hydrochloric acid hydrogen + iron chloride

hydrogen burns with a pop sound dilute hydrochloric acidFigure 4.05.08 Testing for hydrogen Figure 4.05.09 Evaporating a salt solution zincsalt solutionevaporating dishheat

Reactivity of Metals with Oxygen

1. Most metals react with oxygen to form metal oxides.

Metal + oxygen metal oxide

Examples

a. Aluminium + oxygen aluminium oxide

b. Magnesium + oxygen magnesium oxide

c. Zinc + oxygen zinc oxide

d. Copper + oxygen copper oxide

e. Iron + oxygen iron oxide

2. Metals which are reactive react more vigorously with oxygen than metals which are lessreactive.

3. When the Earths crust was very hot, many types of metals reacted with oxygen in the air to

form metal oxides. This explains why there are so many kinds of metal oxides in the Earth.

4. Potassium and sodium react vigorously with oxygen in the air to form oxides and must be

stored in paraffin oil.

5. Metals which are very inactive such as gold and platinum do not react with oxygen.

F igu re 4.05.10 Iron rust is iron oxide


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Reactivity Series of Metals

1. Scientists have experimented with many metals and have prepared a list of the metals in a

reactivity series.

2.

The common metals and some of their reactions are given in Figure 4.05.11.

MOST REACTIVE React with waterand releasehydrogenReact with steamand dilute acids,and set freehydrogen

Figure 4.05.11 Reactivity series of metals Metals extractedfrom their ores byelectrolysisReact very slowlywith dilute acidsand releasehydrogenCARBON Metals extractedby heating theirores with carbonMetal obtainedby heating theoxide directlyUsually found asfree elementsLEAST REACTIVE Do not react withdilute acidsPotassiumSodiumCalciumMagnesium

AluminiumZincCadmium

Iron (ferum)NickelTin (stanum)LeadCopperMercurySilver (argentum)PlatinumGold

The Position of Carbon in the Reactivity Series of Metals

1. Carbon is a non-metallic element.2. It combines easily with oxygen to form carbon dioxide.


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Carbon + oxygen carbon dioxide

3. Carbon is placed in the reactivity series of metals because of its tendency to combine with

oxygen.

4. Carbon is placed between aluminium and zinc in the reactivity series.

5. This means that carbon cannot displace aluminium from aluminium oxide, but can displace

zinc from zinc oxide.

Carbon + zinc oxide zinc + carbon dioxide

4.05.04 Applications of the Reactivity Series ofMetals

Importance of the Reactivity Series of Metals

1. Reaction of a metal.

Since this is a series of metals arranged in order, starting with the most reactive metal, we

can use it to predict reactions which involve metals e.g. magnesium will react more

vigorously with an acid than iron, because magnesium is higher in the reactivity series.

2. Displacement of metal from a solution.

The reactivity series enables us to know whether a metal can displace another in a solution.

copper sulphatesolution

3. Figure 4.05.12 Displacement of a metal

4. iron nail

5. Iron + copper sulphate solution copper + iron sulphate

6. Iron can displace copper from the solution because iron is more reactive than copper.

7. Displacement of metal from an oxide

The series enables us to predict whether a metal can displace another from its oxide.

heat

Zinc + lead oxide

lead + zinc oxide

This reaction occurs because zinc is more reactive than lead.


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8. Extraction of metal using carbon

The series enables us to predict whether a metal can be extracted from its ore by heating it

with carbon.

heat

Carbon + copper oxide

copper + carbon dioxide

In this reaction, carbon is more reactive than copper.

9. Terminals of a simple cell

The series allows us to identify the terminals of a simple cell. The more reactive metal forms

the negative terminal.

electronsFigure 4.05.13 A simple cell ironelectronsleaddilute sulphuric acid

Principle of Metal Extraction from the Ores

1. Metals are usually extracted from their ores.2. Most ores are in the form of metallic oxides, metallic sulphides and metallic carbonates.

