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Page 1 Pure Chemistry Syllabus Paper Topic 1: Experimental Chemistry 5072 01/02 1. The International System of Units is a system of different units used to measure quantities of different things. Quantity SI unit Instrument Time Second (s) Minute (min) and hour (hr) Temperature Kelvin (K) Mercury thermometer Mass Kilogram (kg) Electronic balance Volume m 3 Beaker Estimate volume of liquid Measuring cylinder Measure to nearest cm 3 Burette Measures to nearest 0.1cm 3 Pipette Measures fixed volumes of liquids (25.0 cm 3 ) 2. Collecting and Measuring Volumes of Gases a. How a gas is collected depends on solubility and density. b. Solubility and density of common gases: Ammonia Extremely soluble Hydrogen Chloride Denser Hydrogen chloride Very soluble Sulfur Dioxide Denser Sulfur Dioxide Very soluble Carbon Dioxide Denser Oxygen Very slightly soluble Chlorine Denser Carbon Dioxide Slightly soluble Oxygen Slightly denser Chlorine Soluble Ammonia Less Dense Hydrogen Not Soluble Hydrogen Less Dense c. There are 3 methods for collecting gases: Displacement of water: Suitable for collecting gases insoluble or slightly soluble in water. (e.g. Carbon dioxide, hydrogen, oxygen) Downward delivery: Used to collect gases that are soluble in water and denser than air such as chlorine and hydrogen chloride. Upward delivery: Used to collect gases soluble in water and less dense than air, such as ammonia. d. To collect a dry sample of a gas, we pass it through a drying agent. Some examples include:

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Page 1: Chemistry Finalized Notes

Page 1 Pure Chemistry Syllabus PaperTopic 1: Experimental Chemistry 5072 01/02

1. The International System of Units is a system of different units used to measure quantities of different things.

Quantity SI unit InstrumentTime Second (s) Minute (min) and hour (hr)Temperature Kelvin (K) Mercury thermometerMass Kilogram (kg) Electronic balanceVolume m3 Beaker Estimate volume of liquid

Measuring cylinder Measure to nearest cm3

Burette Measures to nearest 0.1cm3

Pipette Measures fixed volumes of liquids (25.0 cm3)

2. Collecting and Measuring Volumes of Gasesa. How a gas is collected depends on solubility and density.b. Solubility and density of common gases:

Ammonia Extremely soluble Hydrogen Chloride DenserHydrogen chloride Very soluble Sulfur Dioxide DenserSulfur Dioxide Very soluble Carbon Dioxide DenserOxygen Very slightly soluble Chlorine DenserCarbon Dioxide Slightly soluble Oxygen Slightly denserChlorine Soluble Ammonia Less DenseHydrogen Not Soluble Hydrogen Less Dense

c. There are 3 methods for collecting gases:

Displacement of water: Suitable for collecting gases insoluble or slightly soluble in water. (e.g. Carbon dioxide, hydrogen, oxygen)

Downward delivery: Used to collect gases that are soluble in water and denser than air such as chlorine and hydrogen chloride.

Upward delivery: Used to collect gases soluble in water and less dense than air, such as ammonia.

d. To collect a dry sample of a gas, we pass it through a drying agent. Some examples include:i. Concentrated sulphuric acid (Note: Cannot be used

to dry ammonia as it reacts with acid),ii. Quicklime (calcium oxide), andiii. Fused calcium oxide.iv. Apparatus required to dry a gas

e. A gas syringe is used to measure the volume of a gas.

3. A pure substance is made up of only 1 substance and is not mixed with other substance.a. A mixture is a substance that contains 2 or more substances not chemically combined.b. A solid is pure if it has an exact and constant (or fixed) melting point.c. Impurities can affect the melting point:

i. Impurities lower the melting point. The greater the amount of impurities, the lower the temperatureii. They cause melting to occur over a range of temperatures. iii. Impurities increase the boiling point.

Drying a gas.

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Page 2 Pure Chemistry Syllabus PaperTopic 1: Experimental Chemistry 5072 01/02

iv. They cause boiling to occur over a range of temperatures.d. If the pressure acting on a liquid is increased, the boiling point of the liquid is raised.

4. Chromatography is the technique of using a solvent to separate a mixture into its components.a. Paper chromatography can be used to separate dyes in ink, identify poisons or drugs and detect traces of

banned substances in food.b. Method for paper chromatography:

i. A spot of food colouring is applied to the chromatography paper.ii. Once the chromatography paper is dipped in the solvent, it soaks up the solvent which dissolves the dyes.iii. The solvent continues to travel up the paper and carry the dyes along.iv. A dye which is very soluble will be carried far along the paper.

c. The chromatography paper with the separated components is known as a chromatogram.i. If a substance is pure, there is only 1 spot on the chromatogram.

d. The Rf value of a substance is known as the ratio between the distance travelled by the substance and the distance travelled by the solvent. It does not change as long as the chromatography is carried out with the same solvent and the same temperature.

e. Using paper chromatography to analyse a sample:i. Paper chromatography is used to separate the dyes in the sampleii. Each dye can be identified by comparing its position in the chromatogram with that of a known dye.iii. Each dye can also be identified by comparing its Rf value with that of a known dye.iv. Chemists can check if the dyes are permitted for use in food.v. If the dots on the chromatogram corresponds to the dots of any poisonous sample, the substance is

poisonous and should not be consumed.f. Chromatography can so be used for colourless substances like amino acids. A locating agent is sprayed

on the chromatogram.g. Chromatography can be used for:

i. Separate the components in a sample,ii. Identify the number of components in a sample,iii. Identify the components present in a sample, iv. Determine whether a sample is pure.

5. Separation Techniques refer to methods used to separate a mixture into pure substances.

Decanting Used to separate a dense, insoluble solid from a liquid.Filtration Used to separate small solid particles/precipitates from a liquid. Upon filtration, the solid that

remains on the filter paper is called the residue, the liquid or solution that pass through the filter paper is called the filtrate.

Evaporation to Dryness

Used to separate a soluble salt from a liquid. The mixture is heated to dryness to recover salt. However many substances decompose when they are heated strongly.

Crystallization Water is removed by heating the solution. Heating is stopped at the stage whereby a hot saturated solution is formed. The resulting solution is then allowed to cool to room temperature to form pure crystals.

6. Separating solids:a. Filtration can be used to separate a mixture of 2 solids if one is soluble in a solvent and the other is insoluble.

i. The insoluble solid is removed by filtration first.ii. The filtrate is then heated to dryness / crystallized to get the salt.

b. A magnet can be used to separate magnetic materials like iron, steel, nickel and cobalt.c. Sublimation can be used to separate a substance that sublimes from one that does not.

7. Simple Distillation can be used to separate a liquid from a solution. Distillation is the process of boiling the liquid and condensing the vapour. Steps for distillation:a. Thermometer should be placed beside the side arm of the distillation flask to measure the boiling point of the

substance.b. Cold running water should enter the water jacket from the bottom of the condenser and leave at the top.c. Condenser slopes downwards to allow pure solvent to run into the receiver.

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Page 3 Pure Chemistry Syllabus PaperTopic 1: Experimental Chemistry 5072 01/02

d. If the distillate is volatile, the receiver can be put in a large container filled with ice.e. The receiver should not be covered to prevent pressure from acting up in the container.

8. Separating immiscible liquids: Liquids that do not dissolve in each other are described as immiscible. A separating funnel is used to separate immiscible liquids.

9. Miscible liquids are separated by fractional distillation (e.g. water and ethanol). a. Many glass beads are placed in the fractionating column to provide a large surface vapour to condense the

vapour. b. During factional distillation, the liquid with the lowest boiling point will distil over the condenser first and the

vapour of liquids with higher boiling points condenses along the fractionating column and re-enter the round-bottomed flask.

c. Fractional distillation is used in industries to obtain nitrogen, argon and oxygen from air, separate mixtures from crude oil and separate ethanol from glucose solution after fermentation.

Separation of crude oil into its various components (Memorise the order!)10. Identification of cations and anions:

Cation Sodium Hydroxide solution, NaOH Aqueous Ammonia, NH3

Cu2+ A light blue precipitate is formed which is insoluble in excess.

A light blue precipitate is formed which dissolves in excess to form a deep blue solution.

Fe2+ A green precipitate is formed which is insoluble in excess.

A green precipitate is formed which is insoluble in excess.

Fe3+ A reddish-brown precipitate is formed which is insoluble in excess.

A reddish-brown precipitate is formed which is insoluble in excess.

Ca2+ A white precipitate is formed which is insoluble in excess. No precipitate.Zn2+ A white precipitate is formed which dissolves in excess to

form a colourless solution.A white precipitate is formed which dissolves in excess to form a colourless solution.

Pb2+ A white precipitate is formed which dissolves in excess to form a colourless solution.

A white precipitate is formed which is insoluble in excess.

Al3+ A white precipitate formed which dissolves in excess to give a colourless solution.

A white precipitate is formed which is insoluble in excess.

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Page 4 Pure Chemistry Syllabus PaperTopic 1: Experimental Chemistry 5072 01/02

NH4+ No precipitate is formed. On heating, ammonia gas is given

off. Ammonia turns moist red litmus paper blue.

a. N.B. To distinguish Pb2+, add aqueous potassium iodide to the original salt solution. The solution containing Pb2+ gives a yellow precipitate.

b. The precipitate in each of the reactions is the hydroxide of the metal ion. c. Some of the compounds dissolve in excess sodium hydroxide or aqueous ammonia. This is due to the

formation of compounds that are soluble in water.d. Identifying Anions:

Anion Test Observations and InferenceCO3

2- Add dilute acid. Pass the gas given off into limewater.

Effervescence is observed. Gas given off forms a white precipitate with limewater.

Cl- Add dilute nitric acid and then add silver nitrate. A white precipitate of silver chloride is formed.I- Add dilute nitric acid and then add silver nitrate. A yellow precipitate of silver iodide is formed.NO3

- Add dilute sodium hydroxide, add aluminium foil, warm.

Ammonia gas is formed. Most red litmus turn blue.

SO42- Add dilute nitric acid and then add barium nitrate. A white precipitate of barium sulphate is formed.

e. Identifying Gases:

Gases Colour/Odour Test ObservationsHydrogen Colourless and

OdourlessPlace a lighted splint at mouth of test tube.

The lighted splint is extinguished with a ‘pop’ sound.

Oxygen Colourless and Odourless

Insert a glowing splint into the test tube. The glowing splint is rekindled (catches fire)

Carbon Dioxide

Colourless and Odourless

Bubble gas through limewater A white precipitate is formed.

Chlorine Greenish-yellow gas with a pungent smell

Place a piece of moist blue litmus at the mouth of the test tube

The moist blue litmus paper turns red, and then bleaches.

