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LC Chemistry
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•Equilibrium
Change in Pressure
An increase in pressure favours the side with fewer moles of gas. A decrease favours the side with more moles.
•Equilibrium
Change in Temperature
Increasing temperature favours the endothermic reaction. Decreasing temperature favours the exothermic reaction.
•Equilibrium
Change in Concentration of Products
Increasing product concentration favours the reverse reaction. Decreasing it favours the forward reaction.
•Equilibrium
Change in Concentration of Reactants
Increasing reactant concentration favours the forward reaction. Decreasing it favours the reverse reaction.
•Equilibrium
Kc is affected by...
Only temperature affects the value of the equilibrium constant Kc. Concentration and pressure do not.
•Equilibrium
Le Châtelier's Principle states...
If a system at equilibrium is stressed, the system will shift to oppose the stress and restore equilibrium.
•Equilibrium
Dynamic Equilibrium
Both forward and reverse reactions continue to occur at equal rates. Concentrations remain constant.
•Equilibrium
Chemical equilibrium
Occurs when the rate of the forward reaction equals the rate of the reverse reaction in a closed system.
•Analytical Techniques
Chromatography principle and applications
Components move at different rates due to attraction to mobile and stationary phases. Used to test drug purity, alcohol levels, and dye components.
•Analytical Techniques
Mass spectrometry principle and applications
Positive ions are separated by mass in a magnetic field. Identifies isotopes and compounds by mass spectra.
•Organic Chemistry
Where are esters found?
Esters occur naturally in fats and oils. They are formed in condensation reactions.
•Organic Chemistry
Use of ethyl ethanoate
Used as a solvent in perfumes and as a flavouring agent due to its sweet smell.
•Organic Chemistry
Ketones suffix
Ketones use the suffix '-anone'. Ex: propanone is a common ketone.
•Organic Chemistry
Propanone use
Used as a solvent in nail varnish remover due to its ability to dissolve oils and plastics.
•Organic Chemistry
Aldehyde solubility
Aldehydes are soluble in water and cyclohexane due to their polar carbonyl group.
•Organic Chemistry
Unsaturated
A compound is unsaturated if it contains at least one carbon-carbon double or triple bond.
•Organic Chemistry
Use of methanol
Used as a denaturant in industrial alcohols to prevent consumption. It is toxic.
•Organic Chemistry
Why is ethanol an important solvent
Ethanol dissolves both polar and non-polar substances. Used in perfumes, drinks, and as a biofuel.
•Organic Chemistry
Alcohols solubility
Alcohols are very soluble in water due to hydrogen bonding. Solubility drops as carbon chain length increases.
•Organic Chemistry
Why do alcohols have high boiling points?
Strong hydrogen bonds between molecules increase boiling points compared to alkanes.
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•Organic Chemistry
Primary alcohol
Has the hydroxyl group on a carbon attached to only one other carbon. Ex: propanol.
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•Organic Chemistry
Homologous series
A family of compounds with the same functional group and general formula, differing by CH₂.
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•Fuels & Energy
Catalytic converter
Converts harmful gases into safer ones. Uses platinum, palladium, and rhodium as catalysts.
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•Rates of Reaction
Surface Adsorption Theory
Heterogeneous catalysts like nickel adsorb reactants onto their surface, increasing reaction rate. Ex: catalytic converter.
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•Rates of Reaction
Catalysts work by...
Catalysts lower the activation energy required for a reaction, increasing the rate.
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•Rates of Reaction
Catalysis
The increase in rate of a reaction due to the presence of a catalyst. The catalyst remains unchanged.
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•Rates of Reaction
Activation energy
The minimum energy that reacting particles must have to successfully collide and form products.
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•Rates of Reaction
Enzymes
Biological catalysts that speed up reactions in living organisms. They are highly specific.
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•Rates of Reaction
Factors affecting rate of reaction
Factors include concentration, temperature, particle size, nature of reactants, and presence of a catalyst.
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•Rates of Reaction
Rate of reaction
The change in concentration of a reactant or product per unit time. Measured in mol L⁻¹ s⁻¹.
