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Chemistry

Questions &

AnswersBooklet

2018

Prepared by: Zulfigar and Zeek (10S2-2018)

Checked by: Ibrahim Waleed

ContentsI)Questions & answers 5

1) Experimental chemistry 5

2) Tests for ions and gases 7

3) The nature of matter 9

4) Mole concept 12

5) Electrolysis 13

6) Chemical energy 14

7) Chemical reactions 15

8) Use of acids, bases and salts 17

9) The periodic table 20

10) Metals 23

11) The Atmosphere 25

12) Organic chemistry 27

II)Industrial processes 32

13) Haber process 32

14) Contact process 32

15) Extraction of aluminium 33

16) Extraction of iron 34

17) Fractional distillation of liquid air 34

18) Water purification 35

III) Tables 36

19) Test for ions 36

20) Test for cations 36

21) Test for gases 37

22) Fractions of crude oil and their use 37

23) Reactivity series of metals 38

24) Common indicators 38

25) Types of oxides 38

26) Solubility rules 39

27) Redox reactions 39

28) Reducing and oxidising agents 39

29) Pollution table 39

P a g e | 5

Questions & answersExperimental chemistry

1) Which apparatus maybe used to: -

a. Hold approximate volume of 100 or 25cm3 of liquid: Beaker or conical flask.

b. Measure liquid with accuracy of 1cm3: Measuring cylinder.

c. Measure liquid with a long scale of 0-50cm3 with accuracy of 0.1cm3: Burette.

d. Measure exact volumes (25, 50cm3): Pipette and burette respectively.

e. Measure volume of gas up to 10cm3: Gas syringe.

2) State the function of gas syringe in the collection of gas: Measure small volumes of dry gas.

3) State the function of upward delivery tube (downward displacement of air): Collection of dry gases that are less dense than air (Mr < 32).

4) State the function of downward delivery tube (upward displacement of air): Collection of dry gases that are denser than air (Mr > 32).

5) State the use of displacement of water: To collect insoluble or slightly soluble gas in water.

6) What are the types of thermometers used in the laboratory:

a. Mercury in glass. To measure high temperatures.

b. Alcohol in glass. To measure low temperatures.

7) Define pure substances: Single element or compound not mixed with anything else.

8) Define element: A substance having one type of atom only.

9) Define compound: A substance that contains two or more different types of atoms chemically bonded together.

10) Define mixture: Two or more substances physically mixed together.

11) Define purification: The separation process of mixtures into pure substances by using physical methods without chemical reactions.

12) Define filtration. Give 1 example: The method of separating an insoluble solid from a liquid or solution. Separation of sand from water.

Imaduddin School Chemistry Department 2018 Organic chemistry

P a g e | 613) Define crystallization. Give 1 example: Separation of soluble solid from a solution.

Separation of sodium chloride from saltwater.

14) Define sublimation. Give 1 example: Separation of a mixture of solids which one of it sublimes. Separation of iodine from sand.

15) Define simple distillation. Give 1 example: Separation of a pure liquid (solvent) from a solution by condensing vaporised liquid. Water from seawater.

16) Define fractional distillation. Give 1 example: Separation of mixture of miscible liquids with varying B.P. Ethanol and water.

17) Define miscible liquids with examples: Liquids that are soluble with each other (ethanol and water).

18) Define immiscible liquids with examples: Liquids that don’t mix with each other (oil & water).

19) Define solution: The liquid formed when a substance (solute) disappears (dissolves) into another substance (solvent).

20) Define saturated solution: A solution in which no more solute can be dissolves at specific temperature.

21) Define soluble: The ability for a solute to be dissolved in a solvent.

22) Define solvent and give examples: The liquid that the solute has dissolved in. (example: Water and ethanol).

23) Define solute: The substance that dissolves in the solvent.

24) Describe the use of separating funnels with examples: Separating two immiscible liquids. Oil and water.

25) Define chromatography. Give 1 example: A method of separating and identifying mixtures. Separation of dyes in ink.

26) Define locating agent: A substance used to identify colorless substances on a chromatogram.

27) Define rf value (retention factor): It is the distance travelled by the substance per distance travelled by solvent.

28) How to check the purity of substance: Check if it has fixed B.P and M.P.


P a g e | 7Tests for ions and gases

29) Describe the tests for each of the following anion with the results:

a. Carbonate (CO32-): Add dilute acid. Effervescence.

b. Chloride (Cl-): Add dilute nitric acid then aq. Silver nitrate. White ppt.

c. Bromide (Br-): Add dilute nitric acid then add aq. Silver nitrate. Cream ppt.

d. Iodide (I-): Add dilute nitric acid then aq. Lead (ii) nitrate. Yellow ppt.

e. Nitrate (NO3-): Add aq. Sodium hydroxide then aluminium foil; warm.

Ammonia produced.

f. Sulphate (SO42-): Add dilute nitric acid then aq. Barium nitrate. White ppt

insoluble in excess.

g. Sulphite (SO32-): Add dilute nitric acid then aq. Barium nitrate. White ppt

soluble in excess.

30) Describe the effect for each of the following cation when aq. Sodium hydroxide is added:

a. Aluminium (Al3+): White ppt. Soluble in excess.

b. Ammonium (NH4+): Ammonia produced on warming.

c. Calcium (Ca2+): White ppt. Insoluble in excess.

d. Copper (ii) (Cu2+): Light blue ppt. Insoluble in excess.

e. Iron (ii) (Fe2+): Green ppt. Insoluble in excess.

f. Iron (iii) (Fe3+): Red-brown ppt. Insoluble in excess.

g. Zinc (Zn2+): White ppt. Soluble in excess.

h. Chromium (Cr3+): Green ppt. Soluble in excess.