Metal ore Main constituent

BauxiteCassiteriteHaematite

Aluminium oxideTin oxideIron oxide

ArgentiteGalenaIron pyrite

Zinc blend

Silver sulphideLead sulphideIron sulphide

Zinc sulphideCalciteMalachiteMagnesite

Calcium carbonateCopper carbonateMagnesium carbonate

Tabl e 4.05.06

3. Metallic oxides , except for mercury oxide and silver oxide, do not decompose when heated

alone.

4. Most metallic sulphides decompose into their oxides and sulphur dioxide when heated.

Metallic sulphide + oxygen


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metallic oxide + sulphur dioxide Lead sulphide + oxygenlead oxide + sulphur dioxide

5. Most metallic carbonates decompose into their oxides and carbon dioxide when heated.

Metallic carbonate metallic oxide + carbon dioxide Copper carbonatecopper oxide + carbon dioxide

6. Metallic sulphides and carbonates are roasted into their oxides before the metals are

extracted.

7. Metals which are less reactive than carbon are obtained by heating the oxides

with carbon .

heat Metallic oxide + carbon metal + carbon heatLead oxide + carbonlead + carbon dioxide

8. Metals which are more reactive than carbon are obtained by the process

of electrolysis of their molten oxides.

electrolysis

Molten metallic oxide metal + oxygen electrolysisMolten magnesium oxidemagnesium + oxygen

Extraction of Tin from Its Ore

1. Tin is extracted from its ore cassiterite (tin oxide) by heating the ore with carbon in a blast

furnace under very high temperature.

heatTin oxide + carbontin + carbon dioxide

waste gasesFigure 4.05.14 Blast furnace for extracting tin

hot airtin oxide + coke (carbon) + limestonewaste gaseshot airslag from impuritiesmolten tin

2. Once the blast furnace has started, it continues to operate.


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3. Before pouring the tin ore into the blast furnace, it is washed with water to remove any soil,

and then it is roasted in air to remove foreign substances such as sulphur and arsenic.

4. Then the raw materials tin oxide, coke (poor quality carbon) and limestone are added from

the top of the furnace.

5. The coke reduces the tin oxide into tin and carbon dioxide.6. The limestone first decomposes into quicklime (calcium oxide), which combines with foreign

substances such as silica and forms slag. Most of the slag is calcium silicate.

7. Hot air is blown in from the bottom of the furnace to help combustion in the furnace.

8. Slag and molten tin ore are run off from the bottom of the furnace when necessary.

9. The molten tin is made into ingots for sale.

10. Tin is used to make containers for food because it does not rust.

Extraction of Iron from Its Ore

waste gasesFigure 4.05.15 Smelting iron ore in a blast furnace

hot airiron ore + coke + limestonewaste gaseshot airslagmolten iron

1. Iron is extracted from haematite (non-magnetic iron oxide) or magnetite (magnetic iron oxide)

by heating the ore with carbon.

heatIron oxide + carboniron + carbon dioxide

2. The iron ore, coke and limestone are added into the furnace from the top.

3. The coke (carbon) reduces the iron oxide to iron, and sets free waste gases such as carbon

dioxide and carbon monoxide.

4. The limestone is first decomposed by heat into quicklime (calcium oxide) which then

combines with the impurities to form slag (mainly calcium silicate).

5. Slag is less dense than molten iron and so floats on it. The slag is tapped off from time totime.

6. The molten iron is tapped off into large moulds called pigs , and the iron that is obtained is

called pig iron .


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F igur e 4.05.16 A blast furnace

4.05.05 Electrolysis

What is Electrolysis?

1. Electrolysis is the process of using electrical energy to decompose a molten or an aqueous

(watery) solution of an ionic compound.

2. Electrolysis involves the movement of ions . An ion is a charged atom or a charged group of

atoms.

3. The diagram below shows a typical set of apparatus used to electrolyse an aqueous solution.

ammeterFigure 4.05.17 Principle of electrolysis

batteryrheostatcarbon electrode(cathode)carbon electrode(anode)cationelectrolyteanion

4. The battery supplies the electrical energy for electrolysis to take place. It must supplya direct current .

5. The ammeter shows the amount of electric current passing through the circuit.

6. The rheostat is used to control the current flowing through the circuit.

7. The anode is the carbon electrode connected to the positive terminal.

8. The cathode is the carbon electrode connected to the negative terminal.

9. An anion is a negatively charged ion which is attracted to the anode.

10. A cation is a positively charged ion which is attracted to the cathode.

11. The electrolyte is the liquid (molten substance or aqueous solution) used in electrolysis. It

contains anions and cations. It conducts electricity and will undergo chemical change during

electrolysis.


