Electron Orbitals: 2s and 2p in the Second Shell

The transcript references the s orbital and its electrons in the context of the second energy level (n = 2).
In the second energy level, the s-sublevel is 2s and can hold up to 2 electrons; notation: 2s^2\.
The phrase "first two columns" suggests the s-block of the periodic table, which contains elements whose valence electrons reside in an s orbital.
After the 2s sublevel is filled, the 2p sublevel begins filling within the same energy level (n = 2).
The 2p sublevel contains three degenerate orbitals: 2px, 2py, 2p_z, collectively capable of holding up to 6 electrons; notation in full can be 2p^6 for a filled sublevel.
In the second energy level, the total capacity of the outer shell combining 2s and 2p is $2 + 6 = 8$ electrons.
Example shell completion: for Neon, the second shell is filled with 2s^2 2p^6, contributing to the total electron count along with the first shell.
The electron configuration for Neon: $1s^2\ 2s^2\ 2p^6$ .
The 2p orbitals can start with fewer electrons in partially filled cases (e.g., 2p^1, 2p^2, etc.), depending on the element.

2s is the s-sublevel in the second energy level (n = 2).
Capacity: $2$ electrons (one orbital with two possible spins).
Notation: $2s^2$ when filled.
Location in periodic table: s-block elements (groups 1–2) have valence electrons in s orbitals when in the second period as applicable.

The 2p sublevel appears in the second energy level after 2s is filled.
There are three degenerate orbitals: 2px, 2py, 2p_z.
Each orbital can hold up to 2 electrons, so the full 2p sublevel holds up to $6$ electrons; notation: 2p^6 when the sublevel is full.
Hund’s rule implication: electrons will singly occupy the 2p orbitals before any pairing occurs (e.g., 2p^3 can be represented as 2px^1 2py^1 2p_z^1).

Aufbau principle: electrons fill from the lowest energy levels upward (2s before 2p in the second energy level).
Pauli exclusion principle: each orbital can contain at most 2 electrons, and those two electrons must have opposite spins.
Hund’s rule: for degenerate orbitals (like the three 2p orbitals), electrons occupy separate orbitals with parallel spins before any pairing occurs.

Neon example (fully filled second shell): $1s^2\ 2s^2\ 2p^6$
Nitrogen example (partially filled 2p): $1s^2\ 2s^2\ 2p^3$ with the 2p electrons occupying the three 2p orbitals singly: $2p<em>x^1\ 2p</em>y^1\ 2p_z^1$ .
Sodium example (outer electron in the next shell): $1s^2\ 2s^2\ 2p^6\ 3s^1$ , illustrating the transition from the second to the third energy level.

s-block vs p-block: elements in the first two groups (and second period examples) illustrate valence electrons in s orbitals; the p-block begins after the s-block in the same period.
Chemical properties largely depend on the valence electron configuration, especially the presence and arrangement of 2s and 2p electrons in the second energy level.
Real-world relevance: understanding electron configurations helps explain chemical reactivity, bonding, and periodic trends (ionization energy, electronegativity, etc.).

Maximum electrons in an s subshell (l = 0): $2$
Maximum electrons in a p subshell (l = 1): $6$
General subshell capacity formula: $N_{ ext{subshell}}^{\max} = 2(2l + 1)$
Example applications:
- For s: $N_{ ext{subshell}}^{\max} = 2(2\cdot 0 + 1) = 2$
- For p: $N_{ ext{subshell}}^{\max} = 2(2\cdot 1 + 1) = 6$
Common electron-configuration snippets:
- Neon: $1s^2\ 2s^2\ 2p^6$
- Nitrogen: $1s^2\ 2s^2\ 2p^3$
- Sodium: $1s^2\ 2s^2\ 2p^6\ 3s^1$