Science (Chemistry) 5088/03 Paper 3, Oct/Nov 2026

“A chemical formula is like a recipe that shows the ratio of elements that are combined in one unit of a substance. It is made up of: • chemical symbol(s) that indicate the element(s) present; and • subscript(s) indicating the number of atoms of each element present.”

“These subscripts indicate that: • there are two oxygen atoms in one molecule of carbon dioxide; and • there are two hydrogen atoms in one molecule of water.”

“Name of Substance Chemical Formula Atoms Present ammonia NH₃ 1 nitrogen and 3 hydrogen copper(II) sulfate CuSO₄ 1 copper, 1 sulfur and 4 oxygen sodium hydroxide NaOH 1 sodium, 1 oxygen and 1 hydrogen hydrochloric acid HCl 1 hydrogen and 1 chlorine ethanoic acid CH₃COOH 4 hydrogen, 2 carbon and 2 oxygen”

“When there is one carbon atom in the alkane molecule (n = 1): • The number of hydrogen atoms will be 2(1) + 2 = 4. • Hence, the formula of the alkane is CH₄ (Figure 16.7).”

“Photosynthesis Plants take in carbon dioxide and water to make glucose and oxygen in the presence of sunlight. 6CO₂(g) + 6H₂O(l) → C₆H₁₂O₆(s) + 6O₂(g) Glucose is used in other processes that sustain plant life.”

“The element carbon is essential for life. All organic compounds are made of carbon skeletons that originate from carbon dioxide and water during photosynthesis.”

“Carboxylic acids are organic compounds with the carboxyl (–COOH) functional group. All carboxylic acids have this functional group.”

“The formulae of the first three members of the carboxylic acid homologous series are shown in Table 17.3.”

“propanoic acid When n = 2, C₂H₅COOH C₃H₆O₂ CH₃CH₂COOH”

“When a compound is made up of only covalently-bonded carbon and hydrogen atoms, we call it a hydrocarbon.”

“Fossil fuels are mainly organic compounds that contain only carbon and hydrogen, known as hydrocarbons. When fossil fuels are burnt in the presence of oxygen, the main products formed are carbon dioxide and water.”

“Natural gas is a gaseous fossil fuel consisting mainly of methane.”

“Alkanes are hydrocarbons that contain only carbon–carbon single bonds and carbon–hydrogen bonds.”

“In alcohol fermentation, yeast acts on carbohydrates such as glucose to produce ethanol and carbon dioxide. This process takes place in the absence of oxygen.”

“The fermentation of glucose can be represented by the following equation: glucose → ethanol + carbon dioxide C₆H₁₂O₆(aq) → 2C₂H₅OH(aq) + 2CO₂(g)”

“Ethanol and carbon dioxide gas are the products of the fermentation reaction.”

“2 Ethanol can be obtained from the fermentation of glucose. (a) What are the conditions required in this reaction? (b) What other product is formed in the fermentation of glucose?”

“Most reactive potassium (K) sodium (Na) calcium (Ca) magnesium (Mg) carbon (C) zinc (Zn) iron (Fe) lead (Pb) hydrogen (H) copper (Cu) silver (Ag) Least reactive Figure 12.2 The reactivity series”

“In the reactivity series, metals are arranged from the most reactive to the least reactive.”

“From Table 11.4, we can see that the reactivity of alkali metals increases down the group (Figure 11.10). These reactions also explain why they are called alkali metals. All the metals react with water to form soluble bases (alkalis). Order of reactivity: lithium < sodium < potassium < rubidium < caesium < francium”

“There are two main methods for extracting metals from their ores: • reducing the metal compound (ore) to the metal using carbon • electrolysis — using electricity to decompose the molten metal compound (ore) to the metal The method to extract the metals from their ores is determined by the position of the metal in the reactivity series. The more reactive the metal is, the harder it is to extract the metal from its ore. The oxides of metals lower in the reactivity series are more easily reduced by carbon than the oxides of more reactive metals.”

“Metals high up in the reactivity series are reactive. Their compounds are very difficult to break down. These metals can only be extracted by electrolysis.”

“Metals which are less reactive, i.e. from the middle to the bottom of the reactivity series can be extracted by reducing their oxides with carbon.”

“Metals above carbon in the reactivity series can be extracted from their ores using electrolysis. Metals below carbon can be extracted by reduction with carbon.”