P a g e | 831) Describe the effect for each of the following cation when aq. Ammonia is added:

a. Aluminium (Al3+): White ppt. Insoluble in excess.

b. Calcium (Ca2+): No ppt or very slight white ppt.

c. Copper (ii) (Cu2+): Light blue ppt. Insoluble dark blue ppt in excess.

d. Iron (ii) (Fe2+): Green ppt. Insoluble in excess.

e. Iron (iii) (Fe3+): Red-brown ppt. Insoluble in excess.

f. Zinc (Zn2+): White ppt. Soluble excess.

g. Chromium (Cr3+): Green ppt. Insoluble in excess.

32) Describe the tests for the following gases with the results:

a. Ammonia (NH3): Turns damp red litmus paper blue.

b. Carbon dioxide (CO2): Turns limewater milky.

c. Chlorine (Cl2): Bleaches damp litmus paper.

d. Hydrogen (H2): Pops with a lighted splint.

e. Oxygen (O2): Relights a glowing splint.

f. Sulfur dioxide (SO2): Turns acidified aqueous potassium manganate (vii) from purple to colourless.

g. Water vapour (H2O):

i. Turns blue cobalt (ii) chloride paper pink.

ii. Turns anhydrous copper (ii) sulfate from white to blue.


P a g e | 9The nature of matter

33) Define diffusion: Random movement of particles from higher concentration to lower concentration until evenly distributed.

34) What are the factors affecting the rate of diffusion?

a. Temperature: The higher the temperature, the greater the diffusion.

b. Mass of particles: The higher the mass, the slower the diffusion.

35) Explain the following changes of states of matter:

a. Melting: Change from solid to liquid at a fixed temperature.

b. Boiling: Change from liquid to gas a fixed temperature.

c. Condensing: Change from gas to liquid at a fixed temperature.

d. Evaporation: Change from gas to liquid at any temperature.

e. Freezing: Change from liquid to solid at a fixed temperature.

f. Subliming: Change from solid to gas directly.

36) State what happens according to the kinetic particle theory for the following changes of matter:

a. Melting: As the kinetic energy of the particles increase they start to move faster and push each other apart. At the melting point particles breaks away from the attractive forces holding them.

b. Boiling: As the kinetic energy of the particles increase they start to move faster and push each other apart and bounce apart. At the boiling point particles break away from the attractive forces holding them close to each other.

37) State the differences between evaporation and boiling:

a. Evaporation occurs at any temperature but boiling occurs at only fixed temperature.

b. Evaporation occurs at only surface of a liquid but boiling occurs throughout the liquid.

38) State why there is no increase in temperature when melting or boiling: The energy is used to break the bonds between particles.

39) State what is the nucleon number of an atom: The total number of protons and neutrons in the atom. It is also known as the mass number.


P a g e | 1040) Define isotopes: Atoms of the same element with different number of neutrons.

41) Define allotropes: An element that exist as two or more different forms in the same physical state. Eg - graphite and diamond are allotropes of carbon.

42) Define radioactive isotopes: Isotopes with unstable nuclei which break up or disintegrate to give off one or more types of radiation.

43) Define valence electrons: The number of outermost electrons of an atom.

44) Define element with example: A substance that cannot be broken down into simpler substances. Oxygen, O2.

45) Define compounds with example: Substance containing two or more elements chemically joined together. Carbon dioxide, CO2.

46) Define mixtures with example: Substance containing two or more substances physically joined together. Air (mixture of mostly nitrogen (N2), oxygen (O2) and carbon dioxide (CO2)).

47) State the difference between a molecule and an ion: A molecule is a group of two or more atoms joined together. An ion is a charged particle.

48) Define ionic bonding with example: Transfer of electrons from one atom to another to achieve an inert gas configuration. Sodium chloride, NaCl.

a. Metals: Lose electrons. b. Non-metals: Gain electrons.

49) State the properties of ionic compounds:

a. Hard crystalline solids.

b. High B.P & M.P.

c. Strong electrostatic force of attraction.

d. Doesn’t conduct electricity when solid but do in molten/aqueous.

e. Soluble in water.

50) Why ionic compounds have high melting point: Due to strong electrostatic force of attraction between ions, that requires strong heat to break.

51) Define covalent bonding with examples: Sharing a pair/pairs of electrons to gain an electronic configuration of an inert gas. (occurs only between non-metals). Oxygen, O2.

52) Define intermolecular bonds: Attractive forces between covalent molecules.


P a g e | 1153) Why covalent compounds have low melting point: Due to weak intermolecular

force that requires low heat to break.

54) Describe differences between ionic and covalent compounds:

a. Ionic compounds have high M.P/B.P, while covalent has low M.P/B.P.

b. Ionic compound conduct electricity in molten and aqueous state while covalent compunds does not conduct (except graphite).

c. Most ionic compounds soluble in water, but most covalent compounds are soluble in organic solvents like ethanol.

55) Name three compounds with giant covalent structures: Diamond, silicon dioxide and graphite.

56) Why diamond has high melting point: Due to strong covalent bond throughout the giant covalent structure, that requires strong heat to break.

57) Why graphite has high melting point: Due to strong covalent bond within the giant covalent structure, that requires strong heat to break.

58) Why graphite is soft: Contain layers bonded with weak Van Der Waal’s forces that can slide over each other.

59) Why diamond is hard: Contain strong covalent bond throughout the giant molecular structure, arranged in a tetrahedral shape, making it rigid and firm.

60) Describe differences between diamond and graphite.

a. Diamond has tetrahedral shape, while graphite has hexagonal shape.

b. Diamond does not conduct electricity while graphite does.

c. Graphite has layers of molecules, while diamond doesn’t have.

61) Define metallic bond: An electrostatic force of attraction between the mobile ‘sea’ of electrons and the regular array of positive metal ions within the solid metal.

62) Why do metals have high melting point compared to non-metals: Metals have strong electrostatic force of attraction between metal ions and delocalized electrons, which require strong energy, while non-metals have weak intermolecular force.


P a g e | 12Mole concept

63) Define moles: Amount of substance in an object.

64) Define relative atomic mass (RAM): Average mass of an atom.

