A Level Chemistry Revision Notes

Atomic Structure

  • Atoms consist of protons, neutrons, and electrons.

Subatomic Particles

  • Protons:

    • Location: Nucleus

    • Mass: 1

    • Charge: +1

  • Neutrons:

    • Location: Nucleus

    • Mass: 1

    • Charge: 0

  • Electrons:

    • Location: Outer shells

    • Mass: Approximately 1/1836 the mass of a proton (very small)

    • Charge: -1

  • Charges are relative; actual charges are measured in coulombs but using +1, 0, -1 is simpler.

  • Masses are determined relative to carbon-12.

Atomic Structure Observations

  • Atoms are mostly empty space.

  • Nucleus Diameter: 10−1510^{-15} meters

  • Atom Diameter: 10−1010^{-10} meters

  • Atomic models have evolved over time with new evidence.

Periodic Table

  • Boxes contain numbers; the larger number is the mass number (A).

  • Mass number: Average mass of naturally occurring isotopes; can be a decimal due to averaging.

  • Atomic number (Z): Number of protons (smaller number).

  • Mass number (A) = number of protons + number of neutrons

Isotopes

  • Different versions of an element (e.g., carbon-12 and carbon-14).

  • Same atomic number (same number of protons).

  • Different mass number (different number of neutrons).

  • Same electronic structure, hence same chemical properties but may differ in physical properties due to mass differences.

Definitions

  • Relative Molecular Mass (Mr): Average mass of a molecule compared to 1/12 the mass of one carbon-12 atom.

  • Relative Atomic Mass (Ar): Average mass of one atom compared to 1/12 the average mass of one carbon-12 atom.

Mass Spectrometry

  • Used to determine the average mass of naturally occurring isotopes.

  • Example: Boron

    • 20% Boron-10

    • 80% Boron-11

    • Relative atomic mass of boron = (20×10)+(80×11)100=10.8\frac{(20 \times 10) + (80 \times 11)}{100} = 10.8

Ionization Energy

  • First Ionization Energy: Energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous +1 ions.

    • Equation: H(g)→H+(g)+e−H(g) \rightarrow H^+(g) + e^-

    • Equation: Na(g)→Na+(g)+e−Na(g) \rightarrow Na^+(g) + e^-

  • Second Ionization Energy: Energy required to remove one electron from each ion in one mole of gaseous +1 ions to form one mole of gaseous +2 ions.

    • Equation: He+(g)→He2+(g)+e−He^+(g) \rightarrow He^{2+}(g) + e^-

    • Equation: Na+(g)→Na2+(g)+e−Na^+(g) \rightarrow Na^{2+}(g) + e^-

Factors Affecting Ionization Energy

  • Atomic Radius: Larger distance between nucleus and outer electrons reduces attraction.

  • Electron Shielding/Repulsion: Inner electrons repel outer electrons, reducing attraction.

  • Nuclear Charge: More protons increase attraction between nucleus and outer electrons.

Trends in Successive Ionization Energies

  • Ionization energy increases as electrons are removed.

  • Less repulsion among remaining electrons.

  • Big jump in ionization energy indicates a change in electron shell.

Trends in Ionization Energies Across Periods

  • Ionization energy increases across a period.

  • Sharp drop between the end of one period and the beginning of the next provides evidence for electron shells.

Ionization Energy and Electron Structure

  • Small drops in ionization energy between groups 2 and 3 (e.g., Be to B) and between groups 5 and 6 (e.g., N to O) provide evidence for electron configuration.

  • Boron's fifth electron is the first in the 2p subshell, making it easier to remove.

  • Oxygen’s eighth electron is the first to be paired in the 2p subshell, making it easier to remove.

Electronic Structure

  • Atoms have shells, subshells, and orbitals.

  • Periodic table blocks (s, p, d, f) correspond to subshells and orbitals.

Filling Orbitals

  • Example: Calcium (20 electrons)

  • Electrons fill from the bottom up, singly first, with opposite spins in each orbital (box).

  • Electron configuration for calcium: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²

Shells, Subshells, and Orbitals

  • Shell: Energy level.

  • Subshell: s, p, d, f

  • Orbital: Region within a subshell that can hold up to two electrons.

Shapes of Atomic Orbitals

  • s: sphere

  • p: dumbbell

  • Each orbital holds two electrons.

  • First shell: 2 electrons (1s²)

  • Second shell: 8 electrons (2s² 2p⁶)

Shorthand Electronic Configuration

  • Argon (18 electrons): 1s² 2s² 2p⁶ 3s² 3p⁶ or [Ar]

  • Calcium (20 electrons): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² or [Ar] 4s²

Periodic Table Blocks

  • s-block, p-block, d-block, f-block

Period 3 Trends

  • Similar trends occur in other periods (e.g. period 2).

Atomic Radius

  • Decreases across a period.

  • Increased number of protons leads to greater positive charge in the nucleus, attracting outer electrons inwards.

First Ionization Energy

  • Decrease between Mg and Al as the s-shell fills and p-orbitals start filling.

