Molecular mass is the sum of the masses of all the elements present in a molecule.

• Gas particles are in constant random motion
• Gas particles constantly collide with each other and with the walls of the container
• Energy is not lost during collisions
• The total volume of gas particles is much smaller than the volume of the container
• There are no attractive or repulsive forces between gas particles
• Particles move faster as temperature increases
• Kinetic energy: KE = ½mv²
• The average speed of particles depends on molecular mass

Matter can be classified into pure substances and mixtures.

Pure Substance — Has a fixed and unchanging chemical composition.

• Element — Cannot be separated into simpler substances
• Compound — Made of two or more elements in fixed proportions and can only be separated by chemical methods

Mixture — Consists of two or more pure substances that retain their individual properties and can be separated by physical methods.

• Homogeneous (Solution) — Uniform composition and properties throughout
• Heterogeneous — Non-uniform composition and properties throughout

• Do not scatter light and therefore do not show the Tyndall effect
• Solute particles never settle
• Cannot be separated using filtration

Alloys are manmade substances formed by combining multiple metals or a metal with a non-metal to increase strength or resistance to corrosion.

A suspension is a mixture in which solid particles do not dissolve and remain suspended in a liquid or gas.

Colloids are mixtures where one substance is dispersed into another as very small particles. The particle size ranges from 1 nm to 1000 nm.

An emulsion is a type of colloid where tiny droplets of one liquid are dispersed in another liquid. These liquids normally do not mix unless an emulsifier is used.

Example: Mayonnaise — oil mixed in water

A sol is a colloid where a solid is dispersed in a liquid. It appears smooth but contains tiny solid particles.

Example: Blood — solid cells dispersed in plasma

A gel is a mixture where liquid is trapped inside a solid, giving it a soft, jelly-like structure.

Examples: Jelly, cheese

Emulsifiers are FDA-approved food additives that allow substances which normally do not mix to remain combined.

Example: Oil and water in mayonnaise stay mixed due to an emulsifier

The Tyndall effect describes the scattering of light by particles in a mixture.

• True solutions do not show the Tyndall effect because particles are very small
• Colloids show the Tyndall effect because particle size is comparable to or larger than photon size

Example: When light from a flashlight is passed through milk, it appears scattered-- whilst the same is not true for a saltwater solution

Paper chromatography is a simple technique used to separate mixtures of substances (usually coloured) using chromatography paper and a solvent.

Separation occurs because dissolved substances move at different speeds across the paper.

Separation happens because different substances travel at different rates on the paper depending on their attraction to the paper (stationary phase) and the solvent (mobile phase).

Chromatography paper is porous, allowing the solvent to move upward by capillary action.

1. Sample preparation

The mixture is dissolved in a suitable solvent so it can travel with the mobile phase.

2. Placing the sample (spotting)

A tiny drop of the sample is placed on the paper using a capillary tube or micropipette. Correct placement helps produce clear and readable results.

3. Developing the chromatogram

The bottom of the paper is dipped into the solvent. The solvent rises up the paper by capillary action, carrying the sample and separating it into spots.

4. Drying and detecting

Once the solvent has travelled far enough, the paper is removed and dried. A detecting solution may be sprayed to make the spots visible. The visible spots form the chromatogram.

• Amino acids and organic acids
• Alkaloids
• Polysaccharides
• Proteins and peptides
• Natural and artificial pigments
• Inorganic cations
• Plant extracts
• Forensics
• Pharmaceuticals
• Foods
• Isolation and purification

Simple distillation is used to separate miscible liquids with significantly different boiling points.

Example: Obtaining pure water from salt water.

The liquid with the lower boiling point evaporates first, then cools and condenses to form a pure liquid.

Fractional distillation is used when the boiling points of liquids are close together.

Example: Separating crude oil into petrol.

Vapours rise through a fractionating column. Higher boiling liquids condense first, while lower boiling liquids rise further up and are collected separately.

Centrifugation separates particles from a liquid by spinning the mixture at high speed.

