CIE IGCSE Co-ordinated Sciences Chemistry Master Notes

Kinetic Theory and States of Matter

• The kinetic theory describes matter as being composed of tiny particles in constant motion.
• This model explains the differing physical properties of solids, liquids, and gases based on particle arrangement and movement.

Solids

• Properties: Fixed volume, fixed shape, and high density.
• Arrangement: Particles are packed very closely together in a fixed, regular pattern.
• Movement: Particles vibrate in place about fixed positions but cannot change their location.
• Energy: Particles possess low energy.

Liquids

• Properties: Fixed volume but take the shape of their container; generally less dense than solids (water is an exception) but denser than gases; can flow freely.
• Arrangement: Particles are close together but arranged randomly.
• Movement: Particles move and slide past each other.
• Energy: Particles possess greater energy than those in solids.

Gases

• Properties: No fixed volume; take the shape of their container; very low density; easily compressed due to significant space between particles.
• Arrangement: Particles are far apart and arranged randomly.
• Movement: Particles move randomly and quickly in all directions, reaching speeds around $500\,m/s$. Pressure is created by particles colliding with each other and the walls of the container.
• Energy: Particles possess the highest energy of the three states.

State Changes and Phase Transitions

• State changes involve a transition in the energy, arrangement, and movement of particles.
• Heating Curve: When a substance is heated, thermal energy is converted to kinetic energy. On a heating curve, horizontal sections indicate a change of state where the temperature remains constant.
• Cooling Curve: The mirror image of a heating curve, representing the loss of energy as a gas turns back into a solid.

Definitions of State Changes

• Melting: The transition from solid to liquid. It occurs at a specific temperature called the melting point ($m.p.$). Heat energy is transformed into kinetic energy, allowing particles to move.
• Freezing: The transition from liquid to solid. It is the reverse of melting and occurs at the same temperature. Pure water melts and freezes at $0\,^\circ C$.
• Boiling: The transition from liquid to gas. Requires heat to form bubbles of gas below the surface, allowing particles to escape from both the surface and within the liquid. It occurs at a specific temperature called the boiling point ($b.p.$).
• Evaporation: The transition from liquid to gas occurring over a range of temperatures below the boiling point. It only takes place at the surface. Rate is increased by larger surface area and warmer surface temperatures.
• Condensation: The transition from gas to liquid upon cooling. Gas particles lose energy; when they collide, they lack the energy to bounce away and instead group together.
• Sublimation/Desublimation: The direct transition between solid and gas without passing through the liquid phase.

Pressure and Temperature in Gases

• Temperature Effects: Increasing the temperature increases the kinetic energy of gas particles. Particles move faster and spread out more. In flexible containers (like hot air balloons), this causes expansion. In rigid containers, it increases collision frequency and pressure.
• Pressure-Volume Relationship: Increasing gas pressure decreases its volume (compression). For example, a bicycle pump uses high pressure to inflate a tire. Decreasing pressure allows the gas to occupy a larger volume.

Diffusion

• Definition: The process where particles move from an area of high concentration to an area of low concentration due to random motion.
• Diffusion requires no energy input and eventually results in an even concentration throughout the space.
• Temperature Effect: Diffusion occurs faster at higher temperatures because particles possess more kinetic energy.

Diffusion in Liquids and Gases

• Liquids: Demonstrated by adding potassium manganate(VII) ($KMnO_4$) crystals to water. The purple color slowly spreads until the solution is uniform.
• Gases: Diffusion is faster in gases because particles have more energy and move quicker. An example is the spread of orange-brown bromine gas in air.

Diffusion and Molecular Mass

• At a constant temperature, gases with lower relative molecular mass ($M_r$) diffuse faster and further than heavier gases.
• Ammonia and Hydrogen Chloride Experiment:
  • Reaction: $NH_3 (g) + HCl (g) \rightarrow NH_4Cl (s)$
  • $M_r$ of Ammonia ($NH_3$) = $17$
  • $M_r$ of Hydrogen Chloride ($HCl$) = $36.5$
  • Because ammonia is lighter, it travels faster. A white "smoke" ring of solid ammonium chloride forms closer to the hydrogen chloride end of the tube.