  • Another drop between group 5 and 6 (P and S) as pairing starts, increasing repulsion.

Bonding and Melting/Boiling Points

  • Na, Mg, and Al: strong metallic bonds, high melting/boiling points.

  • Si: strong covalent bonds, giant structure, high melting/boiling points.

  • Cl, S, and P: simple covalent molecules with weak intermolecular bonds, low boiling points.

  • Ar: monoatomic, very low boiling point.

Ionic Bonding

  • Occurs between metals and nonmetals.

  • Involves the transfer of electrons from a metal to a nonmetal, forming positive metal ions (cations) and negative nonmetal ions (anions).

  • Electrostatic attraction between ions forms the ionic bond.

Magnesium Chloride Example

  • Magnesium (Mg) has two outer shell electrons, chlorine (Cl) has seven.

  • Magnesium transfers one electron to each chlorine atom, forming Mg²⁺ and two Cl⁻ ions.

  • Formula: MgCl₂

Electronic Configurations

  • Magnesium loses 3s² electrons, chlorine gains one electron to achieve a stable 3p⁶ configuration.
    *

Properties of Ionic Compounds

  • Ions are attracted to all surrounding ions, not just the ones they gained or lost electrons with.

  • Electrostatic attraction is stronger with higher charges and smaller ions.

  • Ions are atoms that have lost or gained electrons.

Predicting Ionic Charges

  • Group 1: +1 ions

  • Group 2: +2 ions

  • Transition metals: variable oxidation states

  • Group 7: -1 ions

  • Group 6: -2 ions

  • Elements aim for noble gas configurations.

  • Example: Calcium (Ca): loses 4s² electrons to form Ca²⁺

Trends in Ionic Radii

  • Increases down a group due to increasing number of electron shells.

Isoelectronic Ions

  • Ions with the same electronic structure (e.g., ions with the same electronic structure as neon).

  • As the number of protons increases in isoelectronic ions, the nuclear attraction to the electrons increases, causing the ionic radii to decrease.

Properties of Ionic Structures

  • Ionic lattices: Large, 3D structures.

  • Example: Sodium Chloride (NaCl): Each Na⁺ ion is surrounded by six Cl⁻ ions, and vice versa.

General Properties

  • High Melting and Boiling Points: Strong electrostatic attractions require large amounts of energy to overcome.

  • Solubility in Water: Water is polar and interacts with the ions, pulling them apart and dissolving them.

  • Electrical Conductivity: Conduct electricity when molten or dissolved because the ions can move freely but not in solid-state.

  • Hardness and Brittleness: Ions are fixed in place and cannot move around, making the solids hard but brittle.

Covalent Bonding

  • Occurs between nonmetals through the sharing of electrons.

Types of Covalent Bonds

  • Single Bond: One pair of shared electrons.

  • Double Bond: Two pairs of shared electrons (four total).

  • Triple Bond: Three pairs of shared electrons (six total).

  • Lewis dot structures can be used to represent covalent compounds.

Examples

  • Oxygen (O₂): double bond

  • Nitrogen (N₂): triple bond

Dative Covalent Bonding

  • One atom donates both electrons for the bond.

  • Example: Ammonium ion (NH₄⁺): the nitrogen in ammonia (NH₃) donates both electrons to the hydrogen ion (H⁺).

  • The dative covalent bond is represented with an arrow indicating the direction of electron donation.

Molecular Shapes

  • Molecules have specific shapes determined by the number of bonding pairs and lone pairs of electrons around the central atom.
    *

Molecular Shapes and Bond Angles

  • Linear: Examples: CO₂, HCN, BeF₂ Bond angle: 180° No lone pairs.

  • Trigonal Planar: Examples: SO₃, BCl₃, AlCl₃. Bond angle: 120° No lone pairs.

  • Tetrahedral: Examples: CH₄, NH₄⁺. Bond angle: 109.5° No lone pairs.

  • Trigonal Pyramidal: Examples: NH₃, ClO₃. Bond angle: 107° One lone pair.

  • Bent: Examples: H₂O. Bond angle: 104.5° Two lone pairs.

  • Trigonal Bipyramidal: Example: PCl₅. Bond angles: 120° and 90° No lone pairs. Expanded octet.

  • Octahedral: Example: SF₆. Bond angles: 90° No lone pairs. Expanded octet.

Valence Shell Electron Pair Repulsion (VSEPR) Theory

  • Electrons in the outer shell arrange themselves as far apart as possible due to electrostatic repulsion.

Examples and Impact of Lone Pairs:

  • Molecules with four pairs of electrons in the outer shell:

    • CH₄: 4 bonding pairs, bond angle 109.5°

    • NH₃: 3 bonding pairs, 1 lone pair, bond angle 107°

    • H₂O: 2 bonding pairs, 2 lone pairs, bond angle 104.5°

  • Lone pairs are more repulsive than bonding pairs.

  • As the number of lone pairs increases, the bond angle decreases.