Sedimentation is the process of separating insoluble solids from a liquid. When left undisturbed, heavier particles settle at the bottom due to gravity.

Decantation involves carefully pouring off the clear liquid after sedimentation has occurred.

A separating funnel is used to separate two immiscible liquids, such as oil and water.

Due to different densities, the heavier liquid settles at the bottom while the lighter liquid stays on top. A tap allows the lower liquid to be released first.

Isotopes are atoms of the same element that have the same number of protons but a different number of neutrons.

• Different numbers of neutrons result in different atomic masses
• Isotopes have nearly identical chemical properties because they share the same electron configuration

Atomic weight represents the weighted average mass of all naturally occurring isotopes of a given element.

It is calculated by combining the mass of individual isotopes with their relative abundance in nature.

Elements behave consistently in chemical reactions, making a weighted average a practical way to represent atomic mass.

Atomic weight accurately reflects the natural distribution of isotopes and provides a realistic representation of an element’s mass in nature.

Percentage abundance describes the proportion of each isotope present in a natural sample of an element. These values are essential for accurately calculating the average atomic mass.

Atomic Weight = (Mass₁ × Abundance₁) + (Mass₂ × Abundance₂) + …

Where abundance is written as a decimal:
Decimal abundance = Percentage abundance ÷ 100

Atomic radius refers to the size of an atom, measured as the distance from the nucleus to the outermost electron shell.

• Increases down a group due to the addition of electron shells
• Decreases across a period because increased protons and electrons increase nuclear attraction

Ionization energy is the amount of energy required to remove an electron from an atom.

• Smaller atoms have higher ionization energy because electrons are closer to the nucleus
• Decreases down a group due to weaker nuclear pull on outer electrons
• Increases across a period as nuclear charge increases

Electronegativity is a measure of how strongly an atom attracts electrons.

• Measured on a scale of 0–4
• Fluorine has the highest electronegativity at 3.98
• Decreases down a group due to increased atomic size
• Increases across a period

Shielding occurs when inner electron shells reduce the attraction between the nucleus and the outermost electrons.

Word equations describe chemical reactions using words.

Example: Oxygen + Hydrogen → Water

Chemical equations use symbols and formulae to represent reactions.

Example: 2H₂ + O₂ → 2H₂O

• Equations must be balanced according to the law of conservation of mass
• Balancing is done by adding coefficients in front of formulae

Metals

• Reactivity increases down a group and across a period
• Metals react by losing electrons

Non-metals

• Reactivity decreases down a group and across a period
• Non-metals react by gaining electrons

• Good conductors of electricity
• High melting and boiling points
• Lustrous (shiny)
• Have a lattice structure

• Soft, shiny, and highly reactive
• Low melting points, boiling points, and densities

• Reactive, but less than alkali metals
• Form ionic compounds

• Form salts with metals
• Highly electronegative
• Poor conductors of heat and electricity

• Monatomic
• Colorless, odorless, tasteless
• Low density and solubility
• Gases at room temperature

Ionic bonding occurs between a metal and a non-metal.

• The metal loses an electron and becomes a cation
• The non-metal gains an electron and becomes an anion
• The bond is held by strong electrostatic attraction

Example: NaCl

• Large electronegativity difference (greater than 1.7)
• Hard, brittle, crystalline solids
• High melting and boiling points
• Do not conduct electricity in solid state due to rigid lattice
• Often soluble in polar solvents
• Have a crystal lattice structure

Covalent bonding occurs between two non-metals when electrons are shared.

• Examples of pure covalent molecules include H₂ and O₂
• Do not normally conduct electricity
• Have very few ions
• Low melting and boiling points
• Generally insoluble in water

• Single bonds — Form simple molecules that are essential in organic chemistry and act as building blocks for more complex compounds.
• Double bonds — Important for forming more stable and rigid molecular structures.
• Triple bonds — Create very strong bonds, contributing to materials and polymers with unique physical and chemical properties.

Polarity refers to the difference in electronegativity between two bonded atoms.