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•Fuels & Energy
Why is it not a good fuel
Hydrogen is hard to store and transport. It is highly flammable and explosive.
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•Fuels & Energy
Why is hydrogen not found naturally
Hydrogen is too light and reactive to exist freely in nature. It bonds readily with other elements.
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•Fuels & Energy
Hydrogen production
Produced by electrolysis of water or steam reforming of natural gas. Electrolysis is cleaner.
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•Fuels & Energy
Reference fuels
2,2,4-trimethylpentane (octane no. 100) and heptane (octane no. 0) are used to compare knocking.
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•Fuels & Energy
Catalytic cracking
Breaks long alkanes into short alkanes (fuels) and alkenes (plastic feedstock) using heat and a catalyst.
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•Fuels & Energy
Reforming
Increases octane number by using catalysts to convert straight-chain alkanes to aromatic hydrocarbons.
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•Fuels & Energy
Isomerisation
Raises octane number by converting straight-chain alkanes into branched isomers using a catalyst.
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•Fuels & Energy
Octane number
A number that shows how well a fuel resists knocking. Higher octane = less knocking.
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•Fuels & Energy
Knocking
Premature ignition of the fuel-air mix in an engine. It reduces efficiency and damages the engine.
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•Fuels & Energy
Auto ignition
The spontaneous ignition of the petrol-air mix before the spark plug fires. Causes knocking.
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•Thermochemistry
Hess's Law
Total enthalpy change is the same whether a reaction occurs in one or several steps.
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•Thermochemistry
Heat of formation
The heat change when one mole of a compound is formed from its elements in their standard states.
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•Thermochemistry
Kilogram calorific value
The heat change when 1 kg of a substance is burned completely in excess oxygen. Used to compare fuels.
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•Thermochemistry
Heat of combustion
The heat released when one mole of a substance is burned in excess oxygen. Exothermic and measured in kJ/mol.
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•Thermochemistry
Bond energy
The average energy needed to break one mole of covalent bonds into separate atoms in the gas phase.
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•Thermochemistry
Heat of reaction
The heat change when the amounts in a balanced equation react completely. Units are usually kJ.
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•Thermochemistry
Endothermic reaction
A reaction that absorbs heat from its surroundings. Positive ΔH. Ex. photosynthesis.
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•Thermochemistry
Exothermic reaction
A reaction that releases heat to the surroundings. Negative ΔH. Ex. combustion.
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•Organic Chemistry
Test for unsaturation
Add bromine water. A colour change from yellow/orange to colourless indicates a double or triple bond.
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•Organic Chemistry
Alkanes are...
Non-polar molecules due to only single bonds and similar electronegativities.
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•Organic Chemistry
Alkane saturation
Alkanes are saturated hydrocarbons containing only single bonds.
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•Organic Chemistry
Saturated compounds
Molecules with only single carbon-carbon bonds. Ex. ethane, propane.
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•Organic Chemistry
Alkane structure
Each carbon in an alkane is tetrahedral, with a bond angle of 109.5°.
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•Organic Chemistry
Alkane solubility
Alkanes are insoluble in water but soluble in non-polar solvents like cyclohexane.
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•Organic Chemistry
Functional group
A specific group of atoms that determines the chemical properties of a homologous series. Ex. –OH in alcohols.
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•Organic Chemistry
3 homologous series
Examples: alkanes, alkenes, alkynes. Each differs by a CH₂ unit.
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•Organic Chemistry
Aliphatic compounds
Molecules that consist of straight or branched chains of carbon atoms. Not aromatic.
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•Organic Chemistry
Homologous series
A series of compounds with the same functional group and similar properties, differing by CH₂.
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•Fuels & Energy
Natural gas is almost pure...
Methane. CH₄ is the main component of natural gas.
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•Fuels & Energy
Why is methane a hazard in coal mines, slurry pits and refuse dumps?
Methane forms explosive mixtures with air. Proper ventilation is required to prevent explosions.
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•Fuels & Energy
Methane is a...
Methane is a good fuel and also a greenhouse gas that contributes to global warming.
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•Fuels & Energy
How is methane produced?
