Comprehensive CAIE IGCSE Chemistry Study Guide: States, Atoms, Bonding, and Reactions

Distinguishing Properties and Structures of Matter

The three states of matter are solid, liquid, and gas. In a solid, particles are closely packed and held in a fixed, regular arrangement; they possess the least amount of energy and vibrate in fixed positions. Solids cannot be compressed. In a liquid, particles are close together but have a random arrangement; they are able to move past one another and take the shape of their container. Like solids, liquids cannot be compressed. In a gas, particles are spread apart and arranged randomly. Gas particles possess the most energy, move randomly in all directions, and the substance can be compressed.

Changes of State

Changes of state are physical changes involving forces between particles.

• Melting: Transition from solid to liquid. This occurs at the melting point.
• Boiling: Transition from liquid to gas where bubbles form throughout the liquid, rising to the surface to evaporate. This occurs at the boiling point.
• Evaporation: Transition from liquid to gas which occurs only at the surface of the liquid.
• Freezing: Transition from liquid to solid. This occurs at the melting point.
• Condensation: Transition from gas to liquid. This occurs at the boiling point.
• Sublimation: Transition from solid to gas.

Gas volume is affected by temperature and pressure. As temperature increases, the volume of a gas increases. As pressure increases, the volume of a gas decreases. Conversely, as pressure decreases, the volume of a gas increases.

Kinetic Particle Theory and Curves

Kinetic particle theory models matter as small solid spheres to explain state changes. For melting, boiling, and evaporation, thermal energy is transferred into kinetic energy, allowing particles to overcome intermolecular forces. For freezing and condensing, energy is lost, and intermolecular forces form to hold particles closer. The state changes from $(s)$ to $(l)$ to $(g)$ as kinetic energy increases.

• Heating Curve: A graph of temperature against time. As temperature rises, particles gain energy. The line plateaus at the melting point while the solid turns to liquid. Once melted, temperature rises again until the boiling point, where the line plateaus again as the liquid turns to gas.
• Cooling Curve: A graph showing temperature decreasing as energy is lost. The line plateaus during condensation as bonds form between gas particles. After turning to liquid, the temperature drops until the freezing point plateau, where the liquid becomes solid.

In gases, increasing temperature increases kinetic energy, causing particles to move faster and collide with the container more frequently, spreading them apart. Increasing pressure forces gas particles closer together, decreasing volume.

Diffusion

Diffusion is the net movement of particles from an area of high concentration to an area of low concentration. It requires particles to be able to move; therefore, it occurs in liquids and gases but not in solids. The rate of diffusion is affected by relative molecular mass ($M_r$). The larger the relative molecular mass, the "heavier" the particles and the slower they move, resulting in a slower diffusion rate. Smaller relative molecular masses result in "lighter," faster-moving particles and a higher rate of diffusion.

Elements, Compounds, and Mixtures

• Element: A pure substance made from only one type of atom that cannot be broken down further (e.g., $O_2$, $H_2$, $Cu$).
• Compound: A substance made from two or more elements chemically bonded together. A chemical reaction is required to break these bonds (e.g., $H_2O$, $CO_2$, $NaCl$).
• Mixture: A substance made of two or more elements or compounds not chemically bonded. They can be separated by physical techniques like filtration, simple distillation, and crystallisation. Chemical properties of components remain unchanged (e.g., air, saltwater).

Atomic Structure and the Periodic Table

An atom consists of a central, positively charged nucleus containing protons and neutrons, surrounded by negatively charged electrons in shells.

Subatomic Particle Properties:

• Proton: Relative charge = $+1$, Relative mass = $1$

• Neutron: Relative charge = $0$, Relative mass = $1$

• Electron: Relative charge = $-1$, Relative mass = $ext{~}\frac{1}{2000}$

• Atomic (Proton) Number: The number of protons in the nucleus.

• Mass (Nucleon) Number: The total number of protons and neutrons in the nucleus.