“Some metals such as zinc and iron do not react with cold water but react with steam. More reactive metals react more violently with steam. metal + steam → metal oxide + hydrogen”

“magnesium Reacts very slowly with cold water to form magnesium hydroxide and hydrogen gas. A test tube of hydrogen gas is produced only after a few days. Mg(s) + 2H₂O(l) → Mg(OH)₂(s) + H₂(g) Hot magnesium reacts violently with steam to form magnesium oxide (a white solid) and hydrogen gas. A bright white glow is produced during the reaction. Mg(s) + H₂O(g) → MgO(s) + H₂(g) zinc No reaction occurs. Hot zinc reacts readily with steam to produce zinc oxide and hydrogen. Zinc oxide is yellow when hot and white when cold. Zn(s) + H₂O(g) → ZnO(s) + H₂(g) iron* No reaction occurs. Red-hot iron reacts slowly with steam to form iron oxide. The iron must be constantly heated for the reaction to progress. 3Fe(s) + 4H₂O(g) ⇌ Fe₃O₄(s) + 4H₂(g)”

“zinc (Zn) no reaction with cold water reacts readily with steam reacts moderately fast with dilute hydrochloric acid”

“1 Name a metal that reacts with steam but not with cold water to give an oxide and hydrogen.”

“The presence of both oxygen (in air) and water are necessary for rusting to occur.”

“Metals may react with oxygen in the air and water and corrode. When an object made of iron is left in moist air for some time, a red-brown substance slowly forms on the surface of the metal (Figure 12.11). This substance is called rust and has the chemical name hydrated iron(III) oxide.”

“Rusting or the corrosion of iron is the process that produces rust. The equation for rusting is: iron + oxygen + water → hydrated iron(III) oxide (rust)”

“From the results of Let’s Investigate 12A, we can conclude that: • both oxygen and water are needed for rusting to occur; and • the presence of sodium chloride increases the rate of rusting. Other than sodium chloride, acidic substances such as sulfur dioxide and carbon dioxide also speed up the rusting process.”

“Table 12.5 Results of Let’s Investigate 12A Test Tube Conditions Observation Is Water Present? Is Oxygen Present? A yes no no rusting B no yes no rusting C yes yes nails rusted D yes yes nails rusted heavily”

“The barrier methods prevent rusting by keeping iron and steel away from oxygen (in air) and water (Figure 12.14).”

“The paint protects the iron gate from moisture and oxygen in the air. It prevents the gate from rusting.”

“Rusting only occurs in iron and steel. It involves both water and oxygen.”

“The relative formula mass (Mr) of an ionic compound is the average mass of one unit of that compound relative to 1/12 the mass of an atom of carbon-12.”

“Relative formula mass is calculated the same way as relative molecular mass and also has the symbol Mr (Table 7.2).”

“Ionic compounds do not exist as molecules. We use the term relative formula mass to refer to the relative mass of an ionic compound.”

“To find the relative molecular mass (Mr) of a substance, we add together the relative atomic masses (Ar) of all the atoms in its chemical formula (Table 7.1).”

“The relative molecular mass (Mr) of a molecular substance is the average mass of one molecule of that substance relative to 1/12 the mass of an atom of carbon-12.”

“Define relative molecular mass, Mr, and calculate relative molecular mass (and relative formula mass) as the sum of relative atomic masses.”

“The relative atomic mass (Ar) of an element is the average mass of one atom of that element relative to 1/12 the mass of an atom of carbon-12.”

“To find the relative molecular mass (Mr) of a substance, we add together the relative atomic masses (Ar) of all the atoms in its chemical formula (Table 7.1).”

“Relative formula mass is calculated the same way as relative molecular mass and also has the symbol Mr (Table 7.2).”

“Calculate the mass, in grams, of 4.5 moles of lead(II) sulfate (PbSO₄). Answer Molar mass of PbSO₄ in g/mol = 207 + 32 + 4(16) = 303 g/mol Mass of PbSO₄ in g = number of moles of PbSO₄ × molar mass of PbSO₄ in g/mol = 4.5 × 303 = 1360 g”

“Since the relative atomic mass (Ar), relative molecular mass and relative formula mass (Mr) are comparisons of mass, we do not assign them with a unit.”

“One mole of any substance has a mass equal to its relative atomic mass (Ar), relative molecular mass or relative formula mass (Mr) in grams. This mass is the molar mass of the substance. The unit for molar mass is g/mol (Table 7.4).”

“Number of moles of a substance = mass of the substance in g / molar mass of the substance in g/mol”

“Step 1 Calculate the number of moles of the reactant or product using the given information. In this case, we find the number of moles of MgCO₃.”