65) Define relative molecular mass (RMM): Average mass of a molecule of a substance.

66) Define molar mass: The mass of one mole of any substance.

67) Define molecular formula: Shows the actual formula and kinds of atoms present.

68) Define empirical formula: Shows the simplest ratio of the atoms present.

69) Define structural formula: Shows how atoms are joined in the molecule.


P a g e | 13Electrolysis

70) Define electrolysis: Decomposition of an ionic compound with the use of electricity.

71) Define electrolytes: Ionic compound which conducts electric current in molten or aq. Solution being decomposed in the process.

72) Define electrode: A rod or plate where electricity enters or leaves the electrolyte during electrolysis. (reactions occur here)

73) Anode: Electrode at which oxidation occurs.

74) Cathode: Electrode at which reduction occurs.

75) Define anions and cations:

a. Anion: Negative ion.

b. Cation: Positive ion.

76) Write down the reactivity series of cations:K+, Na+, Ca2+, Mg2+, Al3+, (carbon), Zn2+, Fe2+/3+, Sn2+, Pb2+, (hydrogen), Cu2+, Ag+, Au2+, PtPlease Stop Calling Me A (Cute) Zebra In The Library, (Homie). Call Me Something Good Please. Ease of discharge increases down the series

77) Write down the reactivity series of anions:OH-, I-, Br-, Cl-, NO3

2- SO42-, F-

Owls In Britain Can Not Stop FlyingEase of discharge increases down the series

78) Define inert electrodes: Electrodes that do not react with the electrolyte or products during electrolysis (e.g. Platinum and graphite).

79) Define active electrodes: Electrodes that react with the electrolyte or products during electrolysis (e.g. Copper).

80) How do electric cells work? The two electrodes are made up of two different metals with different reactivity. The cathode is more reactive. The further apart they are in reactivity series the greater voltage is produced.


P a g e | 14Chemical energy

81) Define the term endothermic in terms of energy: Absorbing energy.

82) Define the term exothermic in terms of energy: Releasing energy (mostly as heat).

83) Define the term endothermic in terms of bond: Bond breaking.

84) Define the term exothermic in terms of bond: Bond forming.

85) What is the enthalpy change? The energy difference between reactant and products. The symbol is ∆h and measured in kJ/mol (kilojoules per mol).

86) What is the activation energy? Minimum energy needed to start a reaction. It is the energy needed to break the reactant bonds before new bonds are formed. The symbol is Ea and measured in kJ/mol (kilojoules per mol).

87) State two ways in which hydrogen may be produced:

a. Cracking of hydrocarbons.

b. Electrolysis of water.

88) Define fuel cell: A cell that converts chemical energy directly to electrical energy.

89) Describe oxygen-hydrogen fuel cell: Production of electricity by using hydrogen and oxygen as a fuel.

90) Disadvantage of oxygen-hydrogen fuel cell

a. Fuel is difficult to store.

b. Obtaining hydrogen and oxygen is expensive.

91) Advantages of oxygen-hydrogen fuel cell

a. More energy efficient.

b. Does not produce any harmful gas.

c. Electricity generated continuously until fuel depleted.

92) Define photosynthesis: Production of glucose and oxygen by plants, using carbon dioxide, water in the presence of sunlight and chlorophyll.

93) Define respiration: Release of energy, carbon dioxide and water by using oxygen and glucose.


P a g e | 15Chemical reactions

94) Define rate of reaction: Speed for a reactant to be used up or a product to be formed.

95) State the ways to find the rate of reaction:

a. Measuring volume of a gas evolved from a reaction during different time intervals.

b. Measure mass of content of a reaction in a period of time.

96) State the factors affecting the rate of reaction and describe and explain their effect:

a. Surface area of reactant, rate increases when area increases. Because more particles are exposed to the reaction so more frequent successful collisions.

b. Concentration of reactant, rate increases when concentration increases. Because more particles per volume that increase more frequent successful collisions.

c. Pressure of reactant, rate increases when pressure increases. Because same number of particles in less volume results in the increase of successful collisions.

d. Temperature of reactant, rate increases when temperature increases. Because particles gain more kinetic energy that increase more frequent successful collisions.

e. Effect of catalyst, rate increases when catalyst is added. Because it reduces the activation energy without being directly involved in the reaction.

97) Define catalysts: They are chemical substances which alters the rate of reaction without itself being used up in the process. Most catalysts are transition metals or compounds of transition metals.

98) State what is meant by oxidation state: Result of a substance when electrons are lost or gained.

99) Define oxidation: Increase of oxidation state by a substance.

100) Define reduction: Decrease of oxidation state by a substance.

101) Define oxidation and reduction in terms of oxygen loss/gain

a. Oxidation is the gain of oxygen.

b. Reduction is the loss of oxygen.


P a g e | 16102) Define oxidation and reduction in terms of hydrogen loss/gain:

a. Oxidation is the loss of hydrogen.

b. Reduction is the gain of hydrogen.

103) Define oxidation and reduction in terms of electron loss/gain:

a. Oxidation Is the Loss of electron.

b. Reduction Is the Gain of electron.

c. It may be remembered as “OIL RIG”.

104) Define redox reaction: A reaction which involves the two processes of reduction and oxidation simultaneously.

105) Define reducing agent: A substance which lose electron or increase its oxidation number so the other reactant gains the electron/hydrogen or loses oxygen.

106) Define oxidising agent: A substance which gain electron or decrease its oxidation number so the other reactant loses the electron/hydrogen or gains oxygen.

107) Define dynamic equilibrium: State of a reaction when the rate of forward reaction is the same as the rate of backward reaction.

108) Conditions for a reaction possessing dynamic equilibrium:

a. All reactants and products must be in gaseous state.

b. Must be in a closed container.