Atomic Structure and the Periodic Table

Classification of Matter

• Elements: Substances made of atoms containing the same number of protons. They cannot be split by chemical means. There are approximately $118$ elements in the Periodic Table.
• Compounds: Pure substances made of two or more elements chemically combined (e.g., $CuSO_4$, $CaCO_3$, $CO_2$). They cannot be separated by physical means.
• Mixtures: Combinations of two or more substances not chemically combined (e.g., sand and water, oil and water). They can be separated by physical methods like filtration or evaporation.

Subatomic Particles

• Atoms consist of protons, neutrons, and electrons.
• Relative Mass and Charge Table:
  • Proton: Relative Mass = $1$, Charge = $1+$
  • Neutron: Relative Mass = $1$, Charge = $0\,(\text{neutral})$
  • Electron: Relative Mass = $\frac{1}{1840}\,(\text{negligible})$, Charge = $1-$

Atomic Definitions

• Atomic Number ($Z$): The number of protons in the nucleus. In a neutral atom, this also equals the number of electrons.
• Nucleon Number / Mass Number ($A$): The total number of protons and neutrons in the nucleus.
• Calculation: $\text{Number of Neutrons} = A - Z$.
• Symbolism: Written as ${^A_Z}X$.

Isotopes

• Different atoms of the same element with the same number of protons but different numbers of neutrons.
• Chemical Properties: Identical because they have the same number/arrangement of outer-shell electrons.
• Physical Properties: Differ (e.g., density, melting point) due to mass differences.

Electronic Configuration

• Electrons occupy shells (energy levels) around the nucleus. Shells fill from the inside out.
• Capacities:
  • First Shell: $2$ electrons
  • Second Shell: $8$ electrons
  • Third Shell: $8$ electrons (for atoms up to atomic number $20$)
• Periodic Table Correspondence:
  • Number of occupied shells = Period number.
  • Number of outer-shell (valence) electrons = Group number (for Groups I to VII).

Ions and Ionic Bonding

Ion Formation

• Ions are charged atoms formed by losing or gaining electrons to achieve a full outer shell (noble gas configuration).
• Cations: Positive ions formed by metals losing electrons.
• Anions: Negative ions formed by non-metals gaining electrons.
• Specific Group Charges:
  • Group 1: $1+$
  • Group 2: $2+$
  • Group 6: $2-$
  • Group 7: $1-$

Ionic Bonds and Structures

• Definition: The strong electrostatic force of attraction between oppositely charged ions.
• Lattice Structure: Ionic compounds form a giant lattice of alternating positive and negative ions in a regular, 3D repeating pattern.
• Properties:
  • High melting and boiling points due to strong electrostatic forces in all directions.
  • Conduct electricity when molten or in aqueous solution (ions are free to move).
  • Poor conductors when solid (ions are fixed in position).

Covalent Bonding and Simple Molecules

• Definition: Formed by the sharing of electron pairs between non-metal atoms.
• Molecules: Groups of two or more atoms covalently bonded together.
• Bond Types:
  • Single Bond: One pair of shared electrons (e.g., $H_2, Cl_2, H_2O, NH_3, CH_4$).
  • Double Bond: Two pairs shared electrons (e.g., $O_2, CO_2$).
  • Triple Bond: Three pairs shared electrons (e.g., $N_2$).
• Properties of Simple Molecular Compounds:
  • Low melting/boiling points due to weak intermolecular forces between molecules (while covalent bonds within molecules are strong).
  • Poor electrical conductivity (no free ions or electrons).

Giant Covalent Structures

Graphite

• Structure: Each carbon atom is bonded to three others in layers of hexagons. Contains delocalised electrons.
• Properties: Soft and slippery (layers slide due to weak intermolecular forces); conducts electricity (delocalised electrons); high melting point.
• Uses: Pencils, industrial lubricants.