“Step 2 Find the number of moles of the substance we are solving for using the mole ratio from the balanced equation and the result from step 1. In this case, we need to find the number of moles of CO₂. From the equation, 1 mol of MgCO₃ will form 1 mol of CO₂. Number of moles of CO₂ = number of moles of MgCO₃ = 2 mol”

“A balanced chemical equation shows the mole ratio of the reactants and products in a chemical reaction. We can use these ratios to calculate certain measurements such as mass and volume.”

“This ratio between the quantities of the different substances in the reaction is the mole ratio. Combining this ratio with the mole concept allows us to find the masses, volumes and concentrations of substances in a chemical reaction.”

“Iron(III) oxide, Fe₂O₃(s) reacts with carbon monoxide gas to produce iron and carbon dioxide gas. Calculate the mass of iron produced when 2000 g of iron(III) oxide is reacted completely. Fe₂O₃(s) + 3CO(g) → 2Fe(s) + 3CO₂(g)”

“From the equation, 1 mol of Fe₂O₃ produces 2 mol of Fe. Number of moles of Fe = 2 × 12.5 = 25.0 mol Mass of Fe = number of moles of Fe × Ar of Fe = 25.0 × 56 = 1400 g”

“Mass of PbSO₄ in g = number of moles of PbSO₄ × molar mass of PbSO₄ in g/mol”

“6 Acids react with carbonates to form a salt, water and carbon dioxide gas. The general equation for the reaction of carbonates with acids is: carbonate + acid → salt + water + carbon dioxide For example, when hydrochloric acid is added to calcium carbonate, the following reaction occurs: calcium carbonate + hydrochloric acid → calcium chloride + carbon dioxide + water CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g) We can test for carbon dioxide gas by bubbling the gas through limewater.”

“A salt is an ionic compound that can be obtained from acid reactions. Different salts are formed when various acids are used in the reactions: • Sulfates are formed from sulfuric acid. • Nitrates are formed from nitric acid. • Chlorides are formed from hydrochloric acid.”

“Isotopes are atoms of the same element that have the same proton (atomic) number but different nucleon (mass) numbers. This means that they have different numbers of neutrons.”

“We have learnt that: • protons and neutrons are found clustered together in an atom’s nucleus; and • the number of protons is fixed for all atoms of a given element. However, the number of neutrons in the nucleus can vary. Neutrons do not affect the electrical charge of the atom, although they affect the total mass.”

“Isotopes of the same element have the same number of protons. Hence, they also have the same number of electrons. This means that they undergo the same types of chemical reactions even though they do not have the same mass. However, isotopes may have different physical properties such as density, melting point and boiling point.”

“This number is also called the atomic number. It tells us what element the atom comes from so that we know what type of atom it is. All atoms of the same element have the same proton number.”

“The proton number (Z) of an atom is the number of protons in its nucleus.”

“The nucleon number (A) is the total number of protons and neutrons in the nucleus. Nucleon number (A) = proton number (Z) + number of neutrons”

“We can use proton and nucleon numbers to calculate the number of neutrons in an atom. Number of neutrons = Nucleon number (A) – Proton number (Z)”

“There are 6 protons and 6 neutrons in the nucleus. Hence, carbon has a proton number of 6 and a nucleon number of 12 (6 protons + 6 neutrons). Since a carbon atom is neutral, it will have the same number of electrons as protons (6 electrons).”

“We can represent the proton and nucleon numbers of an atom using nuclide notation (Figure 3.8).”

“Thought Process • This atom has a nucleon number of 184. There must be a total of 184 protons and neutrons. • The proton number of 74 tells us that it has 74 protons. • We can subtract the proton number from the nucleon number to find the number of neutrons. • The atom is electrically neutral, so the number of electrons is the same as the number of protons. Answer Number of neutrons in tungsten = 184 – 74 = 110 An atom of tungsten has 74 protons, 110 neutrons and 74 electrons.”

“Atoms are electrically neutral. The number of protons in an atom is always the same as the number of electrons. Number of protons in an atom = number of electrons in that atom”

“A negative ion or anion is formed when an atom or combination of atoms gains one or more electrons. The ion formed is negatively charged, because it has more electrons than protons. Most non-metal atoms will gain electrons and form negative ions.”

“After gaining 1 electron, the chlorine atom is no longer electrically neutral. It now has 17 protons and 18 electrons, so it has a charge of –1. The negative ion formed is known as a chloride ion. It has the symbol Cl⁻.”

