Comprehensive Form 2 Chemistry Study Notes for Zambian O-Level Preparation

Chemical Bonding and Atomic Stability

Chemistry serves as the central science because it acts as a bridge between the disciplines of physics and biology, providing the fundamental understanding for everything from the composition of drinking water to the chemical reactions within batteries that power mobile devices. The core reason atoms engage in chemical bonding is to achieve stability. With the exception of noble gases, individual atoms are inherently unstable due to incomplete outer energy levels, known as valence shells. To rectify this, atoms react to achieve a stable electronic configuration, specifically a duplet containing $2$ electrons or an octet containing $8$ electrons.

Types of Chemical Bonding and Structures

There are three primary types of chemical bonding: ionic, covalent, and metallic. Ionic bonding, also referred to as electrovalent bonding, occurs through the complete transfer of electrons from a metal atom to a non-metal atom. This process results in the formation of positive cations and negative anions, which are held together by strong electrostatic forces. Examples include sodium chloride ($NaCl$), calcium oxide ($CaO$), and magnesium chloride ($MgCl_2$). In salt crystals like $NaCl$, billions of $Na^+$ and $Cl^-$ ions lock into a rigid three-dimensional giant ionic lattice, keeping the substance solid. When dissolved in water, these ions break free, which allows an electric current to flow.

Covalent bonding involves the sharing of valence electron pairs between non-metal atoms. These can form single, double, or triple shared electron pairs. Common examples include hydrogen gas ($H_2$), oxygen gas ($O_2$), nitrogen gas ($N_2$), water ($H_2O$), and carbon dioxide ($CO_2$). Metallic bonding is characterized by the electrostatic attraction between positive metal cations and a sea of delocalised electrons. This type of bonding occurs in pure metals and alloys, such as copper ($Cu$) and iron ($Fe$), and accounts for properties like high electrical and thermal conductivity, malleability, and ductility.

Chemical structures are broadly categorized into simple molecular and giant molecular structures. Simple molecular structures consist of small molecules held together by strong internal covalent bonds, but weak intermolecular forces, known as van der Waals forces, exist between the molecules. These substances typically have low melting and boiling points and do not conduct electricity; examples include $CO_2$, $H_2O$, and methane ($CH_4$). Giant molecular or macromolecular structures consist of continuous three-dimensional networks of covalently bonded atoms and possess extremely high melting points. In diamond, each carbon atom is bonded tetrahedrally to $4$ other carbon atoms, making it extremely hard and an electrical insulator. In graphite, carbon atoms are bonded in hexagonal layers with free delocalised electrons, making it a soft, slippery lubricant that conducts electricity.

Polyatomic Radicals and Valencies

A radical is a group of covalently bonded atoms that carries an overall electric charge and functions as a single unit during chemical reactions. The Ammonium radical has the formula $NH_4^+$ with a charge of $+1$ and a valency of $1$, found in compounds like ammonium chloride ($NH_4Cl$). The Nitrate radical is $NO_3^-$ with a charge of $-1$ and a valency of $1$, found in nitric acid and sodium nitrate ($NaNO_3$). The Carbonate radical is $CO_3^{2-}$ with a charge of $-2$ and a valency of $2$, occurring in calcium carbonate ($CaCO_3$). The Sulphate radical is $SO_4^{2-}$ with a charge of $-2$ and a valency of $2$, as seen in copper(II) sulphate ($CuSO_4$). The Phosphate radical is $PO_4^{3-}$ with a charge of $-3$ and a valency of $3$, found in calcium phosphate ($Ca_3(PO_4)_2$).

Chemical Reactions and Equations

A chemical reaction is a process where starting substances, known as reactants, undergo chemical change to form brand-new substances called products. There are five core types of reactions. Direct combination, or synthesis, occurs when two or more substances combine to form a single product, such as $2Mg\text{(s)} + O_2\text{(g)} \rightarrow 2MgO\text{(s)}$. Decomposition involves a single compound breaking down into simpler substances via heat or light, such as $CaCO_3\text{(s)} \rightarrow CaO\text{(s)} + CO_2\text{(g)}$. Single displacement occurs when a more reactive element replaces a less reactive element in a compound, exemplified by $Zn\text{(s)} + CuSO_4\text{(aq)} \rightarrow ZnSO_4\text{(aq)} + Cu\text{(s)}$. Double displacement involves the exchange of positive ions between two reacting ionic compounds, such as $AgNO_3\text{(aq)} + NaCl\text{(aq)} \rightarrow AgCl\text{(s)} + NaNO_3\text{(aq)}$. Finally, redox or chain reactions are continuous self-sustaining reactions driven by free radicals or energy cycles.

