Samoa School Syllabus - Chemistry

This document outlines the topics covered in the Chemistry curriculum. It draws primarily from the Samoa Chemistry Curriculum Year 12 Student Learning Guide prepared by Dr. David Salter.

Introduction to Chemistry

Chemistry is the study of matter, its properties, and how various substances can be beneficial. It explores the properties of substances and how they change in different conditions. Understanding chemical and physical changes, such as the difference between a chemical reaction (like metal rusting) and a physical change (like ice melting), helps in interpreting observations of the world.

Over time, people have studied and categorized numerous substances. This led to the development of classifications based on observable properties, such as elements, compounds, metals, nonmetals, and so on. Learning chemistry involves understanding these categories and their properties. It also involves learning the specific language and symbols chemists use to describe substances and reactions. The periodic table is a key tool, summarizing a wealth of information about the elements.

Chemistry can be seen as having several interconnected components. Chemists classify substances according to observable properties (macroscopic). They also create models to explain these observations at a level too small to be seen (submicroscopic). Finally, they use symbols and formulas (symbolic) to convey information. For example, water is described by its observable properties (clear, colorless liquid at room temperature), by its molecular structure (H₂O), and by the interactions between its molecules.

Atomic Structure

• Describe the model of an atom and the structures that make up an atom.
• Identify the atomic structure (protons, neutrons, electrons, isotopes) of a given atom or ion.
• Explain how an atom changes to an ion and the combining power of an atom that causes that change.
• Describe the formation of an ion.
• Draw Lewis structures of simple molecules.
• Describe the shapes of simple molecules.
• Describe the formation of ionic, covalent (both polar and non-polar), and metallic bonding.

Atoms are composed of smaller particles: protons, neutrons, and electrons. The protons and neutrons are located in a small region at the center of the atom called the nucleus. The number of protons in an atom is its atomic number. The mass of any atom is defined relative to the mass of a carbon atom, which is 12. This is referred to as its relative atomic mass. The relative molecular mass of a molecule is the sum of the relative atomic masses of its atoms. The mole mass is the mass in grams of one mole of particles and has the units gram per mole (g/mol). The mass in grams of one mole of any atom is the same number as the relative atomic mass of that atom, usually shown beneath the symbol on the periodic table. For a particular substance, the amount in moles and mass in grams are related by the mole mass. Since atoms and molecules are far too small to count individually, chemists have devised a way to determine the number of particles present in a pure substance of its mass and chemical formula is known.

Chemical bonds link atoms or ions together. Depending on the nature of the atoms, the bonds are classified as covalent, ionic, or metallic. The strength of attractions between molecules in a substance determines many of its physical properties.

Quantitative Chemistry

• Define the terms molar, relative atomic mass, relative molecular mass, and mole mass.
• Make connections between mass and moles.
• Study the law of conservation of mass, which leads to the balancing of simple equations.
• Calculate percentage composition of elements in compounds.
• Deduce empirical formula.

Chemical equations show the relative amounts of atoms involved in a chemical reaction. Since no atoms are lost or gained in a chemical reaction, the same number and type of atom must appear on both sides of a chemical equation for it to be balanced.

The amounts of substances are measured in chemistry by the mole. A mole is the amount of substance that contains the same number of particles as there are atoms in exactly 12 g of carbon isotope ¹²C, an enormous number. The mass of one mole of any atom is defined relative to the mass of an atom of carbon, which is 12. This is referred to as its relative atomic mass. The relative molecular mass of a molecule is the sum of the relative atomic masses of its atoms. The mole mass is the mass in grams of one mole of particles. For a particular substance, the amount in moles and mass in grams are related by the mole mass. Since mass can be measured using a balance, the amount in moles and the number of particles present can be calculated. Chemical equations show the relative amounts of atoms involved in a chemical reaction. Since no atoms are lost or gained in a chemical reaction, the same number and type of atom must appear on both sides of a chemical equation for it to be balanced. The percentage composition of a compound shows how much each element contributes to the total mass of the compound. The empirical formula shows the simplest whole number ratio of atoms in a compound. The molecular formula shows the actual number of atoms present in a molecule.