“Negative ions (anions) have a net negative charge and usually have a noble gas electronic configuration.”

“Recall that electrons have a charge of –1. Hence, when an atom gains electrons, it will have more negative charges, –1 for each electron gained.”

“An oxide ion has the symbol O²⁻. This means that it has 8 protons and 8 neutrons (just like an oxygen atom), but with a net charge of –2. Thus, it has a total of 2 more electrons than protons. It has 8 + 2 = 10 electrons. Therefore, the charge numbers on an ion indicates • the number of electrons that have been lost in a positive ion; and • the number of electrons that have been gained in a negatively charged ion.”

“Protons • Carry a positive charge (+1) Neutrons • Neutral particles with no charge (0) • Found in the nucleus between protons Electrons • Extremely tiny particles that move very quickly • Carry a negative charge (–1)”

“Halogen Melting Point / °C Boiling Point / °C Appearance at r.t.p. chlorine –101 –34 yellow-green gas bromine –7 59 red-brown liquid iodine 114 184 purple-black solid”

“From Table 11.5, we can see that going down the group, • the melting and boiling points of the halogens increase; and • the colours become darker (colour intensities increase).”

“Figure 11.12 The melting points, boiling points and colour intensities of the halogens increase down the group.”

“Unlike the alkali metals, the reactivity of halogens decreases down the group.”

“Order of reactivity: fluorine > chlorine > bromine > iodine > astatine > tennessine Halogens undergo displacement reactions with halide solutions.”

“In a displacement reaction, 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.”

“Chlorine displaces bromine from a bromide solution. Chlorine displaces iodine from an iodide solution. Bromine displaces iodine from an iodide solution.”

“Figure 11.13 Going down Group 17, the halogens become less reactive.”

“The general molecular formula of alkenes with one carbon−carbon double covalent bond is CnH2n, where n is the number of carbon atoms in each molecule (n = 2, 3 and so on). When there are two carbon atoms in each molecule (n = 2): • The number of hydrogen atoms will be 2(2) = 4. • Hence, the formula of the alkene is C₂H₄.”

“Alkenes are hydrocarbons that contain carbon–carbon double (C=C) bonds.”

“Addition of Hydrogen (Hydrogenation) Alkenes react with hydrogen to form alkanes under the following conditions: • a temperature of 150 °C • in the presence of a catalyst such as nickel This reaction is known as hydrogenation. For example, ethene reacts with hydrogen as follows: ethene + hydrogen → ethane C₂H₄(g) + H₂(g) → C₂H₆(g)”

“Addition Reactions The carbon–carbon double bonds in alkenes are reactive. Hence, alkenes will readily undergo addition reactions. The general equation is as follows:”

“• Other atoms can add to the carbon atoms in the C=C bond. • The carbon–carbon double bond breaks to form new single bonds. Thus, each carbon atom which was originally part of the C=C bond forms four single covalent bonds.”

“A carbon–carbon single bond is a single covalent bond between two carbon atoms. Figure 16.6 shows the structure of ethane.”

“Full structural formula of ethane, an example of an alkane.”

“Combustion Like alkanes, alkenes burn in excess oxygen to produce carbon dioxide and water. For example, ethene burns in excess oxygen as follows: ethene + oxygen → carbon dioxide + water C₂H₄(g) + 3O₂(g) → 2CO₂(g) + 2H₂O(g)”

“Compared to alkanes, the percentage by mass of carbon in alkene molecules is higher. Alkenes are more likely to undergo incomplete combustion. Incomplete combustion produces carbon monoxide and carbon. Hence, alkenes usually burn with a sootier flame than alkanes (Figure 16.15).”

“It is easier to deduce the structure of an addition polymer if the structural formula of its alkene monomer is written in the following way: [diagram: C=C with a box on each of its four bonds] (where [box] = hydrogen, alkyl group, halogen etc.) For example, the formula of propene can be written as [diagram]. The polymer formed, which is poly(propene), can then be represented as [diagram].”

“Addition Polymerisation When monomers join together without losing any molecules or atoms, the polymer formed is called an addition polymer. This process is called addition polymerisation. Addition polymerisation occurs when unsaturated monomers join together without losing any molecules or atoms to form addition polymers. The monomer used for addition polymerisation is usually unsaturated. This means that it contains either a carbon–carbon double bond or a carbon–carbon triple bond. Thus, alkenes can undergo addition polymerisation. At high pressure and temperature, thousands of alkene molecules join together to form polymers.”