Sulfur dioxide

Colourless gas with a pungent smell

Place a piece of filter paper soaked with acidified potassium dichromate(VI) at the mouth of test tube

The orange potassium dichromate(VI) turns green.

Ammonia Colourless gas with a pungent smell

Place a piece of moist red litmus paper at the mouth of test tube.

The moist red litmus paper turns blue.

f. Testing for presence of water.i. Test with anhydrous copper (II) sulphate: Changes colour from white blue.ii. Test with anhydrous cobalt (II) chloride: Changes colour from blue pink

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Page 5 Pure Chemistry Syllabus PaperTopic 2: The Particulate Nature of Matter 5072 01/02

2. The kinetic particle theory states that all matter is made up of tiny particles and these particles are in constant, random motion. The Kinetic Particle Theory describes the states of matter, explains the differences in the properties of solids, liquids and gases and explains the changes of state.

3. When describing a particular state of matter, you must state thea. Arrangement of Particles / Densityb. Forces between Particlesc. Movement (Particle Motion) of Particles d. Kinetic Energy of Particles

4. The description for the different states of matter is listed below. You need to include the important keywords.

Arrangement Movement Attractive Forces EnergySolid Close together in an orderly

arrangementParticles vibrate about their fixed positions.

Very Strong Very Low

Liquid Close together in a disorderly arrangement

Particles slide over each other Strong Low

Gas Far apart in a random arrangement.

Particles move at great speeds. Very Weak High

Table 2.1: Description for the different states of matter

5. A substance can change its state through 5 different processes:a. Melting [Endothermic]b. Boiling [Endothermic]c. Sublimation [Endothermic] [E.g. Dry Ice]d. Condensation [Exothermic]e. Freezing [Exothermic]

6. The main factor that causes a substance to change its state is the kinetic energy of the particles.a. When matter is heated/cooled, the heat taken in/given out causes the kinetic energy of the particles to

change.b. As a result, they overcome the forces of attraction between them, causing the substance to change its state.

7. Diffusion: Diffusion occurs as air molecules are in constant, random motion.a. Diffusion has 2 definitions:

i. Diffusion is the movement of particles from a region of higher concentration to a region of lower concentration.

ii. Diffusion is the process whereby particles move freely to fill up any available space.b. Brownian Motion can be used to demonstrate diffusion.c. Diffusion can be demonstrated in the laboratory.

i. A cover is placed between 2 gas jars, one containing colourless air and the other containing brown bromine vapour.

ii. After the cover is removed, the gas in both gas jars look the same [The particles of each gas are evenly spread throughout both gas jars] and a homogenous mixture of air and bromine is formed.

d. Factors affecting rate of diffusion:i. The relative molecular mass of the gas (smaller hydrogen particles can diffuse faster compared to larger

air particles)ii. The state of the substance (gases diffuse much faster than liquids as the particles can move about more

freely in gases compared to liquids)iii. The temperature of the substance: As the particles gain more energy when the temperature increases,

thus they can move faster and increase the rate of diffusion.

8. Atomic Structurea. Atoms are made up of protons, neutrons and electrons.

Proton Electron NeutronCarries one positive charge Carries one negative charge Carries no electric chargeRelative mass of 1 Have a neligible mass of 1/1840 Has a relative mass of 1

b. All atoms are electrically neutral. An atom contains an equal number of positively charged protons and negatively charged electrons which cancel each other out.

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Page 6 Pure Chemistry Syllabus PaperTopic 2: The Particulate Nature of Matter 5072 01/02

c. The number of protons in an atom is known as the proton number or atomic number. i. As atoms are neutral, no of protons = no. of electrons.ii. It is the thumbprint of an element as it distinguish one element from another.

d. The total number of protons and neutrons in an atom is called the nucleon number or mass number. i. The mass number is not a whole number as every element consists of many different isotopes with

different number of neutrons, and the mass number is an average based on the mass number and the relative abundance of the isotopes.

ii. The element can be represented using only the nucleon number e.g. sodium-23 or 23Na.

The nucleon/proton numbers can be included. E.g. .e. Isotopes are atoms with the same number of protons but different number of neutrons.

i. Isotopes have the same chemical properties but slightly different physical properties.ii. Chemical properties are similar as chemical reactions involve only the electrons and not protons/neutrons.iii. Physical properties differ as the relative masses of the isotopes differ. For instance, hydrogen-2 has a

slightly higher boiling point and density than hydrogen-1.iv. Isotopes that emit high-energy radiation are called radioisotopes and are classified as radioactive

substances. Radiation emitted is dangerous as it can damage living cells and cause cancer. But they have important applications [e.g. heart pacemakers].

f. Arranging Electrons in atoms. i. The way electrons are arranged in an atom is called the electronic configuration of the atom.ii. The shell furthest away from the nucleus is known as the valence shell. Electrons in this shell are known

as valence electrons.iii. The chemical properties of an element depend on the number of valence electrons.

g. deduce the numbers of protons, neutrons and electrons in atoms and ions given proton and nucleon numbersi. Note: Proton number gives the number of protons and electrons.ii. To find number of neutrons take mass/nucleon number – proton number.

9. Elements, Compounds and Mixturesa. An element is a pure substance that cannot be split up into two or more simpler substances by chemical

processes or by electricity.i. Chemical symbols are used to represent elements.ii. Metalloids are elements that have properties of both metals and non-metals.iii. Physical properties of metals and non-metals (IMPT!)

Metals Non – metalsShiny appearance (lustrous) Dull appearance if solidSolids at r.t.p. [except mercury] Either gases, volatile liquids or solids with low

melting points at r.t.p. [except carbon] Malleable (can be hammered into different shapes

without breaking) Sonorous (make a ringing sound when struck) Ductile (can be drawn into wires)

Brittle if solid (easily broken when hammered)

High melting points and boiling points. [Except for sodium, potassium and mercury]

Low melting points and boiling points [except for carbon and silicon]

Good conductors of heat Poor conductors of heat [except carbon – diamond and graphite]

Good conductors of electricity in all states of matter Poor conductors of electricity [except graphite]

b. Atoms are the smallest particles of an element that have the chemical properties of that element.c. Helium, Neon, Argon, Krypton, Xenon and radon are the only elements that exist as individual atoms and

are known as monatomic elements. [Their atoms are not joined together chemically]d. A molecule is a group of two or more atoms that are chemically joined together.

i. Molecules joined by the combination of 2 atoms are called diatomic molecules.ii. Molecules joined by the combination of 3 atoms are called triatomic molecules.iii. Molecules joined by the combination of 4 or more atoms are called polyatomic molecules.

e. A compound is a pure substance that contains 2 or more elements chemically combined.i. The smallest particle of a compound that can exist independently is called the molecule.ii. A compound can be represented by the chemical formula, which states the types of atoms [elements]

present in the compound and the ratio of the different atoms present in the compound.iii. Heat and/or electricity can be used to form compounds and break down compounds into

elements/simpler compounds known as thermal decomposition.f. A mixture is formed when 2 substances are added together without chemical bonds being formed.

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Page 7 Pure Chemistry Syllabus PaperTopic 2: The Particulate Nature of Matter 5072 01/02

i. A mixture of 2 elements: Neon and hydrogen.ii. A mixture of 2 compounds: Water vapour and carbon dioxideiii. A mixture of 1 element and 1 compound: Nitrogen and Ammonia.iv. An alloy is a mixture of elements with other elements. They are widely used as they tend to eb stronger

than pure metals.E.g. Steel consist of iron and carbon.E.g. Stainless steel consists of iron, chromium, nickel and carbonE.g. Brass consists of copper and zincE.g. Bronze consists of copper and tin.

v. Alloys have similar chemical properties but have different physical properties.

10. Ionic Bondinga. An ion is a charged particle formed from an atom or a group of atoms by the gain/loss of electrons.b. Metals form positively charged ions (cations) and non-metals form negatively-charged ions (anions).c. An ionic compound is formed between a metal and a non-metal.

i. The strong forces holding the cations and anions in the compound are known as ionic bonds (electrostatic forces of attraction) or electrovalent bond.

ii. “Dot and cross” diagrams can be used to show ionic bonding.iii. Ionic compounds form giant ionic structures. The cations and anions are arranged in a giant lattice

structure or a crystal lattice.d. Physical Properties of Ionic Compounds:

i. Volatility (melting/boiling points): The lattice of an ionic compound is held together by strong ionic bonds between the ions. A large amount of energy is hence required to overcome these strong bonds, as a result ionic compounds are solids at r.t.p. and are non-volatile. Many ionic compounds have high melting points and high boiling points.

ii. Solubility: Ionic compounds are usually soluble in water. [exceptions: LS chlorides, BCL sulfates and all carbonates except SPA.] However ionic compounds are insoluble in organic compounds.

iii. Electrical Conductivity: Ionic compounds do not conduct electricity in the solid state as there are no mobile ions, whereas when an ionic compound is dissolved in an aqueous solution it can conduct electricity as there are mobile ions in an aqueous solution.

11. Covalent Bondinga. The bond formed between atoms that share electrons is called a covalent bond.b. A molecule is a group of 2 or more atoms joined together by covalent bonds.

i. The sharing of 2 electrons form a single covalent bond.ii. The sharing of 4 electrons form a double covalent bond or a double bond.

c. Molecules formed by covalent bonding are known as covalent compounds.d. Covalent Molecules may exist as simple molecules or giant structures.e. Simple Molecular Structures : Iodine. Within each iodine molecule, the iodine atoms are held together by

strong covalent bonds. Between the iodine molecules there are weak van der Waals’ forces holding the molecules together.i. Volatility – Melting and boiling points. Many simple covalent substances (esp. small molecules) are

liquids or gases at r.t.p. as they are volatile. Little heat energy is required to overcome the intermolecular forces, hence these covalent structures have low melting and boiling points.

ii. Solubility – Most covalent molecules are insoluble in water and soluble in covalent compounds. Exceptions: Alcohol, sugar, chlorine, hydrogen.

iii. Electrical Conductivity: Most covalent elements/compounds do not conduct electricity in all states. This is so as they do not have mobile ions/electrons to conduct electricity.Exceptions: Carbon [Graphite] conducts electricity. Hydrogen chlorine, sulphur dioxide and ammonia react with water to form solutions that conduct electricity.

f. Giant Molecular Structures: 3 case studies – Diamond, Graphite, Silicon(IV) oxide.g. Diamond is an allotrope of carbon. [Allotrope – different forms of the same element].