All chemical equations must include state symbols in lower case: $\text{(s)}$ for solid, $\text{(l)}$ for liquid, $\text{(g)}$ for gas, and $\text{(aq)}$ for aqueous solutions dissolved in water. Balanced equations must strictly obey the Law of Conservation of Mass.

Chemical Kinetics and Reaction Rates

Chemical kinetics is the study of the speed or rate at which a chemical reaction occurs. The rate of reaction is defined as the change in the concentration of reactants or products per unit of time. The formula used is $\text{Rate} = \frac{\text{Amount of Reactant Used or Product Formed}}{\text{Time Taken}}$. According to Collision Theory, a reaction only occurs if particles collide with sufficient energy, which must be greater than or equal to the Activation Energy ($E_a$), and in the correct orientation.

Several factors affect reaction rates. Increasing the temperature increases the rate because particles gain kinetic energy, move faster, and collide more frequently with energy $\text{≥} E_a$. Increasing the concentration or pressure increases the rate by providing more particles per unit volume, leading to a higher frequency of successful collisions. Increasing the surface area of a solid by powdering it exposes more surface particles to collisions, thereby increasing the rate. A catalyst increases the rate by providing an alternative reaction pathway with a lower activation energy ($E_a$) without being consumed itself. Finally, light intensity increases specific photolysis rates as photons provide the energy required to initiate photochemical reactions, such as photosynthesis or the decay of silver bromide ($AgBr$).

Redox Reactions and Laboratory Tests

A redox reaction involves oxidation and reduction occurring simultaneously. Oxidation is defined as the gain of oxygen, the loss of hydrogen, the loss of electrons (described by the acronym OIL: Oxidation Is Loss), or an increase in oxidation state. Reduction is defined as the loss of oxygen, the gain of hydrogen, the gain of electrons (described by the acronym RIG: Reduction Is Gain), or a decrease in oxidation state.

Laboratory tests identify these agents through specific color changes. Oxidising agents, such as acidified potassium manganate(VII) ($KMnO_4$) or acidified potassium dichromate(VI) ($K_2Cr_2O_7$), can turn potassium iodide ($KI$) starch paper dark blue or black. Specifically, acidified $KMnO_4$ changes from purple to colourless when it is reduced. Reducing agents, such as sulphur dioxide ($SO_2$) or hydrogen sulphide ($H_2S$), turn acidified potassium dichromate ($K_2Cr_2O_7$) from orange to green.

Acids, Bases, and Salts

An acid is a chemical substance that produces hydrogen ions ($H^+$) as the only positive ions when dissolved in water. The basicity of an acid refers to the number of ionisable $H^+$ ions per molecule. Monobasic acids like hydrochloric acid ($HCl$) produce $1$ $H^+$ ion ($HCl \rightarrow H^+ + Cl^-$). Dibasic acids like sulphuric acid ($H_2SO_4$) produce $2$ $H^+$ ions ($H_2SO_4 \rightarrow 2H^+ + SO_4^{2-}$). Tribasic acids like phosphoric acid ($H_3PO_4$) produce $3$ $H^+$ ions ($H_3PO_4 \rightarrow 3H^+ + PO_4^{3-}$). Acids react with reactive metals to produce a salt and hydrogen gas ($H_2$), and they react with bases or alkalis in a neutralisation reaction to produce a salt and water. They also react with carbonates to produce a salt, water, and carbon dioxide ($CO_2$).

A base is a metal oxide or hydroxide that reacts with an acid to produce only a salt and water. A soluble base is termed an alkali and produces hydroxide ions ($OH^-$) in water. Oxides are classified into four groups. Acidic oxides are non-metal oxides like $CO_2$, $SO_2$, and $NO_2$ that react with water to form acids or alkalis to form salts. Basic oxides are metal oxides like $CaO$, $CuO$, and $MgO$ that react with acids to form salt and water. Amphoteric oxides, such as aluminium oxide ($Al_2O_3$), zinc oxide ($ZnO$), and lead(II) oxide ($PbO$), show both acidic and basic properties by reacting with both acids and bases. Neutral oxides are non-metal oxides like water ($H_2O$), carbon monoxide ($CO$), and dinitrogen oxide ($N_2O$) that show neither acidic nor basic properties.