“Ethene is unsaturated as it contains a carbon–carbon double bond. Poly(ethene) is saturated as it contains only carbon–carbon single bonds.”

“A polymer is usually made up of similar units joined together. These units are known as repeat units.”

“The formula of a polymer can be written using its repeat unit. The letter n represents a large number. Many ethene molecules are required to form each poly(ethene) molecule.”

“2 Figure 18.14 represents part of a polymer chain. Draw the structural formula of its monomer.”

“Physical Method The physical method of recycling plastics is also known as mechanical recycling. Using this method, the chemical composition of the recycled plastic remains unchanged. This means that if waste poly(ethene) bottles are collected, the recycled plastic products will still be made of poly(ethene). After pre-treatment, small pieces of plastics such as poly(ethene) are melted, cooled, pulled into long, thin strands, and cut into pellets (Figure 18.18). These pellets can then be made into new products.”

“Recycling of plastics can be carried out using the physical method or the chemical method. Before recycling is carried out, the plastic waste needs to undergo pre-treatment (Figure 18.17). This is because not all plastics can be recycled. The recycling method for different types of plastic can be different.”

“Chemical Method The chemical method of recycling plastics converts plastic waste into different raw materials through chemical reactions such as cracking. Plastic waste, such as poly(ethene) waste, can undergo cracking to form short chains of alkanes and alkenes. The short-chain alkanes can be used as fuel, while the short-chain alkenes can be used to make other useful chemicals. Cracking requires a high temperature and the presence of a catalyst.”

“Recall from Chapter 16: Cracking can be used to break down long-chain hydrocarbons into smaller hydrocarbon molecules.”

“Cracking is a process in which larger hydrocarbon molecules, usually alkanes, are broken down into smaller hydrocarbon molecules.”

“Particle Arrangement | very closely packed in an orderly manner | closely packed in a disorderly manner | very far apart in a disorderly manner”

“Particle Movement | vibrate and rotate about fixed positions | slide past each other freely throughout the liquid | move quickly and randomly in any direction”

“During sublimation, the particles gain sufficient thermal energy from the surroundings to overcome the strong forces of attraction. Particles in the solid change from vibrating about fixed positions to moving freely and randomly in the gas.”

“The three physical states of matter (i.e. solid, liquid and gas) could be differentiated based on the arrangement and movement of the particles.”

“For example, pure water boils at 100 °C. Adding salt to water raises its boiling point. Similarly, pure water freezes at 0 °C, but adding salt to water lowers its freezing point.”

“Mixtures containing two or more substances will melt or boil over a range of temperatures. The greater the amount of impurities, the larger the change in melting or boiling points, and the wider the range.”

“A pure substance has a specific melting and/or boiling point under fixed conditions.”

“For example, pure water boils at 100 °C. Adding salt to water raises its boiling point.”

“The greater the amount of impurities, the larger the change in melting or boiling points, and the wider the range.”

“Figure 5.17 shows the arrangement of atoms in pure copper and in bronze. Bronze is an alloy that is about 90 % copper and 10 % tin. It contains atoms of different sizes, as each tin atom is about 1.6 times the size of a copper atom. Therefore, the regular arrangement of copper atoms is disrupted.”

“Alloys are harder and stronger than pure metals. The irregular arrangement of atoms in alloys leads to the difference in properties between alloys and pure metals (Figure 5.18).”

“Pure metals have a regular structure, as shown in Figure 5.19. … metal atoms of the same size”

“Metals can exist as pure elements. They can also be mixed with other elements to produce alloys. … An alloy is a mixture of a metal with one or more other elements.”

“Brass | copper and zinc”

“Describe the general physical properties of metals as solids having high melting and boiling points, malleable, and good conductors of heat and electricity.”

“Metals usually have high melting and boiling points. With the exception of mercury, metals are solids at room temperature. … Both metals and alloys are good conductors of heat. … Metals and alloys usually have high melting and boiling points. They are good conductors of heat.”

“Pure metals have a regular structure, as shown in Figure 5.19. If enough force is applied, the layers of atoms can slide over one another easily. Thus, when hammered, pure metals can be bent or flattened into a thin sheet (malleable). They can also be pulled into a wire without breaking (ductile). … Pure metals are malleable and ductile.”

“Metal atoms have valence electrons that are free to move throughout the structure. Hence, most metals and alloys can conduct electricity. … Metals and alloys are good electrical conductors.”

