AECHE Task 1 Cheatsheet

01/12/2025

AECHE Task 1 Cheatsheet

1. Atomic Theory & History

The Structure of the Atom

An atom is the smallest unit of an element retaining its chemical properties. It consists of two distinct regions:

The Nucleus: A dense, positively charged center containing protons and neutrons (nucleons). It is 10,000–100,000 times smaller than the atom but contains 99.97% of the mass.
The Electron Cloud: Mostly empty space where electrons orbit. It determines the atom’s volume.

Forces at Play:

Electrostatic Force: The attraction satisfying the rule that opposite charges (+ and −) attract. This holds electrons in orbit around the positive protons.
Strong Nuclear Force: Since protons naturally repel each other (like charges repel), this force overcomes electrostatic repulsion to hold the nucleus together. It only acts over extremely short distances.

Historical Models (Timeline)

You must know the scientist, their model, and the experimental evidence.

Scientist	Model Name	Key Discovery/Concept	Flaw/Correction
John Dalton	Solid Sphere	1. All matter is atoms. 2. Atoms of an element are identical. 3. Atoms combine in whole ratios.	Incorrect: Atoms can be divided (subatomic particles). Atoms of one element are not strictly identical (isotopes exist).
J.J. Thomson	Plum Pudding	Discovered the Electron. Modeled atom as a positive sphere with negative electrons embedded like “plums in a pudding”.	Proved wrong by Rutherford because the mass isn’t evenly distributed.
Ernest Rutherford	Nuclear Model	Gold Foil Experiment: Fired alpha particles (+) at gold foil. Most passed through (atom is empty space), but some deflected back (hit a dense, positive center). Discovered the Nucleus.	Could not explain why electrons didn’t spiral into the nucleus (energy loss).
Niels Bohr	Planetary Model	Quantized Energy Levels. Electrons orbit in fixed shells. They cannot exist between shells.	Works well for Hydrogen but struggles with complex electron interactions.
James Chadwick	Neutron Discovery	Discovered the Neutron. Explained why the nucleus was heavier than just protons alone.	N/A

2. Isotopes & Mass Spectrometry

This typically involves calculation questions and “describe the process” extended responses.

Isotope Definitions

Isotopes: Atoms of the same element (same Z, protons) but different mass numbers (A) due to a different number of neutrons.
Chemical vs. Physical: Isotopes have identical chemical properties (same electron configuration) but different physical properties (mass, density, boiling point).
Stability: If the neutron-to-proton ratio is unbalanced, the nucleus becomes unstable and radioactive.

Mass Spectrometry (The 5 Steps)

Memorize this sequence. It is the standard “5-mark” question.

Vaporization: The sample is heated until it becomes a gas to ensure particles can separate ,.
Ionization: An electron beam bombards the vapor, knocking off electrons to form positive ions (usually +1, rarely +2).
Acceleration: An electric field (negative plates) accelerates the positive ions into a focused beam.
Deflection: A magnetic field deflects the ions. The curve depends on the mass-to-charge ratio (m/z). Lighter ions deflect more; heavier ions deflect less.
- Note: Heavier isotopes = Larger radius (deflect LESS). Lighter isotopes = Smaller radius (deflect MORE).
Detection: A detector counts the ions hitting it at specific positions, generating a spectrum of relative abundance vs. mass.

Exam tip: Check units and significant figures when calculating relative abundances — answers should be consistent with the given precision.

Calculations

Relative Atomic Mass (Ar): The weighted average mass of an atom compared to 1/12th the mass of Carbon-12. It has no units.

Ar = (Σ (Isotope Mass × % Abundance)) / 100

3. Electronic Structure & Spectroscopy

Key concept: Energy is discrete (quantized), not continuous.

Shells, Subshells, and SPDF

Bohr described shells (n = 1, 2, 3…) where max electrons = 2n². However, the modern quantum model introduces subshells and explains specific Periodic Table blocks ,.

s-block: Groups 1 & 2 (spherical orbital).
p-block: Groups 13–18 (dumbbell orbital).
d-block: Transition metals.
f-block: Lanthanides/Actinides.

Example Configuration (Calcium, Z = 20):

Bohr: 2, 8, 8, 2.
SPDF: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s².

Important: Electron filling order follows the Aufbau principle, Hund’s rule, and the Pauli exclusion principle. Note the 4s orbital fills before 3d but may be lost first when ionised.

Identifying Excited States

Ground State: Electrons fill the lowest energy levels first (e.g., 2, 8, 8, 1).
Excited State: An electron has jumped to a higher level before the lower level is full.
Example: Potassium Ground state is 2, 8, 8, 1. An excited state might look like 2, 8, 7, 2 (one electron jumped from the 3rd to the 4th shell).

A+ Paragraph: Absorption vs. Emission Spectra

Use this exact phrasing for extended response questions:

Absorption Spectrum: When white light passes through a cool gas, electrons in the atoms absorb specific photons of energy that correspond exactly to the energy difference between their ground state and a higher energy shell. This causes the electrons to jump to an excited state. The frequencies absorbed appear as black lines on a continuous background ,.

Emission Spectrum: When excited electrons in a hot gas or flame return from a higher energy level (excited state) to the ground state, they release the absorbed energy as photons of light. Because electron energy levels are quantized (fixed), only specific wavelengths are emitted. This appears as colored lines on a black background. This spectral fingerprint is unique to every element.

Flame Test: Follows the emission principle. Heat excites the electrons; as they relax, they emit visible light, changing the flame color.

4. Periodic Table Trends

Do not just memorize the arrows. You must explain it using Core Charge and Shielding.