Physical Chemistry

• Describe exothermic and endothermic reactions.
• Investigate the factors affecting the rate of reactions.
• Discuss a simple introduction to equilibrium with reference to reversible reactions.

Chemical and physical change is accompanied by changes in the energy content of the substances involved. For example, when wood burns in air, energy is released as heat and light. Energy from the sun is absorbed when ice melts in sunlight. The energy changes accompanying a process, the speed at which the process occurs, and the extent to which reactants are converted to products are all important aspects of chemical reactivity.

For a chemical reaction to occur, particles must collide. Chemical reactions are faster if the particles collide more frequently or with greater force. Reaction rate can be increased by increasing the surface area of solid reactants, increasing the concentration of reactants, increasing the temperature of the reaction mixture, or adding a catalyst. Changes in the temperature of the surroundings indicate that the energy content of reactants and products differs. Reactions that release energy to the surroundings, which gets warmer, are called exothermic reactions. Reactions that absorb energy from the surroundings, which gets colder, are called endothermic reactions. Every reaction is reversible; that is, the products may react to reform the reactants. At equilibrium, the rates of the forward and reverse reactions are equal, and the relative amounts of reactants and products do not change further.

Redox Reactions

• Define oxidation and reduction in terms of transfer of oxygen, electrons, and change in oxidation state.
• Work out the oxidation state of atoms and ions.
• Write balanced equations to describe redox reactions which lead to balanced equations.
• Describe the role of oxidizing agents and reducing agents.

Any chemical reaction that has an element as a reactant or a product is a redox reaction. For example, the reactions that generate electricity in batteries are redox reactions.

The driving force for redox reactions is the transfer of electrons from one atom to another atom that has a greater attraction for electrons. Oxidation is the loss of electrons, and reduction is the gain of electrons. Both oxidation and reduction occur in every redox reaction. A useful method for identifying the atom or ion that has lost electrons and the atom or ion that has gained electrons in a redox reaction is to assign oxidation numbers to each atom using a set of rules. Oxidizing agents or oxidants are substances that gain electrons in redox reactions. Reducing agents or reductants are substances that lose electrons in redox reactions. The number of electrons lost by the reducing agent in a redox process must equal the number of electrons gained by the oxidizing agent.

Acid-Base Reactions

• Define an acid as a proton donor and a base as a proton acceptor.
• Outline the chemical and physical properties of acids and bases.
• Carry out an investigation of the chemical reactions of acids with metals, bases, and carbonates.
• Carry out simple tests for acids and bases using common indicators.
• Use the pH scale to classify substances as acidic, neutral, or basic.
• Explain the difference between strong and weak acids and bases.
• A strong acid and a strong base react to form water and an ionic metal compound called a salt.

Acid-base reactions occurring in aqueous solutions are important reactions in living organisms as well as in many industrial processes. Neutralization reactions, in which acids and bases react, have many important applications.

An acid is a substance that reacts with water to produce H⁺ ions. More generally, an acid is a proton (H⁺) donor. A base is a substance that produces OH⁻ ions when dissolved in water. More generally, a base is a proton (H⁺) acceptor. The acidity or basicity of a solution is indicated by its pH. A neutral solution has a pH of 7. Acidic solutions have pH values less than 7, and basic solutions have pH values greater than 7. Strong acids and strong bases react completely with water. Weak acids and weak bases react only to a small extent with water. A strong acid and a strong base react to form water and an ionic metal compound called a salt. Acids react with certain metals to form hydrogen gas as one product and react with metal carbonate compounds to form carbon dioxide gas as one product.

