2 Periodic Trends

The causes of periodic trends

Periodic trends are recurring changes in atomic properties across the periodic table. Their main causes are changes in nuclear attraction, distance between the nucleus and outer electrons, and repulsion among electrons. These causes connect the trends in size, ionization energy, electron affinity, and electronegativity.

Effective nuclear charge and shielding

The nucleus has a positive charge determined by its number of protons, $Z$. Electrons, especially those in inner shells, shield outer electrons from some of this attraction. The resulting Effective nuclear charge can be approximated by

Here, $S$ represents shielding. This approximation helps explain trends, but it is not an exact calculation for every electron.

Across a period, each step adds a proton and usually adds an electron to the same main energy level. Shielding does not increase enough to cancel the added nuclear charge, so effective nuclear charge on valence electrons generally rises. Down a group, added shells increase both the distance of outer electrons from the nucleus and shielding. As a result, outer electrons are less tightly held, even though the nucleus has more protons.

Takeaway: Across a period, increasing effective nuclear charge strengthens the nucleus’s attraction to electrons in the same general shell; down a group, increased distance and shielding weaken the attraction experienced by outer electrons.

Atomic and ionic size

An atom has no sharp outer boundary, so Atomic radius is defined using measurement conventions. One example is covalent radius: half the distance between the nuclei of two identical atoms joined by a covalent bond.

• Across a period, atomic radius generally decreases. Increasing effective nuclear charge pulls electrons in the same general shell closer to the nucleus.

• Down a group, atomic radius generally increases. Valence electrons occupy shells with larger principal energy levels and are farther from the nucleus.

Ions differ in size from their neutral atoms. A Cation is usually smaller than its parent atom because removing electrons reduces electron–electron repulsion; losing an entire outer shell can make the ion much smaller. An Anion is usually larger than its parent atom because the added electron increases repulsion among electrons.

In an isoelectronic series, the species have the same number of electrons. Their radii decrease as proton number increases, because the same number of electrons is attracted by a greater nuclear charge. For example, all the species in the series below have $10$ electrons, while proton number increases from left to right:

Takeaway: Atomic size generally decreases across a period and increases down a group; ion formation can change size substantially, especially when an outer shell is lost.

Removing electrons

First ionization energy is the energy required to remove an electron from a ground-state gaseous atom:

Removing an electron always requires energy. First ionization energy generally increases across a period as effective nuclear charge rises. It generally decreases down a group because the outer electron is farther from the nucleus and more shielded.

The pattern has exceptions. Subshell energies and electron pairing can make particular electrons easier to remove than a simple left-to-right trend suggests. Successive ionization energies for the same element increase. A particularly large jump occurs after all valence electrons have been removed: the next electron would come from a more tightly bound inner shell.

Takeaway: Stronger attraction to outer electrons generally means more energy is needed to remove one, but electron arrangement affects the details.

Gaining electrons

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

With the energy-change convention, a negative value means energy is released. Electron gain can be either exothermic or endothermic. Across a period, adding an electron generally becomes more favorable, but groups 2, 15, and 18 are important exceptions. The pattern down a group is irregular rather than perfectly consistent.

For example, chlorine’s electron affinity is more favorable than fluorine’s. In chlorine’s larger valence shell, the added electron experiences less repulsion. When comparing tables, check the sign convention: some report energy released as a positive “affinity,” while others report the energy change with the opposite sign.

Takeaway: Electron affinity is not a perfectly regular trend, and its reported sign depends on the convention being used.

Attraction within bonds

Electronegativity is the relative ability of an atom in a chemical bond to attract shared electrons. It is not the energy change of a single process involving an isolated atom, and its values depend on the scale used.

On the commonly used Pauling scale, electronegativity generally increases across a period and decreases down a group. Fluorine is the most electronegative element on this scale. Noble gases are often omitted from introductory comparisons because many do not form ordinary bonds under typical conditions.

For example, chlorine attracts shared bonding electrons more strongly than sodium does. This helps explain why electrons in a $\mathrm{Na-Cl}$ bond are drawn toward chlorine, but electronegativity alone does not determine every feature of a bond.

Takeaway: Electronegativity describes attraction for shared electrons in a bond; it is related to, but distinct from, the energy changes involved in ionization and electron gain.