Electronic Configuration: Electrons fill the innermost shells first. For proton numbers $1$ to $20$, the shells hold a maximum of:

• Shell 1: $2$ electrons
• Shell 2: $8$ electrons
• Shell 3: $8$ electrons
  For example, Chlorine (proton number $17$) is written as $2, 8, 7$. Sodium (proton number $11$) is $2, 8, 1$, but a Sodium ion ($Na^+$) loses an electron to become $2, 8$.

Periodic Table Organization:

• Periods: Horizontal rows. The number of occupied shells equals the period number.
• Groups: Vertical columns. For Groups I to VII, the number of outer shell electrons equals the group number.
• Group VIII: Noble gases. These have full outer shells (ending in $2$ for Helium or $8$ for others).

Isotopes

Isotopes are atoms of the same element with the same number of protons and electrons but different numbers of neutrons (e.g., Carbon-12 vs. Carbon-13). They have the same chemical properties because they share the same electronic configuration.

Relative Atomic Mass ($A_r$) Calculation:$A_r = \frac{(isotope_1 ext{ mass} imes abundance) + (isotope_2 ext{ mass} imes abundance)}{100}$ Example: Carbon-14 (20 ext{%}) and Carbon-12 (80 ext{%}):
$\frac{(14 imes 20) + (12 imes 80)}{100} = 12.4$

Ions and Ionic Bonding

An ion is an atom/molecule with a charge due to the loss or gain of electrons.

• Anions: Negative ions formed when non-metals gain electrons (e.g., Cl + e^- ightarrow Cl^-).
• Cations: Positive ions formed when metals lose electrons (e.g., Na ightarrow Na^+ + e^-).

An ionic bond is a strong electrostatic attraction between oppositely charged ions. This involves the transfer of electrons. In dot-and-cross diagrams, square brackets and charges indicate the ion.

Ionic Compound Properties:

• Structure: Giant lattice of alternating positive and negative ions.
• Melting/Boiling Points: High, due to strong electrostatic forces.
• Electrical Conductivity: Good when aqueous or molten (ions free to move), but poor when solid (ions fixed).

Covalent Bonding

A covalent bond forms when a pair of electrons is shared between two non-metal atoms to achieve noble gas configurations.

• Simple Molecules: Examples include $H_2$, $Cl_2$, $H_2O$, $CH_4$, $NH_3$, and $HCl$.
• Multiple Bonds: $CO_2$ has double bonds ($2$ pairs shared); $N_2$ has triple bonds ($3$ pairs shared).

Simple Molecular Compound Properties:

• Melting/Boiling Points: Low, due to very weak intermolecular forces (despite strong intramolecular covalent bonds).
• Electrical Conductivity: Poor, as there are no ions or charged particles to carry charge.

Giant Covalent Structures

• Diamond: Each carbon is covalently bonded to $4$ others in a tetrahedral 3D shape. It is very hard and does not conduct electricity (no delocalised electrons). Used in cutting tools.
• Graphite: Each carbon is bonded to $3$ others in hexagonal layers. Layers are held by weak intermolecular forces and can slide, making it soft (lubricant). One delocalised electron per atom allows it to conduct electricity (electrodes).
• Silicon(IV) Oxide ($SiO_2$): Main component of sand. Each Silicon is bonded to $4$ oxygens, and each oxygen to $2$ silicons. It has a tetrahedral arrangement similar to diamond, resulting in high melting points and extreme hardness.

Metallic Bonding

Metallic bonding is the electrostatic attraction between positive metal ions in a giant lattice and a "sea" of delocalised electrons.

• Conductivity: High, as electrons move freely through the structure.
• Malleability/Ductility: Layers of atoms can slide over each other. Alloys are harder because different-sized atoms distort layers, preventing sliding.
• Melting/Boiling Points: High, due to strong attraction between ions and delocalised electrons.

Chemical Formulae and Equations

• Molecular Formula: Actual number of atoms of each element in a molecule (e.g., $H_2O$).
• Empirical Formula: Simplest whole-number ratio of atoms in a compound (e.g., $C_2H_6$ has an empirical formula of $CH_3$).
• State Symbols: $(s)$ solid, $(l)$ liquid, $(g)$ gas, $(aq)$ aqueous.
• Ionic Equations: Equations where the number of atoms and the total charge are balanced on both sides.