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Page 8 Pure Chemistry Syllabus PaperTopic 2: The Particulate Nature of Matter 5072 01/02

i. Diamond consists of a carbon atom joined to 4 other carbon atoms by strong covalent bonds. [The entire molecular structure is made up of covalent bonds]

ii. Physical Properties: Diamond is a very hard substance and is solid at r.t.p. It has a high melting and boiling point. It does not conduct electricity and is insoluble in water.

iii. Lots of heat energy are required to break the strong covalent bonds, hence diamond is hard and difficult to melt. There are also no valence electrons that move through the structure, hence diamond cannot conduct electricity.

iv. Uses of diamond: Gemstones in jewellery, tips of drills/cutting tools.h. Graphite is another allotrope of carbon. It is made up of layers of carbon atoms which are held together by

weak van der Waals’ forces lying on top of each other. i. The bonds within each layer are strong and difficult to break.ii. However the forces of attraction between the layers of carbon are weak and can slide over each other;

hence graphite is soft and slippery.iii. In graphite, each carbon atom has a valence electron not used to form covalent bonds. They are

delocalised electrons and can move along the layers and carry electric charges, hence graphite is a good conductor of electricity.

iv. The uses of graphite depend on several properties:

Slippery, doesn’t decompose at high temperatures, doesn’t attack rubber

Used as a dry lubricant to lubricate machine parts containing rubber.

Soft [layers of carbon atoms can slide off carbon easily]

Graphite is baked with clay and made into pencil lead which flake off and stick to paper when writing.

Good conductor of electricity, fairly unreactive

Graphite is used to make inert electrodes for electrolysis.

i. Silicon(IV) oxide, commonly known as silica has a giant molecular structure as well. It consists of silicon and oxygen atoms linked together in a 3-D structure.i. Each silicon atom is bonded to 4 other oxygen atoms in a tetrahedral structure. Each oxygen atom is

bonded to 2 silicon atoms. Hence the formula is SiO2 [Silicon dioxide]ii. It melts only at high temperatures as all the silicon and oxygen atoms are held by strong covalent bonds.iii. It cannot conduct electricity as there are no mobile ions or electrons present in the structure.iv. It is insoluble in water as well.

12. Metallic Bonding: Metals are strongly held to one another by metallic bonding. In the metal lattice the atoms lose their valence electrons and become positively charged. The valence electrons become delocalised and can move freely between the metal ions.a. The metallic lattice is hence described as a lattice of positiveions surrounded by a ‘sea of mobile electrons’.b. A metallic bond is defined as the force of attraction between positive metal ions and the ‘sea of delocalised

electrons’.c. Physical Properties of Metals:

i. Electrical Conductivity: Metals are good conductors of electricity due to the mobility of the valence electrons within the metal lattice.

ii. Malleability and Ductility: If sufficient force is applied to the metal, one layer of atoms can slide over another without disrupting the metallic bonding. Hence metallic bonds are strong but flexible, so matters are malleable and ductile.

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Page 9 Pure Chemistry Syllabus PaperTopic 3: Formulae, Stoichiometry and the Mole Concept 5072 01/02

Syllabus Outlines:a. state the symbols of the elements and formulae of the compounds mentioned in the syllabusb. deduce the formulae of simple compounds from the relative numbers of atoms present and vice versac. deduce the formulae of ionic compounds from the charges on the ions present and vice versad. interpret chemical equations with state symbolse. construct chemical equations, with state symbols, including ionic equationsf. define relative atomic mass, Arg. define relative molecular mass, Mr, and calculate relative molecular mass (and relative formula mass) as the

sum of relative atomic massesh. calculate the percentage mass of an element in a compound when given appropriate informationi. calculate empirical and molecular formulae from relevant dataj. calculate stoichiometric reacting masses and volumes of gases (one mole of gas occupies 24 dm3 at room

temperature and pressure); calculations involving the idea of limiting reactants may be set (The gas laws and the calculations of gaseous volumes at different temperatures and pressures are not required.)

k. apply the concept of solution concentration (in mol/dm3 or g/dm3) to process the results of volumetric experiments and to solve simple problems (Appropriate guidance will be provided where unfamiliar reactions are involved.)

l. calculate % yield and % purity

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Page 10 Pure Chemistry Syllabus PaperTopic 4: Electrolysis 5072 01/02

1. Electrolysis is the process of using electricity to break down or decompose a compound. The compound is usually molten/dissolved in water.a. Electrolysis takes place in an electrolytic cell which consists of a battery, electrodes and an electrolyte.b. Electrolyte:

i. A molten ionic compound/aqueous solution that conducts electricity.ii. Dissociates to form cations and anions.iii. Ions present in the electrolyte allow electricity to flow through it.

c. Electrodes conduct electric current.i. They are usually carbon (graphite) rods.ii. Electrode connected to the positive terminal is the anode, oxidation always occur here.iii. Electrode connected to the negative terminal is the cathode, reduction always occur here.

d. Battery: It acts as an electron pump, drawing electrons away from the anode [which becomes positively charged] and supply it to the cathode [which becomes negatively charged]

e. When electricity is passed through an electrolyte, chemical reactions take place at the electrode and the electrolyte is decomposed. Reactions taking place at the electrodes are called electrolytic reactions.

f. Electrolytes conduct electricity as they have mobile ions.

Electrical conduction by metals/graphite Electrolytic conduction by electrolytesElectricity is conducted by the flow of electrons and vibration of atoms layer by layer from one end of the conductor to the other end.

Electricity is conducted by the movement of cations and anions across the molten/aqueous electrolyte.

Metals and graphite remains unchanged when an electric current passes through them.

The electrolytes are decomposed to form new substances when they conduct electricity.

2. Explaining Electrolysis:

a. Electrical Energy is converted to chemical energy as chemical/redox reactions take place at the electrodes.b. The process of gaining or losing electrons at the electrodes is called discharge. When ions are discharged at

the electrodes, they form atoms (metals) or molecules (gases).

3. Electrolysis of Molten Ionic Compound: When a molten ionic compound is being electrolysed, the cation goes to the cathode whereas the anion goes to the anode, there is no selective discharge of ions.

4. Inert electrodes are electrodes that do not take part in any chemical reaction during electrolysis. They are used to prevent reactions from occurring between the electrodes/products of electrolysis.a. E.g. Carbon and Platinum Electrodes.

5. Electrolysis of Aqueous solution of compounds: There is a selective discharge of cations and anions.a. In general, when an aqueous solution of an ionic compound is electrolysed, a metal [cation lower than

hydrogen in R.S.] /hydrogen is produced at the cathode. At the anode, a non-metal, oxygen or halogen [if anion is not sulphate, carbonate or nitrate and is concentrated] is given off.

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Page 11 Pure Chemistry Syllabus PaperTopic 4: Electrolysis 5072 01/02

6. Selective Discharge of ions: In electrolysis when there is more than one type of cation or anion present in a solution, only one cation and one anion are preferentially discharged. If inert electrodes are used, the ions discharged and the products formed depend on 3 factors:a. Position of the metal (producing the cation) in the

reactivity series. [If metal is above hydrogen, hydrogen gas is released, else the metal is being deposited on the cathode]

b. The relative ease of discharge of an anion. [Sulfates, nitrates and carbonates will never be discharged]

c. The concentration of the anion in the electrolyte. [If the solution is concentrated, the anion will be discharged preferentially to hydroxide ions]

7. Electrolysis of Aqueous Copper(II) Sulfate:a. Cathode is coated with a layer of reddish-brown solid copper.b. The blue colour of the solution fades gradually as more copper is deposited.c. The electrolyte becomes increasingly acidic as hydroxide ions are being discharged at the anode, hence

resulting in a higher concentration of hydrogen ions left in the electrolyte.d. Anode equation: 4OH- (aq) 2H2O (l) + O2 + 4e-

e. Overall equation: 2CuSO4 (aq) + 2H2O (l) 2Cu (s) + O2 (g) + 2H2SO4 (aq)

8. Electrolysis of Concentrated Sodium Chloride solution:a. One volume of hydrogen gas is given off at the cathode, one volume of chlorine gas produced at anode.b. The resulting solution becomes alkaline as there is a higher concentration of hydroxide ions than hydrogen

ions left in the solution.

9. Electrolysis of Watera. Pure water itself is a poor conductor of electricity as it has few ions in it.b. However if a small amount of an ionic compound is added it becomes a good conductor of electricity.c. The products of the electrolysis of water are always two volumes of hydrogen at the cathode and one volume

of oxygen at the anode.

10. Industrial Applications of Electrolysis:a. Electrolytic purification: Electrolysis can be used to purify metals.

i. Copper electrodes is a reactive electrodes. These reactive electrodes undergo oxidation.ii. When aqueous copper(II) sulphate is electrolysed using copper electrodes, no gas is evolved at the

cathode, blue colour of solution remains unchanged.The copper anode decreases in mass.The copper cathode increases in mass as copper is deposited.

iii. Anode: Cu(s) Cu2+(aq) + 2e-. Anode dissolves to form Cu2+ in aqueous solution.iv. Cathode: Cu2+(aq) + 2e- Cu(s). The cathode becomes coated with a layer of reddish-brown copper.v. Net result: Copper is transferred from the anode to the cathode. Hence to refine copper, impure copper

is placed at the anode of the electrolytic cell. The cathode is a pure sheet of copper and the electrolyte is aqueous copper(II) sulphate.

vi. During electrolysis, impure copper anode dissolves, impurities like silver/platinum fall to the bottom of the cell forming an ‘anode slime’.

b. Electroplating: The process of depositing a layer of metal on another substance using electrolysis.i. Electroplating is used to give a metal object a good decorative finish and to prevent rusting.ii. Object to be plated is placed at the cathode whereas the anode is the source of the plating metal.

11. Simple Cell: It is a device that converts chemical energy to electrical energy. [also known as electric cell]a. It is made by placing 2 different metals in contact with an electrolyte [sulphuric acid].b. The electrons flow out from the cathode [more reactive metal] and flow into the anode [less reactive metal].

Bubbles of hydrogen gas are release from the anode. The electrons which flow in the external circuit constitute the electric current.

c. Overall Equation: Zn(s) + H2SO4(aq) ZnSO4(aq) + H2(g)d. Oxidation and Reduction occur together to cause the flow of electrons, hence electrical energy is produced by

redox reactions in a simple cell.e. The flow of electrons is always from the more reactive metal [cathode] to the less reactive metal [anode].

This is so as more reactive metals tend to give up electrons and form ions more readily.f. If the electrolyte is copper(II) sulphate,

i. Overall ionic equation: Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

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Page 12 Pure Chemistry Syllabus PaperTopic 4: Electrolysis 5072 01/02

ii. The copper anode gains in mass as copper(II) ions from the solution receive electrons to form copper.g. The further apart the 2 metals are in the Reactivity Series, the higher the voltage produced. No voltage is

produced if both electrodes are of the same metal.h. A salt bridge completes the circuit, while keeping the electrolytes separate.