Methane is produced from decaying plant and animal material in anaerobic conditions.
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•Organic Chemistry
Hydrocarbons
Compounds made only of hydrogen and carbon. Can be saturated or unsaturated.
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•Volumetric Analysis
Why do we use H₂SO₄ in permanganate titration?
H₂SO₄ is stable and ensures Mn⁷⁺ reduces fully. HCl and HNO₃ interfere or oxidise themselves.
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•Volumetric Analysis
Why is potassium permanganate used in a weak solution?
It is a strong oxidising agent and has limited solubility. Too concentrated = rapid uncontrolled reaction.
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•Volumetric Analysis
Why is potassium permanganate not a primary standard?
It is not very soluble and decomposes over time. Cannot be weighed out accurately.
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•Volumetric Analysis
Mn⁷⁺ (purple) is reduced to...
Mn²⁺ (colourless). The purple colour disappears during redox titration.
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•Volumetric Analysis
Standardised
Means concentration was determined using a titration with a primary standard.
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•Volumetric Analysis
Indicator for Na₂CO₃ titration
Methyl orange is used because carbonates are weak bases and it changes in acidic range.
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•Volumetric Analysis
Indicator for ethanoic acid titration
Phenolphthalein, as it changes colour in the pH range for weak acid–strong base titrations.
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•Volumetric Analysis
Primary standard used to standardise HCl
Anhydrous sodium carbonate (Na₂CO₃) is used. It is pure, stable, and soluble.
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•Volumetric Analysis
Methyl orange colour change
Orange in neutral solution. Turns red/pink in acid after endpoint.
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•Volumetric Analysis
Phenolphthalein colour change
Pink in base. Turns colourless in acid at the titration endpoint.
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•Volumetric Analysis
Why dilute the vinegar to 500 cm³?
Concentrated vinegar would require too much NaOH and titrate too quickly, increasing error.
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•Acids & Bases
Conjugate acid-base pair
An acid-base pair differs by one proton. The acid has the proton; the base has lost it.
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•Acids & Bases
Difference between strong and weak base
Strong bases fully accept protons. Weak bases only accept a small fraction. Ex. NaOH vs NH₃.
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•Acids & Bases
Difference between strong and weak acid
Strong acids fully donate protons. Weak acids donate partially. Ex. HCl vs CH₃COOH.
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•Acids & Bases
Brønsted-Lowry base
A species that accepts a proton (H⁺). Ex. OH⁻ or NH₃.
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•Acids & Bases
Brønsted-Lowry acid
A species that donates a proton (H⁺). Ex. HCl or H₂SO₄.
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•Acids & Bases
Limitations of Arrhenius Theory
Doesn’t explain H₃O⁺ formation, excludes bases like NH₃, and doesn’t apply to non-aqueous reactions.
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•Acids & Bases
Arrhenius base
Substance that produces OH⁻ ions in aqueous solution. Ex. NaOH.
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•Acids & Bases
Arrhenius acid
Substance that produces H⁺ ions in aqueous solution. Ex. HCl.
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•Everyday Chemistry
Household bases
Common examples: oven cleaner (NaOH), baking soda (NaHCO₃).
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•Everyday Chemistry
Household acids
Examples include vinegar (ethanoic acid) and Coca-Cola (carbonic and phosphoric acid).
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•Acids & Bases
Neutralisation
Reaction of an acid with a base to form salt and water. Ex. HCl + NaOH → NaCl + H₂O.
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•Acids & Bases
Salt
Formed when the H⁺ in an acid is replaced by a metal or ammonium ion. Ex. NaCl from HCl + NaOH.
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•Acids & Bases
Base
A base dissociates in water to form OH⁻ or accepts protons. Can be strong or weak.
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•Acids & Bases
Acid
An acid produces H⁺ (or H₃O⁺) in water or donates a proton. Tastes sour and turns indicators red.
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•Gas Laws
Under what conditions of temperature and pressure would a real gas come closest to ideal behaviour?
A real gas behaves most ideally at low pressure and high temperature. Under these conditions, intermolecular forces and molecular volume become negligible.
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•Gas Laws
Why is the kinetic theory wrong?