“An element is a pure substance that cannot be broken down into two or more simpler substances by chemical methods. Elements can appear as: atoms, if they are not chemically bonded to each other (e.g. helium); or molecules, if they are covalently bonded (e.g. oxygen molecules, O₂).”

“A compound is a pure substance containing two or more elements that are chemically combined in a fixed ratio. The elements are chemically bonded together to form either a covalent compound or an ionic compound.”

“What Is It Made Of? | only one element | two or more elements that are chemically combined | two or more elements and/or compounds that are not chemically combined”

“Compounds can be separated into their constituent elements using chemical methods. The separation can be achieved by: exposing the compound to strong heat (thermal decomposition); or passing an electric current through the compound (electrolysis).”

“Add dilute nitric acid, then add aqueous barium nitrate. | A white precipitate of barium sulfate, BaSO₄, is formed.”

“From Table 9.2, we can see that dilute nitric acid was added in the test to identify chloride and sulfate ions. This is to remove the presence of any carbonate and hydroxide ions, ensuring that there is no error in the analysis obtained.”

“Different salts are formed when various acids are used in the reactions: Sulfates are formed from sulfuric acid. Nitrates are formed from nitric acid. Chlorides are formed from hydrochloric acid.”

“Steel | iron and carbon | to increase hardness | warehouse shelves”

“Stainless steel is an alloy of iron with small amounts of other elements such as chromium and carbon.”

“Green precipitate of iron(II) hydroxide, Fe(OH)₂ | Precipitate is insoluble in excess. The precipitate turns brown on standing.”

“Precipitation is the formation of insoluble solids from two solutions. The precipitate formed in each of the reactions in Table 9.1 is the hydroxide of the metal ion.”

“The reagents used to identify cations are aqueous sodium hydroxide, NaOH(aq) and aqueous ammonia, NH₃(aq). All cations (except Na⁺, K⁺ and NH₄⁺) give precipitates with these alkalis.”

“For example, when iron filings are added to a solution of copper(II) sulfate (Figure 12.7), copper metal is displaced out of the solution as a pink solid (when freshly formed) which turns into a red-brown solid; and the solution turns green. … green aqueous iron(II) sulfate”

“Acids react with reactive metals to form a salt and hydrogen gas. … We can test for hydrogen gas by placing a burning splint at the mouth of the test tube. Hydrogen gas extinguishes the burning splint with a “pop” sound (Figure 8.3).”

“colourless and odourless | Place a burning splint at the mouth of the test tube. | The burning splint is extinguished with a “pop” sound.”

“Metals above hydrogen in the reactivity series react with hydrochloric acid to produce hydrogen gas. … iron | Reacts slowly to give hydrogen gas.”

“We can test for carbon dioxide gas by bubbling the gas through limewater. Carbon dioxide gas forms a white precipitate with limewater (Figure 8.6).”

“Bubble the gas through limewater. | A white precipitate is formed. The precipitate dissolves upon further bubbling.”

“The gas given off forms a white precipitate of calcium carbonate, CaCO₃, with limewater. Carbon dioxide gas is produced.”

“A balanced chemical equation shows the expected ratio of reactants to products. Every atom on the left-hand side of the arrow must also appear on the right-hand side of the arrow. However, it may be bonded to different atoms.”

“Write the chemical formulae of the reactants and products in an unbalanced chemical equation. … Add molecule(s) of the unbalanced species to balance the equation. Do not change the chemical formulae or subtract atoms. … Add the state symbols into the balanced chemical equation.”

“Common states are: solid (s) liquid (l) gas (g) aqueous (aq)”

“When magnesium ribbon is added to dilute sulfuric acid, the following reaction occurs: magnesium + dilute sulfuric acid → magnesium sulfate + hydrogen Mg(s) + H₂SO₄(aq) → MgSO₄(aq) + H₂(g)”

“iron | Reacts slowly to give hydrogen gas. | Fe(s) + 2HCl(aq) → FeCl₂(aq) + H₂(g)”

“iron(III) ion + hydroxide ion → iron(III) hydroxide Fe³⁺(aq) + 3OH⁻(aq) → Fe(OH)₃(s)”

“Elements in the same group have the same number of valence electrons. Elements in Groups 1 and 2 (e.g. lithium and beryllium) have 1 and 2 valence electrons respectively. Elements in Groups 13 to 18 (except helium) have 3 to 8 valence electrons respectively. The number of valence electrons is the group number minus 10 (e.g. chlorine has 17 – 10 = 7 valence electrons).”