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Page 13 Pure Chemistry Syllabus PaperTopic 5: Energy from Chemicals 5072 01/02

11. Energy cannot be created or destroyed, but can be changed from one from to another.a. Reactions that give out heat energy to the surroundings are called exothermic reactions. During a reaction,

i. Heat is liberated and transferred from the chemicals to the surroundings, andii. The temperature of the reaction mixture rises, container feels hot.iii. Examples of such reactions: Combustion, rusting, corrosion, neutralisation, respiration.

b. Reactions that absorb heat from the surroundings are called endothermic reactions.i. Heat energy is absorbed and is transferred from the surroundings to the reactants,ii. The temperature of reaction mixture falls; container feels cold.iii. Examples of such reactions: Photosynthesis, action of light on silver bromide in photographic flim, thermal

decomposition of carbonates.c. The amount of energy involved in a reaction is called the entalphy change of the reaction.

i. It is represented by the symbol ∆H and is measured in kilojoules (kJ).ii. For an exothermic reaction, ∆H is negative as the chemicals have lost energy to the surroundings.iii. For an endothermic reaction, ∆H is positive as the chemicals have gain energy from the surroundings.

d. Energy changes can be shown by drawing energy level diagrams.i. The difference between the energy levels of the products and the reactants is equal to the amount of

energy given out or absorbed during the reaction.

12. Bond Breaking and Bond Makinga. Bond Breaking requires energy to be absorbed, thus the process is endothermic.b. Bond Making requires energy to be released, thus the process is endothermic.c. A process is endothermic if the energy given out during bond breaking is more than the energy taken in

during bond breaking.d. A process is exothermic if the energy taken in during bond breaking is more than the energy given out

during bond breaking.e. The energy required to break a chemical bond = energy released to form the chemical bond. This is also

known as the bond energy of the bond.f. The stronger the bond, the more energy required to break a bond, the higher the bond energy.

13. Activation Energy, EA refers to the minimum energy that reacting particles must possess in order for a chemical reaction to occur.a. Particles which are able to break their bonds on collision and form new bonds are particles that possess the

activation energy of the reaction.b. Energy profile diagrams can be used to show the activation energy of a reaction.c. A catalyst can help to reduce the activation energy for a reaction, thus increasing the rate of reaction.

14. Combustion of Fuels: Fuels are substances that burn easily in air to give out energy. a. The most commonly used fuels include fossil fuels, petroleum and natural gas.b. When fuels burn, carbon dioxide, water and heat energy is released (exothermic reaction) c. If a limited supply of air is used, carbon particles (soot) and carbon monoxide is produced. This process is

known as incomplete combustion.d. A good fuel must:

i. Release lots of heat energy when burnt,ii. Safe to use and convenient to store,iii. Reasonably cheap and readily available, andiv. Must not produce smoke, unpleasant or poisonous gases.

e. A chemical cell in which reactants (usually fuel and oxygen) are continuously supplied to produce electricity directly is called a fuel cell. Overall reaction is the conversion of oxygen and hydrogen to water.

f. Hydrogen can be a fuel of the future. It is renewable and pollution-free.

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Page 14 Pure Chemistry Syllabus PaperTopic 6: Chemical Reactions 5072 01/02

15. Oxidation and Reductiona. Oxidation occurs when a substance gains oxygen, loses hydrogen, loses electrons or increases its oxidation

state after a reaction.b. Reduction occurs when a substance loses oxygen, gains hydrogen, gains electrons or decrease its

oxidation state after a reaction.c. The oxidation state is the charge an atom of an element have existed as an ion in a compound (even if it is

actually covalently bonded)d. The combined process of oxidation and reduction is known as a redox reaction.e. A substance that causes another substance to be oxidized is called an oxidizing agent. E.g. Group VII

elements, potassium manganate(VII) and potassium dichromate(VI).i. It removes electrons from another substanceii. It is reduced during a reaction.

f. A substance that causes another substance to be reduced is called a reducing agent. E.g. Carbon, Carbon Monoxide, Metals, Potassium Iodide, Sulfur dioxide. i. It gives up electrons to another substance.ii. It is oxidized during a reaction.

g. Aqueous Potassium Iodide (KI) is used to test for the presence of an oxidizing agent. It will turn from colourless to brown as the iodide ion is oxidized to iodine by the oxidizing agent.i. 2I- (aq) I2 (aq) + 2e-

ii. Starch-iodide paper can also be used. It changes colour from white to blue in the presence of an oxidizing agent.

h. Acidified Potassium Dichromate(VI) is used to test for the presence of a reducing agent. It will turn from orange to green in the presence of a reducing agent.i. Half Equation: Cr2O7

2-(aq) + 14H+(aq) + 6e- 2Cr3+(aq) + 7H2O(l)ii. The dichromate(VI) ion is reduced to the chromium ion. It loses oxygen and the oxidation state of

chromium decreases from +6 to +3.\

16. The speed of a reaction can be measured by:a. Amount of reactant used per unit timeb. Amount of product obtained per unit timec. Measuring volume of gas per unit time (if a gas is produced)d. The speed of reaction at any instant is indicated by the gradient of the graphs.e. For 2 particles to collide with each other effectively (a.k.a. effective collisions)

i. The reacting particles must collide with one anotherii. The particles must collide with activation energy.

17. There are 5 factors that can affect the Speed of Reaction: (In general, when any factor increases the rate of effective collisions between reacting particles, it will increase the speed of reaction)a. Concentration of particles: When the solution becomes more concentrated, there are more particles of the

reactant occupying a given volume. The reactant particles can thus collide with one another more frequently. With more collisions, the more effective collisions there are between the particles, hence increasing the speed of reaction.

b. Pressure: The higher the pressure, the closer the gaseous molecules, thus collisions between the gaseous molecules become more frequent, increasing the number of effective collisions, hence increasing the speed of reaction.

c. Particle Size/Surface Area of particles: When a surface area of a solid is reduced, the small particles have a larger total surface area, thus allowing more reactions between the particles, increasing the number of effective collisions, hence increasing the speed of reaction.

d. Temperature: The higher the temperature, the more energy the particles contain, thus more particles have energy greater than the activation energy, increasing the number of effective collisions and thus increasing the rate of reaction.Catalyst: A catalyst is a substance which increases the speed of a chemical reaction and remains chemically unchanged at the end of a reaction.i. Characteristics of catalyst

Only a small amount is needed to speed up the reactionIt lowers the activation energy of a reaction.It increases the speed of the reaction and not the yield, thus the same amount of products is formed regardless of the presence of the catalyst.The physical appearance of the catalyst may change, but it remains chemically unchanged after a reaction.Impurities can prevent catalysts from working, poisoning the catalyst.The catalyst is not used up during the reaction.A catalyst is selective in its action; different catalysts speed up different reactions.Transition metals usually acts as catalysts.

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ii. Examples of catalystsFinely divided iron in the Haber process for manufacturing AmmoniaVanadium(V) oxide in the manufacture of sulphuric acidPlatinum in catalytic convertersAluminium oxide or silicon(IV) oxide in the cracking process for producing hydrogenConcentrated Sulphuric Acid during esterification.Nickel used in hydrogenation (alkenes reacting with hydrogen to form alkanes)

iii. Enzymes are substances that catalyst the chemical reactions in palnts and animals. They are often known as biological enzymes. Properties:

They are proteinsThey are very specific in their actionsThey can be made inactive by heating.They are sensitive to pH changes.

iv. Catalysts work by providing an alternative pathway for the reaction to proceed, which has a lower activation energy from the original reaction, allowing a greater number of effective collisions between the reacting particles, hence increasing the speed of reaction.

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Page 16 Pure Chemistry Syllabus PaperTopic 7: Acids, Bases, Salts 5072 01/02

18. An acid is a substance which produces H+ ions when dissolved in water.a. All acids contain hydrogen but not all compounds containing hydrogen are acids.b. It is the hydrogen ions produced that are responsible for the properties of acids, which are:

i. Acids have a sour tasteii. Acids dissolve in water to form solutions which conducts electricityiii. Acids turn blue litmus rediv. Acids react with reactive metals to form hydrogen and a salt.

Salts are called sulfates when formed from sulphuric acid, nitrates from nitric acid, chlorides from hydrochloric acid and ethanoates from ethanoic acid.A lighted splint is used to test for hydrogen gas.Unreactive metals below hydrogen (copper/silver) does not react with dilute acids.Concentrated nitric acid reacts with metals like copper but gives nitrate, salt and nitrogen dioxide.Lead does not react with dilute hydrochloric/sulphuric acid as an insoluble lead salt is formed and coats around the lead metal, preventing further reaction between the metal and acid.

v. Acids react with carbonates to form a salt, carbon dioxide and water.To test for carbon dioxide, bubble the gas through limewater. A white precipitate is formed in the presence of limewater.Ionic equation: 2H+ + CO3

2- H2O

vi. Acids react with metal oxides and hydroxides to form a salt and water. The salt depends on the type of acid used (see point iv)

c. Acids only show the properties of acids when they are dissolved in water. This is so as acids dissociate in water to produce hydrogen ions which are responsible for the acidic properties.

d. Uses of acids:i. Sulfuric acid: Making detergents, making fertilisers, car batteries.ii. Ethanoic acid: (In vinegar) to preserve food, making adhesives (glue)iii. Hydrochloric acid: In leather processing, for cleaning metals.

19. A base is any metal oxide or hydroxide.a. A base is a substance that reacts with an acid to give a salt and water only.b. The oxide or hydroxide ions from the bases react with the hydrogen ions from the acids to form water.c. An alkali is a soluble base. Some common alkalis include Sodium hydroxide, potassium hydroxide, calcium

hydroxide, barium hydroxide and aqueous ammonia.d. Properties of alkalis:

i. Have a bitter taste and soapy feel.ii. Turn red litmus paper blueiii. Produce hydroxide ions when dissolved in water. E.g. when Ammonia dissolves in water, ammonium ions

and hydroxide ions are formed.

NH3 + H2O NH4+ + OH-

iv. All alkalis react with acids to form a salt and water only. This process is known as neutralisation.Ionic equation: H+ + OH- H2O

e. Alkalis heated with ammonium salts give off ammonia gas.General Equation: Alkali + Ammonium Salt Ammonia + Water + Salt

When alkalis react with ammonium salts, the hydroxide ions and the ammonium salts react to produce ammonia gas. The ionic equation is:OH- + NH4

+ NH3 + H2Of. Alkalis can react with the solution of one metal salt to give a metal hydroxide and another metal salt.

Basically the 2 ions exchange partners.g. Uses of bases and alkalis:

i. Ammonia solution: In window cleaning solutions, fertilisersii. Calcium oxide: Neutralising acidic soil, making iron, concrete and cement.iii. Magnesium hydroxide: In toothpaste to neutralise acid, in antacids to relieve indigestioniv. Sodium hydroxide: In making soaps and detergents, industrial cleaning detergents.

20. Concentration refers to how much a substance is dissolved in 1dm 3 of the solution . The concentration can be increased by adding more substance or reduced by adding more water.