Precipitation Reactions

• Demonstrate what happens to a compound that dissolves in water.
• Investigate solubility properties of chlorides, sulfates, nitrates, carbonates, and hydroxides of metals.
• Use the deduced solubility rules to predict the name of the formed precipitate and its color.
• Carry out the tests for chosen anions.
• Carry out the tests for cations.

In precipitation reactions in water, two soluble ionic compounds react to form an insoluble ionic product known as a precipitate. Coral reefs and kidney stones are the result of this type of chemical reaction.

Water is a good solvent for ionic substances. When an ionic compound dissolves in water, it is separated into hydrated positive and negative ions. Precipitation of an ionic solid occurs upon mixing solutions containing hydrated ions if the attraction between the positive and negative ions is stronger than the attraction of either ion for water molecules. Chemists have studied many precipitation reactions and have generated a set of solubility rules that help in predicting whether a precipitate will be observed upon mixing two solutions of ionic compounds. Chosen cation and anion tests can be used to confirm the presence of particular ions in aqueous solution.

Chemistry in the Ocean

Coral reefs are formed by tiny invertebrate animals that secrete a hard, coral structure. Coral animals absorb calcium ions, Ca²⁺, from seawater and convert carbon dioxide from cell respiration into hydrogen carbonate ions, HCO₃⁻. These ions are secreted by the coral animals under their skin around the lower half of their bodies in a substance that hardens to form calcium carbonate, CaCO₃. While precipitation of calcium carbonate in seawater does not support much marine life, corals can cover large areas close to the seashore. Coral animals also host microscopic plants called zooxanthellae in their tissues, which provide simple sugars by photosynthesis. Therefore, corals do not have to move much to obtain food and can survive in tropical waters that contain little food.

Coral reefs are complex and fragile ecosystems. When solid calcium carbonate is heated strongly at about 800°C, solid calcium oxide, CaO, also called lime, is produced. When mixed with water, to form a slurry, it is used by many people outside their homes as a deodorant. When large blocks of coral are heated and crushed, the resulting calcium oxide reacts with water to form calcium hydroxide, Ca(OH)₂. This is used by many people as a deodorant outside their homes. When solid calcium carbonate is strongly heated at about 800°C, solid calcium oxide, CaO, also called lime, is produced. Large blocks of coral were heated and crushed to produce lime for building houses.

Metals

• Carry out an investigation to identify and compare the physical properties of metals and nonmetals.
• Investigate the properties of metals through their reactions with air, oxygen, water, and dilute acids.
• Deduce the activity series of metals using reaction between metals and air, water, and dilute acids, and the displacement of metal ions in metal salt solutions.
• Investigate metal corrosion and how to prevent it.
• Define the term alloy.
• Compare the physical properties of the pure elements in an alloy with the properties of the alloy.

About three-quarters of the elements are metals and have many common properties. Nonmetals show properties that are very different from the properties of metals.

Metals are generally shiny solids at room temperature (mercury is the only liquid), are good conductors of heat and electricity, and can be hammered into thin sheets (malleable) and pulled into wires (ductile). Nonmetals are generally gases or dull brittle solids at room temperature (bromine is the only liquid). Nonmetals do not conduct heat or electricity well. Metals react as reducing agents and lose electrons in chemical reactions to form positive ions. Nonmetals react as oxidizing agents and gain electrons in chemical reactions to become more negatively charged. Reactive metals undergo chemical reaction with air, water, and dilute acids. Chemists list metals in order of their reactivity with these substances in the activity series of metals. The most reactive metal is listed at the top of the series and the least reactive at the bottom. A more reactive metal will displace a less reactive metal from a solution of its ions. Corrosion is the deterioration of a metal by reaction with substances present in the environment. Rusting of iron is a common example of corrosion. Rusting can be prevented or limited by painting the metal, coating it with a less reactive metal, or attaching a more reactive metal (sacrificial protection). An alloy is a mixture of two or more elements, with at least one being a metal. Alloys often have enhanced properties compared to their component elements.