“Elements with the same number of valence electrons have similar chemical properties. We can thus deduce that elements in the same group have similar chemical properties (See Section 11.2).”

“Let us look at the element with a proton number of 7 in Table 11.1(a). We can see that the element would be in Period 2, Group 15 (i.e. nitrogen).”

“A negative ion or anion is formed when an atom or combination of atoms gains one or more electrons. The ion formed is negatively charged, because it has more electrons than protons. Most non-metal atoms will gain electrons and form negative ions.”

“Negative ions (anions) have a net negative charge and usually have a noble gas electronic configuration.”

“Describe the formation of (a) positive ions by losing electron(s); (b) negative ions by gaining electron(s); and that these ions usually have the electronic configuration of a noble gas.”

“Positive ions seldom have a net charge higher than +3 because it gets increasingly harder for atoms to lose more electrons. Negative ions seldom have a net charge lower than –3 because it is difficult for atoms to gain and keep a large number of electrons.”

“Ions of opposite charges have strong electrostatic forces of attraction between them. A large amount of energy is needed to overcome these forces. Hence, ionic compounds tend to have very high melting and boiling points. They are solids at room temperature. … Ionic compounds have high melting and boiling points.”

“When an ionic compound is in the solid state, the ions in the lattice are only able to vibrate about their fixed positions. Hence, the ions are not mobile. However, when the compound is molten (in the liquid state) or in the aqueous state, the ions are mobile and can conduct electricity. … Ionic compounds conduct electricity when molten or in aqueous solution, but not in the solid state.”

“The strong forces of attraction between oppositely charged ions make ionic compounds resistant to deforming. Ionic compounds tend to be hard, as these forces cause the ions to resist motion. … Hence, ionic compounds are hard but brittle. … Ionic compounds are physically hard but brittle.”

“Most ionic compounds are soluble in water and insoluble in organic solvents.”

“Ionic compounds have giant ionic crystal lattice structures (Figure 5.6). The lattice consists of alternating positive and negative ions. These ions are held together by electrostatic forces of attraction.”

“Generally, we can determine how atoms will bond covalently by comparing the number of electrons each atom shares. Suppose one atom shares x electrons to achieve a noble gas electronic configuration. This atom will form a covalent compound with one or more other atoms that can “match” this sharing. For example, ammonia is made up of hydrogen and nitrogen atoms. Let us compare the number of electrons each hydrogen atom and nitrogen atom shares.”

“Hydrogen atom: It has 1 electron in its outermost shell. It shares 1 electron to have a duplet electronic configuration. Nitrogen atom: It has 5 electrons in its outermost shell. It shares 3 electrons to have an octet electronic configuration.”

“Therefore, three hydrogen atoms will share electrons with one nitrogen atom (Figure 4.20).”

“The overlap shows that this pair of electrons is shared between the two atoms. These electrons are called a bonding pair. A single covalent bond has been formed.”

“Each electron shell can only hold a certain number of electrons. The first electron shell can hold a maximum of 2 electrons. The second and third shells can usually have a maximum of 8 electrons each.”

“Let us look at the element with a proton number of 7 in Table 11.1(a). We can see that the element would be in Period 2, Group 15 (i.e. nitrogen).”

“Methane is the simplest covalent compound formed between carbon and hydrogen. Draw a dot-and-cross diagram for methane. (Only the outermost electrons need to be shown.)”

“Exothermic changes: release energy to the surroundings; and increases the temperature of the surroundings.”

“The neutralisation reaction between sodium hydroxide and hydrochloric acid is exothermic (Figure 13.10). When sodium hydroxide and hydrochloric acid react, energy is released to the surroundings and the temperature of the surroundings increases.”

“Exothermic changes involve the release of energy into the surroundings. Hence, the temperature of the surroundings increases.”

“Describe exothermic changes as processes and chemical reactions that release energy to the surroundings.”

“At higher temperatures, the rate of reaction is affected in two ways: … At higher temperatures, reactant particles have more kinetic energy. This means that they move and vibrate faster, increasing the frequency of effective collisions. … At higher temperatures, more reactant particles possess energy that is greater than or equal to activation energy. This increases the frequency of effective collisions. Hence, higher temperatures result in faster reaction rates (Figure 14.19).”

“An increase in the temperature of reactants results in an increase in the rate of reaction.”