21. Strength refers to how easily an acid or alkali dissociates when dissolved in water.a. A strong acid like hydrochloric acid dissociates easily in water.b. A weak acid like ethanoic acid does not fully dissociates.

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c. The strength of an acid cannot be changed. For e.g. ethanoic acid remains a weak acid, regardless of whether it is diluted or concentrated.

d. The strength of an acid can be shown using the pH scale.e. The pH scale is a set of numbers used to indicate whether a solution is acidic, neutral or alkaline.

i. Acidic: pH < 7. E.g. Hydrochloric acid, nitric acid, sulphuric acid, ethanoic acid.ii. Netural: pH = 7. E.g. water.iii. Alkaline: pH > 7. E.g. Aqueous ammonia, sodium hydroxide.

f. The pH of a given solution can be determined using an indicator or a universal indicator.i. It is calculated based on the number of hydrogen ions or hydroxide ions present in a solution.ii. Acids with a smaller pH value has a higher concentration of H+ ions, it will turn UI red colour.iii. Alkaline solutions with a larger pH value as a higher concentration of OH- ions. It will turn UI purple.iv. For neutral solutions, the Universal Indicator remains green.v. Therefore, pH can be used to control the strength of acids and alkalis of the same concentration.

g. Importance of pH: The pH of the soil is important as it can affect the growth and development of plants.i. Plants will not grow well in acidic plants, which can happen when too much fertiliser is added to the soil or

due to acid rain.ii. Chemicals are often added to the soil to adjust its pH. In areas where the soil is too acidic, it can be

treated with bases like quicklime (calcium oxide) or slaked lime (calcium hydroxide). It will react with the acids and raise the pH to allow the plants to grow healthy.

22. An oxide is a compound of oxygen and another element.a. There are 4 types of oxides: Acidic oxides, basic oxides, amphoteric oxides and neutral oxides.b. Acidic oxides are formed by non-metals. Most acidic oxides dissolve in water to form an acid.

i. Acidic oxides do not react with acids. ii. They react with alkali to form a salt and water.iii. Examples: Carbon dioxide, sulphur trioxide, phosphorus(V) oxide, silicon(IV) oxide.

c. Basic oxides are formed by metals.i. Most basic oxides are insoluble in water. Only alkalis (soluble oxides) are soluble in water.ii. Basic oxides are solids in r.t..p.iii. Basic oxides react with acids to form a salt and water.

d. Amphoteric oxides are metallic oxides that react with both acids and bases to form salt and water.i. E.g. Aluminium Oxide, Lead (II) Oxide and Zinc Oxide.ii. If a question consists of any of these oxides, it is likely it refers to this section.

e. Neutral oxides are non-metals that show neither basic nor acidic properties.i. They are insoluble in water.ii. E.g. Water, carbon monoxide and nitric acid.

23. Properties and uses of Sulfur Dioxide:a. As a bleaching agent: A bleaching agent decolourises coloured compounds, causing them to become pale or

white. Sulfur dioxide bleaches coloured compounds by removing oxygen from them. It is a reducing agent.b. As a food preservative: In the food industry, sulphuric dioxide is added to food in small amounts to prevent

the growth of moulds and bacteria.

24. Properties and uses of Sulfuric Acid:a. Manufacture of fertilisers: Sulfuric acid can be used to manufacture fertilisers like ammonium sulphate

(formed by reacting sulphuric acid and ammonia).b. Manufacture of detergents: Sulfuric acid is used to treat hydrocarbons to form an organic acid which is then

neutralised with sodium hydroxide to produce the detergent.[c.] As battery acid in cars: Sulfuric acid is used in batteries for cars. Lead plates and lead(IV) oxide are fitted in

the batteries. When sulphuric acid, lead and lead(IV) oxide react, electrical energy is produced.

25. A salt is formed when a metallic ion or an ammonium ion replaces one or more hydrogen ions of an acid.a. A salt contains a positive metal cation and a negative non-metal ion.b. Many salts combine with water molecules to form crystals. These water molecules are known as water of

crystallization. Salts that contain these water of crystallization are known as hydrated salts and those without are known as anhydrous salts.i. The dot after the formula of the hydrated salt means that when the substance is heated, everything after

the dot will be given off first.ii. When a hydrated salt is heated, water of crystallization is given off.

c. Solubility table: MEMORISE.i. All SPA (Sodium, Potassium, Ammonium) salts are soluble.ii. LS Chlorides (Lead Chloride and Silver Chloride) are insoluble.

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iii. BCL Sulfates (Barium, Calcium and Lead Sulfate) are insoluble.iv. SPA Carbonates/Hydroxides (Sodium, Potassium, Ammonium) salts are soluble.v. To analyse if a salt is soluble,

Check if it is sodium, potassium, ammonium or a nitrate. If it is, the salt is confirm soluble.Check if is chloride/sulfates. LS chlorides and SPA sulfates are insoluble.If pass both test, the salt is soluble.

26. Preparing salts: There are 2 factors to consider – the solubility of the salt and the starting materials.a. If a salt is soluble, it is usually prepared by reaction with acids.

i. Method 1: Reacting an acid with insoluble metal, base or carbonate: The acid is reacted with an excess of the substance to ensure that the acid is all used up. The substance must also be insoluble in water so as to be able to filter it from the salt solution later.

Steps: Add metal, filter mixture, collect filtrate, concentrate (heat), crystallise, filter.This method is not suitable for very reactive metals (Potassium, Calcium and Calcium) or unreactive metals (copper and silver). For these metals, use reactions with bases/carbonates.

ii. Method 2: Titration (used to prepare SPA salts as the starting materials are all soluble!).This is done as no residue is seen when the acid reacts with a soluble substance. A pH indicator is used to see exactly how much alkali or carbonate is required to neutralise the acid completely.Steps: Volumes determined by titration, concentrate, crystallise, filter.

iii. Method 3: Precipitation (used to prepare insoluble salts)This process involves mixing 2 soluble solutions that contains the positive cations of the salt with another solution that contains the anions of the salt to be prepared.As all nitrates are soluble, they are often used for making the cations of the substance. Steps: Add the 2 solutions together, filter, wash the precipitate and dry it.

27. Qualitative Analysis: It is the process of identifying cations and anions to identify an unknown salt.a. Cations can be identified using sodium hydroxide and aqueous ammonia. All cations except SPA (Sodium,

Potassium and Ammonium) give precipitates with these salts as they are soluble.b. A cation can be identified by noting:

i. Colour of precipitateii. Soubility in an excess of the reagentiii. Whether ammonia gas is liberated on addition of sodium hydroxide (NH4

+ only)c. Qualitative Analysis table (MEMORISE!)

28. A reversible reaction is a reaction that can go in either direction. a. The reaction from left to right is known as the forward reaction.b. The reaction from right to left is known as the reverse reaction.

29. Ammonia can be manufactured from the Haber Process.a. The raw materials required are nitrogen and hydrogen.

i. Nitrogen is extracted from air.ii. Hydrogen is produced from the cracking of petroleum.

b. Operating conditions in the Haber Process:i. Temperature of 450oCii. Pressure of 250atm.iii. Presence of iron catalyst

c. Equation for the Haber process: N2 + 3H2 2NH3.

30. Ammonia is displaced from its salt when it is heated with an alkali (e.g. Sodium Hydroxide or Calcium Hydroxide)a. The undesirable effect when calcium hydroxide is used in soil treatment is that it form ammonia gas which

escapes to the atmosphere. This causes the loss of nitrogen already added to the soil by farmers.

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31. The Periodic Table is a list of elements arranged in order of their increasing proton (atomic) numbers.a. A period is a horizontal row of elements, which also represents the number of electron shells.b. A group is a vertical column of elements. The group number represents the number of valence electrons of

the element, which determines the charge of the ion formed by the element.c. The block of elements between Groups II and III are tradition elements. They have multiple oxidation

states, coloured compounds and low densities.d. Metals located close to the bold line are known as metalloids and have both the properties of a metal and

non-metal. From left to right across a period, there is a decrease in metallic properties and an increase in non-metallic properties.

e. When going down the Periodic Table, the proton number increases, atoms become bigger and the properties of elements become more metallic (when going down the group, elements lose electrons more easily)

f. Types of ion formed by elements:i. Elements in Group I, II and III are metals which lose electrons to form cations. The charge of the ion is the

same as the group number of the element forming it. ii. Elements in Group IV and V of the Periodic Table are more likely to share electrons to form covalent

compounds and have a maximum oxidation state that is the same as the group number of the element.iii. Elements in Group VI and VII are non-metals and tend to gain electrons and form negative ions.iv. Elements in Group 0 have stable electronic configurations and do not form compounds.

32. Group I metals – Alkali metals: Group I elements react with water to form alkalis (soluble oxides/hydroxides)a. Physical Properties of Alkali Metals: Soft, cut easily, low melting and boiling points and low densities.b. When going down the group, the melting and boiling decreases and the density generally increases.c. Chemical Properties: Alkali metals are reactive metals, thus they are stored in oil to prevent them from

reacting with air and water.i. Alkali metals react with water to form hydrogen and an alkali.ii. All Group 1 elements form ions with a charge of +1 by losing 1 electron from the outer shell. Since they

give away all their electrons readily, they behave as powerful reducing agents.iii. As the size of the atom increases, it is easier to lose the valence electron from bigger atoms, causing the

reactivity to increase when going down the Group.iv. Compounds of alkali metals are ionic, soluble in water and have similar chemical formulae and are

white in colour.

33. Group VII Elements – Halogensa. Physical Properties: Halogens are non-metals and exist as diatomic covalent molecules. They have low

melting and boiling points and are coloured.i. Upon going down the group, the melting and boiling points of the halogens increase. (Chlorine is a gas,

bromine is a liquid whereas iodine is a solid at r.t.p.)ii. The colours of the halogens become darker.

b. Chemical Properties: Halogens react with most metals to form salts called halides.i. A displacement reaction is a reaction whereby one element takes the place of another element in the

compound. A more reactive halogen will displace a less reactive halogen from its halide solution.ii. A less reactive halogen cannot displace a more reactive halogen from its halide solution.iii. When going down the group, the reactivity of the halogens decrease.iv. During chemical reactions, halogen atoms gain electrons to form halide ions, thus the halogens are also

known as powerful oxidising agents.v. The displacement reactions between halogens and halides are also known as redox reactions. The more

reactive halogen acts as the oxidising agent whereas the less reactive halogen acts as the reducing agent.

34. Group 0 Elements – Noble Gasesa. Group 0 elements are inert gases (unreactive) and rare gases.b. They are monatomic, colourless at r.t.p., low melting and boiling points that increase down the group,

insoluble in water, unreactive.c. They do not lose, gain or share electrons to form compounds, thus they are unreactive.d. E.g. Argon is used to fill light bulbs to prevent the filament from oxidation.