  • Lone pair - lone pair repulsion > lone pair - bonding pair repulsion > bonding pair - bonding pair repulsion

Electronegativity and Polarity

Electronegativity

  • A measure of how much an element attracts electrons in a chemical bond.

  • Fluorine (F) is the most electronegative element.

Trends in Electronegativity

  • Across a Period: Electronegativity increases as the number of protons increases.

    • Number of electron shells remains the same, atomic radius decreases as electrons are pulled in.

  • Down a Group: Electronegativity decreases.
    -Increasing number of shells increases shielding around the nucleus.

Spectrum of Bonding

  • Pure Covalent Bonding: Occurs when two elements are the same; electron cloud is shared evenly.

  • Polar Covalent Bonding: Occurs when different elements are covalently bonded; electron cloud is shifted more towards the more electronegative element.

  • Ionic Bonding: Occurs when the difference in electronegativity is large enough that electrons are transferred from one atom to another.

  • Partial dipoles are set up within the bond due to differences in electronegativity.

  • If the electronegativity difference is greater than 2.0, the bond is generally considered ionic.

  • Sharing electrons is a continuous spectrum with covalent bonding at one end and ionic bonding at the other end.
    *

Examples

  • Hydrogen Gas (H₂): Equal electronegativity, electrons shared equally.
    HCl: Chlorine is more electronegative, attracts electrons more, setting up a dipole.

Intermolecular Forces

  • Forces between molecules.

  • Permanent dipole-permanent dipole forces: Attraction between delta negative and delta positive parts of polar molecules. Higher melting and boiling points due to stronger intermolecular bonding. For asymmetrical molecule, the forces are going to balance each other out and it will not be polar.

  • Induced Dipoles (London Dispersion Forces): Occur in all substances (except ionic compounds) due to random movement of electrons.

    • Strength depends on: Number of electrons and Shape of the molecule. Straight isomers will have more contact points, meaning the intermolecular forces will be stronger and they will have a higher boiling point.

  • Hydrogen Bonding: Special case of dipole-dipole interactions occurring when hydrogen is bonded to either nitrogen, oxygen, or fluorine.
    Consequences of Hydrogen Bonding

  • Anomalously high boiling points.

  • Ice is less dense than liquid water.

  • High specific heat capacity of water.

Chain Length and Intermolecular Forces in Alkenes

  • Long-chain alkanes: As chain length increases, so does the opportunity for intermolecular forces, thus increasing the boiling point.

  • Branching: Branching decreases the opportunity for intermolecular forces, decreasing the boiling point compared to a straight chain of the same length.

Alcohols

  • Polar due to the electronegative oxygen atom, leading to stronger intermolecular bonds.

  • Lower volatility compared to similar-length alkanes.

Hydrogen Halides

  • Boiling points generally increase as you move down the group from HCl to HI (due to larger halogens and increasing strength in intermolecular bonds).

  • Hydrogen fluoride (HF) exhibits hydrogen bonding, resulting in a higher boiling point than would otherwise be expected.

  • These are very strong intermolecular bonds and requires a lot of energy to break them.

Solvents and Intermolecular Forces

  • Choice of solvent is dependent on whether it is polar or nonpolar, and whether the compound to be dissolved is polar or nonpolar.

  • "Like dissolves like" rule.

  • Water: good solvent for ionic compounds due to hydration of ions and for alcohols due to hydrogen bonding.

  • Capsaicin (in chilies): nonpolar compound; milk is a suitable nonpolar solvent to remove the burning sensation, while water is not.

Metallic Bonding and Properties

  • Model: Positive blue metal ions and delocalized green electrons.

  • Bonding: Electrostatic attraction between delocalized electrons and positive ions.

  • Factors affecting bond strength: smaller ions, more delocalized electrons or a more positive nucleus have stronger attractions.

Properties of Giant Metallic Lattices

  • High Melting and Boiling Points: Strong electrostatic attractions require lots of energy to overcome.

  • Insoluble in Water

  • Conductivity: Conduct electricity and heat when solid or molten because electrons can move freely.

  • Ductility and Malleability: Metal ions can slide over each other.

Giant Covalent Macromolecular Structures

Silicon Dioxide (SiO₂)

  • High melting and boiling points: Strong intramolecular bonds require large amounts of energy to overcome.

  • Insoluble in water.

  • Poor electrical conductors: Electrons are not freely available to move around.

Allotropes of Carbon

Diamond

  • Each carbon makes four carbon-carbon bonds.

  • Very hard: Lattice structure.

  • Does not conduct electricity: All electrons are involved in bonding, no free electrons.

Graphite

  • Each carbon makes three carbon-carbon bonds.

  • Soft: Atoms are arranged in layers that can slide over each other.

  • Conducts electricity: One free electron that is not involved in bonding can move around.

Simple Molecular Structures

Common Examples

  • Water (H₂O)

  • Ammonia (NH₃)

  • Nitrogen Gas (N₂)

  • Carbon Dioxide (CO₂)

  • Oxygen Gas (O₂)

Properties