109) Factors that affect dynamic equilibrium:

a. Pressure b. Temperature c. Concentration

110) Do catalysts affect the dynamic equilibrium? No

111) State what happens to the dynamic equilibrium when:

a. Pressure increases: The reaction favours the side with less moles.

b. Pressure decreases: The reaction favours the side with more moles.

c. Temperature increases: The reaction favours endothermic direction.

d. Temperature decreases: The reaction favours exothermic direction.

e. Concentration of reactant(s) increase: The reaction favours opposite side.

f. Concentration of reactant(s) decreases: The reaction favours the same side.


P a g e | 17Use of acids, bases and salts

112) Physical properties of acids:

a. Sour in taste.

b. Change moist blue litmus to red.

c. Corrosive.

113) Chemical properties of acids:

a. Acid and metal produce salt and hydrogen.

b. Acid and base produce salt and water.

c. Acid and metal carbonates produce salt, water and carbon dioxide.

114) Reaction of acids:

a. Metal hydroxide + acid salt + water

b. Metal oxide + acid salt + water

c. Metal + acid salt + hydrogen

d. Metal carbonate + acid salt + carbon dioxide + water

115) Define strong acids with example: Acid that completely ionises in water. E.g. Hydrochloric acid, HCl.

116) Define weak acids with example: Acids that partially ionises in water. E.g. Ethanoic acid, CH3COOH.

117) State uses for the following acids:

a. Sulfuric acid: Car batteries, manufacture of ammonium sulphate in fertilisers.

b. Hydrochloric acid: Remove rust.

c. Ethanoic acid: Preservation of food.

118) Define bases: Oxides or hydroxides of metals.

119) Define alkalis: Bases soluble in water.

120) Physical properties of alkali:

a. Bitter taste.

b. Slippery feel.

c. Turn hydrated red litmus paper to blue.

d. Corrosive.


P a g e | 18121) Chemical properties of alkali:

a. Alkali and acid produce salt and water.

b. Alkali and ammonium salt produce salt, water, ammonia.

122) Reactions of alkalis

a. Acid + alkali salt + water

b. Alkali + ammonium compounds salt + ammonia + water

c. Alkali + metal ions salt + metal hydroxide

123) State uses for alkalis

a. Neutralise acids in teeth (in toothpaste).

b. In soap and detergents to dissolve grease.

c. Floor and oven cleaners.

124) Define neutralisation: Neutralisation of an acid by a substance to form water.

125) State the solubility rules for the following salts:

a. Group 1 and ammonium salts: All are soluble.

b. Nitrates: All are soluble.

c. Halides: All are soluble except silver halide and lead halide.

d. Sulphates: All are soluble except BaSO4, PbSO4 and CaSO4.

e. Carbonates: All are insoluble except group 1 and ammonium salts.

f. Hydroxides: All are insoluble except group 1, Sr(OH)2 and Ba(OH)2.

126) Define indicators: Substance used to determine whether a substance is acidic, alkaline or neutral.

127) Define universal indicator: Solution that determines the pH of a substance by the verity of colour change.

128) State some common indicators along with their colour change:

a. Litmus (neutral: Purple, acid: Red, alkali: Blue).

b. Methyl orange (neutral: Orange, acid: Red, alkali: Yellow).

c. Phenolphthalein (neutral: Colourless, acid: Colourless, alkali: Pink).

129) Define pH: It is a measure of acidity or alkalinity of a solution.


P a g e | 19130) State the colour of universal indicator in neutral solution with pH respectively:

Green (pH 7).

131) State the colour of universal indicator in acids with pH respectively: Red (pH 1,2,3), orange or yellow (pH 4,5,6).

132) State the colour of universal indicator in alkali with pH respectively: Blue (pH 8,9,10,), indigo (pH 11), violet (pH 12,13,14).

133) What’s meant by an ionic equation: Equation involving ions in aqueous solution, showing formation and changes of ions during the reaction.

134) Define the following oxides

a. Acidic: Oxides of non-metals.

b. Basic: Oxides of metals.

c. Amphoteric: Oxides that react with both acids and alkalis.

d. Neutral: Oxides that don’t react with either acids/alkalis.

135) Dehydrating agent: Substance that removes molecules of water from compounds.


P a g e | 20The periodic table

136) What is a period in the periodic table: A horizontal (--) row of elements in the periodic table. It represents the number of shells an element has.

137) What is a group in the periodic table: A vertical (|) column of elements in the periodic table. It represents the number of valence electrons the element has.

138) What happens to the reactivity of metals down the group: Increases as the size of the particle increases so the distance between nucleus and valence electron increase thus it is easier to lose it.

139) What happens to the reactivity of non-metals down the group: Decreases due to increase of size and distance between nucleus and valence electrons so it is difficult to attract electrons.

140) What happens to the reactivity of metals along the period (left to right): Decreases due to more valence electrons so greater electrostatic force of attraction between nucleus and valence electron making it difficult to lose electrons.

141) What happens to the reactivity of non-metals along the period: Increases due to more valence electrons so greater electrostatic force to attract electrons.

142) Physical properties of alkali metals:

a. Soft (can be cut with a knife).

b. Have low M.P and B.P.

c. Low density compared to other metals.

d. Silvery shine.

143) Chemical properties of alkali metals:

a. Reacts with water to form strong alkali and hydrogen gas.

b. Reacts with acid to form soluble salt and hydrogen gas.

c. Burns in oxygen to form metal oxides.

d. Reacts with halogens to form soluble salts

144) State observation when:

a. Lithium reacts with water: Floats, metal disappears, effervescence.

b. Sodium reacts with water: Metal melts, fizzes, darts around on the surface, yellow sparks.

c. Potassium reacts with water: Cracks, fizzes, pops and burns with a lilac flame.


P a g e | 21145) State the properties of group 7 elements (halogens)

a. Very reactive.

b. 7 valence electrons.

c. Diatomic.

d. Become darker down the group.

e. Low M.P & B.P.

f. Poisonous.

146) What is the state and colour of the following halogens:

a. Fluorine at RTP: Gas. Pale yellow.

b. Chlorine at RTP: Gas. Yellowish green.

c. Bromine at RTP: Liquid. Reddish brown.

d. Iodine at RTP: Solid. Black.