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12. The following substances may be used as electrolytes.

a. Acids, e.g. hydrochloric acid and sulphuric acid

b. Alkalis, e.g. sodium hydroxide and ammonia solutions

c. Salt solutions, e.g. sodium chloride and copper sulphate solutions

13. Molecular substances cannot be used as electrolytes because they do not contain ions, e.g.kerosene, petrol and alcohol.

Principle of Electrolysis

1. a. During electrolysis, positively charged ions (cations) are attracted to the negativeelectrode (cathode). At the cathode the positively charged ions receive electrons,

become discharged, and form atoms.Positive ions + electronsneutral atoms

b. These atoms are usually hydrogen or metals.2. a. During electrolysis, negatively charged ions (anions) are attracted to the positive

electrode (anode). At the anode the negatively charged ions release electrons, becomedischarged and form atoms.

Negative ions electronsneutral atoms

b. These atoms are non-metallic substances.

Application of Electrolysis in Industry

Electrolysis is used in industry for the following purposes.

a. To extract metals which are more reactive than carbon, e.g. to obtain calcium, magnesium

and aluminium

b. To electroplate metallic objects, e.g. to silver plate trophies

c. To purify metals, e.g. to purify silver.

Extraction of Aluminium by Electrolysis

carbon cathodeFigure 4.05.18 Electrolysis of molten aluminium oxide to extract aluminium carbon anode

electric supplymolten aluminium oxide and molten cryolitemolten aluminiumsteel container (electrolytic cell)

1. Aluminium is extracted from its ore, bauxite (aluminium oxide), by electrolysis.

2. Bauxite (boiling point 2015C) is heated with cryolite in a blast furnace until they melt.

3. Cryolite lowers the boiling point of bauxite to 900C.

4. The molten mixture is transferred to a steel container called an electrolytic cell for

electrolysis.

5. In the molten state, the aluminium oxide splits up into positively charged aluminium ions andnegatively charged oxide ions. These ions move about freely in the molten aluminium oxide.


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6. At the anode

The negatively charged oxide ions are attracted to the anode. They surrender electrons to

the anode, become discharged and form oxygen atoms. So oxygen gas is released at the

anode.

7. At the cathodeThe positively charged aluminium ions are attracted to the cathode. They receive electrons

from the cathode, become discharged and form aluminium atoms. So aluminium is deposited

in the electrolytic cell.

8. Finally, aluminium oxide is electroysed into aluminium and oxygen.

Aluminium oxidealuminium + oxygenelectrolysis

Electroplating

1. Metals easily react with oxygen in damp air and corrode. For example, iron corrodes easily in

damp air and forms rust.

Iron + oxygeniron oxide (rust)

2. Sometimes metals which corrode easily are coated with a layer of metal which does not

corrode e.g. tin, silver, chromium and gold. This process, called electroplating , is done by

electrolysis.

3. The electroplated metal becomes shiny and more attractive. It is also protected from

corrosion.

4. For electroplating an article:

a. the article e.g. badge is used as the cathode.

b. the metal chosen for electroplating e.g. silver, is used as the anode.

c. the electrolyte used must be a salt solution of the metal chosen for electroplating e.g.

silver nitrate may be used in silver plating.

Purification of Metals

1. Metals may contain foreign substances or impurities.

2. These impure metals can be purified by electrolysis.

3. For purifying an impure metal:

a. the impure metal is used as the anode, e.g. impure silver;

b. a pure piece of the metal is used as the cathode, e.g. silver;

c. a salt solution of the metal to be purified is used as the electrolyte, e.g. silver nitrate

solution.

4. During electrolysis:

a. The impure metal loses mass because positively charged metal ions from it areattracted to the cathode.