“When a chemical reaction occurs between particles, the collision is said to be an effective collision. The activation energy is the least amount of energy that colliding reactant particles must possess to react with each other. For an effective collision to occur, the reactant particles must collide with energy that is greater than or equal to the activation energy (Figure 14.10).”

“As the temperature increases, reactant particles have more kinetic energy. More reactant particles possess energy that is greater than or equal to activation energy. This increases the frequency of effective collisions.”

“In this test, dark purple aqueous potassium manganate(VII) that has been acidified with a little sulfuric acid is used. A few drops of the purple solution are added to a solution of the unknown substance. … In a positive test, manganate(VII) ions are reduced to manganese(II) ions by the reducing agent (Figure 10.16).”

“The test for reducing agents can also be applied to gases. Use a strip of filter paper with a spot of acidified potassium manganate(VII) solution. If the spot decolourises when exposed to the gas, the gas is a reducing agent. … Positive test (colourless solution)”

“Place a piece of filter paper soaked with acidified potassium manganate(VII) at the mouth of the test tube. | The purple acidified potassium manganate(VII) turns colourless.”

“A reducing agent is able to: reduce another substance; decrease the oxidation state of an element in another substance; and lose electrons to the other substance. Thus, a reducing agent causes reduction while being oxidised itself.”

“From Table 8.1, we can see that all the acids produce hydrogen ions, H⁺, in aqueous solutions. When acids are dissolved in water, the acid molecules ionise, or dissociate, to produce H⁺ ions. This H⁺ ion is mainly responsible for the behaviour of acids. … An acid is a substance that produces hydrogen ions, H⁺, in aqueous solutions.”

“pH < 7: The concentration of H⁺ ions is higher than that of OH⁻ ions, so the solution is acidic.”

“A successful test for acids (e.g. blue litmus paper) indicates the presence of H⁺ ions.”

“Universal Indicator is a mixture of coloured dyes. It may come in the form of a solution or paper. It gives different colours depending on how acidic or alkaline a solution is (Figure 8.11).”

“A set of numbers from 0 to 14 is used to indicate whether a solution is acidic, neutral or alkaline. This is called the pH scale. The pH of a solution may be determined by using: an indicator; a pH meter; or a pH sensor connected to a data logger (Figure 8.12).”

“The colours of the different pH values in Universal Indicator mimic the colours of the rainbow.”

“The pH of a solution is related to the concentration of hydrogen ions or hydroxide ions present in a solution. An acid with a lower pH value is more acidic. It has a higher concentration of hydrogen ions.”

“Number of moles of KOH reacted = concentration of KOH × volume of KOH = 0.100 × (30.0 ÷ 1000) = 0.003 mol”

“Concentration in mol/dm³ = number of moles of solute [over] volume of solvent in dm³”

“1 dm³ is equivalent to 1 litre. There are 1000 cm³ in 1 dm³.”

“From the equation, 2 mol of KOH will react with 1 mol of H₂SO₄. Number of moles of H₂SO₄ reacted = 0.003 ÷ 2 = 0.0015 mol”

“This ratio between the quantities of the different substances in the reaction is the mole ratio. Combining this ratio with the mole concept allows us to find the masses, volumes and concentrations of substances in a chemical reaction.”

“Find the number of moles of the substance we are solving for using the mole ratio from the balanced equation and the result from step 1. In this case, we need to find the number of moles of CO₂. From the equation, 1 mol of MgCO₃ will form 1 mol of CO₂.”

“One mole of any gas occupies a volume of 24 dm³ at room temperature and pressure (r.t.p.).”

“Calculate the volume occupied by 2 moles of oxygen gas (O₂), measured at room temperature and pressure (r.t.p.). Answer Volume of O₂ in dm³ = number of moles of O₂× 24 dm³ = 2 × 24 = 48 dm³”

“Lastly, calculate the answer using the number of moles obtained in step 2. In this case, we need to find the volume of CO₂. Volume of CO₂ in dm³ = number of moles of CO₂ × 24 = 2 × 24 = 48 dm³”

“The concentration of a solution shows the number of particles in a given volume. As the concentration of reactants increases, the number of reacting particles per unit volume increases. They collide more frequently and hence, the frequency of effective collisions also increases (Figure 14.11).”

“An increase in concentration of reactants increases the rate of reaction.”

“As the concentration of reactants increases, the number of reacting particles per unit volume increases. The reacting particles collide more frequently, increasing the frequency of effective collisions.”

“When a chemical reaction occurs between particles, the collision is said to be an effective collision. … When more effective collisions occur between reactant particles, the rate of reaction increases.”