Stoichiometry and the Mole

The Relative Atomic Mass ($A_r$) is the average mass of isotopes compared to $\frac{1}{12}$th the mass of $^{12}C$. The Relative Molecular/Formula Mass ($M_r$) is the sum of $A_r$ values.

The Mole Concept: One mole ($1 ext{ mol}$) contains the Avogadro constant: $6.02 imes 10^{23}$ particles. $ext{Mass (g)} = ext{Moles (mol)} imes M_r (g/mol)$$ext{Number of particles} = ext{Avogadro constant} imes ext{Amount of substance (mol)}$

Gases: The volume of $1 ext{ mol}$ of any gas at Room Temperature and Pressure (RTP) is $24 ext{ dm}^3$ (or $24000 ext{ cm}^3$). $ext{Volume (dm}^3) = ext{Moles} imes 24$

Concentration:$ext{Concentration (mol/dm}^3) = \frac{ ext{Moles}}{ ext{Volume (dm}^3)}$ To convert $mol/dm^3$ to $g/dm^3$, multiply by $M_r$. To convert $cm^3$ to $dm^3$, divide by $1000$.

Reacting Masses and Yields:

• Limiting Reactant: The reactant used up first, determining the product amount.
• Percentage Yield = $\frac{ ext{Actual amount}}{ ext{Theoretical amount}} imes 100$
• Percentage Composition by Mass = $\frac{ ext{Total } A_r ext{ of element}}{M_r ext{ of compound}} imes 100$
• Percentage Purity = $\frac{ ext{Mass of pure substance}}{ ext{Total mass of sample}} imes 100$

Electrolysis

Electrolysis is the decomposition of an ionic compound (molten or aqueous) by an electric current.

• Anode: Positive electrode; attracts anions (oxidation occurs).
• Cathode: Negative electrode; attracts cations (reduction occurs).
• Electrolyte: The compound undergoing electrolysis containing mobile ions.

Products of Electrolysis:

• Molten Salts: Metal forms at cathode; non-metal at anode.
• Aqueous Solutions:   - Cations compete with $H^+$: If the metal is more reactive than hydrogen, $H_2$ gas forms. If less reactive, the metal forms.   - Anions compete with $OH^-$: If halide ions ($Cl^-$, $Br^-$, $I^-$) are present, the halogen forms. Otherwise, oxygen and water form.

Specific Processes:

• Brine (Concentrated Aqueous $NaCl$): Cathode = $H_2$, Anode = $Cl_2$, Solution left = $NaOH$.
• Dilute Sulfuric Acid: Cathode = $H_2$, Anode = $O_2$.
• Aqueous Copper(II) Sulfate: With inert electrodes, Cathode = $Cu$ (brown deposit), Anode = $O_2$. With copper electrodes, the anode dissolves (Cu ightarrow Cu^{2+} + 2e^-) and the cathode builds copper plating.

Electroplating: Purpose is appearance and corrosion resistance. The cathode is the object to be plated; the anode is the plating metal; the electrolyte contains ions of the plating metal.

Hydrogen-Oxygen Fuel Cells

These produce electricity with water as the only product. Overall Reaction: 2H_2 (g) + O_2 (g) ightarrow 2H_2O (l). Advantages: Less $CO_2$ emissions (if hydrogen is renewable), reduces fossil fuel use. Disadvantages: Expensive production/transport, difficult/dangerous storage, less durable than petrol engines.

Energy Changes in Reactions

• Exothermic: Transfers thermal energy to surroundings (temp increases). $ext{ΔH}$ is negative. Examples: Combustion, neutralisation.
• Endothermic: Takes in thermal energy (temp decreases). $ext{ΔH}$ is positive. Examples: Thermal decomposition.
• Activation Energy ($E_a$): Minimum energy colliding particles need to react.