Using the Downs Cell of Iron

Rusting is a redox reaction in which iron is oxidized to iron(III) oxide, Fe₂O₃, and oxygen in air is reduced. Water is also necessary for rusting to occur. Rusting can be controlled in several ways by preventing oxygen and water from coming into contact with the iron. Painting the iron or coating it with a less reactive metal or plastic are common methods. Another method, called cathodic protection, involves attaching a more reactive metal such as zinc or magnesium to the iron. The more reactive metal corrodes in preference to the iron and provides electrons to prevent the iron from rusting. Fuel for vehicles and planes arrives in Samoa in tankers that moor off the coast. Harbor fuel tanks that have a large amount of iron surface are brittle and weak. Therefore, it is important to paint iron roofs and fuel storage tanks as well as the inside of ships to prevent rusting occurring. Colorless roofing paint that contains zinc metal is often used. As the paint forms a physical barrier to water and oxygen, the zinc in the paint provides sacrificial protection to prevent rusting.

Oxygen and Hydrogen and their Compounds

• Carry out the laboratory preparation of oxygen and hydrogen.
• Investigate the occurrence, properties, and uses of oxygen and hydrogen.
• Discuss the ozone layer and its importance to the earth.
• Describe the effect of CFCs on the ozone layer.
• Describe the industrial preparation of oxygen and hydrogen.
• List the uses of hydrogen.

Oxygen

Symbol: O₂ Atomic Number: 8 Electron Arrangement: 2,6

• The most abundant element on Earth's surface.
• Present as elemental oxygen in the atmosphere (21% by volume of the atmosphere).
• Present as compounds with metals and silicon and as a component of water H₂O.
• About 2/3 of the mass of Earth's crust consists of compounds with O.
• Essential constituent of all living matter with C, H, and N.
• Required for respiration, the process by which organisms obtain energy.

• Only slightly soluble in water but enough oxygen dissolves to support aquatic life.
• Required for combustion, the process by which fuels burn in pure oxygen to produce oxides.

Physical Properties:

• Occurs as diatomic molecules O₂ with a double bond O=O linking two oxygen atoms.
• Colorless and odorless gas, slightly denser than air, and freezes at -219°C and boils at -183°C.

Chemical Properties:

• Reacts directly with most other elements and forms oxide compounds with all elements except He, Ne, Ar, and Kr.

Uses of oxygen:

• In the manufacture of steel. Iron that has a large amount of carbon impurity is brittle and weak. Steel consists mostly of iron and is much harder and less brittle than iron. Oxygen is blown into molten iron to reduce the amount of impurities, mainly carbon, present.
• For cutting and welding metals. Oxyacetylene torches use pure oxygen and ethyne, C₂H₂, gas in a combustion reaction that gives a very hot flame, about 3000°C. This is hot enough to melt metals for welding and cutting.
• In breathing apparatus. Hospitals use oxygen in life support systems and for patients with breathing difficulties. Also used in high-altitude mountaineering and deep-sea diving.
• In spacecraft. Rocket fuel uses liquid oxygen to burn fuels such as liquid hydrogen.

Industrial production of oxygen: Obtained from the fractional distillation of liquid air. Air that has had dust, carbon dioxide, and water vapor removed is compressed and cooled to liquefy most of it at -200°C. As liquid air boils, nitrogen, N₂, boils off first at -196°C leaving liquid oxygen, which boils at -183°C.

Hydrogen

Symbol: H₂ Atomic Number: 1 Electron Arrangement: 1

• The most abundant element making up about 75% of the universe.
• The lightest element and occurs on Earth combined in water H₂O, natural gas methane CH₄, and living matter in proteins, carbohydrates, and fats and oils.

Physical Properties:

• Occurs as diatomic molecules H₂ with a single bond H-H linking two hydrogen atoms.
• Colorless and odorless gas, much less dense than air, and freezes at -259°C and boils at -253°C.

Chemical Properties:

• Chemistry dominated by tendency to share its electron with other nonmetal atoms.
• Reacts directly with most other elements and forms oxide compounds with all elements except the noble gases.