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b. The pure metal increases in mass because pure metal atoms are deposited on it.

c. The electrolyte does not change because the metal lost by the anode is gained by the

cathode.

d. Impure substances are deposited at the bottom of the electrolytic cell.

4.05.06 Production of Electrical Energy FromChemical Reactions

Simple Cells

1. Simple cells produce electrical energy from chemicalreactions.

Chemical energy electrical energy

2. A simple cell consists of:a. two different metal plates or a carbon plate and a

metal plate as electrodes;b. a dilute acid, an alkali or a salt solution as an

electrolyte.3. The more reactive metal in the pair forms

the negative electrode and sets free electrons. The lessreactive metal or carbon forms the positive electrode.

4. The flow of electrons round the circuit forms an electriccurrent.

5. The weakness in a simple cell is that hydrogen bubbles areformed on the positive electrode. This gradually slowsdown the production of electrical energy. (Thisphenomenon is called polarisation.)

6. Another weakness is that the negative electrode reactswith an acid electrolyte even when the cell is not in use.(This corrosion of the negative electrode in the acid iscalled local action.)

7. One advantage of a simple cell is that it is easy to set upand its voltage can be increased by using more plates.

8. The disadvantage is that its electrolyte is easily spilled.(Also, polarisation and local action take place.)

The Chemical Reactions in a Simple Cell


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1. A simple cell can produce electrical energy to light up a bulb of low voltage.

2. In the cell above, zinc forms the negative electrode because it is more reactive than copper.

3. Zinc releases electrons and forms zinc ions. So zinc dissolves in the acid.

4. The electrons set free flow through the bulb and light it up, and then flow to the copper

electrode.5. At the copper electrode, the electrons from the zinc combine with the positive hydrogen ions

in the acid to form hydrogen gas.

6. Hydrogen gas is set free at the copper electrode.

A Simple Cell Using a Salt Solution as Electrolyte copperFigure 4.05.21 A simple cell using a salt solution aselectrolyte electronszinc

coppersulphatesolution

1. Zinc forms the negative electrode in the simple cell because it is more reactive than copper

and sets free electrons more easily.

2. Zinc forms zinc ions and loses mass.

3. The electrons set free form an electric current and flow through the lamp to the copper

electrode. Thus the lamp becomes lighted.

4. On reaching the positive copper electrode, the electrons from the zinc combine with the

positive copper ions from the solution and form copper atoms.

5. These copper atoms are deposited on the copper electrode, which increases in mass.

6. It this way, chemical reactions in a simple cell change chemical energy into electrical energy.

Other Cells and Their Uses

1. There are many types of cells being used in daily life.

2. Cells may be classified as:

a. primary cells;

b. secondary cells.3. Primary cells cannot be recharged, e.g. dry cells.

4. Secondary cells can be recharged, e.g. lead-acid accumulators.

Dry Cells

sealingmaterialFigure 4.05.22 Vertical section through a dry cell metal capammonium chloride

pastecarbon rod (+)


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manganese oxide+ carbon powderzinc container ( )

1. A dry cell supplies about 1.5 V.

2. The carbon rod in the centre together with its metal cap serve as the positive electrode.3. Surrounding the carbon rod is a mixture of manganese oxide and carbon powder.

a. The manganese oxide oxidises the hydrogen produced around the carbon rodinto water.

b. The carbon powder reduces resistance in the cell.

4. The ammonium chloride paste acts as the electrolyte. Sometimes it is mixed with zincchlorid

5. The zinc container serves as the negative electrode.

6. a. The zinc atoms give out electrons and form zinc ions.

b. The electrons flow to the carbon rod.

c. These electrons combine with the ammonium ions in the electrolyte, formingammonia and hydrogen.

7. Advantages a. A dry cell is light and easy to carry it about.

b. It does not contain any liquid electrolyte which can spill.

c. It supplies a steady high current.

8. Disadvantages a. It cannot last long. In time it leaks and gives out a very corrosive liquid.

b. It cannot be charged (primary cell).

9. Uses a. It is widely used in torches, wall clocks and toys.

b. It is usually joined in series to get a higher voltage.