Real gases have forces between molecules and their particles occupy space. These violate two assumptions of the ideal gas model.
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•Gas Laws
Assumptions of Kinetic Theory
1. No forces between molecules. 2. Volume of molecules is negligible. 3. Collisions are elastic. 4. Particles move randomly in straight lines. 5. Average kinetic energy is proportional to temperature.
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•Gas Laws
Ideal gas
A gas that obeys all kinetic theory assumptions under all conditions of temperature and pressure. No real gas is perfectly ideal.
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•Gas Laws
Molar volume of a gas
At STP (0°C, 1 atm), one mole occupies 22.4 L. At room temperature and pressure (RTP), it occupies 24 L.
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•Mole Concept
A mole
A mole contains 6 × 10²³ particles. It is the SI unit for the amount of substance.
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•Mole Concept
Avogadro's number
Avogadro’s number is 6 × 10²³ particles per mole. It applies to atoms, ions, and molecules.
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•Gas Laws
Avogadro's Law
Equal volumes of gases at the same temperature and pressure contain equal numbers of molecules.
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•Gas Laws
Gay-Lussac's Law of combining volumes
When gases react at constant temperature and pressure, their volumes and the volumes of any gaseous products are in simple whole number ratios.
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•Gas Laws
Charles' Law
At constant pressure, the volume of a gas is directly proportional to its temperature in Kelvin: V₁/T₁ = V₂/T₂.
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•Gas Laws
Boyle's Law
At constant temperature, the pressure of a gas is inversely proportional to its volume: P₁V₁ = P₂V₂.
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•Bonding & Shapes
2 bonding pairs + 2 lone pairs
V-shaped molecule. Ex: H₂O. Bond angle approx. 104.5° due to lone pair repulsion.
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•Bonding & Shapes
3 bonding pairs + 1 lone pair
Pyramidal shape. Ex: NH₃. Bond angle approx. 107° due to one lone pair.
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•Bonding & Shapes
4 bonding pairs
Tetrahedral shape. Ex: CH₄. Bond angle is 109.5°.
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•Bonding & Shapes
3 bonding pairs
Trigonal planar shape. Ex: BCl₃. Bond angle is 120°.
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•Bonding & Shapes
2 bonding pairs
Linear shape. Ex: BeCl₂. Bond angle is 180°.
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•Bonding & Shapes
Hydrogen bonding
Strong dipole-dipole force between H and highly electronegative atoms (F, O, or N). Ex: H₂O, NH₃.
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•Bonding & Shapes
Dipole-dipole forces
Attractive forces between opposite poles of polar molecules. Ex: Methanal has stronger forces than ethane.
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•Bonding & Shapes
Van der Waals forces
Weak temporary attractions due to induced dipoles. Increase with molecule size. Ex: O₂ has a higher boiling point than H₂.
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•Bonding & Shapes
Electron pair repulsion theory
Electron pairs around a central atom repel as far as possible. Lone pairs repel more strongly than bonding pairs.
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•Periodic Table
Down a group: Electronegativity
Electronegativity decreases. Atoms have more shells and greater shielding, so attraction for electrons drops.
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•Periodic Table
Across a period: Electronegativity
Electronegativity increases. More protons and smaller atomic radius cause stronger attraction for bonding electrons.
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•Bonding & Shapes
Electronegativity and bonding
0 = pure covalent, 0–0.4 = non-polar covalent, 0.4–1.7 = polar covalent, >1.7 = ionic bond.
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•Bonding & Shapes
Electronegativity
The relative attraction an atom has for a shared pair of electrons in a covalent bond.
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•Bonding & Shapes
Non-polar bond
A covalent bond where electrons are shared equally. Ex: Cl₂, O₂.
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•Bonding & Shapes
Polar covalent bond
A covalent bond where electrons are unequally shared. One atom is slightly negative, the other slightly positive.
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•Bonding & Shapes
3 characteristics of covalent substances
1. Don’t conduct electricity. 2. Made of neutral molecules. 3. Often low melting and boiling points.
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•Bonding & Shapes
2 examples of non-polar substances
Helium and neon are non-polar due to symmetrical distribution of electrons.