Uses of hydrogen:

• In the manufacture of steel. Iron that has a large amount of carbon impurity is brittle and weak. Steel consists mostly of iron and is much harder and less brittle than iron. Hydrogen is blown into molten iron to reduce the amount of impurities, mainly carbon, present.
• For cutting and welding metals. Oxyhydrogen torches use pure oxygen and hydrogen gas in a combustion reaction that gives a very hot flame, about 3000°C. This is hot enough to melt metals for welding and cutting.
• In breathing apparatus. Hospitals use hydrogen in life support systems and for patients with breathing difficulties. Also used in high-altitude mountaineering and deep-sea diving.
• In spacecraft. Rocket fuel uses liquid hydrogen to burn fuels such as liquid oxygen.

Industrial production of hydrogen: Obtained from the reaction of methane with steam at 800°C using a nickel catalyst.

Laboratory preparation of oxygen: Several methods can be used in a laboratory to produce oxygen gas. A test for oxygen gas is that a stick with a glowing ember relights when placed in oxygen.

Laboratory preparation of hydrogen: Several methods can be used in a laboratory to produce hydrogen gas. A test for hydrogen gas is that a lighted splint makes a squeaky pop when placed in hydrogen.

The ozone layer: Ozone O₃ is an allotrope of oxygen. Allotropes are different forms of an element in the same state. Oxygen gas O₂ and ozone O₃ are gases at room temperature. Ozone is formed in the upper atmosphere, about 20 km above Earth's surface, when oxygen molecules absorb high-energy ultraviolet (UV) radiation from outer space and split into oxygen atoms. These oxygen atoms react with other oxygen molecules to form ozone. Ozone absorbs more UV radiation and decomposes back to oxygen atoms and oxygen molecules. Therefore, the ozone layer acts as a filter to prevent harmful UV radiation from reaching Earth's surface.

CFCs: Chlorofluorocarbons (CFCs) are compounds containing carbon, fluorine, and chlorine that were widely used as refrigerants in air conditioners and refrigerators, as propellants in aerosol cans, and in the manufacture of foamed plastics. CFCs are chemically unreactive in the lower atmosphere but in the upper atmosphere, they decompose to form chlorine atoms when exposed to UV radiation. Chlorine atoms react with ozone and convert it into oxygen. This has caused a decrease in the amount of ozone in the upper atmosphere, particularly over Antarctica. It is predicted that a decrease in the ozone layer will allow an increase in the amount of harmful radiation reaching Earth's surface and consequently contribute to an increase in the incidence of skin cancer.

Water

• Describe the following properties of water such as melting point and boiling point and solvent properties.
• Explain the purification of water by the process of filtration and chlorination.
• Define the terms soft and hard water.
• Investigate the role of soap in water.
• Define the terms dissolved oxygen, hydrogen, biological, and chemical oxygen demand.

Water is all around us. It is a unique substance that has some unusual properties that are vital for life on Earth. These properties arise from the nature of the water molecule, which is made up of two H atoms and an O atom.

Formula: H₂O

• Most abundant liquid on Earth about 70% of Earth's surface covered by water. Living organisms are composed mostly of water.
• Boils at 100°C and freezes at 0°C.
• Very poor conductor of electricity unless it contains large amounts of dissolved ions.
• Water molecule has an angular or bent structure with a H-O-H angle of 105°. Bonding electron pairs are attracted more strongly to O atom and partially negative charge δ⁻ and partially positive charges δ⁺ appear on H atoms. This uneven distribution of electrical charges causes water to be a polar molecule.

Solvent properties: Water molecules are attracted to the charged ions in an ionic solid and cause the crystal structure to break down as each ion becomes surrounded by several water molecules. This explains why water is a good solvent for ionic compounds.

