Learners' AlleyDesigned by Khow Willard

Atomic Structure

Sub-atomic Particles

1.1Protons, neutrons and electrons

Particle	Relative charge	Relative mass
proton	+1	1
neutron	0	1
electron	−1	1/1836

Protons and neutrons reside together in the nucleus, which holds almost all the mass and all the positive charge.
Electrons are found in regions of space around the nucleus, and are attracted electrostatically by the protons.

1.2Proton number and nucleon number

DefinitionNucleon number: total number of protons and neutrons in the nucleus of an atom.

DefinitionProton number: number of protons in the nucleus of an atom. It determines the identity of the element.

DefinitionNuclide: a specific combination of protons and neutrons in a nucleus.

3517Cl

35 nucleon number (mass number) = p + n 17 proton number (atomic number) = p

Atomno. of p = no. of e⁻

Cationloses e⁻: p > e⁻

Aniongains e⁻: e⁻ > p

1.3Isotopes

DefinitionIsotopes are atoms of the same element with the same number of protons but different numbers of neutrons.

³⁵Cl	17 p, 18 n
³⁷Cl	17 p, 20 n

Same electronic configurationSame chemical properties

Different massesDifferent physical properties

Name an isotope with its nucleon number, e.g. ²H. In every stable isotope except ¹H and ³He, neutrons ≥ protons.

1.4Beams of p, n and e⁻ in an electric field

angle of deflection ∝ chargemass

Protons are deflected towards the negatively charged plate.
Electrons are deflected towards the positively charged plate. They are much lighter than protons, so are deflected to a larger extent.
Neutrons have no charge and are undeflected.

Worked exampleA beam of ²⁴¹Am⁺ ions is deflected by +2°. Under the same conditions, find the angle of deflection of ³²S⁻ ions.

θ = k × chargemass Am⁺: +2 = k × +1241 → k = 482 S⁻: θ = 482 × −132 = −15.1°

2Topic 01 · Atomic Structure

Relative Masses of Atoms

2.1Relative isotopic mass

DefinitionRelative isotopic mass: the mass of one mole of atoms of an isotope relative to 1/12 the mass of one mole of ¹²C atoms.

2.2Relative abundance

Relative abundance: the percentage of atoms of each isotope in a naturally occurring sample of the element.

e.g. chlorine:
75% ³⁵Cl, 25% ³⁷Cl

2.3Relative atomic mass, A_r

DefinitionRelative atomic mass: the weighted average mass of one mole of atoms of an element relative to 1/12 the mass of one mole of ¹²C atoms.

A_r = Σ (relative isotopic mass × relative abundance)total abundance (100) Cl: (35 × 75) + (37 × 25)100 = 35.5

Shells, Subshells and Orbitals

3.1Principal quantum shells, subshells and orbitals

DefinitionOrbital: the region of space where there is a high probability (95%) of finding an electron. Each orbital can accommodate 2 electrons.

Atom→principal quantum shells (n)→subshells (s, p, d, f)→orbitals→max. 2 e⁻ each

A larger n gives a larger, more diffused orbital further from the nucleus, and hence a higher energy.

n	Subshell	No. of orbitals	Types of orbitals	Total (n²)	Max e⁻ (2n²)
1	1s	1	1s	1	2
2	2s	1	2s	4	8
2	2p	3	2p_x, 2p_y, 2p_z	4	8
3	3s	1	3s	9	18
	3p	3	3p_x, 3p_y, 3p_z
	3d	5	3d_xy, 3d_xz, 3d_yz, 3d_x²−y², 3d_z²
4	4s 4p 4d (4f)	1, 3, 5	4f not in syllabus	16	32

3.2Relative energies of the orbitals

Within a shell: s < p < d < f, as an s electron is more likely to be near the nucleus than a p, and a p than a d.
Levels get closer as n increases. First overlap: 3d and 4s (4s fills first).

3.3Shapes of orbitals

p_x

p_y

p_z

d_xy

d_xz

d_yz

d_x²−y²

d_z²

s spherical, p dumb-bell. Orbitals within a subshell have the same energy: degenerate.
d_x²−y², d_z²: along the axes. d_xy, d_xz, d_yz: between the axes.

Topic 01 · Atomic Structure3

Learners' AlleyDesigned by Khow Willard

Electronic Configuration

4.1Rules for filling orbitals

1. Aufbau principleElectrons fill the orbitals of lowest energy first. After 3p, 4s is filled before 3d.

2. Pauli exclusion principleAn orbital holds a maximum of two electrons, of opposite spins.

3. Hund's ruleOrbitals of a subshell are filled singly first, with parallel spins, before pairing. This gives less inter-electronic repulsion.