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•Bonding & Shapes
2 examples of polar molecules
Water and ammonia are polar due to lone pairs and asymmetric shape.
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•Bonding & Shapes
Double bond
Two pairs of electrons shared between atoms. Ex: O=C=O in carbon dioxide.
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•Bonding & Shapes
Covalent bond
A bond formed when two atoms share a pair of electrons to achieve noble gas configuration.
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•Bonding & Shapes
Anions
Negatively charged ions formed when atoms gain electrons. Ex: Cl⁻.
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•Bonding & Shapes
Cations
Positively charged ions formed when atoms lose electrons. Ex: Na⁺.
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•Bonding & Shapes
Ion
An atom or molecule that has gained or lost electrons, giving it a charge.
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•Bonding & Shapes
Ionic bond
Attraction between oppositely charged ions. Forms between metals and non-metals. Ex: NaCl.
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•Bonding & Shapes
Octet rule
Atoms bond to achieve eight electrons in their outer shell. Exceptions: H, He, transition metals.
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•Bonding & Shapes
Valency
The number of hydrogen atoms or other monovalent atoms an element can combine with. Based on outer electrons.
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•Chemical Reactions
Compound
A substance formed when two or more elements chemically bond together in fixed proportions. Ex: H₂O.
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•Oxidation & Reduction
Common reducing agents
Sulfur dioxide (bleaching), carbon (metal extraction), carbon monoxide (metal extraction), hydrogen (organic reduction).
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•Oxidation & Reduction
Common oxidising agents
Oxygen (combustion), potassium permanganate (titration), sodium dichromate (organic prep), hydrogen peroxide (bleaching).
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•Oxidation & Reduction
Reducing agent
A substance that donates electrons or reduces another species. It itself is oxidised.
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•Oxidation & Reduction
Oxidising agent
A substance that gains electrons or oxidises another species. It itself is reduced.
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•Oxidation & Reduction
Reduction
1. Gain of electrons. 2. Decrease in oxidation number. Ex: Fe³⁺ → Fe²⁺.
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•Oxidation & Reduction
Oxidation
1. Loss of electrons. 2. Increase in oxidation number. Ex: Fe²⁺ → Fe³⁺.
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•Atomic Structure
First ionisation energy down a group
Decreases due to increased atomic radius and more inner shells that shield outer electrons.
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•Atomic Structure
First ionisation energy across a period
Increases due to more protons in the nucleus and smaller atomic radius. Electrons are held more tightly.
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•Atomic Structure
First ionisation energy
The minimum energy needed to remove the most loosely bound electron from a neutral gaseous atom in its ground state.
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•Atomic Structure
Atomic radius down a group
Increases due to the addition of extra shells and increased shielding effect.
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•Atomic Structure
Atomic radius across a period
Decreases as the number of protons increases, pulling electrons closer with no extra shielding.
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•Atomic Structure
Atomic radius
Half the distance between nuclei of two atoms of the same element bonded by a single covalent bond.
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•Atomic Structure
Ions
Charged atoms or groups of atoms. Cations are positive, anions are negative.
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•Analytical Techniques
Uses of the AAS
Used to detect metals in water (e.g., lead, cadmium) and in blood samples (e.g., lead exposure).
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•Atomic Structure
When an electron moves from a lower to a higher energy level...
Energy is absorbed from surroundings. This leads to excitation of the electron.
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•Atomic Structure
When an electron moves from a higher to a lower energy level...
Energy is emitted in the form of light (a photon). This gives line spectra.
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•Atomic Structure
Excited state
When electrons occupy higher energy levels than in the ground state. Temporary condition.
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•Atomic Structure
Ground state
When electrons occupy the lowest available energy levels. Most stable arrangement.
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•Atomic Structure
Ionisation energy
The energy needed to remove the most loosely held electron from a neutral gaseous atom.
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•Atomic Structure
Heisenberg uncertainty principle
It is impossible to know both the position and velocity of an electron at the same time.
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•Atomic Structure
Orbital
Region of space around the nucleus where there is a high probability of finding an electron.