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Page 20 Pure Chemistry Syllabus PaperTopic 9: Metals 5072 01/02

1. Properties of Metals and their structurea. Atoms in a metal are packed tightly into layers and are held together by strong metallic bonds. These bonds

require lots of energy to break, thus causing metals to have high densities, melting and boiling points.b. Atoms in a metal are packed regularly in layers which are of the same size, thus allowing them to slide over

each other when a force is applied. This makes metals soft, ductile (drawn into fine wires without breaking) and malleable (can be beaten into thin seats)

c. Metals have a structure whereby there are positive metal cations surrounded by a “sea of mobile electrons”. The mobile electrons make metals good conductors of heat and electricity.

d. Pure metals have useful properties but are not widely use as they are soft and may react with air and water and corrode easily. Alloys are used instead.

e. An alloy is a mixture of a metal with one or few other elements. For example,i. Bronze is an alloy of copper and tin,ii. Brass is an alloy of copper and zinc,iii. Steel is an alloy of iron and carbon.iv. Stainless steel is an alloy of iron, chromium, nickel and carbon.

f. Advantages of Alloys:i. Metals can be made harder and stronger by alloying them with other elements. This is so as the

added element breaks the regular arrangement of atoms in the pure metal. The atoms of the different sizes cannot slide over each other easily, making the alloy harder and less malleable.

ii. Alloying can improve the appearance of metals.iii. Alloys are more resistant to corrosion than pure metals.iv. Alloying is used to lower the melting points of metals.

2. The Reactivity Series: In the Reactivity Series, metals are listed from the most reactive to the least reactive. The order of reactivity is determined by the reactions of the metal.a. Order of the Reactivity Series:

Most Reactive Least ReactivePlease Stop Calling Me A Zebra I Like Coloured Stripes

Potassium Sodium Calcium Magnesium Aluminium

Zinc Iron Lead Copper Silver

b. Reaction with water: Metals react with water to produce metal hydroxide and hydrogen gas .i. Metal + Water Metal Hydroxide + Hydrogenii. Reactions of the different metals:

Potassium Reacts very violently/explosively.Heat produced ignites hydrogen gas produced.Hydrogen gas burns with lilac flame.

Sodium Reacts violently.Hydrogen gas may catch fire and burn with yellow flame.

Calcium Reacts readily.Magnesium Reacts very slowly. Hydrogen gas produced after a few days.No reaction for Zinc, Iron, Lead, Copper, Silver.

c. Reaction with steam: Metals react with steam to produce metal oxide and hydrogen gas .i. Metal + Steam Metal oxide + Hydrogenii. Reactions of the different metals:

Magnesium Reacts violently/explosively.A bright white glow is produced during reaction

Zinc Reacts readily. Zinc Oxide is yellow and white when cold

Iron Reacts slowly.The iron must be heated constantly for the reaction to proceed.

No reaction for Lead, Copper and Silver.d. Reaction with Dilute Hydrochloric Acid:

i. Metal + Dilute Hydrochloric Acid Metal Chloride + Hydrogenii. Reactions of the different metals:

Potassium, Sodium

Reacts explosively.

Calcium Reacts violently.

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Magnesium Reacts rapidly.Zinc Reacts moderately fast.Iron Reacts slowly.No reaction for Lead, Copper, Silver.

e. Reduction of Metal Oxides with Carbon: The reactivity of the different metals can be compared by studying how easily the metal oxides decompose. The more reactive the metal, the more difficult it is to decompose its oxides (reduce the oxide to the metal).

i. Metal + Carbon Metal + Carbon Dioxide/Monoxideii. Reduction reactions for the different metals:

Potassium, Sodium,Calcium,Magnesium

Oxides are not reduced by carbon. This is so as the oxides are too stable to be reduced (Carbon is below the metals in the Reactivity Series and is less reactive)

Zinc, Iron, Lead, Copper

Oxides are reduced by carbon. (Carbon is above the metals in the Reactivity Series and is more reactive)

Silver Oxide is reduced by heating (200oC). No carbon required.Equation: Silver Oxide Silver + OxygenChemical Eqn: 2Ag2O 4Ag + O2

f. Reduction of Metal Oxides with Hydrogen: Hydrogen gas is passed over the metal oxide. It acts as a reducing agent and reduces the oxides of some metals.

i. Metal oxide + Hydrogen Metal + Steamii. Reduction reactions for the different metals:

Potassium, Sodium, Calcium, Magnesium, Zinc.

Oxides are not reduced by hydrogen.

Iron, Lead, Copper , Silver Oxides are reduced by hydrogen.

g. Displacement Reactions of Metals: More reactive metals can displace less reactive metals from their salt solutions.

i. More reactive metal + less reactive metal salt More reactive metal salt + Less reactive metalii. Ionic equation (X is more reactive, Y less reactive): X + YSO4 XSO4 + Y.iii. A metal higher up in the Reactivity Series will displace a metal that is lower in the Reactivity Series from

its salt solution. No change occurs if the less reactive metal tries to displace the more reactive metal.iv. This is a redox reaction. The more reactive metal is oxidized whereas the less reactive metal is reduced

h. Reaction between a metal and the oxide of another metal: A more reactive metal has a higher tendency to form positive ions compared to a less reactive metal, thus it can reduce the oxide of a less reactive metal.

i. The more reactive the metal, the more readily it forms compounds.ii. E.g. Zinc + Copper(II) oxide Zinc Oxide + Copper

i. Action of heat on metal carbonates: Some compounds are more difficult to decompose and are more stable than other carbonates.

i. The more reactive the metal, the more stable the carbonate, the less likely/more difficult the carbonate will decompose.

ii. Reduction reactions for the different metals:

Potassium, Sodium,

Unaffected by heat, very stable.

Calcium, Magnesium, Zinc, Iron, Lead, Copper

Decompose into metal oxide and carbon dioxide upon heating.Eqn: Metal Carbonate Metal Oxide + Carbon Dioxide

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Silver Decomposes into silver and carbon dioxide upon heating.

j. The reactivity series is useful for:i. Predicting the behaviour of a metal from its position in the Reactivity Series,ii. Predicting the position of an unfamiliar metal in the Reactivity Series from a given set of experimental

results.

3. Extracting Metals: Most metals react with other elements to form ores (a compound of the metal mixed with large amounts of Earth and rock). a. There are 2 main methods for extracting metal from their ores. The position of a metal in the Reactivity

Series determines the method use for extraction.i. Reducing the metal compound (ore) to the metal using carbon (Metals below carbon)ii. Using electricity to decompose the molten metal compound (ore) to the metal. (Electrolysis) Used for

extracting metals above carbon in the Reactivity Series.iii. Generally, the more reactive the metal, the harder it is to extract the metal from its ore.

b. Extracting iron from Haematitei. Haematite contains iron(III) oxide mixed with impurities such as sand and clay. Iron is extracted from

haematite in a blast furnance.ii. Haematite, coke (mainly carbon) and limestone (calcium carbonate) are added at the top of the blast

furnace. Blasts of hot air are blown into the furnace near the bottom.iii. Firstly, Carbon Dioxide is produced.

a) The carbon in coke burns in oxygen to produce carbon dioxide. It is highly exothermic.Equation: Carbon + Oxygen Carbon Dioxide Symbol Equation: C (s) + O2 (g) CO2 (g)

b) The limestone is thermally decomposed (it undergoes thermal decomposition) by heat at the same time to produce carbon dioxide and calcium oxide (quicklime)Equation: Carbon + Carbon Dioxide Carbon Monoxide.Symbol Equation: CaCO3 (s) CaO (s) + CO2 (g)

iv. Secondly, Carbon Monoxide is produced. Carbon monoxide is the reducing agent. As the carbon dioxide rises up the furnance, it reacts with more coke to form carbon monoxide.Equation: Carbon + Carbon Dioxide Carbon Monoxide.Symbol Equation: C (s) + CO2 (g) 2CO (g)

v. Thirdly, Haematite is reduced to iron. The Carbon Monoxide reduces the iron(III) oxide in haematite to iron. Equation: iron(III) oxide + carbon monoxide molten iron + carbon dioxide.Symbol Equation: Fe2O3 (s) + 3CO (g) 2Fe (l) + 3CO2 (g)The iron formed is molten and runs to the bottom of the furnace. Hot waste gases containing carbon monoxide, carbon dioxide and nitrogen escape through the top of the furnance.

vi. Lastly, Impurities are removed. Calcium oxide from the calcium carbonate broken down earlier reacts with silicon(IV) oxide and other impurities to form a molten slag. The slag floats on top of the molten iron and are tapped off separately at the bottom of the furnace.Equation: CaO (s) + SiO2 (s) CaSiO3 (l)Solidified slag is used mainly for road surfacing.

4. Uses of Iron and Steela. Steel is an alloy of iron with carbon/other metals.b. Steel-making involves oxidation, following by alloying. Carbon and other small amounts of meteals are added

to molten iron to produce different types of steel with special properties. c. Removing Impurities by oxidation

i. Cast iron is melted and pure oxygen is passed through the molten iron to remove impurities.ii. The impurities form gases like sulphur dioxide and carbon dioxide which are blown out of the surface.iii. Calcium carbonate is later added to the furnace. At high temperatures, calcium carbonate decomposes

to form calcium oxide which reacts with silicon(IV) oxide and phosphorous oxide to form slag.iv. The slag is poured off and pure iron (wrought iron) is left behind).

d. Adding other elements to make steel: Carbon and other metals are then stirred into the motlen iron in the correct amounts to make steel with special properties.

e. Uses of steel:

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i. Carbon Steels: General purpose steels which contain mainly iron and carbon.ii. Mild Steel / Low carbon steel is hard, strong and malleable and is used to make car bodies and

machinery. iii. High carbon steel is strong but brittle. It is used to make cutting and boring tools (things used to hit hard

objects)iv. Alloy steel: Stainless steel is an alloy of iron, chromium, nickel and carbon. It does not rust and is used

extensively for making cutlery and surgical instruments.

5. Rusting: Rusting is the process that produces rust. E.g. When iron rusts, it forms hydrated iron(III) oxide (rust).a. Iron + Oxygen + Water Hydrated Iron(III) oxideb. Both air and water are required for rusting to occur.c. The presence of sodium chloride increases the speed of rusting. This is so as sodium chloride gives of ions

and accelerates electron transfer.d. Acidic substances such as sulphur dioxide and carbon dioxide also speed up the rusting process.e. There are 3 general ways for rust prevention:

i. Coating with a protective layer (not a permanent method ): The metal is coated with paint/grease, covering it with plastic, electroplating it or dip plating it. E.g. Galvanising is the coating of the metal to with zinc.

ii. Using a sacrificial metal: A more reactive metal is attached to the metal to be protected. The more reactive metal will then corrode instead of the less reactive metal. This is known as sacrificial protection as the more reactive metal is sacrificed to protect the less reactive metal.a) The more reactive metal cannot be too reactive (react with water, etc.)b) Metals less reactive will accelerate the rusting process.

iii. Using Alloys : The best known rust-resistant alloy is stainless steel. On exposure to air and moisture, a hard coating of Chromium(III) Oxide forms on the surface of stainless steel, protecting it from further corrosion.