147) State the uses of the following halogens:

a. Fluorine: Added to toothpaste as fluoride ions.

b. Chlorine: Added to swimming pools/water treatment to kill pathogens.

c. Bromine: Preparation of drugs and dyes.

d. Iodine: Used as mild antiseptic for wounds.

148) State why a halogen higher up in the group can displace a halogen from an ionic compound: Because it’s more reactive.

149) State why noble gases don’t form compounds: They have a stable electronic configuration.

150) State two other properties of noble gases: Coloured gases, low M.P and B.P.

151) State the properties of transition metals:

a. High M.P & density.

b. Variable oxidation states.

c. Form coloured compounds.

d. Form complex ions.

e. Act as catalysts.


P a g e | 22152) State three uses of transition metals in industrial chemistry

a. Iron: Haber process (Manufacture of ammonia).

b. Vanadium(v) oxide: Contact process to manufacture sulfuric acid.

c. Nickel: Hydrogenation of alkenes to form saturated fats (e.g. Margarine).

153) State the advantages of using transition metals in industrial chemistry:

a. Save time in manufacture due to catalysing reactions.

b. Lower cost as it lowers activation energy for reactions


P a g e | 23Metals

154) State the properties of metals:

a. Malleable and ductile.

b. Conduct electricity.

c. Conduct heat easily.

d. High M.P and B.P.

155) What is meant by malleable and ductile:

a. Malleable: Can be bent.

b. Ductile: Can be stretched.

156) Define alloys: A mixture of metallic compounds with metallic or non-metallic compounds.

157) State some alloys and the elements present in them:

a. Steel: Mixture of iron, little carbon and trace elements.

b. Brass: Copper and zinc. c. Bronze: Copper and tin.

158) State what is meant by a displacement reaction: Displacement of ions from compounds of elements lower in reactivity series by elements higher in reactivity series.

159) Define minerals: Elements/compounds that make up rocks.

160) Define metal ore: Rock containing metal.

161) State why aluminium is used for window frames and airplanes: Light weight, non-toxic, cheap, resist corrosion and strong.

162) What are the benefits of recycling:

a. Help conserve metals.

b. Saves cost of extraction.

c. Benefit environment.

163) Problems with recycling:

a. Some metals are cheaper to extract than to recycle. E.g. Iron.

b. May damage environment due to fumes produced by smelting.

c. Separation of metals from waste is high.

d. Transport cost.


P a g e | 24164) State the iron ore used in the blast furnace: Haematite (it consists of Fe2O3 and

other impurities such as sand).

165) Why may iron from blast furnace not be good for some applications:

a. Contains impurities making it brittle.

b. Cannot be bent or stretched.

166) Advantages of steel:

a. Strong and tough.

b. Can be bent and stretched.

167) Define rusting with its conditions: Corrosion of iron and steel. Air and water are required.

168) What is rust: A brown solid product formed consisting of hydrated iron (ii) oxide.

169) How is rusting prevented:

a. Surface protection.

b. Sacrificial protection.

c. Use of stainless steel.

170) State how surface protection prevents rusting: The surface is covered so it stops air and water from reaching the iron or steel underneath.

171) State how sacrificial protection prevents rusting: A more reactive metal block is attached to the iron surface so the more reactive reacts rather than the iron. (usually magnesium or zinc is attached).


P a g e | 25The Atmosphere

172) What is the atmosphere? A layer of air containing a mixture of several gases.

173) What is combustion: Substances reacting with oxygen gas in exothermic reactions.

174) State 3 uses of oxygen:

a. As rocket fuel.

b. In steel making, to burn off impurities.

c. In oxy-acetylene cutting and welding.

175) Define pollutants: Substances in the atmosphere which are harmful for living things and environment.

176) Name 7 pollutants and their sources:

a. Carbon monoxide: Exhaust fumes, unburnt hydrocarbons, forest fires.

b. Sulfur dioxide: Volcanic eruptions, combustion of fossil fuels containing sulfur impurities.

c. Oxides of nitrogen: Lightning, forest fires, internal combustion of engines, power stations.

d. Methane: Decomposition of organic matter, natural gas, mines.

e. Unburnt hydrocarbons: Internal combustion of engines, incomplete combustion of hydrocarbons.

f. Ozone: Reaction of nitrogen oxides with volatile organic compounds in uv radiation.

g. Lead compounds: Combustion of leaded petrol.

177) Name the effects of the 7 pollutants:

a. Carbon monoxide: Respiratory poison.

b. Sulfur dioxide: Leads to acid rain, irritates eyes/lungs.

c. Oxides of nitrogen: Eutrophication, lung damage, acid rain.

d. Methane: Global warming.

e. Unburnt hydrocarbons: Carcinogenic, form photochemical smog.

f. Ozone: Reacts with unburn hydrocarbons to form photochemical smog.

g. Lead compounds: Lead poisoning, brain damage.


P a g e | 26178) Define photochemical smog: It is a hazy brown air, which is a mixture of fog and

smog, that reduces visibility, causes eye irritation and breathing difficulties.

179) How is photochemical smog formed: It is produced by reaction between NO2 and O2

in the presence of sunlight to form NO, O and O2. This reaction is called photochemical reaction. The oxygen atom then reacts with another oxygen to form ozone which can react with unburnt hydrocarbons to produce eye-irritating substances.

180) Why does acid rain occur? Acidic gases in the air dissolves into the rain, making it acidic.

181) State 3 effects of acid rain: Damage trees, increase soil acidity, acidifies lakes.

182) Define flue gas: Exhaust gas that contain high percentage of sulfur dioxide gas.

183) Define flue gas desulfurization: Removal of sulfur dioxide from flue gas by using calcium carbonate or calcium hydroxide.

184) Define the greenhouse effect: Trapping of heat from sun by greenhouse gases to regulate earth temperature so that not all heat is reradiated back to space.