✓ ✗

Electrons-in-boxes: label the boxes with the subshell only (2p, not 2p⁶). Ground state: all e⁻ in the lowest energy levels. Excited state: an e⁻ promoted, e.g. O: 1s² 2s² 2p³ 3s¹.

4.2Writing configurations of atoms

From the Periodic Table

Period = n of the valence shell. A new period starts when e⁻ enter a new principal quantum shell
Block: the subshell filled last. The d-block is one shell behind: 3d in Period 4
Position in the block: the number of e⁻ in that subshell

e.g. Mn: Period 4, 5th in the d-block → [Ar] 3d⁵ 4s²

e.g. Se: Period 4, 4th in the p-block → [Ar] 3d¹⁰ 4s² 4p⁴

Elements in the same Group have similar valence shell configurations.

Noble gas configurationInner electrons shown as the previous noble gas: Li: [He] 2s¹. Exam tipWrite the full configuration in order of increasing n: …3d⁶ 4s².

Anomalies: Cr and Cu

Cr [Ar] 3d⁵ 4s¹ 3d⁴ 4s²

Cu [Ar] 3d¹⁰ 4s¹ 3d⁹ 4s²

A half-filled (d⁵) or fully-filled (d¹⁰) 3d subshell is more stable.

4.3Configurations of ions

Write the configuration of the atom first.
Cation: electrons are removed from the outermost shell of the ground state configuration (4s before 3d).
Anion: electrons are added to the empty or partially-filled valence orbitals.

S²⁻	1s² 2s² 2p⁶ 3s² 3p⁶
Na⁺	1s² 2s² 2p⁶
Fe²⁺	1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁶

Ionisation Energy

DefinitionIonisation energy: the energy required to remove 1 mole of electrons from 1 mole of gaseous atoms or ions.

Definition2nd IE: the enthalpy change when 1 mole of electrons is removed from 1 mole of gaseous, singly charged cations.

X(g) → X⁺(g) + e⁻ΔH = 1st IE of X X⁺(g) → X²⁺(g) + e⁻ΔH = 2nd IE of X X⁽ⁿ⁻¹⁾⁺(g) → Xⁿ⁺(g) + e⁻ΔH = n^th IE of X

All IEs are endothermic. Always show (g) in the equation.

4Topic 01 · Atomic Structure

Ionisation Energy (continued)

5.1Trends in first IE Z_eff = Z − shielding constant

Across periods 2 and 3, 1st IE generally ↑

Nuclear charge ↑, while there is a negligible ↑ in inner shell shielding effect, as e⁻ are added to the same shell
→ effective nuclear charge ↑, atomic radius ↓
→ valence e⁻ more strongly attracted by the nucleus → more energy required to remove them

Down a group, 1st IE ↓

Atomic radius ↑
No. of principal quantum shells ↑ → shielding effect ↑, which largely cancels out the ↑ in nuclear charge
→ valence e⁻ less attracted by the nucleus → less energy required to remove them

First IE of elements 1 to 20 / kJ mol⁻¹

Dip: Be → B also Mg → Al

Be
B

The 2p e⁻ removed from B is higher in energy than the 2s e⁻ removed from Be → less energy required to remove it → 1st IE of B < Be

Dip: N → O also P → S

N
O

Coulombic repulsion between the paired 2p e⁻ in O → easier to remove one of them than an unpaired 2p e⁻ from N → less energy required → 1st IE of O < N

5.2Successive ionisation energies

Answer structureReading a successive IE graph

General trend: successive IEs ↑ as each e⁻ is removed from an ion of increasing positive charge
Gradual increase (1st to 7th IE): removal of e⁻ from the same principal quantum shell
Sharp increase from 7th to 8th IE: 8th e⁻ removed from an inner principal quantum shell, nearer to the nucleus, with less shielding

Note: the small jump from 5th to 6th IE is an e⁻ from a lower energy subshell (2s after 2p) of the same shell.

Worked exampleThe successive IEs of element X are shown below. Predict the Group of X.

IE / kJ mol⁻¹

870

1800

3000

3600

5800

7000

13000

Successive IEs ↑ as each e⁻ is removed from an ion of increasing positive charge
1st to 6th IE ↑ gradually: removal of the outermost e⁻ only
Sharp ↑ from 6th to 7th IE: 7th e⁻ removed from an inner principal quantum shell, nearer to the nucleus, with less shielding
∴ 6 e⁻ in the outermost shell (ns² np⁴) → X is in Group 16. Big jumps + 1 = occupied shells.

Topic 01 · Atomic Structure5