MYP Chemistry 1.5 Electronic Configuration and Valency Notes

Electrons Occupy Energy Levels in Atoms

1. Electrons in atoms do not move randomly anywhere around the nucleus.
2. They occupy specific energy levels, also called shells or principal energy levels, labelled with numbers: $$n=1,2,3,4, \ldots$$
3. These shells are arranged around the nucleus of an atom.
   1. Electrons in shells closer to the nucleus (smaller nn) have lower energy.
   2. Electrons in shells further from the nucleus (larger nn) have higher energy.

Maximum Number of Electrons in a Shell

Each shell can hold a maximum number of electrons given (in theory) by: $$ \text { Maximum electrons in shell } n=2 n^2$$

Orbitals Are Regions That Describe Electron Positions

Each energy level (shell) is divided into subshells (s, p, d, f), and each subshell contains one or more orbitals.

1. We often show orbitals in diagrams, but they are not fixed paths like planets around the Sun.
2. They are probability clouds showing where an electron is likely to be found.

Electrons and Spin

1. Each orbital can hold a maximum of two electrons.
2. When two electrons share an orbital, they must have opposite spins.
3. This is a consequence of the Pauli exclusion principle: no two electrons in an atom can have exactly the same set of quantum numbers.

Types of Orbitals

Orbitals come in different shapes, linked to different subshells:

1. s orbitals
   1. Shape: spherical
   2. Present in every energy level (1s, 2s, 3s, …)
2. p orbitals
   1. Shape: dumbbell-shaped
   2. First appear from the second energy level (2p, 3p, …)
   3. There are three p orbitals in each p subshell, oriented along the x, y, and z axes.

Subshells, Orbitals and Electron Capacity

1. Subshells are labelled s, p, d and f.
2. Each type of subshell contains a fixed number of orbitals, according to the table below.

Electronic Configuration: Distributing Electrons

1. The electronic configuration of an atom or ion shows how its electrons are arranged in shells, subshells and orbitals.
2. It tells us how many electrons are in each part of the atom.
3. The electrons in the outermost shell (valence electrons) largely determine the element’s chemical properties and reactivity.

We usually write electron configurations using a notation such as: $$1 s^2 2 s^2 2 p^6 3 s^2 3 p^4$$

1. The number (1, 2, 3, …) is the energy level (shell).
2. The letter (s, p, d, f) is the subshell.
3. The superscript (², ⁶, ⁴, …) is the number of electrons in that subshell.

How Electrons Fill Orbitals in Atoms and Ions

To write electron configurations for atoms and ions, we need to know how electrons fill orbitals.

The filling pattern follows three main rules:

1. Aufbau principle – electrons fill the lowest energy orbitals first.
2. Hund’s rule – electrons occupy orbitals singly before pairing.
3. Pauli exclusion principle – no two electrons in an atom can have the same set of quantum numbers (in one orbital they must have opposite spins).

Aufbau Principle – Lowest Energy First

The Aufbau principle states that:

Electrons fill the lowest available energy orbitals first, then move to higher energy orbitals.

Hund’s Rule – Spread Out Before Pairing

Hund’s rule states that:

Electrons fill degenerate (same-energy) orbitals one at a time with parallel spins before any orbital gets a second electron.

This reduces electron–electron repulsion and gives a more stable arrangement.

Pauli Exclusion Principle – Opposite Spins in One Orbital

The Pauli exclusion principle states that:

No two electrons in the same atom can have the same set of four quantum numbers.

In simpler terms for MYP Chemistry:

1. Each orbital can hold a maximum of two electrons.
2. When two electrons share an orbital, they must have opposite spins.

Valency and Electron Configuration

1. Valency is the combining power of an element.
2. For many main-group elements, valency is related to the number of electrons in the outermost shell (valence electrons).
   1. Metals in Group 1 usually lose 1 electron → valency 1.
   2. Metals in Group 2 usually lose 2 electrons → valency 2.
   3. Non-metals in Group 17 usually gain 1 electron → valency 1 (for forming −1 ions).
   4. Non-metals in Group 16 usually gain 2 electrons → valency 2, etc.

Transition Metals and Exceptions

Some transition metals have slightly different (more stable) electron configurations than you would predict from a simple Aufbau pattern.

Common Exceptions

1. Chromium (Cr)
   1. Expected: $[A r] 4 s^2 3 d^4$
   2. Actual: $[A r] 4 s^1 3 d^5$ (half-filled 3d subshell is more stable)
2. Copper (Cu)
   1. Expected: $[A r] 4 s^2 3 d^9$
   2. Actual: $[A r] 4 s^1 3 d^{10}$ (fully filled 3d subshell is more stable)

Which Electrons Are Lost First?

When transition metals form ions, the 4s electrons are lost before the 3d electrons, even though 4s filled first in the neutral atom.

Charges of Ions and Electron Configuration

The electron configuration explains how ions form:

1. Cations are formed when atoms lose electrons
   1. e.g. $\mathrm{Na} \rightarrow \mathrm{Na}^{+}+e^{-}$
2. Anions are formed when atoms gain electrons
   1. e.g. $\mathrm{Cl} +e^{-} \rightarrow \mathrm{Cl}^{-}$

How Does Valency Determine the Formula of Compounds?

1. Valency tells us how many electrons an atom loses, gains or shares when it forms a compound.
   1. Metals usually lose electrons → form positive ions (cations).
   2. Non-metals usually gain electrons → form negative ions (anions).
2. To form a neutral compound, the total positive charge must equal the total negative charge.
3. So valency helps us predict the ratio of atoms in the formula of an ionic compound.