185) State the examples of greenhouse gases: Carbon dioxide, methane.

186) Define global warming: Increase in temperature of earth’s atmosphere due to trapping of heat by greenhouse gases.

187) State how global warming can be reduced:

a. Reduce the use of fossil fuels.

b. Use alternative forms of energy such as wind, tidal and hydroelectric power.

c. Use more electric vehicles.

d. Reduce number of cars on road.

e. Create efficient engines in cars to ensure complete hydrocarbon combustion.

188) What is the carbon cycle? Carbon cycle is the removal of carbon dioxide by plants by photosynthesis and the replacement of these carbon molecules by combustion, respiration and natural processes.


P a g e | 27Organic chemistry

189) State the characteristics of organic compounds:

a. All contain carbon element.

b. Most come with hydrogen.

c. Others with oxygen, nitrogen or a halogen.

190) Hydrocarbons: Compounds containing hydrogen and carbon atoms only.

191) Define homologous series: Special group of atoms available in homologous series compounds which are responsible for the chemical properties of the compound.

a. With same general formula.

b. Similar chemical properties.

c. Follow a trend in physical properties.

d. Each corresponding compound differ by CH2.

192) Define saturated hydrocarbons: A hydrocarbon without any double bonds between carbon molecules.

193) Define unsaturated hydrocarbons: A hydrocarbon containing a carbon to carbon double or triple bonds.

194) State the test and observation for unsaturated hydrocarbons: Mix bromine solution with the hydrocarbon. Reddish-brown colour of bromine disappears if it is unsaturated. But the colour remains if it is saturated.

195) Define substitution reaction: A reaction in which an atom or group of atoms are replaced by another.

196) Define addition reaction: A reaction in which one molecule combines with another to form a larger molecule with no other products.

197) Define hydrogenation: Addition of hydrogen.

198) Define hydration: Addition of water.

199) Define alkanes: Homologous series with the general formula of CnH2n+2.


P a g e | 28200) State the properties of alkanes:

a. Alkanes are insoluble in water but soluble in organic solvents.

b. Alkane density increases down the series.

c. Alkanes become more viscous (uneasy to flow) going down the series as they longer.

d. Become less flammable down the series.

e. Alkanes are unreactive with either metals, water, acids or bases.

201) State three ways in which alkanes may react:

a. Combustion.

b. Halogenation (reaction with halogens).

c. Substitution reaction.

202) Define alkenes: Homologous series with the formula of CnH2n.

203) State the combustion reactions of alkanes/alkenes: Alkane/alkene + oxygen carbon dioxide + water

204) State 4 ways in which alkenes may react:

a. Combustion.

b. Halogenation (reaction with halogens).

c. Substitution reaction.

d. Addition reaction.

205) Define polyunsaturated foods: Food containing C=C bond in their molecules. E.g: Vegetable oil.

206) Define cracking: Breaking down of hydrocarbons to produce smaller hydrocarbons.

207) State the conditions for cracking:

Zeolite as catalyst. High temperature.

208) State why cracking is required: Due to high demand for smaller hydrocarbons.

209) Define isomers: Compounds with same molecular formula but different structural formula.


P a g e | 29210) Define alcohols: Homologous series with general formula CnH2n+1OH.

a. They have -oh functional group (hydroxyl group).

b. Names end with suffix -ol.

211) State two ways in which alcohol is made:

a. Fermentation of sugars with yeast.

b. Reacting alkene with steam.

212) State the condition for fermentation:

a. 37°C of temperature.

b. Anaerobic conditions.

213) State the reactions of alcohols:

a. Combustion.

b. Oxidation.

c. Esterification.

214) State the conditions for reacting alkene with steam to form alcohols

a. Temperature of 300°C.

b. Pressure of 65atm.

c. Phosphoric acid.

215) Define carboxylic acids: Homologous series with general formula CnH2n+1COOH (first series n=0).

216) State why carboxylic acids are considered to be weak acids: They only partially ionise in water.

217) State two ways in which carboxylic acids may be prepared:

a. Oxidation of alcohols.

b. Catalytic reaction of natural gas with air.

218) Define esters: Organic compound made from carboxylic acid and alcohol with the removal of one molecule of water.

219) Define esterification: Process that carboxylic acid reacts with alcohol by a reversible reaction on heating and at the presence of concentrated sulfuric acid catalyst.

220) Define hydrolysis: Breaking down of a compound with the addition of water.


P a g e | 30221) What are the steps for hydrolysis of esters: Add sodium hydroxide and heat.

222) Define macromolecule: A large molecule made by joining together many small molecules.

223) Define polymer: A long-chain macromolecule made by joining together many monomers.

224) Define polymerisation: The addition of monomers to make one large polymer.

225) Define addition polymerisation: Type of polymerisation in which unsaturated monomers join up by addition reaction.

226) Define repeat unit: The simplest part of the polymer which is repeated many times to form the polymer.

227) Define condensation polymerisation: Type of polymerisation in which monomers join up by the elimination of small molecules like H2O, HCl.

228) State how nylon is formed: Dicarboxylic acid and diamine undergo condensation polymerisation with the removal of water.

229) State the type of linkage found between monomers in nylon: Amide linkage.

230) State the elements found in an amide linkage: Carbon, hydrogen, oxygen, nitrogen.

231) State why nylon is called a polyamide: Because of the amide linkage between monomers.

232) State 4 uses of nylon:

a. Manufacture of garments.

b. Manufacture of parachutes/tents.

c. Fishing lines.

d. Rugs and carpets.

233) State how terylene is formed: Dicarboxylic acid (acid with 2 –COOH groups) and diol (alcohol with 2 –OH groups) undergo condensation polymerisation.

234) State the elements involved in an ester linkage: Carbon and oxygen.


