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Ionisation Energy

  • Definition of Ionisation Energy: Ionisation energy measures the energy required to completely remove an electron from an atom of an element to form an ion.

  • First Ionisation Energy: The energy required to remove one electron from each atom in one mole of the gaseous element to form one mole of gaseous 1+ ions.

    • Equation:
      X(g) → X^+(g) + e^-

  • Successive Ionisation Energies: Applies to the removal of electrons after the first ionisation energy.

    • nth Ionisation Energy:
      X^{(n-1)+}(g) → X^n+(g) + e^-

  • Significance: Successive ionisation energies provide evidence for the shell structure of atoms.

    • Within each shell, successive ionisation energies increase due to decreased electron repulsion.

    • Between shells, big jumps in ionisation energies occur, as the electron is removed from a shell closer to the nucleus.

  • Factors Affecting Ionisation Energies:

    • Atomic Radii: Larger atomic radius increases distance of outer electrons from the nucleus, reducing nuclear attraction.

    • Nuclear Charge: Higher nuclear charge increases attractive force on outer electrons.

    • Shielding: Increased number of inner shells creates greater repulsion on outer shell electrons, reducing attraction.

  • Concepts on Ionisation Energies:

    • Greater electron attraction results in higher ionisation energy.

    • Periodicity in Atomic Radii:

    • Across a period: radius decreases.

    • Down a group: radius increases.

    • Trends in Ionisation Energy:

    • Increases across a period due to greater attraction in same shell.

    • Decreases down a group due to increased atomic radius and shielding, reducing attraction.

The Mole & Concentration

  • Definition of Mole: The mole is the unit used to quantify the amount of a substance, applicable to any chemical species (atoms, electrons, molecules, ions).

  • Mole Definition: A mole contains the same number of atoms or particles as 12 g of carbon-12.

    • Avogadro Constant:
      N_A = 6.022 imes 10^{23} ext{ mol}^{-1}

  • Relationship:

    • n = rac{m}{M}
      where:

    • n is number of moles (mol)

    • m is mass (g)

    • M is molar mass (g mol^{-1})

  • Concentration of a Solution: Amount of solute in a known volume of solution.

    • c = rac{n}{V}
      where:

    • c is concentration (mol dm^{-3})

    • n is number of moles in solution (mol)

    • V is volume (dm^3)

  • Conversions:

    • 1 dm^3 = 1000 cm^3

    • 1 m^3 = 1000 dm^3

Gas Equations

  • Molar Volume of Gas: One mole of any gas under standard conditions occupies the same volume.

    • Molar Gas Volume: 24 dm^{-3} mol^{-1} at 298 K and 100 kPa.

  • Number of Moles of Gas Calculation:

    • n = rac{V}{24}
      where V is volume (dm^{-3}).

  • Ideal Gas Assumptions:

    • Intermolecular forces between gas particles are negligible.

    • The volume of particles compared to their container's volume is negligible.

  • Ideal Gas Equation:

    • PV = nRT
      where:

    • P is pressure (Pa)

    • V is volume (m^{3})

    • n is number of moles (mol)

    • R is gas constant (8.314 J K^{-1})

    • T is temperature (K)

Empirical & Molecular Formula

  • Empirical Formula: The simplest whole-number ratio of atoms of each element in a compound.

    • Calculation from mass or percentage by mass.

    • Example: P and O combine to form phosphorous oxide:

    • Mass of O2: 14.2 ext{ g} - 6.2 ext{ g} = 8 ext{ g}

    • Number of moles of each element calculated, giving empirical formula P2O5.

  • Molecular Formula: Shows the number and type of atoms of each element in a molecule.

    • Derived from empirical formula and relative molecular mass.

    • Example: For empirical formula CH2, relative molecular mass of 224:

    • 224/14=16

    • Molecular formula: 16 imes CH2 = C{16}H_{32}

Balanced Equations

  • Chemical Reaction Principle: Atoms are neither created nor destroyed; they rearrange in products and reactants.

  • Balanced Equation: Equal number of atoms of each element in both reactants and products.

  • State Symbols: Show the physical state of each species:

    • Solid (s), Liquid (l), Gaseous (g), Aqueous (aq).

  • Ionic Equations: Include only reacting ions and products.

    • Example: NaCl (aq) + AgNO3 (aq) → AgCl (s) + NaNO3 (aq)

    • Full ionic: Na^+(aq) + Cl^-(aq) + Ag^+(aq) + NO3^-(aq) → AgCl(s) + Na^+(aq) + NO3^-(aq)

    • Net Ionic Equation: Cl^-(aq) + Ag^+(aq) → AgCl(s)

  • Stoichiometry: Expresses molar ratios of reactants and products in a reaction.