Diamond

• Structure: Each carbon atom is bonded to four others in a tetrahedral arrangement.
• Properties: Extremely hard and dense; highest melting point; does not conduct electricity (no free electrons).
• Uses: Cutting tools (drills).

Metallic Bonding

• Structure: A giant lattice of positive metal ions surrounded by a "sea" of delocalised electrons.
• Definition: The strong force of attraction between positive metal ions and delocalised electrons.
• Properties:
  • High melting and boiling points.
  • Conduct heat and electricity (delocalised electrons carry charge/energy).
  • Malleable (can be hammered into shape) and ductile (can be drawn into wires) because layers of ions can slide over each other without breaking the metallic bond.

Chemical Formulae and Equations

Symbols and Diatomic Molecules

• Common diatomic elements: $H_2, N_2, O_2, F_2, Cl_2, Br_2, I_2$.
• Molecular Formula: Shows the type and exact number of atoms (e.g., $H_2SO_4$).

Writing Formulae for Ionic Compounds

• The compound must have no overall charge.
• Swap-and-Drop Method: Used to balance charges. For example, Copper(II) chloride consists of $Cu^{2+}$ and $Cl^{1-}$, resulting in $CuCl_2$.
• Polyatomic Ions: Carbonate ($CO_3^{2-}$), Sulfate ($SO_4^{2-}$), Hydroxide ($OH^-$), Nitrate ($NO_3^-$), Ammonium ($NH_4^+$).

Chemical Equations

• Word Equations: $\text{Reactants} \rightarrow \text{Products}$.
• Symbol Equations: Must be balanced to satisfy the Law of Conservation of Mass.
• State Symbols: Solid ($s$), Liquid ($l$), Gas ($g$), Aqueous solution ($aq$).

Relative Masses and The Mole

Relative Atomic Mass ($A_r$) and Molecular Mass ($M_r$)

• $A_r$ is measured relative to $\frac{1}{12}\text{th}$ the mass of a Carbon-12 atom.
• $M_r$ is the sum of $A_r$ for all atoms in a formula.
• Law of Conservation of Mass: Total mass of reactants equals total mass of products.

The Mole and Avogadro Constant

• The Mole ($mol$): The SI unit of amount of substance.
• Avogadro Constant: $6.02 \times 10^{23}$ particles per mole.
• Molar Mass: The mass in grams for one mole of a substance (equals $A_r$ or $M_r$ in grams).
• Formula: $\text{Moles} = \frac{\text{Mass (g)}}{\text{Molar Mass (g/mol)}}$.

Molar Volume of Gas

• Avogadro’s Law: Equal volumes of gases at the same temperature and pressure contain equal numbers of molecules.
• At RTP ($20\,^\circ C, 1\,atm$): One mole of any gas occupies $24\,dm^3$ ($24,000\,cm^3$).
• Formula: $\text{Volume (dm}^3) = \text{Moles} \times 24$.

Energetics: Exothermic and Endothermic Reactions

• System: The reacting chemicals.
• Surroundings: Everything else.
• Exothermic: Heat is transferred to the surroundings (surroundings get hotter). $\Delta H$ is negative. Examples: Combustion, Neutralization.
• Endothermic: Heat is absorbed from the surroundings (surroundings get colder). $\Delta H$ is positive. Examples: Thermal decomposition, Photosynthesis.
• Activation Energy ($E_a$): The minimum energy required for a successful collision to occur and start a reaction.
• Bond Energy:
  • Bond breaking is endothermic.
  • Bond forming is exothermic.
  • Reaction is exothermic if energy released in making bonds > energy absorbed in breaking bonds.