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•Atomic Structure
Sublevel
Subdivisions of energy levels (s, p, d, f), each with slightly different energies.
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•Atomic Structure
Line spectrum
Spectrum showing light at specific wavelengths. Produced when electrons fall to lower levels.
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•Atomic Structure
Absorption spectrum
A spectrum with dark lines showing absorbed wavelengths. Light is absorbed as electrons are excited.
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•Atomic Structure
Emission spectrum
A spectrum with bright coloured lines. Formed when excited electrons return to lower levels.
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•Radioactivity
Uses of radioisotopes
Used in cancer treatment, medical imaging, archaeological dating, food irradiation, and smoke detectors.
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•Radioactivity
Half-life
The time taken for half of the atoms in a radioactive sample to decay.
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•Radioactivity
Radioisotope
An isotope with an unstable nucleus that emits radiation as it decays.
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•Radioactivity
When an isotope undergoes beta decay
Its atomic number increases by 1. Mass number remains the same.
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•Radioactivity
When an isotope undergoes alpha decay
Its atomic number decreases by 2, and its mass number decreases by 4.
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•Radioactivity
Chemical reaction
Involves changes in electron arrangement. Bonds are broken and formed. Elements do not change.
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•Radioactivity
Nuclear reactions
Involve changes in the nucleus and may change one element into another. Include alpha and beta decay.
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•Radioactivity
Gamma rays are
Electromagnetic waves with no mass or charge. Highly penetrating and not deflected by fields.
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•Radioactivity
Gamma rays
Electromagnetic radiation. Highly penetrating, damages cells, not deflected, blocked by thick lead.
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•Radioactivity
Beta particles are
High-speed electrons emitted from the nucleus during radioactive decay.
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•Radioactivity
Beta particles
Fast, moderately penetrating. Atomic number increases by 1. Stopped by 5 mm of aluminium.
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•Radioactivity
Alpha particles are
Helium nuclei (2 protons, 2 neutrons) with a +2 charge. Low penetration.
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•Radioactivity
Alpha particles
Slow and heavy. Stopped by paper. Atomic number drops by 2, mass by 4. Eg: Americium in smoke alarms.
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•Radioactivity
Measuring radiation
Use a Geiger-Müller tube to detect and count radioactive emissions.
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•Radioactivity
Pierre and Marie Curie
Discovered polonium and radium. Pioneers in radioactivity research.
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•Radioactivity
Henri Becquerel
Discovered radioactivity using uranium salts and photographic plates.
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•Radioactivity
Radioactivity
The spontaneous breakdown of an unstable atomic nucleus with emission of radiation.
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•Analytical Techniques
Mass spectrometer steps
1. Vaporisation 2. Ionisation 3. Acceleration 4. Separation by mass 5. Detection
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•Analytical Techniques
Mass spectrometer uses
Used to identify isotopes, perform drug tests, and determine relative atomic and molecular masses.
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•Atomic Structure
Relative atomic mass
The average mass of an atom of an element, taking into account all its isotopes and their abundances.
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•Atomic Structure
Isotope
Atoms of the same element with the same atomic number but different mass numbers. Different numbers of neutrons.
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•Atomic Structure
Mass number
Sum of protons and neutrons in the nucleus of an atom. Determines the isotope.
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•Atomic Structure
Atomic number
The number of protons in the nucleus of an atom. Determines the element.
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•Atomic Structure
James Chadwick
Discovered the neutron by bombarding beryllium with alpha particles. Neutrons are neutral and found in the nucleus.
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•Atomic Structure
Conclusion of Gold Foil Experiment
Most of the atom is empty space. The positive charge and mass are concentrated in a small dense nucleus.
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•Atomic Structure
Gold Foil Experiment (Rutherford)
Alpha particles were fired at gold foil. Most passed through, some were deflected, and a few reflected, showing the presence of a dense nucleus.
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•Atomic Structure
Thomson's plum pudding model
Proposed atoms as a sphere of positive charge with embedded electrons. Later disproved by Rutherford’s experiment.
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•Atomic Structure
Robert Millikan
Used the oil drop experiment to measure the charge of the electron and calculate its mass.