Atom Economy & Percentage Yield

  • Atom Economy: The theoretical measure of atoms from reactants forming the desired product; requires a balanced chemical equation.

    • % Atom Economy:
      ext{Atom Economy} = rac{ ext{mass of desired product}}{ ext{mass of reactants}} imes 100

  • Importance: Maximizing atom economy leads to:

    • More sustainable practices (fewer raw materials).

    • Minimizing waste products.

    • Higher efficiency and cost savings on separation processes.

  • Limiting Reagent: Determines theoretical yield; it is the reagent not in excess.

  • Percentage Yield: Measures efficiency of reaction route:

    • ext{Percentage Yield} = rac{ ext{actual yield}}{ ext{theoretical yield}} imes 100

  • Reducing Factors: Formation of by-products, unreacted reactants, and losses in extraction impact percentage yield.

Ions

  • Definition of Ion: An atom or molecule with a net charge from gain or loss of electrons.

  • Cations vs Anions:

    • Cation: Formed by loss of electrons.

    • Anion: Formed by gain of electrons.

  • Noble Gases: Do not form ions due to fully filled outer electron shells; hence they're unreactive.

  • Molecular Ions: Groups of covalently bonded atoms gaining/losing electrons. Examples:

    • +1: Ammonium (NH4^+)

    • -1: Hydroxide (OH^-), Nitrate (NO3^-), Hydrogencarbonate (HCO3^-)

    • -2: Carbonate (CO3^{2-}), Sulphate (SO4^{2-}), Sulphite (SO3^{2-}).

Ionic Bonding

  • Definition: Electrostatic attraction between positive and negative ions resulting in a lattice structure.

  • Lattice Structure: A three-dimensional framework where each ion is surrounded by oppositely charged ions.

  • Common Formation: Occurs between metals and non-metals aiming for stable outer shells.

  • Factors Affecting Ionic Bond Strength:

    • Charge of ions: Higher charges result in stronger bonds.

    • Distance between ions: Smaller ions can form stronger bonds due to increased attraction.

  • Properties of Ionic Structures:

    • High melting/boiling points.

    • Soluble in polar solvents.

    • Conducts electricity when molten or dissolved in water.

Covalent Bonds

  • Definition: Strong electrostatic attraction between shared pair of outer electrons and nuclei of bonded atoms.

  • Types of Covalent Bonds:

    • Single Covalent Bond: One shared pair.

    • Multiple Covalent Bond: More than one shared pair.

    • Coordinate (Dative Covalent Bond): One atom donates both electrons.

  • Crystalline Structures: Types of structures and their properties:

    • Diamond: Giant covalent, high melting point, does not conduct.

    • Graphite: Conductor due to delocalized electrons, high melting point.

    • Ice: Molecular structure, low melting point.

    • Iodine: Molecular, low melting point.

    • Sodium Chloride: Ionic, high melting point, conducts when aqueous/molten.

    • Metals: Metallic bonding with lattice of positive ions and delocalized electrons; high melting/boiling points, good conductors.

Simple Molecules

  • Molecular Shape: Determined by electron pairs around the central atom, including bonding and lone pairs.

  • Repulsion: Electron pairs minimize repulsion by arranging far apart.

  • Bond Angles and Types: Each type of molecular geometry has associated bond angles influenced by lone pairs.

    • Examples:

    • Linear: CO2 (180°)

    • Trigonal Planar: BF3 (120°)

    • Bent: SO2 (<120°)

    • Tetrahedral: CH4 (109.5°)

    • Trigonal Pyramidal: NH3 (107°)

    • V-Shaped: H2O (104.5°)

    • Trigonal Bipyramidal: PCl5 (90°, 120°)

    • Seesaw: SF4 (87°, 102°)

    • T-Shaped: ClF3 (88°)

    • Octahedral: SF6 (90°)

    • Square Pyramidal: BrF5 (<90°)

    • Square Planar: XeF4 (90°)

Enthalpy

  • Definition of Enthalpy: The thermal energy stored in a system, symbolized as H.

  • Enthalpy Change (ΔH): Heat energy change at constant pressure, usually under standard conditions (100 kPa, 298 K).

    • Standard Enthalpy Change:
      ΔH^Ɵ = ext{total enthalpy of products} - ext{total enthalpy of reactants}

  • Key Types of Enthalpy:

    • Standard Enthalpy of Formation (ΔfHƟ): Change when one mole of a substance forms from its elements.

    • Standard Enthalpy of Combustion (ΔcHƟ): Change when one mole of a substance completely burns in oxygen.

Exothermic & Endothermic Reactions

  • Exothermic Reactions: Release heat; ΔH is negative.