The Periodic Table Trends

• Periodic Arrangement: Elements ordered by increasing atomic number. Metals on the left, non-metals on the right, separated by metalloids.
• Group I (Alkali Metals): $Li, Na, K, Rb, Cs, Fr$. Soft, low density, low melting points. React with water to produce hydrogen and an alkaline hydroxide. Reactivity increases down the group ($Li$ is least, $Fr$ is most).
• Group VII (Halogens): $F, Cl, Br, I, At$. Diatomic non-metals. Reactivity decreases down the group. Colors darken (Chlorine: pale yellow-green gas, Bromine: red-brown liquid, Iodine: grey-black solid). Melting/boiling points increase down the group.
• Displacement Reactions: A more reactive halogen displaces a less reactive halide from its aqueous solution.
• Transition Elements: Metals in the center of the Periodic Table. Form colored compounds, have variable oxidation states (ions with different charges), and are used as catalysts (e.g., Iron in the Haber process).
• Noble Gases (Group VIII): Monoatomic, unreactive gases with full outer shells.

Properties and Extraction of Metals

Uses of Metals

• Aluminium: Low density, corrosion-resistant (due to oxide layer). Used in aircraft bodies and food cans.
• Copper: Excellent electrical conductor, ductile. Used in wiring and water pipes.
• Alloys: Mixtures of a metal with other elements (e.g., Brass = Copper + Zinc; Stainless Steel = Iron + Chromium + Nickel). Alloys are harder/stronger because atoms of different sizes disrupt the regular lattice, preventing layers from sliding.

Reactivity Series

• Mnemonic: "Please send cats, monkeys and cute zebras into hot countries signed Gordon" (Potassium, Sodium, Calcium, Magnesium, Aluminium, Carbon, Zinc, Iron, Hydrogen, Copper, Silver, Gold).
• Extraction method depends on position relative to Carbon:
  • Above Carbon: Electrolysis (e.g., Aluminium from Bauxite in molten cryolite).
  • Below Carbon: Reduction with Carbon/CO (e.g., Iron from Hematite in the Blast Furnace).

Extraction of Iron (Blast Furnace Reactions)

1. Zone 1: $C (s) + O_2 (g) \rightarrow CO_2 (g)$ (Provides heat).
2. Zone 2: $CO_2 (g) + C (s) \rightarrow 2CO (g)$ (Forms reducing agent).
3. Zone 3: $Fe_2O_3 (s) + 3CO (g) \rightarrow 2Fe (l) + 3CO_2 (g)$ (Reduces iron ore).
4. Impurities: $CaCO_3 \rightarrow CaO + CO_2$; Then $CaO (s) + SiO_2 (s) \rightarrow CaSiO_3 (l)\,(\text{Slag})$.

Corrosion and Rusting

• Rusting: Specifically the oxidation of iron in the presence of water and oxygen.
• Prevention:
  • Barriers: Paint, oil, grease, plastic.
  • Sacrificial Protection: Attaching a more reactive metal (e.g., Zinc on steel ships).
  • Galvanising: Coating iron with zinc; provides both barrier and sacrificial protection.

Water and Air Quality

Water

• Chemical Tests:
  • Anhydrous Cobalt(II) chloride paper: Blue to pink.
  • Anhydrous Copper(II) sulfate: White to blue.
• Purification: Sedimentation, filtration (sand/gravel/carbon), and chlorination (to kill microbes).
• Distilled Water: Used in labs as tap water contains ions that interfere with reactions.

Air

• Composition: $78\% \, N_2$, $21\% \, O_2$, $0.9\% \, Ar$, $0.04\% \, CO_2$.
• Pollutants:
  • Carbon Monoxide ($CO$): From incomplete combustion; toxic, binds to haemoglobin.
  • Sulfur Dioxide ($SO_2$) and Oxides of Nitrogen: Cause acid rain.
  • Methane ($CH_4$) and $CO_2$: Greenhouse gases causing global warming.
• Catalytic Converters: Employ transition metals (Pt, Rh) to convert pollutants via redox: $2NO + 2CO \rightarrow N_2 + 2CO_2$.