• Non-polar covalent: electronegativity difference < 0.4
• Polar covalent: electronegativity difference between 0.4 and 1.8
• Ionic: electronegativity difference > 1.8
• Higher polarity generally leads to higher melting and boiling points

Polarity direction:
δ⁺ ←— bond —→ δ⁻

Hydrogen bonding occurs when hydrogen forms a covalent bond with a highly electronegative element (such as oxygen, nitrogen, or fluorine).

• Creates a dipole within the molecule
• Results in equal and opposite partial charges
• Molecules have a partially positive and partially negative end
• Often produces a bent or asymmetric molecular structure

Metallic bonding occurs between metal atoms.

• Electrons move randomly and become delocalized
• Delocalized electrons form a “sea of electrons”
• Positive metal ions are held together by attraction to these electrons
• Metallic bonds are generally harder to break than non-metallic bonds

Van der Waals forces are weak, short-range electrostatic attractions between neutral atoms or molecules.

• Caused by temporary movement of electrons
• Uneven electron distribution creates a temporary dipole
• One side becomes slightly negative while the other becomes slightly positive

Dipole–dipole forces occur between molecules with permanent dipoles.

These form when an atom bonds with a highly electronegative atom, creating uneven charge distribution.

London dispersion forces are a type of intermolecular force caused by induced dipoles.

• Random electron movement creates a temporary dipole in one atom
• This dipole induces dipoles in neighboring atoms
• Occurs in all atoms and molecules, even those that are electrically symmetrical

The reactivity series arranges metals in order of how easily they lose electrons (their reactivity).

Order is determined by reactions with:

• Water — Potassium reacts violently, iron reacts very slowly
• Dilute acids — Metals above hydrogen produce hydrogen gas
• Air/oxygen — Magnesium oxidizes rapidly, gold does not react

Trend: Metals higher in the series are stronger reducing agents.

Common order:
K > Na > Ca > Mg > Al > Zn > Fe > Sn > Pb > Cu > Ag > Au

• Predicting products of chemical reactions
• Understanding compound stability
• Planning extraction methods and corrosion prevention

A more reactive metal displaces a less reactive metal from its compound.

Example: Zn + CuSO₄ → ZnSO₄ + Cu

• Carbonates: CaCO₃ → CaO + CO₂
• Nitrates: 2KNO₃ → 2KNO₂ + O₂
• Hydroxides: Ca(OH)₂ → CaO + H₂O

1. Concentration: Removal of impurities
2. Roasting: ZnS + O₂ → ZnO + SO₂
3. Reduction: ZnO + C → Zn + CO

1. Blast furnace: Fe₂O₃ + 3CO → 2Fe + 3CO₂
2. Removal of impurities: Flux reacts with Si, S, P to form slag

Rusting reaction:
4Fe + 3O₂ + 6H₂O → 4Fe(OH)₃ → Fe₂O₃·xH₂O

Conditions: Oxygen and water (humidity speeds up rusting)

Prevention methods:

• Painting or coating — physical barrier
• Sacrificial protection — more reactive metal corrodes first
• Cathodic protection — metal attached to negative terminal

Oxygen is blown through molten iron to oxidize impurities such as carbon, sulfur, phosphorus, and silicon, producing purer steel.

Alloying involves mixing metals to improve properties such as strength, hardness, and corrosion resistance.

• Carbon steel: High carbon → hard, low carbon → malleable
• Stainless steel: Chromium and nickel → corrosion resistant
• Percentage composition affects hardness, ductility, and resistance
• Used in construction, tools, and appliances

• Combination: A + B → AB
• Decomposition: AB → A + B
• Displacement: A + BC → AC + B
• Double displacement: AB + CD → AD + CB

• Oxidation: Loss of electrons or gain of oxygen
• Reduction: Gain of electrons or loss of oxygen
• Oxidizing agent: Causes oxidation, itself reduced
• Reducing agent: Causes reduction, itself oxidized

Example: Zn + Cu²⁺ → Zn²⁺ + Cu

Zinc is oxidized (loses electrons) and copper ions are reduced (gain electrons).

Electrolysis uses electricity to break down ionic compounds.