  • Endothermic Reactions: Absorb heat; ΔH is positive.

  • Activation Energy: Minimum energy needed for a reaction.

Calorimetry

  • Definition: Measuring heat changes during chemical reactions.

  • Experimental Variability:

    • Incomplete reactions, incomplete combustion, approximated heat capacities, heat loss, non-standard conditions, and evaporation can lead to discrepancies in reported values.

  • Heat Change Equation:

    • q = mcΔT
      where:

    • q is heat change (J)

    • m is mass of substance (g)

    • c is specific heat capacity (J g^{-1} K^{-1})

    • ΔT is temperature change (K or °C).

  • Coffee Cup Calorimetry: Used to calculate enthalpy changes of neutralisation.

    • Process: Measure volumes and temperatures, insulate reaction; plot temperature against time

Bond Enthalpies

  • Bond Enthalpy Definition: Energy needed to break a bond, averaged over various compounds.

  • Endothermic and Exothermic: Bond breaking is endothermic; bond formation is exothermic.

  • Estimation: Bond enthalpy values are approximate and less accurate than Hess's Law values.

Hess’s Law

  • Hess's Law: Enthalpy change of a reaction is independent of the route taken between reactants and products.

  • Applications: Use combustion and formation enthalpy to find enthalpy changes of a reaction.

  • Standard Enthalpy of Formation: 0 kJ mol^-1 for elements.

Collision Theory

  • Principle: Particles in a gas/liquid move and collide, but not all collisions cause reactions.

    • For a reaction to occur:

    • Sufficient energy to overcome activation energy.

    • Correct orientation for collision.

Maxwell-Boltzmann Distribution

  • Maxwell-Boltzmann Curve: Distribution of molecular kinetic energies in a gas at constant temperature.

  • Characteristics:

    • Area under curve = total molecules.

    • Curve starts at origin; no zero energy molecules.

    • Only molecules with energy > activation energy can react.

    • Peak represents most probable energy; average energy is to the right of the peak.

  • Impact of Temperature and Catalysts:

    • Higher temperature increases kinetic energy and leads to more successful collisions.

    • Catalysts lower activation energy, increasing reaction rate.

Rate of Reaction

  • Definition: Change in concentration of a reactant or product over time.

  • Factors Influencing Rate:

    • Temperature: Higher temperatures increase kinetic energy, leading to more frequent collisions.

    • Pressure: Increases concentration of gaseous reactants, leading to more collisions.

    • Concentration: Higher concentration in solutions increases successful collisions.

    • Surface Area: Increased surface area allows for more contact area, increasing reactions.

    • Catalysts: Provide alternative pathways with lower activation energy, increasing reaction rate.

Chemical Equilibrium

  • Reversible Reactions: Represented with equilibrium symbol (⇌).

  • Dynamic Equilibrium: Conditions with constant concentrations of reactants/products with equal forward and reverse reaction rates in a closed system.

  • Le Chatelier’s Principle: Changes will shift equilibrium to oppose perturbations (concentration, temperature, pressure).

  • Effect of Conditions on Equilibrium:

    • Changing temperature affects reaction direction per heat addition/removal rules.

    • Changing concentration shifts equilibrium to consume excess or replenish.

    • Changing pressure shifts to favor fewer moles of gas in a reaction.

The Haber Process

  • Haber Process Reaction: N2(g) + 3H2(g) ⇌ 2NH3(g) ΔH = -92 kJ mol^-1.

  • Optimum Conditions:

    • High Pressure: Increases yield, favors product side (fewer moles).

    • Low Temperature: Favors exothermic direction, but slows reaction rate.

  • Compromise Conditions: Striking balance between yield rate and economic feasibility in industry:

    • 400–500 °C, 200 atm, Iron catalyst.

Redox Reactions

  • Redox Reaction Composition: Combination of oxidation (electron loss) and reduction (electron gain) half-equations to form full equations.

  • Half Equation Example:

    1. Oxidation: Mg → Mg^{2+} + 2e^-

    2. Reduction: Cu^{2+} + 2e^- → Cu

  • Equilibrium Constant (Kc): Ratio of concentrations of products to reactants at equilibrium (for the general reaction: aA + bB ⇌ dD + eE).

Lattice Enthalpy

  • Definition: Enthalpy change when one mole of ionic lattice forms from gaseous ions under standard conditions.

  • Opposite Expressions: Formation and dissociation enthalpies are inverses.

  • Indicators of Ionic Bond Strength: Lattice enthalpy informs strength of the ionic bonds, examining formation/dissociation extremes; individual bonds and their energies come into play especially in ionic compounds and their crystal structures.