Science Test Revision
Part 1: Atoms and the Periodic Table
Key Definitions
Element: Pure substance made of one type of atom
Atom: Smallest unit of an element
Subatomic particles
Proton: +1, nucleus
Neutron: 0, nucleus
Electron: –1, around nucleus
Atomic number (Z): number of protons
Mass number (A): protons + neutrons
Isotopes: same element, different neutrons (different mass)
Ions: charged atoms
Cations (+): lose electrons
Anions (–): gain electrons
Periodic Table
Arranged by increasing atomic number
Periods (rows): number of electron shells
Groups (columns): number of valence electrons (main group)
Groups
Group 1: Alkali metals
Group 2: Alkaline earth metals
Groups 3–12: Transition metals
Group 17: Halogens
Group 18: Noble gases
f-block: Lanthanides and actinides
Same group → similar chemical properties
Element 118: oganesson
Families and Trends
Group properties
Group 1: +1 ions, very reactive, soft, not found pure
Group 2: +2 ions, reactive, less than Group 1
Group 17: –1 ions, diatomic (F₂, Cl₂, Br₂, I₂), reactivity decreases down
Group 18: unreactive, found in atmosphere
Transition metals: hard, high melting, often coloured, some magnetic (Fe, Co, Ni)
Group 14: changes from non-metal → metalloid → metal down group
Down a group
More shells added
More shielding
Outer electrons further from nucleus
Held less tightly
Reactivity trends
Group 1 and 2 metals: increase down the group
Group 17 halogens: decrease down the group
Reaction Patterns
Group 1 + halogens
2M + X₂ → 2MX
Examples:
2Na + Cl₂ → 2NaCl
2Na + Br₂ → 2NaBr
2Fr + I₂ → 2FrI
Group 2
With halogens: M + X₂ → MX₂
Be + Cl₂ → BeCl₂
Ba + Br₂ → BaBr₂
With water:
Cold: Mg + 2H₂O → Mg(OH)₂ + H₂
Steam: Mg + H₂O → MgO + H₂
Halogens + hydrogen sulfide
X₂ + H₂S → S + 2HX
Reactivity order: F₂ > Cl₂ > Br₂ > I₂
Carbon and Organic Chemistry
Carbon forms most known compounds (>90%)
Allotropes
Diamond: hard, non-conductive
Graphite: layered, conducts electricity
Carbon nanotubes: strong, light, conductive
Amorphous carbon
Organic chemistry: study of carbon compounds
Includes fuels, polymers, medicines
Molecular machines: molecules that perform mechanical tasks
Electron Structure
Shells
Maximum electrons = 2n²
Shell 1: 2
Shell 2: 8
Shell 3: 18
Shell 4: 32
Subshells
s: 2 electrons
p: 6 electrons
d: 10 electrons
f: 14 electrons
Filling order
1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p
Electron configurations
Na: 1s² 2s² 2p⁶ 3s¹ → (2,8,1)
N: 1s² 2s² 2p³ → (2,5)
Key ideas
Period = number of shells
Group = number of valence electrons (main group)
Octet rule: atoms aim for 8 valence electrons (or 2 for H, He)
Chemical Symbols and Isotopes
Symbols: 1–2 letters (some from Latin, e.g. Na)
Atomic mass on table = weighted average of isotopes
Isotopes have same atomic number → same position
Chemical Equations Basics
Subscripts: number of atoms
Coefficients: number of molecules
States:
(s), (l), (g), (aq)
Law of conservation of mass: atoms are rearranged, not created/destroyed
Part 2: Chemical Reactions
Balancing Equations
Do not change subscripts
Use coefficients
Use smallest whole numbers
Check atoms and charges
Examples
2Al + 6HCl → 2AlCl₃ + 3H₂
2Ca + O₂ → 2CaO
Types of Reactions
Decomposition
AB → A + B
Types:
Thermolytic (heat)
Electrolytic (electricity)
Photolytic (light)
Examples:
2NaN₃ → 2Na + 3N₂
2H₂O → 2H₂ + O₂
2AgCl → 2Ag + Cl₂
Synthesis (Combination)
A + B → AB
Examples:
C + O₂ → CO₂
2H₂ + O₂ → 2H₂O
Acid reactions
Acid + metal → salt + H₂
Acid + base → salt + water
Acid + carbonate → salt + water + CO₂
Examples:
Mg + 2HCl → MgCl₂ + H₂
HCl + NaOH → NaCl + H₂O
CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂
Precipitation Reactions
Two aqueous solutions → solid forms
Example:
AgNO₃ + NaCl → AgCl(s) + NaNO₃
Net ionic:
Ag⁺ + Cl⁻ → AgCl(s)
Spectator ions remain in solution
Solubility Rules
Soluble
Group 1 and NH₄⁺
NO₃⁻, CH₃COO⁻
Most Cl⁻, Br⁻, I⁻ (except Ag⁺, Pb²⁺, Hg₂²⁺)
Most SO₄²⁻ (except Ca²⁺, Sr²⁺, Ba²⁺, Pb²⁺, Ag⁺)
Insoluble
Most OH⁻ (except Group 1, NH₄⁺; Ca²⁺, Sr²⁺, Ba²⁺ slightly)
Most S²⁻ (except Group 1, NH₄⁺, Ca²⁺, Sr²⁺, Ba²⁺)
Most CO₃²⁻, PO₄³⁻ (except Group 1, NH₄⁺)
Ionic Compounds
Must be neutral overall
Name: cation + anion
Use brackets for multiple ions
Ca(OH)₂
Transition metals can have different charges
FeCl₃ = iron(III) chloride
Combustion
Complete: hydrocarbon + O₂ → CO₂ + H₂O
Incomplete: → CO or C + H₂O
Example
2C₂H₆ + 7O₂ → 4CO₂ + 6H₂O
Balancing combustion
Balance C → H → O
Allow fractions, then multiply to whole numbers
Part 3: Collision Theory and Rates
Collision Theory
Reaction only occurs if:
Particles collide
Enough energy (≥ activation energy)
Correct orientation
Activation Energy
Minimum energy required
Energy profiles
Exothermic: products lower energy
Endothermic: products higher energy
Maxwell–Boltzmann
Particles have a range of energies
Higher temperature → more particles above activation energy
Rate of Reaction
Rate = change in concentration over time
Rate = −Δ[reactant]/Δt = +Δ[product]/Δt
Units: mol L⁻¹ s⁻¹
Steeper graph = faster reaction
Initial rate = slope at start
Measuring Rate
Mass change (gas lost)
Gas volume
Colour change (spectrophotometer)
Factors Affecting Rate
Temperature
Increases energy and successful collisions
Concentration / Pressure
More particles → more collisions
Surface Area
Smaller particles → more collisions
Catalyst
Lowers activation energy
Not used up
Stoichiometry and Rate
Example:
Mg + 2HCl → MgCl₂ + H₂
HCl reacts twice as fast as Mg
Common Mistakes
Changing subscripts when balancing
Forgetting state symbols
Not balancing properly
Mixing up group trends
Forgetting ion charges
Not using whole-number coefficients