🥼Organic Chemistry Unit 1 Review

Electron configurations describe how electrons are arranged in atoms, and they follow a set of rules that determine which orbitals fill first. Understanding these arrangements is central to organic chemistry because the way electrons are distributed dictates how atoms bond, what shapes molecules take, and how reactive a compound will be.

Valence electrons, the ones in the outermost shell, are especially important. They're the electrons directly involved in forming bonds. Knowing how to identify and count them connects directly to predicting an element's bonding behavior and its position on the periodic table.

Order of Orbital Filling

The Aufbau principle states that electrons fill orbitals starting from the lowest energy and working up. An electron won't occupy a higher-energy orbital until the lower-energy ones are filled.

The energy sequence of atomic orbitals follows this pattern:

$1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p < 7s < 5f < 6d < 7p$

Notice that sublevel energies can overlap between different principal quantum numbers. For example, $4s$ fills before $3d$ because $4s$ is slightly lower in energy. This overlap is a common source of confusion, so pay close attention to the sequence rather than assuming all $n = 3$ orbitals fill before any $n = 4$ orbitals.

Electron configuration notation uses the principal quantum number ( $n$ ) and the sublevel letter ( $s, p, d, f$ ) to label each orbital, with a superscript showing how many electrons are in that sublevel. For example, neon with 10 electrons is written as $1s^2 2s^2 2p^6$ .

Application of Electron Configuration Rules

Three rules work together when you're writing electron configurations:

The Pauli exclusion principle says each orbital can hold a maximum of two electrons, and those two electrons must have opposite spins ( $\uparrow\downarrow$ ). More precisely, no two electrons in the same atom can share an identical set of four quantum numbers. This is what caps each orbital at two electrons.

Hund's rule applies when you have orbitals of equal energy (like the three 2p orbitals). Electrons will spread out into separate orbitals before they start pairing up, and unpaired electrons in the same sublevel all have parallel spins ( $\uparrow\;\uparrow\;\uparrow$ ). Think of it like passengers on a bus: everyone takes their own seat before anyone sits next to a stranger.

Steps for writing an electron configuration:

Determine the total number of electrons (equal to the atomic number for a neutral atom).
Fill orbitals in order of increasing energy following the Aufbau principle.
Place a maximum of two electrons per orbital (Pauli exclusion principle), with opposite spins.
For orbitals of equal energy, place one electron in each before pairing any (Hund's rule), keeping unpaired spins parallel.

Example: Carbon (6 electrons) Configuration: $1s^2 2s^2 2p^2$ . The two 2p electrons go into separate 2p orbitals with parallel spins rather than pairing up in the same orbital. Drawing an orbital diagram makes this easier to see: you'd show two boxes with one up-arrow each and one empty box in the 2p sublevel.

Electrons in the Outermost Shell

Valence electrons are the electrons in the outermost (highest $n$ ) shell of an atom. These are the electrons that determine chemical reactivity and bonding behavior, which is why organic chemistry cares about them so much.

To find the number of valence electrons from an electron configuration:

Identify the highest principal quantum number ( $n$ ) in the configuration.
Add up all the electrons in the $s$ and $p$ sublevels of that highest $n$ .

Example: Nitrogen ( $1s^2 2s^2 2p^3$ ). The highest $n$ is 2. There are 2 electrons in $2s$ and 3 in $2p$ , giving 5 valence electrons.

For a quick shortcut, you can use noble gas (core) notation. Replace the inner-shell electrons with the symbol of the previous noble gas in brackets, then write out only the valence electrons. Potassium, for instance, becomes $[Ar]\,4s^1$ , showing one valence electron beyond the argon core. This notation keeps things clean and puts the focus on the electrons that matter most for bonding.

For main-group elements, the group number on the periodic table tells you the number of valence electrons directly (Group 1 = 1, Group 14 = 4, Group 17 = 7, etc.).

Energy Levels and Electron States

Shells (designated by the principal quantum number $n = 1, 2, 3, \ldots$ ) represent the main energy levels of electrons around the nucleus. Higher $n$ values mean greater average distance from the nucleus and higher energy.

The ground state is the lowest-energy arrangement of electrons in an atom. This is the configuration you write using the Aufbau principle. When an electron absorbs energy and jumps to a higher orbital, the atom is in an excited state, which is less stable and temporary.

Electron spin is an intrinsic property of electrons. Each electron has one of two possible spin states: spin-up ( $\uparrow$ ) or spin-down ( $\downarrow$ ). This property is what makes the Pauli exclusion principle work: two electrons can share an orbital only if their spins are opposite.

🥼Organic Chemistry Unit 1 Review