Asal chemistry cheat sheet
1.1 Particles in the atom and atomic radius
1. Understand that atoms are mostly empty space surrounding a very small, dense nucleus that contains
protons and neutrons; electrons are found in shells in the empty space around the nucleus.
Atoms consist of a central nucleus (tiny and dense) with protons and neutrons, while electrons orbit in energy
levels or shells. The space between the nucleus and the electrons is mostly empty.
2. Identify and describe protons, neutrons, and electrons in terms of their relative charges and relative
masses.
| Particle | Relative Charge | Relative Mass |
|------------|------------------|----------------| | Proton | +1 | 1 | | Neutron | 0 | 1 | | Electron | -1 | ~0.0005 (≈1/1836 of
proton) |
3. Understand the terms atomic and proton number; mass and nucleon number.
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Proton number (atomic number, Z): Number of protons in an atom.
Nucleon number (mass number, A): Total number of protons and neutrons.
4. Describe the distribution of mass and charge within an atom.
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Almost all mass is in the nucleus (protons + neutrons).
Protons give the nucleus a positive charge.
Electrons orbit and balance the charge, but contribute almost no mass.
5. Describe the behaviour of beams of protons, neutrons, and electrons moving at the same velocity in an
electric field.
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Protons: Deflected towards the negative plate (positive charge).
Electrons: Deflected towards the positive plate (negative charge, more deflection due to
smaller mass).
Neutrons: No deflection (neutral).
6. Determine the numbers of protons, neutrons, and electrons present in both atoms and ions given
atomic or proton number, mass or nucleon number, and charge.
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Protons = atomic number
Neutrons = mass number – atomic number
Electrons = protons (if neutral); adjust for charge (e.g. lose 1 electron for +1 ion)
7. State and explain qualitatively the variations in atomic radius and ionic radius across a period and
down a group.
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Across a period: Atomic radius decreases (more protons = stronger pull on electrons).
Down a group: Atomic radius increases (more shells = larger atoms).
Ionic radius: Cations are smaller (lose electrons = smaller size); anions are larger (gain
electrons = more repulsion).
1.2 Isotopes
1. Define the term isotope in terms of numbers of protons and neutrons.
Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons.
2. Understand the notation ᵃZ X for isotopes, where ᵃ is the mass or nucleon number and Z is the atomic
or proton number.
Example:
₆¹²C → Carbon with atomic number 6 and mass number 12
→ 6 protons, 6 neutrons
3. State that and explain why isotopes of the same element have the same chemical properties.
Chemical properties depend on electrons, not neutrons. Since isotopes have the same number of electrons, they
react the same way.
4. State that and explain why isotopes of the same element have different physical properties, limited to
mass and density.
Isotopes have different masses (due to different numbers of neutrons), which affects density and mass-related
properties like diffusion rate or boiling point.
1.3 Electrons, Energy Levels and Atomic Orbitals
1. Understand the terms: shells, sub-shells and orbitals; principal quantum number (n); ground state,
limited to electronic configuration
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Shells are energy levels around the nucleus, labeled n = 1, 2, 3...
Sub-shells are divisions within shells: s, p, d, f.
Orbitals are regions in sub-shells where electrons are found. One orbital holds 2 electrons.
The principal quantum number (n) indicates the shell's size and energy level.
Ground state refers to the lowest-energy arrangement of electrons.
2. Describe the number of orbitals making up s, p and d sub-shells, and the number of electrons that can
occupy the orbitals and sub-shells
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s sub-shell: 1 orbital = 2 electrons
p sub-shell: 3 orbitals = 6 electrons
d sub-shell: 5 orbitals = 10 electrons
3. Describe the order of increasing energy of the sub-shells within the first three shells and the 4s and 4p
sub-shells
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Order of energy: 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p
4. Describe the electronic configurations to include the number of electrons in each shell, sub-shell and
orbital
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Example (Ca, Z = 20): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s²
This shows the full arrangement of electrons in shells and sub-shells.
5. Explain the electronic configurations in terms of energy of the electrons and inter-electron repulsion
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Electrons fill orbitals from lowest to highest energy.
Within a sub-shell, orbitals are filled singly first (Hund’s Rule), because electrons repel each
other.
6. Determine the electronic configuration of atoms and ions given the atomic or proton number and
charge, using either of the following conventions:
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Full configuration: Fe = 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶ 4s²
Shorthand: Fe = [Ar] 3d⁶ 4s²
For ions: Fe²⁺ = [Ar] 3d⁶ (electrons removed from 4s first)
7. Understand and use the electrons in boxes notation, e.g. for Fe: [Ar] ⬜⬜⬜⬜⬜ (each box represents
an orbital, arrows represent electrons)
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Electrons fill each orbital singly before pairing:
For nitrogen (Z = 7): 1s ↑↓, 2s ↑↓, 2p ↑ ↑ ↑
8. Describe and sketch the shapes of s and p orbitals
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s orbital: spherical
p orbitals: dumbbell-shaped, arranged along x, y, and z axes
9. Describe a free radical as a species with one or more unpaired electrons
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Example: Cl• has 7 valence electrons, with one unpaired. Free radicals are highly reactive.