• Molten salts: Produce metal and non-metal
• Aqueous solutions: Products depend on electrode reactivity

Examples:

• Water: 2H₂O → 2H₂ + O₂
• Bauxite (Al₂O₃): Produces Al and O₂ using carbon electrodes
• CuSO₄:
  • Carbon electrodes — copper deposited at cathode
  • Copper electrodes — copper transfers from anode to cathode

Ionic equations show only the ions involved in the reaction.

• Electroplating: Coating metals for appearance and corrosion resistance
• Electrorefining: Purifying metals such as copper

• Simple cell: Produces electricity from spontaneous redox reactions
• Electrochemical series: Ranks metals by tendency to lose electrons
• Batteries: Multiple cells connected to increase voltage

• System: The part being studied
• Surroundings: Everything else
• Open system: Exchanges energy and matter
• Closed system: Exchanges energy only

An exothermic reaction releases energy to the surroundings, causing the temperature of the surroundings to increase.

• ΔH is negative
• Energy released during bond formation > energy required to break bonds

Example: Combustion of methane

CH₄ + 2O₂ → CO₂ + 2H₂O + energy

An endothermic reaction absorbs energy from the surroundings, causing the temperature of the surroundings to decrease.

• ΔH is positive
• Energy required to break bonds > energy released during bond formation

Example: Thermal decomposition of calcium carbonate

CaCO₃ → CaO + CO₂

• Exothermic: Products have lower energy than reactants; activation energy shown as an energy hump
• Endothermic: Products have higher energy than reactants; greater energy input required

ΔH = Energy of products − Energy of reactants

Units: kJ mol⁻¹

Calorimetry measures the heat released or absorbed during a chemical reaction.

Formula:

q = mcΔT

• m = mass of water
• c = specific heat capacity
• ΔT = change in temperature

The rate of reaction is the speed at which reactants are converted into products.

Measured by:

• Change in mass
• Volume of gas produced
• Change in concentration over time

• Concentration: Higher concentration → more frequent collisions → faster rate
• Temperature: Higher temperature → greater kinetic energy → more successful collisions
• Pressure (gases): Higher pressure → particles closer together → faster rate
• Surface area: Greater surface area → more exposed particles → faster rate
• Catalyst: Lowers activation energy and is not consumed

• Biological catalysts made of proteins
• Highly specific to substrates
• Work best at optimum temperature and pH
• Denature at high temperatures

A dynamic equilibrium occurs when forward and backward reactions happen at the same rate.

Occurs only in closed systems.

If a system at equilibrium is disturbed, it will shift to oppose the change.

Concentration:

• Increase reactants → shift right
• Increase products → shift left

Temperature:

• Exothermic forward reaction: increase temperature → shift left
• Endothermic forward reaction: increase temperature → shift right

Pressure (gases only):

• Increase pressure → shift towards fewer gas molecules

Catalyst:

• Does not change equilibrium position, only speeds up equilibrium

N₂ + 3H₂ ⇌ 2NH₃

ΔH = −92 kJ mol⁻¹

Conditions:

• Temperature: ~450°C (rate vs yield compromise)
• Pressure: ~200 atm
• Iron catalyst

Application: Fertilizers

• S + O₂ → SO₂
• 2SO₂ + O₂ ⇌ 2SO₃ (V₂O₅ catalyst)
• SO₃ + H₂SO₄ → H₂S₂O₇
• H₂S₂O₇ + H₂O → 2H₂SO₄

• Giant ionic lattice
• Strong electrostatic forces
• High melting and boiling points
• Conduct electricity when molten or dissolved

• H₂ — linear
• O₂ — linear
• H₂O — bent (104.5°)
• CO₂ — linear
• CH₄ — tetrahedral
• NH₃ — trigonal pyramidal

Diamond

• Each carbon bonded to four others
• Very hard, high melting point
• Electrical insulator

Graphite

• Each carbon bonded to three others
• Layers slide easily → soft
• Conducts electricity due to delocalized electrons

Silicon Dioxide (SiO₂)