6. Recycling Metals: Metals are finite resources (amounts of metals in the Earth are limited)a. Metals can be recycled. E.g. Lead is recovered from car batteries. Aluminium is recycled mainly from drink

cans and food containers.b. Advantages:

i. Land will be free for other uses (Fewer landfills required to dispose used metal objects and waste material, save cost for building landfill sites)

ii. Air and water pollution greatly reducediii. Conserve limited fossil fuel reserves.

c. Economic Issues of recyclingi. Recycling can be extremely costly (transporting the metals, separating the metals, sorting and cleaning

the scrap metals)d. Social Issues

i. Metal recycling programmesii. Being environmentally friendly.

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35. The Earth is surrounded by a blanket of air known as the atmosphere. Composition of air:

Nitrogen 78%Oxygen 20%Carbon Dioxide 0.03%Water Vapour 0 – 5% (varies on location)Noble Gases (Argon) 0.9%Neon/Helium 0.002%

36. Fractional Distillation of Liquid Air (important to get nitrogen and oxygen).a. Air is first cooled and compressed into liquid.b. Liquid air is then separated into its constituents by fractional distillation.c. The liquid with the lowest boiling point distils over first.

37. Air Pollution is defined as the condition in which air contains a high concentration of certain chemicals that may harm living things or damage non-living things. They are caused by air pollutants.a. Air pollutants include poisonous gases such as carbon monoxide, oxides of nitrogen, sulphur dioxide and

other pollutants.b. Carbon Monoxide is a poisonous gas that is colourless and odourless. It is formed by the incomplete

combustion of carbon-containing fuels. E.g. incomplete combustion of petrol in car engines.c. Oxides of Nitrogen can be produced in 2 ways:

i. In a car engine where temperatures are high, nitrogen combines with oxygen to form nitrogen monoxide which further reacts with oxygen to form nitrogen dioxide.N2 (g) + O2 (g) 2NO (g)2NO (g) + O2 (g) 2NO2 (g)

ii. During a thunderstorm, the heat energy released by lightning causes nitrogen and oxygen in the air to react to form oxides of nitrogen, which dissolve in water to form acid rain.

d. Sulfur Dioxide is mainly formed by the combustion of fossil fuels that contain sulphur. When they are burnt, they are converted to sulphur dioxide. They are also released in large quantities during volcanic eruptions. Equation: S (s) + O2 (g) SO2 (g)

e. Other pollutants: These pollutants include unburnt hydrocarbons, methane, lead and ozone. i. Unburnt hydrocarbons and lead are released in car exhaust.ii. Methane is a colourless and odourless gas produced when plant and animal matter/rubbish decay.

f. Effects of air pollution:i. Carbon monoxide can react with haemoglobin in the blood to form carboxyhaemoglobin, preventing

haemoglobin to transport oxygen to the rest of the body. It can cause headaches, fatigue, breathing difficulties and even death!

ii. Sulfur dioxide and oxides of nitrogen irate the eyes and cause breathing difficulties by irritating the lungs. High levels of them can also lead to inflammation of lungs. They also form acid rain which destroys buildings, aquatic life and plants.

38. Acid rain is formed when acidic air pollutants such as sulphur dioxide and nitrogen dioxide dissolve in rainwater.a. Sulfur dioxide dissolves in water to form sulphurous acid (H2SO3). b. In the presence of oxygen, this acid is oxidized to sulphuric acid (H2SO4).c. Nitrogen dioxide react with oxygen and water to form nitric acid.

4NO2 (g) + 2H2O (l) + O2 (g) 4HNO3 (g)d. The pH of normal rainwater is usually slightly below 7 as carbon dioxide in the air dissolves in rainwater to

form carbonic acid (H2CO3) which is a weak acid. However acidic rain is much more acidic with a pH value of 4 or less.

e. Effects of acid rain:i. Acid rain reacts with metals and carbonates in marble and limestone, damaging metal structures.ii. Acid rain reduces the pH of water bodies to below 4, killing aquatic life.iii. Acid rain leaches important nutrients from the soil and destroys plants, stunting plant growth.

39. Depletion of the Ozone Layera. Ozone is formed by the photochemical reaction between oxygen molecules and atoms in the atmosphere.

O (g) + O2 (g) O3 (g) [Oxygen can exist as atom???]b. When oxygen is formed at low altitudes, it can cause severe pollution problems. However in the

stratosphere, it acts as a shield, filtering out the harmful ultraviolent radiation from the sun.

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c. If ultraviolet radiation reaches the surface of the Earth, there could be an increase in number of skin cancer, genetic mutations and eye damage.

d. Ozone is pale blue, almost colourless with an odour. e. The depletion of the ozone layer is caused by chlorofluorocarbons (CFCs) which are compounds

containing the elements carbon, fluorine and chlorine. CFCs are widely used as coolants in refrigerators and air conditioners.

f. CFCs in the atmosphere decompose to form chlorine atoms which react with ozone molecules in the stratosphere to form chlorine oxide and oxygen, destroying the ozone layer.

g. Some countries have agreed to ban the use of CFCs (Kyoto Protocol)

40. Reducing air pollutiona. Some steps taken by the Government to reduce air pollution:

i. Prohibition on the use of open fires for disposal of domestic and industrial wastes, which can release air pollutants that are toxic to humans and can cause irritation and respiratory problems.

ii. Introduction of unleaded petrol, phasing out of lead petrol.iii. Fitting of all petrol-driven vehicles with catalytic converters.

b. Catalytic converters and flue gas desulfurisation can help to reduce the effects of acid rain.c. A catalytic converter is attached to the exhaust system of a car and contains platinum/rhodium.

i. When the hot exhaust gases pass over the catalysts, redox reactions occur. The harmful pollutants are converted to harmless substances.

ii. Carbon monoxide is oxidized to carbon dioxide. Oxides of nitrogen are reduced to nitrogen, and unburnt hydrocarbons are oxidized to carbon dioxide and water.Equation: 2NO (g) + 2CO (g) N2 (g) + 2CO2 (g)

iii. E.g. of unburnt hydrocarbons: octane.Equation: 2C8H18 (g) + 25O2 (g) 16CO2 (g) + 18H2O (g)

d. Flue Gas Desulphurisation refers to the process of removing sulphur dioxide from flue gases (waste gases formed when fossil fuels undergo combustion)

41. The carbon cycle refers to the mechanism that maintains the amount of carbon dioxide in the atmosphere. There are 3 processes involved: Combustion, respiration and photosynthesis.a. Carbon dioxide can be produced by 2 main processes:

i. Respiration: C6H12O6 (aq) + 6O2 (g) 6CO2 (g) + 6H2O (l) + energyii. Combustion of fuels: When fuels are burnt, carbon dioxide and water are produced. For e.g. combustion

of natural gas (consists of mainly methane):CH4 (g) + 2O2 (g) CO2 + 2H2O + heat energy.In a limited supply of oxygen, carbon particles (soot) and carbon monoxide are formed. This is known as incomplete combustion. Equation:4CH4 (g) + 5O2 (g) 2CO + 8H2O + 2C (s) + heat energy.

b. Plants are essential to us as they help remove carbon dioxide from the atompshere through photosynthesis. During photosynthesis, plants convert carbon dioxide and water into glucose and oxygen in the presence of sunlight.i. Equation: 6CO2 (g) + 6H2O (l) C6H12O6 (aq) + 6O2 (g)

42. The greenhouse effect is a warming effect due to the trapping of the infrared radiation from the sun.a. The natural greenhouse effect is crucial in maintaining the proper temperature needed to sustain life on

Earth. b. The effect of carbon dioxide build up is an increase in the Earth’s global temperature. This is known as

global warming.

c. Effects of global warming:i. A decrease in crop yields world-wide

ii. Melting of large quantities of ice in the Poles, which may cause sea levels to rise and flood low-lying countries.

iii. The rapid evaporation of water from the Earth’s surface, causing carbon dioxide dissolved in the oceans to be driven out into the atmosphere, worsening the greenhouse effect.

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43. All organic compounds contain the element carbon.a. Organic compounds that contain hydrogen and carbon only are called hydrocarbons.b. Not all carbon-containing compounds are organic. [E.g. carbon dioxide, carbonates]

44. A homologous series is a family of organic compounds with similar chemical properties. Compounds of the same homologous series contain the same functional group.a. A functional group is an atom or a group of atoms that gives a molecule its characteristic properties.b. Organic compounds in the same homologous series have the same chemical properties as they have the

same functional group.

45. Organic Compounds in the same homologous series have the following properties:a. They have the same functional group.b. They have similar chemical properties.c. They show a gradual change in their physical properties.

46. Naming organic compounds: Note the prefix.

Prefix meth- eth- prop- but-No. of carbon atoms one two three four

47. Petroleum (crude oil) is a dark brown, foul-smelling liquid which is a mixture of different hydrocarbons. Petroleum must be separated into fractions before it is used, which is done through oil refininga. Petroleum can be separated by fractional distillation in an oil refinery. b. It makes use of the fact that a hydrocarbon with more carbon atoms has a higher boiling point than one with

fewer carbon atoms.c. Sequence of hydrocarbons obtained from petroleum [learn uses as well]:

i. Petroleum gas: Fuel for cooking and heating.ii. Petrol: Fuel for car engines.iii. Naphtha: Feedstock for petrochemical industryiv. Kerosene: Fuel for aircraft engines/cooking using oil stovesv. Diesel Oil: Fuel for diesel enginesvi. Lubricating Oil: Used for lubricating machines, making waxes and polishesvii. Bitumen: For paving road surfaces

d. Competing uses for petroleum:i. Petroleum is mainly used as a fuel to generate electricity (90%) and the rest is used as chemical feedstock

for the manufacture for petrochemicals.ii. Naphtha makes up 20% of petroleum. The naphtha fraction is the main source of hydrocarbons used as

the chemical feedstock for the production of a wide range of organic compounds [e.g. plastics, detergents, medicines, fertilizers and pesticides].

e. Issues regarding uses of petroleum:i. Petroleum is a non-renewable resources and world’s petroleum reserves are finite.ii. The use of petroleum can be cut down by reducing the number of motor vehicles on the road, driving cars

that consume less petrol and taking public transport like MRT or using alternative energy sources.f. Fuels: A fuel is a substance that can be burnt in air to give out energy.

i. Our main source of fuel are coal, petroleum and natural gas.ii. General Equation for combustion of fuel: Hydrocarbon + Oxygen Carbon Dioxide + water + energyiii. Natural gas is mainly made up of methane (CH4). Petroleum (crude oil) is a mixture of hydrocarbons

(mainly alkanes) whereas coal consists mainly of carbon.iv. Alternative fuels include palm oil as a substitute for petrol in Malaysia. In Brazil, alcohol mixed with petrol

can be used to run vehicle engines. Biogas (gas produced when organic matter decay in the absence of air) can also be used, it contains about 50% methane.

