Atomic Structure Standard Level

How are electrons arranged in an atom?

  • electrons are arranged in energy levels, sub-levels and orbitals around the nucleus.
Level (n)Sublevel (n)Max. # of e- in a level (2n²)Max # of e- in a sub-level
11 (s)2s → 2
22 (s, p)8p → 6
33 (s, p, d)18d → 10
44 (s, p, d, f)32f → 14

How can we write the electron configuration of an atom?

You start writing the electron configuration by filing from the first arrow.

 

 

Format:
3s2;
→ 3 = level
→ s = sublevel
→ 2 = number of e

 

 

 

11Na: 1s22s22p63s1
           11 – 2 = 9 – 2 = 7- 6 = 1 – 1 = 0

17Cl: 1s22s22p63s23p5

9F: 1s22s22p5

18Ar: 1s22s22p63s23p6

21Sc: 1s22s22p63s23p64s23d1

 

What are the shapes of the orbitals?

Levels → sub-levels (s, p, d, f) → orbitals

  • Each orbital can hold up to 2 electrons only

The ‘s’ orbital has a spherical shape while the ‘p’ orbital has a dumbbell shape or two balloons tied at the nucleus.

  • s sub-level → 1 orbital (sphere)
  • p sub-level → 3 orbitals → Px, Py, Pz
    • the orbitals of the p sub-level are spaced with an angle of 90° between each other.
  • d sub-level → 5 orbitals
  • f sub-level → 7 orbitals

What are the main principles used to write the electron configuration?

  • Pauli Exclusion Principle:
    Each orbital can hold a maximum of two electron with opposite spins.
  • Aufbau Principle:
    Electrons are placed in orbitals with the lowest energy first.

  • Hund’s Rule
    In the p, d and f orbitals the electrons will occupy different orbitals with parallel spins.

How can we write the electron configuration of the ‘d’ block?

Since the energy differences between the 4s and the 3d sub-levels are very small, the first two electrons should be removed from the 4s before the 3d.

Cu: 1s22s22p63s23p64s13d10
Cu+: 1s22s22p63s23p63d10
Cu2+: 1s22s22p63s23p63d9

Fe: 1s22s22p63s23p64s23d6
Fe2+: 1s22s22p63s23p63d6
Fe3+:1s22s22p63s23p63d5

After all the electrons in the 4s sub-level are removed, you start removing from the 3d sub-level.

 

Define the ionisation energy and write a chemical equation to represent it.

Ionisation energy is the minimum amount of energy needed to remove one mole of electron from one mole of gaseous atom, to form a gaseous ion.

Example:

X(g) → X+(g) + 1e

Fe(g) → Fe+(g) + 1e

 

What are the factors that affect the ionisation energy?

There are two factors that affect the ionisation energy:

  1. Size of the atom:
    As the number of shells increases, outer electrons will be further away from the nucleus and will need less energy to remove the electron; therefore, the ionisation energy will decrease. /
  2. The nuclear charge:
    As the nuclear charge increases, the outer electrons will be attracted more to the nucleus, therefore they need more energy to be removed and the ionisation energy will increase.

How does the ionisation energy change across a period?

There is a general increase in trend in the ionisation energy across a period, but this trend shows some discontinuity.

  1. If the electron is removed from the new sub-level, there would be a small drop in the ionisation energy.
    Mg: 1s22s22p63s2                                              Al: 1s22s22p63s23p1
    Aluminium has a lower ionisation energy than than Mg.
  2. If the electrons removed from a sub-level having an electron pair in one of its orbitals, then ionisation energy will be lower than the atom that has no lone pairs due to electron-electron repulsion that makes the size of the atom bigger and the ionisation energy smaller.
    N: 1s22s22p3                                                      O: 1s22s22p4
    Oxygen has a lower ionisation energy than nitrogen.