LIFEF002 Atomic and Electronic Structure

Course Overview

Course: LIFEF002 Atomic and Electronic Structure 2
Learning Outcomes:
- Explain trends in ionisation energy across a period.
- Determine the electron configuration of elements using s, p, d, and f notation.

Electronic Structure: Spectral Lines

Key Observations:
- Spectral lines are fixed, indicating that electrons exist only in certain fixed energy levels.
- As energy levels get closer together, the spectral lines also get closer.
- The frequency of emitted light during an electron's transition from a higher to a lower shell depends on the energy difference between the two levels.

The Bohr Model

Principle Energy Levels:
- Seven principal energy levels are identified: n = 1 to 7.
- The relationship between the number of energy levels and the number of periods in the Periodic Table is relevant.

Successive Ionisation Energies

Definition:
- Successive ionisation energies measure the energy required to remove electrons sequentially from an atom.
Process:
- Electrons are removed one by one.
- Successive ionisation energies provide critical evidence for understanding electron arrangements.

First and Second Ionisation Energies

First Ionisation Energy (IE1):
- Energy needed to remove one electron from each atom in a mole of gaseous atoms:
- Reaction: $X(g) \rightarrow X^+(g) + e^-$
- Enthalpy change (DH): Positive value.
Second Ionisation Energy (IE2):
- Energy needed to remove one electron from each ion in a mole of gaseous 1+ ions:
- Reaction: $X^+(g) \rightarrow X^{2+}(g) + e^-$
- Enthalpy change (DH): Positive value.

Example: Ionisation Energies of Oxygen

Atomic Structure of Oxygen:
- 8 neutrons, 8 protons, 8 electrons.
1st Ionisation Energy:
- From O to O+: 8 neutrons, 8 protons, 7 electrons.
2nd Ionisation Energy:
- From O+ to O2+: 8 neutrons, 8 protons, 6 electrons.
Nucleus Composition:
- No change in nucleus throughout ionisation steps.

Ionisation Energies of Oxygen

Representation of Energies:
- $O(g) \rightarrow O^+(g) + e^-$ ;; DH = +1314 kJ mol^-1
- $O^+(g) \rightarrow O^{2+}(g) + e^-$ ;; DH = +3388 kJ mol^-1

Data on Successive Ionisation Energies of Oxygen

Ionisation energies listed (kJ mol^-1):
- 1st: 1314
- 2nd: 3388
- 3rd: 5300
- 4th: 7469
- 5th: 10990
- 6th: 13327
- 7th: 71330
- 8th: 84078

Evidence for Electron Shells

Ionisation and Electron Removal:
- As electrons are removed, repulsion between remaining electrons decreases, pulling them closer to the nucleus, reducing atomic radius.
Energy Considerations:
- Electrons closer to the nucleus have stronger attraction and require more energy to remove.
- Large jumps in energy suggest the removal of an electron from a different shell closer to nuclear charge.

Ionisation Energy Trends

Down a Group

Trend Observation:
- Overall decrease in ionisation energy as one moves down a group, consistent with the Bohr model of electron shells.

Across a Period

Trend Observation:
- General increase in ionisation energy across a period due to:
- Increase in number of protons increases positive charge.
- Decrease in atomic radius strengthens attraction between nucleus and electrons.
- Shielding effect remains constant across periods.

More Complex Electronic Structure

Anomalies in Ionisation Energies:
- Trends do not apply universally across all elements, especially within transition metals.

Subshells and Atomic Orbitals

Energy Levels:
- Each principal energy level (n) is not strictly defined in energy value.
- Each -level consists of sub-shells: s, p, d, and f.
- Each sub-shell contains a set of orbitals that can each hold two electrons.

Orbital Structure

Principal Quantum Number (n) Correspondence:
- n = 1: 1 s-subshell (1 orbital)
- n = 2: 1 s-subshell (1 orbital) + 1 p-subshell (3 orbitals)
- n = 3: 1 s-subshell + 1 p-subshell + 1 d-subshell (5 orbitals)
- n = 4: 1 s-subshell + 1 p-subshell + 1 d-subshell + 1 f-subshell (7 orbitals)

Orbital Shapes

Orbital Characteristics:
- s: Spherical shape
- p: Dumbbell shape along x, y, z axes
- d and f: More complex geometries not detailed here.

Electron Filling Rules

AUFBAU Principle:
- Electrons are added one at a time starting from the lowest available energy.
- Each energy level must be fully occupied before moving to the next.
HUND'S Rule:
- Each orbital in a sub-shell is singly occupied before any orbital starts to pair up.
PAULI Exclusion Principle:
- An orbital can hold a maximum of two electrons with opposite spins.

Examples of Electron Configurations

Configuration for Helium:
- He: 2 electrons, Configuration = 1s²
Configuration for Beryllium:
- Be: 4 electrons; Configuration = 1s² 2s²
Configuration for Nitrogen:
- N: 7 electrons; Configuration = 1s² 2s² 2p³

Summary of Electron Configurations

Recorded Configurations for Common Elements:
- He: 1s²
- Be: 1s² 2s²
- N: 1s² 2s² 2p³

Anomalies in Ionisation Energy

Decrease between Be and B explained:
- Boron’s 5th electron occupies the 2p-subshell, further from the nucleus than 2s electrons in Be, resulting in reduced nuclear attraction despite higher proton count.
Decrease between N and O explained:
- Oxygen has paired electrons in p-orbitals experiencing greater repulsion than singly occupied p-orbitals in nitrogen, requiring less energy for removal of paired electrons.

Transition Metal Exceptions

Notable Exceptions:
- Elements like chromium and copper exhibit preferential filling of 3d subshells for stability, preferring a half-full rather than fully filled configuration.

Concluding Confidence Assessment

Knowledge Self-Test:
- Can you:
- Explain successive ionisation energies?
- Discuss trends in 1st ionisation energy across periods and groups?
- Use spdf notation for electron configurations?
- Write configurations for ions?
- Identify anomalies in transition elements?
- Identify outer electrons based on periodic table position?

Additional Practice and Review

Practice Quizzes:
- Links and QR codes for quizzes are provided for self-assessment and practice.