Alkaline Cell

outer steel caseFigure 4.05.23 Vertical section through an alkaline cell metal capmanganese oxideand carbon powder (+)metal rod(collects current)potassium hydroxide(electrolyte)zinc powder ( )

1. An alkaline cell supplies about 1.5 V.

2. The metal rod in the centre collects current and sends it out through the metal cap.

3. The electrolyte used is either potassium hydroxide or sodium hydroxide.

4. The zinc powder forms the negative electrode.


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5. The manganese oxide and carbon powder serve as the positive electrode.

6. Advantages

a. It supplies a steady high current, and can last a long time.

b. In the long run, it supplies more electrical energy than a dry cell.

c. It is light and can be carried about easily.7. Disadvantage

Most cannot be recharged (primary cells).

8. Uses

a. It is used in torches, radios, toys and wall clocks.

b. It is also used in electrical appliances which require high energy such as in

flash cameras.

Lead-acid Accumulator

negative lead platesFigure 4.05.24 A lead-acid accumulator

terminalspositive lead plates covered with lead dioxideterminalscelldilute sulphuric acidacidlead plate

Vertical sectionShowing cells inside

1.A lead-acid accumulator usually supplies 12.0 V. It is usually made up of six lead-acidcells.

2. a. The negative terminal is made up of lead plates.

b. The positive terminal consists of lead plates covered with lead dioxide.

c. The electrolyte used is sulphuric acid.

3. a. The accumulator must be charged before it can be used .

b. During charging, the accumulator changes electrical energy into chemical energy andstores it.Electrical energy chemical energy

c. During discharging when it is in use, the chemical energy stored is converted backinto electrical energy.Chemical energy electrical energy

4. Advantages a. A lead-acid accumulator is portable.

b. It can be recharged when its voltage runs low (secondary cell).

c. It can last a very long time (2 years) if properly looked after.

d. It supplies a steady high current.

5. Disadvantages a. A lead-acid accumulator is heavy and expensive.


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b. The acid can be easily spilled.

c. The accumulator cannot store much energy and has to be recharged frequently.

6. Uses

Lead-acid accumulators are used in motor vehicles and for temporary lighting.

Silver Oxide-mercury Cell

zincpowder ( )

Figure 4.05.25 A silver oxide-mercury cell silver oxide or mercury oxide (+)steel casepotassium hydroxide (electrolyte)

1. a. The positive terminal is silver oxide or mercury oxide.

b. The negative terminal is zinc.c. The electrolyte used is usually potassium hydroxide.

d. The cell supplies 1.2 V.

2. Advantages a. It is very small (size of a button).

b. It can last a very long time.

3. Disadvantage

It cannot be charged (primary cell).

4. Uses

It is used in watches, cameras and hearing aids.

Nickel-cadmium Cell

Figure 4.05.26 Nickel-cadmium cells

1. a. Nickel oxide is used for the positive terminal.

b. Cadmium is used for the negative terminal.

c. Potassium hydroxide is used for the electrolyte.

d. The cell supplies about 1.25 V.

2. Advantages a. It can be recharged (secondary cell).

b. It lasts a long time.

3. Disadvantage

It is very expensive.

4. Uses

It is used in mobile handphones and emergency lamps.


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Question 1

Diagram 1 shows the process in the synthesis and secretion of extracellularenzymes in an animal cell.

Diagram 1

a) Based on Diagram 1:

i) Explain the role of nucleus in the synthesis of enzyme. [ 3 marks]

Answer: Roles of nucleus in synthesis of enzyme are: it contains the cells hereditary information that is DNA in nucleolus it manufactures of subunits of ribosomes (RNA) and proteins associated with

ribosome the important structures in the synthesis of proteins it has nuclear envelope with pores that allows ribosomal subunits to be exported

to the cytoplasm where finally assembly of ribosomes takes place.

ii) Name one extracellular enzymes and describe how the different cellularcomponents are involved in the secretion of this enzyme. [10 marks]

Answer: Digestive enzyme that is made by pancreas cell such as lipase is a type of extracellularenzyme. The instruction for making the extracellular enzyme is transcribed

from deoxyribonucleic acid (DNA) to ribonucleic acid (RNA) in the nucleus. The RNA then leaves the nucleus through the nuclear pore and attaches itself to

the ribosome located on the endoplasmic reticulum.