Hard water: Water is described as being hard if it forms a scum with soap and does not readily form a lather. The presence of moderate amounts of calcium ions, Ca²⁺, and magnesium ions, Mg²⁺, makes water hard. These ions react with soap to form insoluble precipitates, scum. Soap will not lather until all the calcium and magnesium ions have been precipitated. Hard water results when rain water flows over rocks containing calcium or magnesium minerals which dissolve in the water.

Purification of water for drinking: Water for drinking must be purified to remove disease-causing microorganisms and suspended matter. The main steps in water purification are:

• Filtration: Water is filtered through beds of sand and gravel to remove suspended matter such as silt, natural organic matter, and some microorganisms.
• Chlorination: Chlorine gas, Cl₂, is bubbled through the filtered water to kill any remaining microorganisms. Chlorine dissolves in water to form hypochlorous acid, HOCl, which kills bacteria.

Soft water: Soft water readily forms a lather with soap. Rainwater is naturally soft but becomes hard as it flows over and dissolves minerals in rocks. Soap: Animal fat heated in a solution of sodium hydroxide, NaOH, is converted into sodium stearate, also known as soap. The long non-polar hydrocarbon end of the stearate ion dissolves in the thin film of oil on dirty fabrics or skin, while the ionic carboxylate end stays in the water. This causes tiny oil droplets surrounded by stearate ions to form and be suspended in the water so they can be rinsed away.

Unusual properties:

• Water is very stable to heat.
• High heat capacity. Large volumes of water such as lakes and oceans can absorb a lot of heat energy without their temperature changing much. Earth's oceans help prevent massive temperature changes in Earth's climate.
• Ice is less dense than liquid water. When water freezes, hydrogen bonding extends throughout the whole structure causing water molecules to be arranged in a regular open network structure. This causes ice to have a lower density than liquid water and explains why ice floats on water. As cold air above the lake cools the surface water, the surface water becomes denser and sinks to the bottom of the lake. This continues until all the water in the lake is at 4°C. As the surface water cools further to below 4°C, it becomes less dense and remains on the surface of the lake. When the surface water freezes, the ice floats on the denser water below. This insulates the water below from further heat loss and allows aquatic plants and animals to survive under the ice during winter.

Nitrogen

• Investigate the occurrence, properties, and uses of nitrogen.
• Describe the properties of nitric oxide, NO₂, and HNO₂.
• Carry out the laboratory preparation of nitrogen dioxide and ammonia.
• Discuss the nitrogen cycle including the use of bacteria and fertilizers and the role of nitrogen in plant and animal life.
• Outline the industrial preparation of nitric acid and ammonia.
• List the uses of ammonia.

Nitrogen

Symbol: N₂ Atomic Number: 7 Electron Arrangement: 2,5

• Main constituent of air by volume (78%).
• Present in all living matter as part of the structure of proteins and nucleic acids (DNA).
• Natural deposits of nitrate compounds, for example, sodium nitrate NaNO₃, occur in dry areas, Chile.
• Gas at room temperature, colorless, odorless, and relatively unreactive.
• Occurs as diatomic molecules N₂ with a triple bond N≡N linking two nitrogen atoms.
• Boils at -196°C and freezes at -210°C.
• Very strong N≡N bond causes nitrogen gas to be chemically unreactive at room temperature.
• Essential for plant growth but most plants cannot use nitrogen gas directly. Instead, they absorb nitrogen compounds from the soil.
• Animals obtain nitrogen by eating plants or other animals.

Laboratory preparation of nitrogen dioxide: Copper metal reacts with concentrated nitric acid to form nitrogen dioxide, a brown poisonous gas.

Ammonia: Formula NH₃

• Occurs as a colorless gas with a characteristic pungent odor.
• Dissolves readily in water to give a basic solution.
• Laboratory preparation of ammonia: Prepared by heating an ammonium salt with a strong base.
• Industrial preparation of ammonia: Produced by reacting nitrogen with hydrogen at high temperature and pressure using a catalyst.

• Uses of ammonia:
• Used to manufacture fertilizers, nitric acid, and explosives.
• Used as a refrigerant gas.

