ocr-a-level-chemistry-cheatsheet-periodic-table-and-energy

Page 1: Periodicity Cheat Sheet

Arrangement: Elements arranged by increasing atomic number.
Groupings:
- Periods: Horizontal rows.
- Groups: Vertical columns.
- Properties of Groups: Similar physical/chemical properties due to the same number of outer shell electrons.
Periodicity: Regularly repeating patterns of atomic properties with increasing atomic number.
Blocks in Periodic Table: Splits into s-, p-, d-, and f- blocks based on highest energy electron orbital.

Electron Filling: 2s subshell filled before 2p in Period 2; 3s before 3p in Period 3.
Melting Points:
- Trend: Increases from Group 1 to Group 14 due to giant structures; decreases from Group 14 to 15 due to simple molecular structure and weak intermolecular forces.
Atomic Radius: Decreases across a period as effective nuclear charge increases; no increase in shielding.

Definition: Energy needed to remove an electron from an atom.
First Ionisation Energy Equation: X(g) → X+(g) + e-.
Successive Ionisation Energies: Involves removing additional electrons, denoted as: X(n-1)+ (g) → Xn+(g) + e-.
Evidence for Shell Structure:
- Successive ionisation energies increase within shells due to less repulsion.
- Large jumps between shell ionisation energies because of electrons closer to nucleus being removed.
Factors Affecting Ionisation Energies:
- Atomic Radii: Larger radius leads to less nuclear attraction for outer electrons.
- Nuclear Charge: Greater nuclear charge results in a stronger attraction.
- Shielding Effect: More inner shells mean greater repulsion of outer electrons, reducing attraction.
Trends:
- Atomic radii decrease across a period, increase down a group.
- Ionisation energy increases across a period, decreases down a group.

Outer Shell: All have 2 electrons in the outer s-subshell.
Trends Down the Group:
- Ionisation Energy: Decreases due to increased atomic radius and shielding.
- Melting Point: Decreases due to weaker metallic bonding as atomic radii increase.
Reactions:
- With Water: Forms hydroxides and hydrogen gas: Mg (s) + 2H2O(l) → Mg(OH)2 (aq) + H2 (g).
- With Oxygen: Forms oxides: Mg (s) + O2 (g) → MgO (s).
- With Acids: Produces salts: Mg (s) + HCl(aq) → MgCl2 (aq) + H2 (g).
Reactivity Trend: Increases down the group as ionisation becomes easier.
Hydroxides Solubility:
- Mg(OH)2 (slightly soluble), Ca(OH)2 (sparingly soluble), Sr(OH)2 (more soluble), Ba(OH)2 (most soluble).

Molecular Form: Exist as diatomic molecules (X2).
Electronegativity: Decreases down the group due to increased atomic radius.
Boiling Points: Increase down the group due to stronger London forces.
Reactivity: Halogens gain one electron (s2p5, forming 1– ions).
- Displacement Reactions: More reactive halogens can displace less reactive ones from solutions.
Testing for Halide Ions:
- Precipitation reactions with silver ions: Ag+ + X– → AgX (s).
Use of Chlorine: Chlorine in water for purification creates HCl and HClO, both with oxidizing properties.
- Advantages: Kills bacteria, prevents diseases; however, has potential health risks.

Definition: Enthalpy (H) is thermal energy stored in a system.
Enthalpy Change: Heat energy change at constant pressure (ΔHƟ under standard conditions).
Standard Enthalpy Change Definitions:
- Reaction (ΔHr ϴ): Enthalpy change for specific molar quantities.
- Formation (ΔfHƟ): Heat change for forming one mole from elements.
- Combustion (ΔcHƟ): Heat change for complete combustion in oxygen.
- Neutralisation (ΔneutHϴ): Heat change of acid-base reactions forming one mole of water.

Key Points: Particles collide but not all result in reactions.
- Must collide with sufficient energy and correct orientation.

Concept: Shows the distribution of kinetic energies at constant temperature.
- Area under the curve = total number of molecules; peak represents most probable energy.

Function: Increases reaction rate without being consumed.
- Types: Homogeneous (same phase) and heterogeneous (different phase).
Benefits: Economic, environmental advantages, and reduced energy consumption.

Temperature: Higher temperatures increase kinetic energy, leading to more collisions.
Pressure: In gaseous reactions, increased pressure brings molecules closer, increasing collisions.
Concentration: More reactant molecules increase collision frequency.
Catalysts: Provide an alternative pathway with lower activation energy.

Reversible Reactions: Denoted by ⇌; in dynamic equilibrium.
Le Chatelier’s Principle: Shifts equilibrium to oppose changes.
- Temperature Change: Affects equilibrium based on reaction's enthalpy.
- Concentration Change: Shifts to either produce or consume reactants/products.
- Pressure Change: Affects gases; shifts to the side with fewer moles.
Catalysts: Speed up approach to equilibrium but do not change its position.

Equilibrium Reaction: N2 + 3H2 ⇌ 2NH3; ΔH = –92 kJ mol–1.
Optimum Conditions:
- High pressure (favors product formation) but cost/risks.
- Low temperature (favoring exothermic reaction) but slow.
Compromise Conditions: Typically used: 400–500 °C and 200 atm with iron catalyst.

Definition: Pressure exerted by each gas in a mixture.
Mole Fraction & Total Pressure Relationship: p(A) = mole fraction x total pressure.

KC and Kp: Indicate equilibrium position based on concentrations or partial pressures.
- Affect of Temperature and Pressure on equilibrium and constants.