Show the distribution of electrons in oxygen atom (atomic number 8) using orbital diagram.

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Q. Show the distribution of electrons in oxygen atom (atomic number 8) using orbital diagram.

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Concept understanding — Effective Nuclear Charge

The Intuition: Why Don't Electrons Just Fly Away?

Imagine you're holding a magnet near a pile of paperclips. The closer the magnet, the stronger the pull. Now imagine you put a sheet of cardboard between the magnet and the paperclips. The pull weakens — the cardboard "shields" the paperclips from the full force of the magnet.

An atom works similarly. The nucleus (positive charge) pulls on the electrons (negative charge). But an electron is not alone — there are other electrons buzzing around between it and the nucleus. Those inner electrons act like the cardboard sheet: they shield or screen the outer electron from feeling the full positive charge of the nucleus.

So an outer electron doesn't "see" the full nuclear charge $Z$ (the atomic number). It sees a smaller, effective charge — the net positive pull after accounting for the repulsion from inner electrons.

That's effective nuclear charge, denoted $Z_{\text{eff}}$ .

The Precise Statement

Z_{\text{eff}} = Z - S

$Z_{\text{eff}}$ is always less than $Z$ (except for hydrogen, which has no other electrons to shield — there $Z_{\text{eff}} = Z$ ).

What Determines the Shielding Constant $S$ ?

Not all electrons shield equally. The key rules:

Inner electrons shield outer electrons very effectively. An electron in the $n=1$ shell completely blocks about 1 unit of charge from an electron in $n=2$ .
Electrons in the same shell shield poorly. They're at roughly the same distance, so they don't block much of the nucleus from each other.
Outer electrons do not shield inner electrons at all. An electron farther out cannot block the nucleus from one closer in.

There are detailed rules (Slater's rules) to calculate $S$ numerically, but the core idea is simple: the more electron shells between an electron and the nucleus, the more shielding, and the lower $Z_{\text{eff}}$ .

Why Does This Matter?

$Z_{\text{eff}}$ explains three fundamental patterns in the periodic table:

Trend	What happens to $Z_{\text{eff}}$	Why
Across a period (left to right)	Increases	Adding protons ( $Z$ up) but electrons go into the same shell (shielding roughly constant). Net pull on outer electrons gets stronger.
Down a group (top to bottom)	Stays roughly constant or decreases slightly	Adding a new shell means much more shielding. The extra protons are almost completely cancelled by the new inner electrons.
Atomic size	Larger $Z_{\text{eff}}$ → smaller atom	Stronger pull pulls electrons closer to nucleus.

Note

This is why fluorine is smaller than lithium, even though fluorine has more protons. The extra protons in fluorine are not fully shielded — the outer electrons feel a much stronger pull.

A Concrete Example: Sodium vs. Chlorine

Sodium ( $Z = 11$ ): Electron configuration $1s^2 2s^2 2p^6 3s^1$

The outermost electron (3s) is shielded by the 10 inner electrons ( $1s^2 2s^2 2p^6$ ). Roughly, $S \approx 10$ , so $Z_{\text{eff}} \approx 11 - 10 = 1$ . The outer electron feels a pull equivalent to just one proton. …

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Oxygen has 8 electrons that fill orbitals in order of increasing energy: $1s^2\,2s^2\,2p^4$ . The orbital diagram shows paired electrons in $1s$ and $2s$ , then four electrons distributed across three $2p$ orbitals with two unpaired.

Why orbital diagrams matter

An orbital diagram is a visual map of where electrons live in an atom. Unlike electron configuration notation, which just counts electrons in each subshell, the diagram shows how electrons occupy individual orbitals within a subshell. This matters because electrons arrange themselves to minimize repulsion—a principle captured by Hund's rule—and the diagram reveals the atom's magnetic properties and bonding behavior.

Oxygen's 8 electrons fill orbitals starting from the lowest energy level, following three rules:

Aufbau principle: electrons occupy the lowest-energy orbital available
Pauli exclusion principle: each orbital holds at most two electrons with opposite spins
Hund's rule: within a subshell, electrons spread out into separate orbitals (all with parallel spin) before pairing up

Building the orbital diagram step by step

Start with the $1s$ orbital The first two electrons enter the $1s$ orbital, one spin-up ( $↑$ ) and one spin-down ( $↓$ ), filling it completely:

1s: \boxed{↑↓}

Move to the $2s$ orbital The next two electrons (3rd and 4th) fill the $2s$ orbital, again as a pair with opposite spins:

2s: \boxed{↑↓}

Distribute the remaining four electrons in the $2p$ subshell

The $2p$ subshell contains three orbitals of equal energy: $2p_x$ , $2p_y$ , and $2p_z$ . We have four electrons left.

By Hund's rule, the first three electrons enter separate orbitals, all with the same spin (say, spin-up):

2p: \boxed{↑}\quad \boxed{↑}\quad \boxed{↑}

The fourth electron must now pair with one of the existing electrons (opposite spin):

$2p: \boxed{↑↓}\quad \boxed{↑}\quad \boxed{↑}$ …

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Q.Show the distribution of electrons in oxygen atom (atomic number 8) using orbital diagram.

Why orbital diagrams matter

Building the orbital diagram step by step

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