Nitric acid: Formula HNO₃

• Industrial preparation of nitric acid: Produced by the catalytic oxidation of ammonia to form nitrogen dioxide, which is then dissolved in water.
• Uses of nitric acid:
• Used to manufacture fertilizers, explosives, dyes, and drugs.

Nitrogen cycle: Nitrogen from the atmosphere is converted into nitrogen compounds that plants can absorb from the soil by several processes. Nitrogen-fixing bacteria living in the root nodules of leguminous plants such as peas and beans convert nitrogen into ammonia. Lightning causes nitrogen and oxygen in the air to react and form nitrogen oxides, which dissolve in rain water to form nitric acid. Nitric acid in rainwater reacts with minerals in the soil to form nitrates. Plants absorb nitrates and convert them into proteins. Animals obtain nitrogen by eating plants or other animals. When plants and animals die, decomposer bacteria convert nitrogen compounds in dead organisms into ammonia. Nitrifying bacteria convert ammonia into nitrites and then into nitrates, which plants can absorb. Denitrifying bacteria convert nitrates into nitrogen gas, which is returned to the atmosphere. Farmers add nitrogen-containing fertilizers to the soil to replace nitrogen removed by plant growth.

Carbon and Sulfur

• Discuss the properties and uses of carbon.
• Describe the preparation and properties of SO₂ and SO₂.
• Outline the Contact Process for the manufacture of sulfuric acid.
• Discuss the properties and uses of sulfuric acid.
• Study chlorine and its compounds.

Carbon

Symbol: C Atomic Number: 6 Electron Arrangement: 2,4

• Occurs naturally as diamond and graphite and in the atmosphere as carbon dioxide CO₂.
• Present in all living organisms as part of the structure of proteins, carbohydrates, fats and oils, and nucleic acids.
• Forms more compounds than any other element.
• Forms strong covalent bonds to itself and other nonmetals.

Sulfur

Symbol: S Atomic Number: 16 Electron Arrangement: 2,8,6

• Occurs naturally as the element and as sulfide and sulfate minerals.
• Extracted from underground deposits by the Frasch process. Superheated water at about 170°C and compressed air are pumped down into the sulfur deposit. The hot water melts the sulfur, and the compressed air forces the molten sulfur to the surface.
• Uses of sulfur:
• Used to manufacture sulfuric acid and other chemicals.
• Used in the vulcanization of rubber.
• Used in fungicides and insecticides.

Sulfur dioxide: Formula SO₂

• Preparation of sulfur dioxide: Burning sulfur in air or roasting sulfide ores.
• Properties of sulfur dioxide:
• Colorless gas with a choking odor.
• Dissolves in water to form sulfurous acid, H₂SO₃.
• Used as a bleach, disinfectant, and food preservative.

Sulfuric acid: Formula H₂SO₄

• Contact process for the manufacture of sulfuric acid: Sulfur is burned in air to form sulfur dioxide. The sulfur dioxide is then oxidized to sulfur trioxide using a catalyst. The sulfur trioxide is dissolved in concentrated sulfuric acid to form oleum, which is then diluted with water to produce sulfuric acid.
• Properties of sulfuric acid:
• Colorless, oily liquid.
• Strong acid.
• Dehydrating agent.
• Oxidizing agent.
• Uses of sulfuric acid:
• Used in the manufacture of fertilizers, detergents, explosives, dyes, and drugs.
• Used in petroleum refining and metal processing.

Halogen Compounds

Chlorine

Symbol: Cl₂ Atomic Number: 17 Electron Arrangement: 2,8,7

• Greenish-yellow poisonous gas with a choking odor.
• Occurs as diatomic molecules Cl₂.
• Fairly reactive and forms compounds with most elements.
• Most common chlorine compounds are ionic chlorides formed with metals.
• Chlorine reacts with hydrogen to form hydrogen chloride, HCl, a colorless gas. Hydrogen chloride dissolves readily in water to form hydrochloric acid.
• Uses of chlorine:
• Used as a disinfectant in water treatment.
• Used to manufacture other chemicals such as PVC plastic.