Key Definitions

Atomic Radius: Distance from nucleus to valence electrons.
Ionization Energy (1st IE): Energy required to remove the outermost electron from a gas-phase atom.
Electronegativity: Ability of an atom to attract electrons in a chemical bond.
Core Charge: Protons − Inner Shell Electrons. This is the effective pull the nucleus has on the valence shell.
Shielding: The repulsion from inner-shell electrons that “blocks” the pull of the nucleus on valence electrons.

Trend Logic (The “Why”)

Trend	Down a Group (↓)	Across a Period (→)
Atomic Radius	Increases. More shells are added. Shielding increases, reducing the pull on valence electrons.	Decreases. Shells stay the same, but protons increase. Higher Core Charge pulls electrons closer.
Ionization Energy	Decreases. Because radius increases (electrons are further away) and shielding increases, electrostatic attraction is weaker. It is easier to steal an electron.	Increases. Core charge increases and radius decreases. Stronger electrostatic attraction holds electrons tighter.
Electronegativity	Decreases. Weak attraction for outside electrons due to distance and shielding.	Increases. High core charge and small radius mean the nucleus attracts external electrons strongly.

Quick note: Watch for exceptions (e.g., between groups 2→13 and 15→16) due to subshell electron pairing and penetration.

Successive Ionization Energy & The “Big Jump”

We can remove electrons one by one. The energy required always increases (as the ion becomes more positive).
The Rule: Look for the massive jump in energy. This jump occurs when you switch from removing valence electrons to removing inner shell electrons (which are closer to the nucleus).
Application: If the jump happens between the 6th and 7th electron, the element has 6 Valence Electrons (Group 16).

Reactivity Trends

Metals (Group 1 & 2): React by losing electrons.
- Trend: Reactivity increases down the group.
- Why: Larger radius + more shielding = Lower Ionization Energy = Easier to lose electrons.
Non-Metals (Group 17): React by gaining electrons.
- Trend: Reactivity decreases down the group.
- Why: Larger radius + more shielding = Lower Electronegativity = Harder to attract electrons.

5. Bonding & Material Properties

This is the core of WACE short answers. You must use the term “Electrostatic Attraction” in every definition.

Metallic Bonding

Definition: The strong Electrostatic Attraction between a lattice of positive metal cations and a sea of delocalized valence electrons.
Properties:
- Conducts Electricity: Delocalized electrons are free to move and carry charge.
- Malleable/Ductile (A+ Explanation): Bonding is non-directional. When force is applied, metal ions slide over each other, but the sea of electrons moves with them, maintaining Electrostatic Attraction so the lattice does not break.
- High Melting Point: Strong forces between cations and the electron sea require high energy to break.

Ionic Bonding

Definition: The Electrostatic Attraction between positive metal ions (cations) and negative non-metal ions (anions) arranged in a 3D crystal lattice. Formed by electron transfer.
Properties:
- Brittle (A+ Explanation): When force is applied, ions of like charge (e.g., + and +) align adjacent to each other. The resultant electrostatic repulsion causes the lattice to shatter.
- Conductivity: Does not conduct as solid (ions fixed). Conducts when molten or aqueous because ions become free-moving carriers of charge.
- High MP/BP: Strong Electrostatic Attraction between oppositely charged ions requires large energy to overcome.

Covalent Bonding

Definition: Electrostatic Attraction between positive nuclei of non-metals and their shared pairs of electrons.

Type 1: Covalent Molecular

Discrete molecules with strong intramolecular bonds but weak intermolecular forces.
Properties: Low MP/BP (easy to break weak intermolecular forces), soft, non-conductors (no free charged particles).

Type 2: Covalent Network (e.g., Diamond, Graphite, Silicon Dioxide)

Giant 3D lattices where atoms are continuously bonded by strong covalent bonds.
Properties: Extremely high MP/BP (requires breaking actual covalent bonds), very hard.
Exception: Graphite conducts electricity due to delocalized electrons between layers.

6. Essential Skills Checklist

Guaranteed questions based on the revision booklet.

Naming & Formulas

Ionic: Balance the charges.
- Ex: Aluminum Oxide. Al³⁺, O²⁻. Swap numbers → Al₂O₃.
Covalent: Use prefixes (mono, di, tri, tetra).
- Ex: Dinitrogen Pentoxide → N₂O₅.

Lewis Dot Diagrams

Show valence electrons only.
Ionic: Put brackets around the anion with its charge; show transfer.
Covalent: Show shared pairs (lines or dots) satisfying the Octet Rule (8 electrons).

Periodic Table Geography

Group 1: Alkali Metals (Reactive, +1 ions).
Group 2: Alkaline Earth Metals (+2 ions).
Group 17: Halogens (Diatomic, −1 ions, very reactive).
Group 18: Noble Gases (Inert, full octet).
Transition Metals: Groups 3–12.

Valence & The Octet Rule

Atoms lose, gain, or share electrons to achieve a full valence shell (usually 8 electrons), mimicking the stability of Noble Gases.

Metals lose electrons → Cations (+).
Non-metals gain electrons → Anions (−).
Isoelectronic Species: Ions that have the same number of electrons. To find the electron count in ions: Protons − Charge.

Additional Quick Facts (Highlight)

Electron charge: −1.602×10⁻¹⁹ C; Proton charge: +1.602×10⁻¹⁹ C.
Masses (approx): Proton ≈ 1.007 u, Neutron ≈ 1.009 u, Electron ≈ 0.00055 u.
Unit reminder: Atomic mass unit = unified atomic mass unit (u); Ar has no units.
When answering exam questions, always: define using key terms (e.g., “Electrostatic Attraction”), show working for calculations, and label units.