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Electron Configuration and Valency

Have you ever wondered why some elements react extremely aggressively, and others don't seem to react at all? Or why some elements can form different compounds and others can only form one?

It all has to do with what is going on in the center of their atoms, with their electrons.

Electrons are far more than just tiny particles that orbit the atoms. They are the drivers of chemical behaviour and chemical bonding. Electron configuration and valency are concepts that offer insight into how different elements will chemically react, and are the bedrock of Chemistry.

What Are Electrons?

Electrons are negatively charged particles that orbit around the nucleus of an atom. They are much less massive than neutrons and protons.

An atom's electrons determine most of the atom's properties.

You can relate electrons to a series of tenants that live in an apartment building (the atom).

Every level of the building is an energy level (or shell), and within each level, there are several suites (or orbitals) where the electrons are located.

Energy Levels and Sublevels

Electrons occupy regions around the nucleus called energy levels (or shells), labeled K, L, M, N… or numerically 1, 2, 3, 4….

• 1st shell (K): can hold up to 2 electrons
• 2nd shell (L): can hold up to 8 electrons
• 3rd shell (M): can hold up to 18 electrons
• 4th shell (N): can hold up to 32 electrons

Each energy level contains sublevels (types of orbitals) labeled s, p, d, f:

• Electrons fill the lowest energy orbitals first; this is the Aufbau principle.
• Each orbital can hold 2 electrons, with opposite spins (Pauli Exclusion Principle)
• Electrons prefer to occupy empty orbitals first before pairing up. (Hund's Rule)

Electron Configurations

Electron configuration is the listing of electrons in their orbitals.

Example: Oxygen (Atomic number 8)

Energy levels to fill: 1, 2

Electron filling: 1s² 2s² 2p⁴

Explanation:

• 1st shell: 1s² → 2 electrons
• 2nd shell: 2s² → 2 electrons
• 2nd shell: 2p⁴ → 4 electrons

So, the configuration is 1s² 2s² 2p⁴.

Another Example: Sodium (Atomic number 11)

Electron configuration: 1s² 2s² 2p⁶ 3s¹

Notice how the outermost electron (3s¹) is responsible for chemical reactions.

Valence Electrons

Valence electrons are the electrons located at the outer shell of an atom.

They are important because they take part in the formation of chemical bonds.

Examples:

• Sodium (Na): has 1 valence electron → highly reactive, loses 1 electron to become Na⁺.
• Chlorine (Cl): has 7 valence electrons → gains 1 electron to become Cl⁻.
• Oxygen (O): has 6 valence electrons → gains 2 electrons to complete the octet.

When an atom has closer to 8 valence electrons, it becomes more stable. This is known as the Octet Rule.

Determining Valency

Valency refers to the combining capacity of an element. It determines how many electrons an atom can lose, gain, or share to achieve a stable configuration.

• An atom that has 1–4 valence electrons will tend to lose electrons → valency = number of electrons lost
• An atom that has 5–7 valence electrons will tend to gain electrons → valency = 8 − number of valence electrons
• Having 8 valence electrons means you are a noble gas → valency = 0 (inert, stable)

Examples:

• Sodium (Na, 1 valence electron): valency = 1 (loses 1 electron)
• Oxygen (O, 6 valence electrons): valency = 2 (gains 2 electrons)
• Chlorine (Cl, 7 valence electrons): valency = 1 (gains 1 electron)

Electron Configuration and the Periodic Table

Elements in a group of the Periodic Table have the same number of valence electrons. Elements in the same period have the same number of electron shells.

For example:

• Group 1 (Alkali metals): Li → 2s¹, Na → 3s¹, K → 4s¹ → all have 1 valence electron
• Group 17 (Halogens): F → 2s² 2p⁵, Cl → 3s² 3p⁵ → all have 7 valence electrons

This is why elements in the same group tend to have similar chemical properties.

Rules to Remember

• Aufbau Principle: Whichever energy level is the lowest will fill first.
• Pauli Exclusion Principle: Each orbital can hold a maximum of 2 electrons.
• Hund's Rule: All the lowest energy unfilled orbitals will gain 1 electron before any pair up.
• Octet Rule: Atoms are only stable if they have 8 valence electrons. The only exceptions are hydrogen and helium.

Orbital Diagrams

Electrons can also be shown using boxes or arrows to represent orbitals.

Each box represents an orbital. Each arrow represents an electron.

If they are up and down arrows, they have opposite spins.

Example: Oxygen (O, 1s² 2s² 2p⁴)

1s: ↑↓

2s: ↑↓

2p: ↑↓ ↑ ↑

This visual representation helps determine which electrons are paired or unpaired. This is important and is used for bonding.

Connection Between Electron Configuration and Chemical Behavior

The number of valence electrons determines how an atom bonds (ionic or covalent), its reactivity, and the types of compounds it can form.

Ionic Bonding:

This is created when electrons are transferred.

Example: Sodium (1 valence electron) + Chlorine (7 valence electrons) → NaCl

Covalent Bonding:

This is created when electrons are shared.

Example: Oxygen (6 valence electrons) shares electrons with hydrogen to form H₂O

Transition Elements and Valency

Transition metals (Fe, Cu, Zn) can have more than one valency. This is because its d orbitals can participate in bonding.

Example: Iron → Fe²⁺ or Fe³⁺

This is the reason why transition metals can form multiple compounds with different properties.

Why Electron Configuration is Important

• Predicts chemical reactions & bonding patterns.
• Describes trends of the period table in:
  • Reactivity
  • Ionization energy
  • Electronegativity
• Essential for modern chemistry, including organic chemistry & materials science.