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•Atomic Structure
JJ Thomson
Discovered the electron using the cathode ray tube. Showed that atoms contain negatively charged particles.
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•Atomic Structure
Discharge tube
A sealed glass tube with metal electrodes used to observe the flow of electric discharge. Cathode rays were discovered using this.
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•Atomic Structure
Cathode rays
Streams of electrons emitted from the cathode in a discharge tube. Negatively charged and deflected by electric fields.
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•Atomic Structure
John Dalton's Atomic Theory
1. Matter is made of atoms. 2. All atoms of an element are identical. 3. Atoms are indivisible. 4. Atoms combine in fixed ratios.
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•Chemical Reactions
Law of Conservation of Mass
Mass is conserved in a chemical reaction. No mass is lost or gained.
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•Periodic Table
Differences in Mendeleev's periodic table
Left gaps for unknown elements, no noble gases, arranged by atomic mass, d-block elements in subgroups.
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•Periodic Table
Why did Mendeleev place tellurium before iodine?
Their chemical properties fit better that way. Later explained by Moseley’s atomic number concept.
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•Periodic Table
Dmitri Mendeleev
Created the periodic table and grouped elements by properties. Left gaps for undiscovered elements like gallium.
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•Periodic Table
Law of Octaves
Every eighth element had similar properties when arranged by atomic mass. Discarded due to inconsistencies.
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•Periodic Table
John Newlands
Proposed the Law of Octaves. Grouped elements by atomic mass in rows of 8.
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•Periodic Table
Law of Triads
Triads are groups of 3 elements with similar properties. The middle element's mass was an average of the other two.
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•Periodic Table
Johann Dobreiner
Formulated the Law of Triads based on similarities in properties of groups of three elements.
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•Periodic Table
Henry Moseley element definition
An element is a substance whose atoms all have the same atomic number.
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•Periodic Table
Henry Moseley
Discovered the atomic number using X-ray spectra. Re-arranged the periodic table by atomic number.
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•Periodic Table
Humphrey Davy
Used electrolysis to isolate reactive elements like potassium, sodium, and calcium.
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•Periodic Table
Robert Boyle
Defined an element as a substance that cannot be broken down into simpler substances. Early definition of elements.
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•Chemical Formulas
What is empirical formula?
The simplest whole-number ratio of atoms of each element in a compound. Ex: CH₂O is the empirical formula of glucose.
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•Radioactivity
Describe beta particles
Beta particles are fast-moving electrons. Penetrate 2–3 mm of aluminium. Ex: used in carbon-14 dating.
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•Radioactivity
Describe alpha particles
Helium nuclei with +2 charge. Low penetration. Stopped by paper. Ex: Americium-241 in smoke detectors.
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•Radioactivity
What are the 3 types of radioactive isotopes?
Alpha particles, beta particles, and gamma rays. Each has different penetration and effects.
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•Atomic Structure
Define an isotope
Atoms of the same element with the same atomic number but different mass numbers. Different number of neutrons.
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•Analytical Techniques
What is the principle of the mass spectrometer?
Positively charged ions are separated based on their mass in a magnetic field.
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•Analytical Techniques
What are the 5 fundamental processes that the mass spectrometer works by?
1. Vaporisation 2. Ionisation 3. Acceleration 4. Separation 5. Detection
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•Analytical Techniques
How does a mass spectrometer work?
Ions are sorted by mass-to-charge ratio. Lighter ions are deflected more by the magnetic field.
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•Atomic Structure
Define relative atomic mass
The average mass of an atom compared to one-twelfth the mass of a carbon-12 atom. Includes all isotopes.
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•Bonding & Shapes
What is a polar covalent bond?
A bond where electrons are shared unequally. One atom is slightly negative, the other slightly positive. Ex: HCl.
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•Bonding & Shapes
Define electronegativity
The relative ability of an atom to attract a shared pair of electrons in a covalent bond.
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•Periodic Table
What trend in electronegativity occurs across the periodic table and why?
Electronegativity increases across a period due to increased nuclear charge and decreasing atomic radius.
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