P a g e | 31235) State the problems associated with plastic

a. Plastics are non-biodegradable.

b. They cannot be decomposed by bacteria.

c. Plastics produce toxic gas when burnt and this contributes to acid rain.

d. Plastics produce carbon dioxide when burnt that increases global warming.

e. Plastics that require cfc during production may contribute to global warming when the cfc is allowed to escape.

236) Carbohydrates: Molecules made up of only carbon, hydrogen and oxygen.

237) State the general formula for carbohydrates: Cn(H2O)n.

238) State the monomers that form starch: Glucose.

239) State how to hydrolyse starch: Heat with sulfuric acid.

240) Proteins: Polymer made up of only amino acids joined together.

241) State why proteins are called polyamides: The monomers are linked by the amide linkage.

242) State how hydrolysis of proteins may occur: Heating with sulfuric acid.

243) Fats: Polymer made up of one molecule of glycerol and 3 molecules of fatty acids with ester linkage.

244) State how the hydrolysis of fats occurs: Heat with an alkali (e.g. Sodium hydroxide).


34) P a g e | 32

Industrial processesHaber process

245) Define Haber process: Process for manufacturing ammonia industrially.

246) State the raw materials required for Haber process and how they are obtained:

a. Nitrogen – fractional distillation of liquid air.

b. Hydrogen – cracking of hydrocarbons or by reacting methane with steam.

247) Write down the equation for the production of ammonia:N2 + 3H ⇌ 2NH3

248) State the conditions for Haber process:

a. Temperature: 450°C.

b. Pressure: 200atm.

c. Catalyst: Iron.

249) State the uses of ammonia: Fertilisers, explosives, nitric acid.

Contact process250) Define contact process: Process for manufacturing sulfuric acid.

251) State the raw materials required for contact process:

a. Sulfur.

b. Air (oxygen).

c. Water.

d. Sulfuric acid.

252) Write down the reactions for contact process

a. S + O2 SO2

b. 2SO2 + O2 ⇌ 2SO3conditions: 450°C, 1-2atm, vanadium pentoxide catalyst.

c. SO3 + H2SO4 H2S2O7

d. H2S2O7 + H2O 2H2SO4

253) State why sulfur trioxide is not dissolved directly in water to produce sulfuric acid: It will produce a corrosive mist of conc. Sulfuric acid which is difficult to deal with.

254) State uses of sulfuric acid: Fertilisers, detergents, dyes, car batteries.Imaduddin School Chemistry Department 2018 Tables

34) P a g e | 33Extraction of aluminium

255) State the ore from which aluminium is extracted from: Bauxite (mainly Al2O3).

256) State how the aluminium ore is purified: Using conc. Sodium hydroxide.

257) State how the electrolysis of the ore is carried out:

a. It is carried out in a steel tank lined with graphite (cathode) with blocks of graphite (anode) dipping in from above.

b. Electrolyte used is: Aluminium oxide dissolved in molten cryolite.

258) State why aluminium oxide is dissolved in cryolite: It decreases the M.P of aluminium oxide.

259) State the half and overall equations of the electrolysis:

a. Cathode: Al3+ + 3e- Al

b. Anode: 2O2- O2 + 4e-

c. Overall: 2Al2O3 4Al + 3O2

260) State why the anode needs to be replaced regularly: The oxygen formed reacts with the anode to form co2. Thus, using it up.

261) State the uses of aluminium:

a. Aircraft body parts.

b. Overhead power cables.

c. Ship superstructures.

Imaduddin School Chemistry Department 2018 Tables

34) P a g e | 34Extraction of iron

262) State the raw materials required for extraction of iron:

a. Haematite (mainly iron (iii) oxide).

b. Limestone (calcium carbonate).

c. Coke (carbon).

d. Air (oxygen).

263) State how the iron is extracted: Reduction of the ore with carbon or carbon monoxide in a blast furnace.

264) State the reactions involved in the extraction of iron:

a. C + O2 CO2

b. CO2 + C 2CO

c. Fe2O3 + 3CO 2Fe + 3CO2 or 2Fe2O3 + 3C 4Fe + 3CO2

265) State how the impurities in the iron made in the blast furnace are removed: By adding limestone.

a. CaCO3 CaO + CO2

b. CaO + SiO2 CaSiO3

slag (CaSiO3) is less dense so it floats on top of the iron preventing the iron from being re-oxidised.

Fractional distillation of liquid air266) State the steps of fractional distillation of liquid air:

a. First air is filtered to get rid of dust particles.

b. Next carbon dioxide and water are removed as they would freeze and block the pipes.

c. Then the air is compressed and expanded repeatedly until it liquefies (at about (-200°C). Helium and neon do not liquefy.

d. The liquid air is passed into a fractionating column where it is slowly warmed up. The gases then boil off one by one starting with the one with the lowest B.P.


34) P a g e | 35Water purification

267) State the steps for water purification:

a. Water from rain and river downstream is collected in reservoir.

b. Water is transported via pipe to a tank where chemicals are added to water so that clay coagulates.

c. Water is moved to sedimentation tank where sedimentation occurs.

d. Water is filtered off in filtration tank, where there are sand particles filter which traps the remaining solid particles in water.

e. Chlorine and carbon are added into the tank.

268) State why carbon is added to the tank: To remove taste and smell (odour).

269) State why chlorine is added to the tank: To kill the pathogens present in the water.

270) Define coagulation: Tiny particles joining together to form large lumps of solid.

271) Define sedimentation: Settling down of particles onto the bottom of a tank.


34) P a g e | 36

Tables Test for ions

Anion Test Result Carbonate (CO3

2-)Add dilute acid. Effervescence, carbon

dioxide produced.Chloride (Cl-) Acidify with dilute nitric acid, then add

aqueous silver nitrate.White ppt.

Bromide (Br-) Acidify with dilute nitric acid, then add aqueous silver nitrate.

Cream ppt.

Iodide (I-) Acidify with dilute nitric acid, then add aqueous lead (ii) nitrate.

Yellow ppt.

