AP Chemistry 1.5: Atomic structure and electron configuration

This is one of the most important and detailed parts of Unit 1. Khan Academy covers shells, subshells, orbitals, electron configurations, Aufbau filling, valence electrons, ions, and 3d transition metals.

An atom contains:

• nucleus

• and

• electron cloud

The nucleus contains:

• protons

• neutrons

Electrons occupy regions of space around the nucleus.

Electrons should not be imagined as tiny planets following fixed circular paths.

• Instead, quantum mechanics describes regions where electrons have a high probability of being found.

• These regions are called orbitals .

Electrons are arranged into main energy levels called shells .

Shells are represented using the principal quantum number:

• n = 1, 2, 3, 4...

Generally:

• larger n → farther from nucleus → higher energy

• Example:

• n = 1 is closer to the nucleus than n = 3.

Each shell contains one or more subshells .

The major subshell types are:

• s

• p

• d

• f

Allowed subshells depend on the shell.

• n = 1:

• 1s

• n = 2:

• 2s, 2p

• n = 3:

• 3s, 3p, 3d

• n = 4:

• 4s, 4p, 4d, 4f

Subshells contain orbitals .

Each orbital can hold a maximum of:

• 2 electrons

Subshell capacities:

You should absolutely memorize this.

An electron configuration describes how electrons are distributed among subshells.

• Example:

Hydrogen:

• 1s¹

The "1":

• first energy level

The "s":

• s subshell

The superscript "1":

• one electron

Oxygen has atomic number 8.

A neutral oxygen atom therefore has:

8 electrons.

Electron configuration:

• 1s² 2s² 2p⁴

Count electrons:

• 2 + 2 + 4 = 8

The Aufbau principle says electrons generally fill the lowest-energy available orbitals first.

Important order:

• 1s

• 2s

• 2p

• 3s

• 3p

• 4s

• 3d

• 4p

• 5s

• 4d

• 5p

• 6s

• 4f

• 5d

• 6p

• 7s

You may have seen the diagonal-arrow diagram for remembering this.

Energy does not depend only on the shell number.

For neutral atoms near the beginning of the transition metals, 4s is generally occupied before 3d.

• So potassium is:

• [Ar]4s¹

• not:

• [Ar]3d¹

Each orbital holds no more than:

• 2 electrons

If two electrons occupy the same orbital, they have opposite spins.

Orbital diagram:

• ↑↓

• is allowed.

But three electrons in one orbital are not.

If several orbitals have the same energy, electrons occupy them one at a time before pairing .

For p orbitals, there are three orbitals.

• Suppose there are three p electrons:

• ↑ ↑ ↑

• not:

• ↑↓ ↑ empty

• Why?

Separating the electrons reduces repulsion and gives a lower-energy arrangement.

Nitrogen has 7 electrons.

Configuration:

• 1s² 2s² 2p³

• Its 2p orbital diagram is:

• ↑ ↑ ↑

• one electron in each p orbital.

Oxygen:

• 1s² 2s² 2p⁴

For the p orbitals:

• ↑↓ ↑ ↑

After placing one electron into each orbital, the fourth begins pairing.

Long configurations can be shortened using the previous noble gas.

• Example:

Sodium:

• full:

• 1s² 2s² 2p⁶ 3s¹

The first 10 electrons are the configuration of neon.

• So:

• [Ne]3s¹

Full:

• 1s² 2s² 2p⁶ 3s² 3p⁵

Shorthand:

• [Ne]3s²3p⁵

Valence electrons are the electrons most directly involved in chemical bonding.

For main-group elements, they are usually electrons in the highest occupied principal energy level .

• Example:

O:

• 1s² 2s² 2p⁴

• highest n = 2

Valence configuration:

• 2s²2p⁴

Number of valence electrons:

• 6

Khan Academy uses oxygen as this kind of example.

For main-group elements:

Exception:

He has only 2 , because the first shell can contain only two electrons.

A cation is a positively charged ion.

Cations form when electrons are lost .

• Example:

Na:

• [Ne]3s¹

Na loses one electron:

Na⁺:

• [Ne]

The atom now has:

• 11 protons

• 10 electrons

Net charge:

• +1

An anion is negatively charged.

Anions form when electrons are gained .

• Example:

Cl:

• [Ne]3s²3p⁵

Add one electron:

Cl⁻:

• [Ne]3s²3p⁶

• which is the same electron configuration as argon:

• [Ar]

Very useful relationship:

• charge = protons − electrons

• Example:

• 12 protons

• 10 electrons

• 12 − 10 = +2

• So the ion has charge:

• 2+

The d-block introduces an important complication.

• Example:

Fe has 26 electrons.

Configuration:

• [Ar]4s²3d⁶

You may also see the order written:

• [Ar]3d⁶4s²

Both describe the occupied subshells.

This is extremely testable.

Although 4s fills before 3d , when forming transition-metal cations, the electrons in the highest principal energy level are generally removed first.

• So 4s electrons are removed before 3d electrons.

• Example:

Fe:

• [Ar]4s²3d⁶

Fe²⁺:

• remove the two 4s electrons

• → [Ar]3d⁶

Fe³⁺:

• remove another electron from 3d

• → [Ar]3d⁵

• Common mistake

Incorrect:

• Fe²⁺ = [Ar]4s²3d⁴

• That removes d electrons before the 4s electrons.

Some transition metals have electron configurations that differ from the simplest predicted Aufbau arrangement.

Two major examples:

Chromium:

• Cr = [Ar]3d⁵4s¹

Copper:

• Cu = [Ar]3d¹⁰4s¹

• Instead of the simple predicted:

• Cr → [Ar]3d⁴4s²

• Cu → [Ar]3d⁹4s²

The actual arrangements are lower in energy.

For AP Chemistry, know that the Aufbau order is a model , and experimental electron configurations can contain exceptions.