Chemistry: Periodic Table Trends

Define atomic radius and describe its trend across a period.

Atomic radius is half the distance between the nuclei of two identical atoms bonded together. It generally decreases across a period due to increasing effective nuclear charge pulling valence electrons closer to the nucleus.

How does atomic radius change down a group, and why?

Atomic radius increases down a group. This is because electrons occupy higher energy levels (shells) further from the nucleus, and increased shielding reduces the attraction of the nucleus.

What is ionization energy, and what is its trend across a period?

Ionization energy is the energy required to remove one electron from a gaseous atom. It generally increases across a period because effective nuclear charge increases, holding valence electrons more tightly.

How does ionization energy change down a group, and why?

Ionization energy generally decreases down a group. The valence electrons are further from the nucleus and experience greater shielding, making them easier to remove.

Define electron affinity and state its general trend for nonmetals.

Electron affinity is the energy change when an electron is added to a gaseous atom. For nonmetals, it generally becomes more negative (more favorable) across a period as effective nuclear charge increases.

What is electronegativity, and how does it trend across a period?

Electronegativity is an atom's ability to attract electrons in a chemical bond. It generally increases across a period because of increasing effective nuclear charge and smaller atomic radius.

Describe the trend of electronegativity down a group.

Electronegativity generally decreases down a group. This is due to increased atomic radius and electron shielding, which reduces the attraction for bonding electrons.

How does metallic character change across a period and down a group?

Metallic character decreases across a period (left to right) and increases down a group. Metals tend to lose electrons, a property enhanced by larger atomic radii and lower ionization energies.

Compare the size of a cation to its neutral parent atom.

A cation is always smaller than its neutral parent atom. This is because the loss of valence electrons often removes an entire electron shell, and the remaining electrons are pulled closer by the same nuclear charge.

Compare the size of an anion to its neutral parent atom.

An anion is always larger than its neutral parent atom. The addition of electrons increases electron-electron repulsion within the same shell, causing the electron cloud to expand.

What is the trend for the reactivity of alkali metals (Group 1) down the group?

The reactivity of alkali metals increases down the group. This is because their single valence electron becomes easier to lose due to decreasing ionization energy and increasing atomic radius.

What is the trend for the reactivity of halogens (Group 17) down the group?

The reactivity of halogens decreases down the group. Halogens react by gaining an electron, which becomes less favorable due to increasing atomic radius and decreasing electronegativity.

Explain the concept of effective nuclear charge.

Effective nuclear charge (Zeff) is the net positive charge experienced by an electron in a multi-electron atom. It generally increases across a period as protons are added and inner shell shielding remains relatively constant for valence electrons.

What is electron shielding, and how does it affect periodic trends?

Electron shielding is the reduction of the attractive force between the nucleus and an outer electron due to the presence of inner electrons. It increases down a group, contributing to larger atomic radii and lower ionization energies.

How does the number of valence electrons change across a period and down a group?

The number of valence electrons increases across a period (from 1 to 8 for main group elements). It remains constant for all elements within the same main group.