Nitrate (NO3-) Add aqueous sodium hydroxide, then

aluminium foil and warm.Ammonia produced.

Sulphate (SO42-) Acidify with dilute nitric acid, then add

aqueous barium nitrate.White ppt insoluble in excess.

Sulphite (SO32-) Add dilute nitric acid then aq. Barium

nitrate.White ppt soluble in excess.

Test for cationsCation Effect of aq. Sodium hydroxide Effect of aq. AmmoniaAluminium (Al3+) White ppt. Soluble in excess. White ppt. Insoluble in excess.Ammonium (NH4

+)Ammonia produced on warming. None.

Calcium (Ca2+) White ppt. Insoluble in excess. No ppt or very slight white ppt.

Copper (ii) (Cu2+) Light blue ppt, insoluble in excess.

Light blue ppt insoluble in excess.

Iron (ii) (Fe2+) Green ppt, insoluble in excess. Green ppt. Insoluble in excess.Iron (iii) (Fe3+) Red-brown ppt, insoluble in

excess.Red-brown ppt. Soluble in excess.

Zinc (Zn2+) White ppt. Soluble in excess. White ppt. Soluble excess.Chromium (Cr3+) Green ppt. Soluble in excess. Green ppt. Insoluble in excess.


34) P a g e | 37Test for gases

Gas Test and resultAmmonia (NH3) Turns damp red litmus paper blue.Carbon dioxide (CO2) Turns limewater milky.Chlorine (Cl2) Bleaches damp litmus paper.Hydrogen (H2) Pops with a lighted splint.Oxygen (O2) Relights a glowing splint.Sulfur dioxide (SO2) Turns acidified aqueous potassium manganate (vii) from purple

to colourless.Water vapour (H2O) -Turns blue cobalt (ii) chloride paper pink.

-Turns anhydrous copper (ii) sulphate from white to blue.

Fractions of crude oil and their useFraction UseRefinery gases Heating fuel.

Solvents. Plastic manufacture.

Gasoline Petrol in cars.Naphtha Feedstock for chemical industry.

Producing high octane gasoline.Kerosene or paraffin Jet fuels.

Heating/lighting fuel. Cracked to form petrol.

Gas oil (diesel) Fuel for diesel engines. Cracked to form petrol.

Heavy oil or lubricating oil Fuel for furnaces and ships. Lubricating oil, wax, polish.

Bitumen Make road surfaces.


34) P a g e | 38Reactivity series of metals

Metals Reaction with water

Reaction with dilute

acids

Reduction of metal oxide by Action of heat

on carbonateCarbon HydrogenK

Reac

ts w

ith

cold

wat

er

to p

rodu

ce

met

al

hydr

oxid

e an

d hy

drog

en

gas

Reac

ts

viol

ently

to

form

salt

and

hydr

ogen

ga

s

Not

redu

ced

Not

redu

ced

Does

not

de

com

pos

eNa

Ca

Mg

Reac

ts w

ith st

eam

to

prod

uce

met

al o

xide

an

d hy

drog

en g

as

Reac

ts to

form

salt

and

hydr

ogen

gas

Deco

mpo

sed

to fo

rm m

etal

ox

ide

and

carb

on d

ioxi

de

Al

Zn

Oxi

des o

f the

se m

etal

s are

re

duce

d to

form

the

met

al a

nd

carb

on d

ioxi

de

Fe

Sn

Pb

Cu

Do n

ot re

act

with

wat

er

Do n

ot re

act

Oxi

des r

educ

ed

to fo

rm m

etal

an

d w

ater

va

pourAg

Deco

mpo

sed

to fr

om

met

al,

carb

on

diox

ide

and Au

Pt

Common indicatorsIndicator Acid Neutral AlkaliLitmus solution Red Purple BlueMethyl orange Red Orange YellowPhenolphthalein Colourless Colourless PinkUniversal indicator Red, orange, yellow Green Blue, indigo, violet

Types of oxidesAcidic oxides Basic oxides Amphoteric oxides Neutral oxidesOxides of non-metals which react with bases to form salts.

Oxides of metals which reacts with acids to form salts.

Oxides of metals which react with both acids and alkalis to form salts.

Oxides of non-metals which do not react with either acids or bases.

E.g. SO2, NO2, CO2 E.g. Na2O, CaO, CuO E.g. Al2O3, ZnO, PbO E.g. H2O, NO, CO


34) P a g e | 39Solubility rules

Compound Soluble InsolubleGroup 1 and ammonium All NoneNitrates All NoneChlorides Most AgCl, Pbcl2

Sulphates Most BaSO4, PbSO4, CaSO4

Carbonates All group 1 and ammonium MostHydroxides All group 1, Sr(OH)2, Ba(OH)2 Most

Redox reactionsOxidation ReductionIncrease of oxidation state Decrease of oxidation stateGaining oxygen Losing oxygenLosing hydrogen Gaining hydrogenLosing electrons Gaining electrons

Reducing and oxidising agentsOxidising agent: Acidified potassium

permanganatePurple to colourless Reduced in

reactionsReducing agent: Potassium iodide Colourless to brown Oxidised in

reactionsPollution table

Pollutant Sources EffectsSulfur dioxide Burning fossil fuels containing

sulfur. Volcanic eruptions.

Irritates the eyes.Cause breathing problems.Causes acid rain.

Oxides of nitrogen Lightning. Forest fires. Internal combustion in engines.

Irritate/damage lungs.Leads to acid rain.Form ozone in presence of sunlight.

Carbon monoxide Incomplete combustion of fossil fuel.

Cause respiratory poison.

Methane Decay of organic matter. Contributes to global warming.

Unburnt hydrocarbons

Internal combustion of engines. Can cause cancer.React to form ozone.

Ozone Photochemical reactions. Irritate eyes, nose & throat.Damage crops.

Lead Use of fuel containing lead. Causes brain damage.CFC Used in refrigerators, air

conditioners, aerosols.Damages ozone layer.


34) P a g e | 40


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