**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 |

1 | 1 (s) | 2 | s → 2 |

2 | 2 (s, p) | 8 | p → 6 |

3 | 3 (s, p, d) | 18 | d → 10 |

4 | 4 (s, p, d, f) | 32 | f → 14 |

… | … | … |

**How can we write the electron configuration of an atom?**

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

Format:

3s^{2};

→ 3 = level

→ s = sublevel

→ 2 = number of e^{–}

_{11}Na: 1s^{2}2s^{2}2p^{6}3s^{1 } 11 – 2 = 9 – 2 = 7- 6 = 1 – 1 = 0

_{17}Cl: 1s^{2}2s^{2}2p^{6}3s^{2}3p^{5}

_{9}F: 1s^{2}2s^{2}2p^{5}

_{18}Ar: 1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}

_{21}Sc: 1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{1}

**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 → P
*x*, P*y*, P*z*- the orbitals of the p sub-level are spaced with an angle of 90° between each other.

- 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: 1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{1}3d^{10}

Cu^{+}: 1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}3d^{10}

Cu^{2+}: 1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}3d^{9}

Fe: 1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}4s^{2}3d^{6}

Fe^{2+}: 1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}3d^{6}

Fe^{3+}:1s^{2}2s^{2}2p^{6}3s^{2}3p^{6}3d^{5}

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:

- 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. / - 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.

- If the electron is removed from the new sub-level, there would be a small drop in the ionisation energy.

Mg: 1s^{2}2s^{2}2p^{6}3s^{2}Al: 1s^{2}2s^{2}2p^{6}3s^{2}3p^{1}

Aluminium has a lower ionisation energy than than Mg. - 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: 1s^{2}2s^{2}2p^{3}O: 1s^{2}2s^{2}2p^{4}

Oxygen has a lower ionisation energy than nitrogen.