Organic Chemistry

• Investigate why there are numerous carbon compounds.
• Define the following terms as important concepts in organic chemistry: homologous series, isomerism in compounds of up to five carbon atoms, functional group, and saturation.
• Source the naturally occurring hydrocarbons.
• Outline the process of fractional distillation of petroleum.
• Outline the physical properties of alkanes.
• Investigate selected reactions of ethene: addition polymerization to form polyethene.
• Discuss the bromine test for unsaturation.
• Discuss the properties of alkenes.
• Investigate selected reactions of ethyne.
• Study brewing in Samoa.
• Investigate proteins, fats, and oils.

Organic chemistry is the study of carbon compounds. Carbon forms strong covalent bonds to itself and other nonmetals such as hydrogen, oxygen, nitrogen, sulfur, phosphorus, and the halogens. This explains why there are so many organic compounds.

Hydrocarbons are organic compounds containing only carbon and hydrogen. Alkanes are saturated hydrocarbons with the general formula CₙH₂ₙ₊₂. Alkanes with only a few carbon atoms are gases at room temperature. As the number of carbon atoms increases, the boiling points of the alkanes increase, and they become liquids and then waxy solids. Alkanes are relatively unreactive but burn in excess oxygen to form carbon dioxide and water. Alkenes are unsaturated hydrocarbons with at least one carbon-carbon double bond. Alkenes have the general formula CₙH₂ₙ. The simplest alkene is ethene, C₂H₄. Alkenes are more reactive than alkanes and undergo addition reactions. Alkynes are unsaturated hydrocarbons with at least one carbon-carbon triple bond. Alkynes have the general formula CₙH₂ₙ₋₂. The simplest alkyne is ethyne, C₂H₂. Alkynes are even more reactive than alkenes and also undergo addition reactions.

Fractional distillation of petroleum: Petroleum is a complex mixture of hydrocarbons formed over millions of years from the remains of marine organisms. Fractional distillation is used to separate petroleum into fractions based on their boiling points. The fractions have different uses depending on their physical properties.

Brewing in Samoa: Vailima Breweries Ltd. produces the popular beer Vailima in Samoa. Brewing involves the fermentation of sugars by yeast to produce ethanol and carbon dioxide.

Proteins, fats, and oils: Proteins are polymers of amino acids. Fats and oils are esters of glycerol and fatty acids.

Hydrocarbon Reactions

Functional groups: A functional group is an atom or group of atoms that is responsible for the characteristic reactions of a homologous series of organic compounds. For example, the functional group of the alkenes is the C=C double bond. The functional group of the alcohols is the -OH group.

Saturation: Organic compounds are saturated if they contain only single bonds between carbon atoms. Alkanes are saturated hydrocarbons. Organic compounds are unsaturated if they contain at least one double or triple bond between carbon atoms. Alkenes and alkynes are unsaturated hydrocarbons.

Bromine test for unsaturation: Bromine reacts with unsaturated hydrocarbons by adding across the double or triple bond. The bromine water, which is brown, becomes colorless when shaken with an unsaturated hydrocarbon.

Addition polymerization: Alkenes can react with each other in addition reactions to form long chains called polymers. Polyethene is formed by the addition polymerization of ethene.

Food Chemistry

Proteins: Polymers of amino acids. Amino acids contain two functional groups, an amine group, -NH₂, and a carboxylic acid group, -COOH. Proteins are essential components of living organisms.

Fats and oils: Esters of glycerol and fatty acids. Glycerol has three -OH groups, and fatty acids have a -COOH group. Fats and oils are important energy sources for living organisms. Fats are solid at room temperature, and oils are liquid. Fats generally contain saturated fatty acid groups, and oils generally contain unsaturated fatty acid groups.