Bond Energies: Bond breaking is endothermic; bond making is exothermic. $ext{ΔH} = ext{Energy In (Breaking)} - ext{Energy Out (Making)}$ Example: If Reactants = $3064 ext{ kJ/mol}$ and Products = $3166 ext{ kJ/mol}$, then $ext{ΔH} = -102 ext{ kJ/mol}$ (Exothermic).

Physical and Chemical Changes

• Physical Changes: Reversible, no new substances made, no significant energy change (e.g., melting).
• Chemical Changes: Usually irreversible, new substances formed, large energy changes, possible colour changes (e.g., burning).

Rate of Reaction

Factors Increasing Rate:

1. Concentration: More particles in the same volume = more frequent collisions.
2. Pressure (Gases): Increases particle density = more frequent collisions.
3. Surface Area (Solids): Fine powder has more exposed particles than a lump.
4. Temperature: Increases kinetic energy; particles move faster and collisions are more frequent and energetic.
5. Catalyst: Lowers activation energy ($E_a$) by providing an alternative pathway. Catalysts are not used up.

Collision Theory: Reaction occurs if particles collide with sufficient energy ($E_a$). Tangents on a graph ($\frac{ ext{Δy}}{ ext{Δx}}$) are used to calculate specific rates at certain times.

Reversible Reactions and Equilibrium

Reversible reactions ( ightleftharpoons) can reach a dynamic equilibrium in a closed system when the rate of the forward reaction equals the rate of the reverse reaction and concentrations remain constant.

Le Chatelier’s Principle:

• Temperature Increase: Shifts toward the endothermic reaction.
• Pressure Increase: Shifts toward the side with fewer gas molecules.
• Concentration Increase: Shifts to consume the added substance.

Industrial Processes:

• Haber Process (N_2 + 3H_2 ightleftharpoons 2NH_3): Conditions = $450^ ext{o}C$, $20,000 ext{ kPa (200 atm)}$, Iron catalyst. Forward is exothermic.
• Contact Process (2SO_2 + O_2 ightleftharpoons 2SO_3): Conditions = $450^ ext{o}C$, $200 ext{ kPa (2 atm)}$, Vanadium(V) oxide catalyst.

Redox Reactions

• Oxidation: Gain of oxygen, loss of electrons (OIL), increase in oxidation number.
• Reduction: Loss of oxygen, gain of electrons (RIG), decrease in oxidation number.
• Oxidising Agent: Gains electrons (gets reduced).
• Reducing Agent: Loses electrons (gets oxidised).

Tests:

• Acidified Aqueous Potassium Manganate(VII): Purple ($Mn^{7+}$) $ ightarrow$ Colourless ($Mn^{2+}$) in the presence of a reducing agent.
• Aqueous Potassium Iodide: Colourless ($I^-$) $ ightarrow$ Brown ($I_2$) in the presence of an oxidising agent.

Acids, Bases, and Salts

• Acids: Proton donors ($H^+$). pH < 7. Turn litmus red, methyl orange red, thymolphthalein colourless.
• Bases: Proton acceptors. Typically metal oxides or hydroxides. Alkalis are soluble bases ($OH^-$).
• Strong vs. Weak: Strong acids (e.g., $HCl$) fully dissociate; weak acids (e.g., ethanoic acid) partially dissociate ( ightleftharpoons).

Oxides:

• Basic Oxides: Formed by metals (e.g., $CuO, CaO$).
• Acidic Oxides: Formed by non-metals (e.g., $CO_2, SO_2$).
• Amphoteric Oxides: React with both acids and bases (e.g., $Al_2O_3, ZnO$).

Salts Preparation:

• Soluble Salts: React acid with excess metal, base, or carbonate. Or use titration for alkalis. Crystallise by evaporating water.
• Insoluble Salts: Prepared by precipitation (mixing two soluble salts).
• Solubility Rules: Nitrates and Sodium/Potassium/Ammonium salts are always soluble. Chlorides are soluble except Silver and Lead. Sulfates are soluble except Barium, Calcium, and Lead. Carbonates and Hydroxides are generally insoluble.