48. Alkanes are saturated hydrocarbons that contain only single covalent bonds between the carbon atoms.g. Alkanes have the general formula CnH2n+2

h. Based on alkanes, members of the same homologous series also have the same general formula and each member differs from the next by a –CH2– unit.

i. A structural formula is the formula which shows how atoms are arranged in a molecule.

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j. Alkanes are saturated hydrocarbons as they contain only single covalent bonds. Each carbon atom in an alkane molecule uses all its valence electrons to form single bonds with 4 other atoms, and no more atoms can be added to an alkane.

k. Isomers are compounds that have the same molecular formula but different structural formula.i. Butane is a straight-chain alkane as all the carbon atoms are arranged in a row.ii. Methylpropane is a branched-chain alkane as it has a branch/side chain.iii. The side chains are called alkyl groups, which is obtained by removing a hydrogen atom from an alkane.

It have the general formula of CnH2n+1.l. Physical Properties of Alkanes: All alkanes are insoluble in water and are flammable.

i. Volatility: The melting and boiling points of the alkanes increase as their molecular sizes increase. This is so as when alkane molecules become bigger, the attractive forces between the molecules become stronger.

ii. Density: The density of alkanes increase as their molecular size increase.iii. Viscosity: The alkanes become more viscous (more difficult to pour out) as their molecular size increase.

This is so as the intermolecular forces are stronger as the alkane molecules become bigger, making it difficult for the liquid to flow.

iv. Flammability: As the molecular size of the alkane molecules increase, the percentage of carbon in the alkanes increase, hence they become less flammable and burn with a smokier flame. The smoky flame is caused by the incomplete combustion of alkane molecules.

m. Chemical Properties of Alkanes:i. Combustion: Combustion of alkanes in excess air/oxygen produces carbon dioxide and water vapour.

When there is an insufficient supply of air/oxygen combustion of alkanes is incomplete, which may form carbon or carbon monoxide as well.Combustion is highly exothermic, hence alkanes are good fuels. [E.g. methane is used in cooking gases]

ii. Cracking: Alkanes with long carbon chains are broken up into smaller molecules in this process.iii. Substitution Reaction: Alkanes are saturated hydrocarbons with single carbon-carbon bonds and carbon-

hydrogen bonds which are strong and difficult to break, making alkanes generally unreactive. However alkanes react with halogens in the presence of UV light.

This reaction is known as a substitution reaction as a hydrogen atom is being substituted by another atom of the halogen.

49. Alkenes are unsaturated hydrocarbons that contain one or more carbon-carbon double bonds.n. Functional Group: C = C bond.o. General Formula: CnH2n

p. Alkanes are obtained by cracking petroleum (crude oil). Cracking is the breaking down of long-chain hydrocarbons into smaller molecules.i. A catalyst may be used to speed up cracking. This process is known as catalytic cracking.ii. In industries, cracking is done passing the petroleum fraction over a catalyst [aluminium oxide or

silicon(IV) oxide] at a high temperature [600oC].iii. Reaction: Long chain alkanes short chain alkene + short chain alkanes or hydrogen gas.iv. Importance of cracking:

Produce petrol: Our need for petrol is greater than our need for diesel oil or lubricating oil [made up of long-chain alkanes]. Hence cracking can convert them into petrol [short-chain alkanes]Produce short-chain alkenes [e.g. ethane and propene which are used as starting materials for making ethanol and plastics.]Produce hydrogen [a by-product]

q. Chemical Properties of Alkenes: Alkenes can take part in combustion and addition reactions.i. Combustion: Alkenes burn in plentiful supply of oxygen to form carbon dioxide and water vapour. Since

alkenes have a relatively higher percentage of carbon than alkanes, carbon particles may be produced. They also burn with a smokier flame.

ii. Addition: The C=C bonds in alkenes are reactive and can undergo addition reactions. An addition reaction is a reaction whereby an unsaturated organic compound combines with another substance to form a single new compound.

C=C bonds become single bonds; hence an unsaturated organic compounds become saturated.Hydrogenation: Addition of hydrogen to alkanes. Temperature: 200oC, catalyst: Nickel. It is used to make margarine from vegetable oil.Bromination: Addition of bromine to alkenes. It serves as a chemical test for the presence of unsaturated hydrocarbons and distinguishes between alkane and alkene. Alkanes do not decolourse bromine solution, however alkenes decolourise bromine solution.Hydration: Alkenes can react with steam to produce an alcohol. The reaction requires a catalyst [Phosphoric(V) acid, a temperature of 300oC and a pressure of 60atm.]

50. Alcohols are organic compounds with the hydroxyl (-OH) functional group.

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r. They contain the elements carbon, hydrogen and oxygen.s. They have the general formula CnH2n+1OH.t. The name of an alcohol ends with ‘-ol’.u. Physical Properties of Alcohols:

i. Alcohols are soluble in water, but their solubility decreases as the molecular size increases.ii. The first 4 alcohols are liquids at r.t.p.

v. Chemical Properties of Alcohols:i. Alcohols are not alkalis (they are neutral) despite having the hydroxyl group.ii. They are more reactive than alkanes because of the bonds.iii. Alcohols take part in combustion and oxidation.iv. Combustion: An alcohol burns in air to produce carbon dioxide and water vapour. It allows alcohols to be

used as a fuel. E.g. Methanol is less volatile than petrol and is a clean fuel. It is also burnt on fruit cake to give it its distinct flavor.

v. Oxidation: Alcohols can be easily oxidized to form a carboxylic acid. In the laboratory, ethanol is oxidized by heating it with a mixture of acidified potassium dichromate(VI) solution. This is a redox reaction. The colour of the potassium dichromate changes from orange to green. It is used in breathalyzers to test the amount of alcohol consumed by drivers.

w. Manufacturing Ethanoli. Hydration of Ethene: Ethanol is manufactured by the catalytic addition of steam to ethene. The mixture is

passed through phosphoric(V) acid at 300oC and 60atm. The acid acts as a catalyst in the reaction.ii. Fermentation of Ethanol: Fermentation is a chemical process whereby microorganisms such as yeast act

on carbohydrates to produce ethanol and carbon dioxide. [E.g. sugar, starch]Yeast contains enzymes that break down starch to glucose.The glucose is further broken down to ethanol and carbon dioxide.The mixture is kept at a temperature of 37oC [If the temperature is raised, the enzymes will die and fermentation stops]The fermentation of sugars only produce a dilute solution of ethanol. [When the alcohol content exceeds the value the yeast dies]. Ethanol can then be obtained by fractional distillation.If alcoholic drinks are left to air, it will turn sour due to the action of bacteria from the air which oxidizes ethanol into ethanoic acid.

x. Uses of Ethanol:i. In alcoholic drinks: Large quantities of ethanol are used in alcoholic drinks.ii. As a solvent: Ethanol is a good solvent and dissolves many subsutances insoluble in water. It also

evaporates quickly, hence it is used in paints, varnishes, perfumes and lotions.iii. As a fuel: Ethanol is the main constituent of methylated spirit which is used in spirit lamps and burners

for cooking.

51. The carboxylic acids are a homologous series of organic compounds with the carboxyl functional group (-COOH). They are weak acids and are a class of organic acids.y. General Formula: CnH2n+1COOH.z. Properties of Ethanoic Acid: It is a colourless liquid at r.t.p. and has a characteristic vinegar smell. It is

completely miscible in water.aa. Physical Properties of carboxylic acids: As the number of carbon atoms increase, the boiling point of a

substance increases.bb. Functional Group: -COOH (Carboxyl).cc. Carboxylic acids are weak acids as they dissociate partially in water to form hydrogen ions, which give the

carboxylic acids their acidic properties.i. Reaction with metals. [Salts of ethanoic acid are ethanoates, and so on]

ii. Reaction with carbonatesiii. Reaction with bases

dd. An ester is formed when a carboxylic acid reacts with an alcohol to form an organic compound. This process is known as esterification.i. During the formation the –OH group of the carboxylic acid is replaced by the CH3-O- group of the ethanol.

ii. Conditions: Concentrated Sulfuric Acid [catalyst], heat.iii. In the name of an ester, the first part tells the ethyl group whereas the second part consists of the

ethanoate group. iv. Uses of esters: Esters are colourless, neutral liquids which are insoluble in water. They have a sweet,

fruity smell, hence they are used in the preparation of perfumes and artificial food flavourings. Some natural occurring esters include animal fats and vegetable oils.

52. A macromolecule is any long-chain molecule that contains hundreds and thousands of atoms joined together by covalent bonds.ee. It is formed by linking together many small repeating units known as monomers.

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ff. Polymerisation: The process of joining together a large number of monomers to form a macromolecule.gg. The macromolecule formed is called a polymer.hh. E.g. of polymers:

i. Natural: Proteins, carbohydrates, fats.ii. Synthetic: Poly(ethene), nylon, terylene, PVC.

ii. Addition Polymerisation occur when monomer units join together without losing any molecules or atoms.i. At high pressure (1000atm) and high temperature (200oC) thousands of alkene molecules join together

to form polymers.ii. E,g. Poly(ethene), also known as polythene. It is mainly used for making bags, toys, buckets, clingfilm

etc. It is flexible but difficult to break.iii. Addition polymers are good insulators of heat and electricity and are resistant to chemical attack. For

instance, PVC is used to make pipes, raincoats. Windscreen of cars are made up of Perspex.jj. Condensation Polymerisation occurs when 2 different types of monomers react to form a polymer and a

small molecule is produced. There are 2 main groups: polyamides (polymer with amide linkage) and polyesters (polymer with ester linkage).i. Nylon – A Synthetic Polyamide: It is made up of the

monomers dicarboxylic acid and diamine. ii. Terylene – A Synthetic Polyester: It is made up of a

dicarboxylic acid and a diol (alcohol with 2 –OH groups).iii. Examples of items made from Nylon and Terylene include

curtains, parachutes, fishing lines and sleeping bags.kk. Plastics from polymerization are highly used as they are cheap, easily moulded into various shapes, light,

tough and waterproof and are durable.i. Plastic are carbon compounds and are flammable. When plastics burn, fires spread quickly and poisonous

gases are produced.ii. Plastics are non-biodegradable and cannot be decomposed by bacteria in the soil. Burying them in

landfills take up space and cause land pollution.iii. Most plastics are destroyed by burning (incineration) to produce carbon dioxide and water. However,

many plastics like PVC produce poisonous gases like hydrogen chloride upon burning. Burning plastics will hence cause air pollution, and the best way to deal with them is to reuse or recycle them!