06 Electron Configurations and Periodic Trends

Electron configurations and orbital filling

An electron configuration describes how an atom’s electrons are distributed among its orbitals. For a neutral atom, the number of electrons equals its atomic number. In configuration notation, the number identifies the principal energy level, the letter identifies the subshell, and the superscript gives the number of electrons in that subshell. For example, oxygen’s configuration is $1s^2\,2s^2\,2p^4$.

Three principles guide orbital filling:

• The Aufbau principle says that electrons occupy available orbitals in order of increasing energy.

• The Pauli exclusion principle limits each orbital to two electrons, which must have opposite spins.

• Hund’s rule says that orbitals of equal energy are singly occupied before any orbital receives a second electron.

Subshell capacities are $s^2$, $p^6$, $d^{10}$, and $f^{14}$. A commonly used filling order is:

Oxygen’s four electrons in the $2p$ subshell occupy three equal-energy orbitals as one pair and two unpaired electrons. Noble-gas shorthand represents inner electrons with the symbol of the preceding noble gas in brackets; potassium is $[\mathrm{Ar}]\,4s^1$. Chromium and copper are examples of elements whose ground-state configurations differ from a simple prediction based on the filling order because the energy differences between subshells are small.

The periodic table’s blocks indicate which type of subshell is being filled. Elements in the same main-group column have similar outer-electron configurations, which helps explain why their chemical properties recur.

The forces behind periodic trends

Shielding and effective nuclear charge help explain why properties change across the periodic table. Inner electrons shield outer electrons from some of the nucleus’s positive charge. The net attraction an electron experiences is its effective nuclear charge, often written as $Z_{\mathrm{eff}}$.

Across a period, protons are added while electrons are generally added to the same principal energy level. As a result, effective nuclear charge tends to increase. Down a group, each step adds an occupied energy level. The outer electrons are then farther from the nucleus and experience more shielding.

These changes in attraction and distance provide a common explanation for several trends: across a period, the electron cloud is generally pulled closer to the nucleus; down a group, outer electrons are generally farther away.

Atomic size

Atomic radius is not based on a sharply defined edge, so chemists compare atomic sizes using specified measures such as covalent radius.

• Across a period: atomic radius generally decreases because increasing effective nuclear charge pulls the electron cloud closer.

• Down a group: atomic radius generally increases because outer electrons occupy higher energy levels farther from the nucleus.

For example, potassium is larger than lithium because its outer electron occupies the fourth principal energy level rather than the second. In period three, sodium is generally larger than chlorine, as the stronger effective nuclear attraction across the period contracts the electron cloud. Overall, atoms tend to be largest toward the lower left of the periodic table.

Removing an electron: first ionization energy

First ionization energy is the energy needed to remove one electron from a gaseous atom. The process can be represented as:

First ionization energy generally increases across a period as the nucleus attracts electrons more strongly. It generally decreases down a group because the outer electron is farther from the nucleus and more shielded. Thus, elements near the upper right tend to have high first ionization energies, while those near the lower left tend to have low values.

The trend has exceptions. Boron’s first ionization energy is lower than beryllium’s because boron’s outermost electron is in a $2p$ subshell, which is higher in energy and easier to remove than beryllium’s $2s$ electron. Pairing and subshell stability also lead to smaller departures from the general pattern.

Gaining an electron: electron affinity

Electron affinity describes the energy change when a gaseous atom gains an electron:

With the energy-change convention, a more negative value means more energy is released. Electron affinity tends to become more negative across a period, but the pattern is irregular. An added electron may begin a new subshell, pair in an occupied orbital, or enter beyond a filled shell, leading to departures from a simple trend.

Chlorine’s electron affinity is more negative than fluorine’s. Although fluorine is higher in its group, the added electron enters its compact $2p$ shell, where electron–electron repulsion is relatively strong. In chlorine, the electron enters the roomier $3p$ shell. Some references instead state electron affinity as the positive amount of energy released, so check the convention before comparing values.

Takeaway: Across a period, increasing effective nuclear charge generally makes atoms smaller and raises first ionization energy. Down a group, added energy levels and shielding generally make atoms larger and lower first ionization energy. Electron affinity follows a less regular pattern, and its sign depends on the convention used.