Periodic Properties of Elements

Position and Nuclear Attraction

An element's position in the periodic table gives clues about atomic size, how easily an electron is removed, and how strongly the atom pulls electrons in a bond.

These trends depend on nuclear attraction toward valence electrons. Valence electrons are electrons in the outer shell. When the number of protons increases, the nuclear pull usually becomes stronger. When the number of shells increases, valence electrons are farther from the nucleus and feel a weaker pull.

Periodic Trend Model

Pick one property, then compare whether its value increases or decreases across a period and down a group.

Start with the period track: the $\mathrm{Na}$ marker is largest, then the markers shrink toward $\mathrm{Cl}$ . In the same period, the nucleus pulls the outer electrons more strongly.

Meaning: Relative atomic size, usually in $\mathrm{pm}$ .
Across a period: From $\mathrm{Na}$ to $\mathrm{Cl}$ , the size decreases.
Down a group: From $\mathrm{N}$ to $\mathrm{Bi}$ , the size increases.
Cause: To the right, $Z$ increases. Downward, the number of shells increases.

Four Periodic Properties

Each property describes a different aspect of an atom. The trend arrows describe changes in atomic size, the energy needed to remove an electron, the energy change when an electron is added, and the attraction of shared electrons in a bond.

Atomic radius measures atomic size. Atoms do not have hard surfaces, so the radius is inferred from distances between nuclei using a stated measurement model.
First ionization energy is the energy needed to remove the first electron from an isolated neutral atom in the gas phase.
Electron affinity is the energy change when an electron is added to an isolated atom in the gas phase.
Electronegativity describes how strongly an atom attracts shared electrons in a bond.

The difference between electron affinity and electronegativity lies in the electron's situation. Electron affinity describes the energy change when an electron is added to a gaseous atom. Electronegativity describes the attraction of shared electrons within a bond.

Trends Across Periods and Groups

Across a period, proton number increases while valence electrons enter the same main shell. Shielding changes less than nuclear charge, so the effective nuclear charge felt by the valence electrons generally increases. Atomic radius therefore tends to decrease, while first ionization energy and electronegativity tend to increase.

Down one group, the number of occupied shells increases. Valence electrons are farther from the nucleus, and inner electrons reduce the attraction they feel from the nucleus. This reduction is called electron shielding.

Two sign conventions are used for electron affinity. The IUPAC Gold Book uses the energy released when an electron attaches, so favorable attachment has $\mathrm{EA}>0$ . When the same process is written as a reaction-energy change, $\Delta E=-\mathrm{EA}<0$ .

For chlorine, attaching an electron releases energy. We write $\mathrm{EA}>0$ for its positive electron affinity and $\Delta E=-\mathrm{EA}<0$ for the negative energy change of the reaction:

\mathrm{Cl(g)}+e^- \rightarrow \mathrm{Cl^-(g)}

Noble gases such as $\mathrm{Ar}$ have filled valence shells, so they do not easily accept extra electrons. Electron affinity has more exceptions to a simple left-to-right trend than atomic radius or first ionization energy.

Direction of Periodic Trends

Use these two questions as a quick guide: are the valence electrons getting closer to or farther from the nucleus? Is the effective nuclear attraction getting stronger or weaker?

If valence electrons are closer and the nuclear pull is stronger, atoms tend to get smaller and electrons become harder to remove. If more shells place valence electrons farther away, atoms tend to get larger and outer electrons become easier to remove. These two relationships explain the main trend arrows.

Position and Nuclear Attraction

An element's position in the periodic table gives clues about atomic size, how easily an electron is removed, and how strongly the atom pulls electrons in a bond.

Periodic Trend Model

Pick one property, then compare whether its value increases or decreases across a period and down a group.

Start with the period track: the $\mathrm{Na}$ marker is largest, then the markers shrink toward $\mathrm{Cl}$ . In the same period, the nucleus pulls the outer electrons more strongly.

Meaning: Relative atomic size, usually in $\mathrm{pm}$ .
Across a period: From $\mathrm{Na}$ to $\mathrm{Cl}$ , the size decreases.
Down a group: From $\mathrm{N}$ to $\mathrm{Bi}$ , the size increases.
Cause: To the right, $Z$ increases. Downward, the number of shells increases.

Four Periodic Properties

Atomic radius measures atomic size. Atoms do not have hard surfaces, so the radius is inferred from distances between nuclei using a stated measurement model.

First ionization energy is the energy needed to remove the first electron from an isolated neutral atom in the gas phase.

Electron affinity is the energy change when an electron is added to an isolated atom in the gas phase.

Electronegativity describes how strongly an atom attracts shared electrons in a bond.

Trends Across Periods and Groups

Two sign conventions are used for electron affinity. The IUPAC Gold Book uses the energy released when an electron attaches, so favorable attachment has

\mathrm{EA}>0

. When the same process is written as a reaction-energy change,

\Delta E=-\mathrm{EA}<0

For chlorine, attaching an electron releases energy. We write

\mathrm{EA}>0

for its positive electron affinity and

\Delta E=-\mathrm{EA}<0

for the negative energy change of the reaction:

\mathrm{Cl(g)}+e^- \rightarrow \mathrm{Cl^-(g)}

Noble gases such as

\mathrm{Ar}

have filled valence shells, so they do not easily accept extra electrons. Electron affinity has more exceptions to a simple left-to-right trend than atomic radius or first ionization energy.

Direction of Periodic Trends

Use these two questions as a quick guide: are the valence electrons getting closer to or farther from the nucleus? Is the effective nuclear attraction getting stronger or weaker?