1.4 Ionisation Energy
1. Define and use the term first ionisation energy, IE
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First ionisation energy is the energy required to remove 1 mole of electrons from 1 mole of
gaseous atoms to form 1 mole of 1⁺ ions.
Example: Na(g) → Na⁺(g) + e⁻
2. Construct equations to represent first, second and subsequent ionisation energies
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1st IE: Mg(g) → Mg⁺(g) + e⁻
2nd IE: Mg⁺(g) → Mg²⁺(g) + e⁻
3rd IE: Mg²⁺(g) → Mg³⁺(g) + e⁻
3. Identify and explain the trends in ionisation energies across a period and down a group of the Periodic
Table
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Across a period: IE increases due to higher nuclear charge and smaller atomic radius.
Down a group: IE decreases due to more shielding and larger atomic radius.
4. Identify and explain the variation in successive ionisation energies of an element
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Each successive electron is harder to remove.
A large jump indicates that a new shell (closer to the nucleus) is being accessed.
5. Understand that ionisation energies are due to the attraction between the nucleus and the outer
electron
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Greater attraction = higher ionisation energy. Influenced by nuclear charge, distance from
nucleus, and electron shielding.
6. Explain the factors influencing the ionisation energies of elements in terms of nuclear charge,
atomic/ionic radius, shielding by inner shells and sub-shells, and spin–pair repulsion
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Higher nuclear charge increases IE.
Larger atomic radius and more shielding decrease IE.
Spin-pair repulsion slightly reduces IE when electrons are paired in the same orbital.
7. Deduce the electronic configurations of elements using successive ionisation energy data
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A large increase in IE shows when an inner shell begins.
Number of low IEs = number of outer (valence) electrons.
8. Deduce the position of an element in the Periodic Table using successive ionisation energy data
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From the number of valence electrons (before the jump in IE), you can find the group number.
Total electrons = atomic number = position in the period.
2.1 Relative Masses of Atoms and Molecules
1. Define the unified atomic mass unit as one twelfth of the mass of a carbon-12 atom
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The unified atomic mass unit (u) is a standard unit for atomic mass.
1 u = 1/12 the mass of a carbon-12 atom ≈ 1.66 × 10⁻²⁷ kg
2. Define relative atomic mass, Aᵣ; relative isotopic mass; relative molecular mass, Mᵣ; and relative
formula mass in terms of the unified atomic mass unit
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Relative atomic mass (Aᵣ): Weighted average of isotopic masses of an element compared to
1 u.
Relative isotopic mass: Mass of a specific isotope compared to 1 u.
Relative molecular mass (Mᵣ): Sum of Aᵣ values of all atoms in a molecule.
Relative formula mass: Same as Mᵣ, but used for ionic compounds and giant structures.
2.2 The Mole and the Avogadro Constant
1. Define and use the term mole in terms of the Avogadro constant
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A mole contains 6.022 × 10²³ particles (Avogadro's constant).
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1 mole of atoms/molecules/ions = Avogadro's number of them.
Example: 1 mol of H₂ O has 6.022 × 10²³ molecules.
2.3 Formulas
1. Write formulas of ionic compounds from ionic charges and oxidation numbers (shown by a Roman
numeral), including:
(a) The prediction of ionic charge from the position of an element in the Periodic Table
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Group 1 = +1, Group 2 = +2, Group 17 = –1, etc.
Transition metals vary; e.g., Fe²⁺ or Fe³⁺ (shown as iron(II), iron(III))
(b) Recall of the names and formulas for the following ions:
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NO₃ ⁻: Nitrate
CO₃ ²⁻: Carbonate
SO₄ ²⁻: Sulfate
OH⁻: Hydroxide
NH₄ ⁺: Ammonium
Zn²⁺: Zinc
Ag⁺: Silver
HCO₃ ⁻: Hydrogencarbonate
PO₄ ³⁻: Phosphate
2. (a) Write and construct equations (which should be balanced), including ionic equations (which should
not include spectator ions)
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Example ionic equation:
Ag⁺(aq) + Cl⁻(aq) → AgCl(s)
(b) Use appropriate state symbols in equations
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(s) = solid, (l) = liquid, (g) = gas, (aq) = aqueous solution
3. Define and use the terms empirical and molecular formula
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Empirical formula: Simplest whole-number ratio of atoms (e.g. CH₂ O)
Molecular formula: Actual number of atoms in a molecule (e.g. C₆ H₁ ₂ O₆ )
4. Understand and use the terms anhydrous, hydrated and water of crystallisation
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Anhydrous: No water present
Hydrated: Contains water molecules
Water of crystallisation: Water chemically bound in a crystal (e.g. CuSO₄ ·5H₂ O)
5. Calculate empirical and molecular formulas using given data
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Use % composition or masses
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Steps:
1. Convert mass to moles
2. Divide by smallest number of moles
3. Get whole-number ratio
4. For molecular formula: divide relative molecular mass by empirical formula mass