Electron Configuration and Subshells

Energy Levels (n): Represented by the principal quantum number $n$ , where $n = 1, 2, 3,…$ . Higher $n$ values indicate higher energy levels. Analogous to lifting a book higher requiring more energy.
Subshells: Each energy level contains subshells, denoted by letters (s, p, d, f). The number of subshells increases with the energy level.
- $n = 1$ : s
- $n = 2$ : s, p
- $n = 3$ : s, p, d
- $n \geq 4$ : s, p, d, f
Subshells have specific shapes, for example, an s subshell is spherical.

The first energy level (Shell 1) has only one subshell (1s).
The second energy level (Shell 2) has two subshells (2s and 2p).
The analogy of a parking deck is used: electrons fill the lowest energy levels (floors) first. If the first level is full, electrons occupy higher levels.
The energy difference between main energy levels is larger than the difference between subshells.
The 2p subshell consists of three orbitals, each shaped like a figure eight.

Periods: Correspond to the highest energy level. The current periodic table ends at period 7 (element 118). Hypothetically, if an element beyond 118 is created, it would begin period 8.
Blocks:
- Columns 1A and 2A fill the s orbital (s-block).

Electron Configuration: Describes the arrangement of electrons within an atom's energy levels and subshells.
Example: Lithium (Li)
- Lithium has 3 protons and 3 electrons.
- The first two electrons fill the 1s orbital (1s²).
- The remaining electron occupies the 2s orbital (2s¹).
- The complete electron configuration is 1s²2s¹.
Example: Boron (B)
- Boron has 5 electrons.
- Electron configuration: 1s²2s²2p¹ (requires three p orbitals)
Example: Carbon (C)
- Carbon has 6 electrons.
- Electron configuration: 1s²2s²2p²
Example: Nitrogen (N)
- Nitrogen has 7 electrons.
- Electron configuration: 1s²2s²2p³
Example: Oxygen (O)
- Oxygen has 8 electrons.
- Electron configuration: 1s²2s²2p⁴
Example: Fluorine (F)
- Fluorine has 9 electrons.
- Electron configuration: 1s²2s²2p⁵
Example: Neon (Ne)
- Neon has 10 electrons.
- Electron configuration: 1s²2s²2p⁶
- Neon is inert because it has a full outer energy level.

Example: Sodium (Na)
- Electron configuration: 1s²2s²2p⁶3s¹
Example: Magnesium (Mg)
- Magnesium has 12 electrons.
- Electron configuration: 1s²2s²2p⁶3s²
Example: Aluminum (Al) (or Aluminium)
- Aluminum has 13 electrons.
- Electron configuration: 1s²2s²2p⁶3s²3p¹
Example: Silicon (Si)
- Electron configuration: 1s²2s²2p⁶3s²3p²

After filling the 3p subshell, the 4s subshell fills before the 3d subshell.
The 4s subshell is lower in energy than the 3d subshell due to interelectronic repulsion and effective nuclear charge.
This is a key exception to the filling order that needs to be remembered. (Think of it as someone cutting in front of you in line.)
The periodic table's structure reflects this filling order: s-block, p-block, d-block.

The number of elements in each block corresponds to the number of electrons that can occupy the subshells:
- s-block: 2 elements
- p-block: 6 elements
- d-block: 10 elements (requires 5 orbitals)
Two n squared can be used to calculate the number of electrons at each energy level. $(2n^2)$

Noble gases have full outer electron shells, making them generally unreactive.
Helium (He): 1s²
Neon (Ne): 2s²2p⁶
Argon (Ar): 3s²3p⁶
Krypton (Kr): 4s²4p⁶
Elements with similar outer electron configurations exhibit similar chemical behavior.

Electrons behave as spinning, electrically charged particles, creating magnetic fields.
These magnetic fields have north and south poles. Electrons want to occupy their own space due to repulsion between negative charges.

Electrons will individually occupy each orbital within a subshell before doubling up in any one orbital. This minimizes electron repulsion.
Example: Potassium (K) has the configuration 4s¹ and Calcium (Ca) has 4s².

The d-block elements (Scandium to Zinc and below) are called transition elements.
The f-block elements are called inner transition elements.