When the enzyme synthesis is completed, it is extruded into the interior of theendoplasmic reticulum.
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The enzyme is then encapsulated in a transport vesicle. The transport vesicle fuses with the Golgi apparatus, releasing the enzyme into

the Golgi apparatus. In the Golgi apparatus, the enzyme is further modified before being packed in a

secretory vesicle. The secretory vesicle transports the enzyme to the plasma membrane. The secretory vesicle membrane fuses with the plasma membrane and the

enzyme is released outside the cell.

b)

"In multicellular organisms, cell specialization allows for division amongtissues, organs and systems to carry out their specific roles."

Using the information given, explain what will happen to a cell if particularcellular components are absent. [8 marks]

Answer: Every cellular component works together for the survival and specific function of the cell.For example in nervous system, nerve cells are adapted for the transmission andreception of nerve impulses. Thus nerve cells requires more energy generated bymitochondria to carry out their functions.

If there is absence of mitochondria, no energy will produce. As a result nerve cells fail tocarry out their functions. Similarly, in green plants, mesophyll cells in the leaves have

chloroplasts to perform photosynthesis. If chloroplast is absent, light energy cannot betrapped for the purpose of photosynthesis.

Epithelial tissues that lining the inner wall of trachea consists of epithelial cells that havemodified with the structure of goblet cells. These goblet cells secrete mucus to trap dustand particles in inhaled air.

A lot of Golgi apparatus with secretory vesicles are needed to secrete mucus. If there areno Golgi apparatus, no mucus will be secreted and the tissues cannot perform efficiently.

Question 2

a) Figure 2(a) and 2(b) shows two different conditions of a plant before andafter fertilizers are added within a week .


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Describe the condition of plant in figure 2(a) before fertilizers are added.Explain how excess fertilizers affected this plant shown in figure 2(b)? Ifthe plant in Figure 2(b) is watered immediately predict the changes that will

occur in the plant cell.

Answer: Plant A has turgid cells because the cell saps of the plant cells are isotonic to the waterin the soil. If too much fertilizer is added to the soil, the concentration of dissolvedminerals is increases and water in the soil becomes hypertonic to the cell saps of theroots. Thus, water diffuses out of the cell to the soil by osmosis.

The cells become plasmolyses and the plant is wilted (Plant B). If this condition isprolonged the plant will die. But if the plant is watered immediately, it will turgid onceagain. This is because by watering the plant it will produce hypotonic condition of water

in the soil to the cell saps. This will cause the water to diffuse into the cells by osmosis.The cells become deplasmolysed.

b) [10 marks]

"Mushrooms, mangoes and fish can be preserved longer using natural preservative such as sugar and salt. A housewife prefer to maintain thefreshness of the mushrooms using suitable salt solution for a week. She alsowanted to preserve the mangoes by using sugar solution and make salted fishby using salt."

i) Explain how the freshness of the mushroom can be maintained.

Answer: The mushroom is placed in a salt solution that is isotonic to the cell sap of themushroom. There will be no net movement of water molecules. The mushroom cellsmaintain it turgidity. Hence, the freshness of the mushroom can be maintained.

ii) Describe how the principle of osmosis is applied when preserving mangoesusing sugar solution and making salted fish by using salt.

Answer:
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The mangoes can be preserved by using sugar solution. The mangoes is places in aconcentrated sugar solution that is hypertonic to the cell sap of mangoes. So watermolecules diffuse from the mangoes into the surrounding hypertonic sugar solution byosmosis. This dehydrated condition is not suitable for the growth of bacteria. So, themangoes is preserved.

Salted fish is made by using salt. When salt is applied on the surface of the fishs skin, itwill form a hypertonic condition. Water will diffuses from from the fishs cells into thesurrounding hypertonic saltsolution by osmosis. Water also diffuses from the cells of thebacteria found on the fish into the hypertonic salt solution. The bacterial cells areplasmolysed. Hence, the fish can be preserved.

Nota Sains Bab 5.Pptx

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