Indicators, pH Scale, and Salt Preparation

Indicators change color based on the pH level. Red litmus paper turns blue in an alkali ($\text{pH} > 7$) and remains red or turns purple in neutral ($\text{pH} = 7$). Methyl orange is red or pink in acid ($\text{pH} < 7$), orange in neutral, and yellow in alkali. Phenolphthalein is colourless in acid and neutral solutions but turns pink or magenta in alkalis. Universal indicator changes through red, orange, or yellow in acids, is green in neutral, and becomes blue or violet in alkalis.

A salt is formed when the replaceable hydrogen ions of an acid are partially or completely replaced by a metal or ammonium ion. According to solubility rules, all nitrates ($NO_3^-$), sodium ($Na^+$), potassium ($K^+$), and ammonium ($NH_4^+$) salts are soluble. All chlorides are soluble except silver chloride ($AgCl$) and lead(II) chloride ($PbCl_2$). All sulphates are soluble except barium sulphate ($BaSO_4$), lead(II) sulphate ($PbSO_4$), and calcium sulphate ($CaSO_4$). All carbonates are insoluble except those of Group 1 metals and ammonium.

Three methods are used for salt preparation. Titration is used for soluble salts from alkalis (Acid + Alkali $\rightarrow$ Salt + Water), followed by evaporation to crystallise. The excess insoluble base or metal method involves adding an excess of the metal or carbonate to an acid, filtering the unreacted residue, and heating the filtrate to crystallise. Precipitation is used for insoluble salts by mixing two soluble salt solutions to produce an insoluble salt precipitate and a soluble solution; the precipitate is then filtered, washed with distilled water, and dried.

Qualitative Analysis of Ions and Gases

Qualitative analysis identifies unknown chemical species using reagent testing. When testing cations with sodium hydroxide ($NaOH$) and aqueous ammonia ($NH_3$), distinct results appear. Aluminium ($Al^{3+}$) forms a white precipitate with $NaOH$ that is soluble in excess to form a colourless solution, while it forms an insoluble white precipitate with $NH_3$. Calcium ($Ca^{2+}$) forms a white precipitate with $NaOH$ that is insoluble in excess, but yields no precipitate or only a slight white one with $NH_3$. Copper(II) ($Cu^{2+}$) produces a light blue precipitate that is insoluble in $NaOH$ but dissolves in excess $NH_3$ to form a deep blue solution. Iron(II) ($Fe^{2+}$) forms a dirty green precipitate insoluble in both. Iron(III) ($Fe^{3+}$) forms a reddish-brown precipitate insoluble in both. Zinc ($Zn^{2+}$) forms a white precipitate soluble in excess of both reagents. Ammonium ($NH_4^+$) produces ammonia gas upon warming with $NaOH$, turning damp red litmus paper blue.

Anions are identified through specific tests. Carbonates ($CO_3^{2-}$) produce effervescence when dilute $HCl$ is added, releasing $CO_2$ that turns limewater milky. Chlorides ($Cl^-$) are tested by acidifying with dilute nitric acid and adding aqueous silver nitrate ($AgNO_3$), yielding a white precipitate. Sulphates ($SO_4^{2-}$) are tested by acidifying with dilute nitric acid and adding barium nitrate ($Ba(NO_3)_2$), also yielding a white precipitate. Nitrates ($NO_3^-$) are identified by adding $NaOH$ and aluminium foil and warming gently to evolve ammonia gas.

Gases have distinct identification profiles. Ammonia ($NH_3$) is colourless with a pungent smell and turns damp red litmus paper blue. Carbon dioxide ($CO_2$) is colourless and odourless and forms a white precipitate in limewater ($Ca(OH)_2$). Chlorine ($Cl_2$) is yellow-green with a choking smell and bleaches damp blue litmus paper white. Hydrogen ($H_2$) is colourless and odourless and puts out a lighted splint with a 'pop' sound. Oxygen ($O_2$) is colourless and odourless and relights a glowing splint. Sulphur dioxide ($SO_2$) is colourless with a suffocating smell and turns acidified potassium dichromate(VI) from orange to green. When conducting exams, it is vital to state the reagent added, the observation (precipitate or color change), and the identity of the product.