• Giant covalent structure
• High melting point
• Hard and brittle

• Positive metal ions arranged in a lattice
• Surrounded by delocalized electrons
• Explains conductivity, malleability, and ductility

Combustion of coke:

C + O₂ → CO₂

Formation of carbon monoxide:

CO₂ + C → 2CO

Reduction of iron oxide:

Fe₂O₃ + 3CO → 2Fe + 3CO₂

Role of limestone (flux):

• CaCO₃ → CaO + CO₂
• CaO + SiO₂ → CaSiO₃ (slag)

Limestone removes acidic impurities such as silicon dioxide by forming molten slag, which floats on top of the molten iron.

Froth flotation:

Used to concentrate the sulfide ore by removing impurities.

Roasting:

2ZnS + 3O₂ → 2ZnO + 2SO₂

Reduction:

ZnO + C → Zn + CO

Purification:

Electrolysis is used to obtain pure zinc.

• Aluminium is too reactive to be reduced by carbon
• Extracted from bauxite (Al₂O₃) using electrolysis
• Cryolite (Na₃AlF₆): Lowers melting point and increases electrical conductivity

Electrode reactions:

Cathode:

Al³⁺ + 3e⁻ → Al

Anode:

2O²⁻ → O₂ + 4e⁻

The oxygen produced reacts with the carbon anodes to form carbon dioxide (CO₂), causing the anodes to wear away.

Food dyes make drinks and candies more attractive, but some dyes may cause health issues. Quinoline dye E104 is suspected of causing rashes in children and has the molecular formula C₉H₇N.

In quinoline, the nitrogen atom is covalently bonded within a solid organic compound and is not present in its elemental gaseous form. By contrast, at room temperature, nitrogen exists as diatomic molecules (N₂) in the gaseous state, where the particles are far apart and exhibit rapid, random motion due to their kinetic energy.

Q1.1 Using the Bohr model, draw and label the electronic structure of a nitrogen atom showing the distribution of its electrons in the respective energy levels.

Q1.2 Compare the movement and arrangement of the molecules in solid nitrogen to those in nitrogen gas.

Q1.3 Using the periodic table, state the group and period of nitrogen.

Nitrogen Isotopes

Q1.4 Calculate the relative atomic mass of nitrogen to two significant figures.

Q1.5 The relative atomic mass of carbon is 12.0, hydrogen is 1.0 and nitrogen is 14. Calculate the relative molecular mass (Mr) of quinoline (C₉H₇N). Show your working clearly.

Q1.6 The diagram shows the electronic configuration of an atom of element X.

Determine the number of protons contained in the nucleus of the atom.

Quinoline Yellow (E104) is a synthetic dye that produces a bright yellow colour. Chromatography can be used to separate mixtures and compare dyes with standard references. Students tested a soft drink using ethanol:water solvent ratios of 60:40, 70:30, and 50:50.

Hypothesis: If the soft drink contains Quinoline Yellow, then the solvent mixtures will produce a yellow spot with an R_f value close to 0.62, because solvent composition affects the movement and separation of dyes on chromatography paper.

Procedure Summary

1. Prepare three ethanol:water solvents (100 mL each).
2. Cut chromatography paper (≈ 2.5 × 10–12 cm) and draw a pencil baseline 1.5 cm from the bottom.
3. Spot the soft drink using a capillary tube and allow to dry.
4. Suspend paper in solvent with baseline above solvent level.
5. Allow solvent to rise to ~9.0 cm and mark solvent front.
6. Circle spots, measure distances, calculate R_f values.
7. Compare results with Quinoline Yellow standard.

Q2.1 Identify the independent variable, dependent variable, and at least two control variables.

Q2.2 State a suitable research question for the experiment.

Q2.3 Formulate your own hypothesis about the effect of different ethanol:water solvent ratios.

Q2.4 List all equipment used with correct quantities and sizes.

Q2.5 Suggest one safety measure and explain how the method is kept safe.

Q2.6 Explain how variables are manipulated and how sufficient, relevant data is collected.

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