Periodic Properties of the Elements: Trends and Explanations

Periodic Properties of the Elements

Development of the Periodic Table

The periodic table organizes elements based on increasing atomic number and recurring chemical properties. This arrangement reveals periodic trends in element properties, which are essential for understanding chemical behavior.

• Periodic Law: The properties of elements are periodic functions of their atomic numbers.

• Groups (columns): Elements with similar chemical properties.

• Periods (rows): Elements with increasing atomic number.

Effective Nuclear Charge (Zeff)

Effective nuclear charge is the net positive charge experienced by an electron in a many-electron atom. Due to electron shielding, the outer electrons do not experience the full nuclear charge.

• Definition: The amount of positive charge "felt" by an electron after accounting for shielding by other electrons.

• Formula:

$Z_{\mathrm{eff}} = Z - S$

• Z: Atomic number (number of protons)

• S: Number of core (shielding) electrons

Trend: Zeff increases across a period (left to right) and increases slightly down a group.

Sizes of Atoms and Ions

Trends in Atomic Radii

The atomic radius is the distance from the nucleus to the outermost electron shell. Atomic size varies predictably across the periodic table.

• Down a group: Atomic radius increases due to the addition of electron shells.

• Across a period: Atomic radius decreases from left to right due to increasing Zeff, pulling electrons closer to the nucleus.

Example: Which of the following atoms is largest: Be, Mg, Ca, or Sr? Answer: Sr is the largest because it is furthest down the group.

Example: Which of the following atoms is largest: K, V, As, or Br? Answer: K is the largest because it is furthest to the left in the period.

Trends in Ionic Radii

• Cations (positive ions): Smaller than their parent atoms due to loss of electrons and decreased electron-electron repulsion.

• Anions (negative ions): Larger than their parent atoms due to gain of electrons and increased electron-electron repulsion.

General Trend: For ions with the same charge, ionic radius increases down a group.

Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom or ion.

• First Ionization Energy (IE1): Energy needed to remove the first electron from a neutral atom.

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

• Second Ionization Energy (IE2): Energy needed to remove the second electron from a +1 ion.

$\mathrm{Na}^+(g) \rightarrow \mathrm{Na}^{2+}(g) + e^-$

Trends in First Ionization Energy

• Down a group: Ionization energy decreases because outer electrons are farther from the nucleus and more shielded.

• Across a period: Ionization energy increases due to increasing Zeff and decreasing atomic radius.

Example: Which element has the highest 1st ionization energy: Be, Mg, Ca, or Sr? Answer: Be (furthest up and right in the group).

Example: Which element has the highest 1st ionization energy: K, V, As, or Br? Answer: Br (furthest right in the period).

Electron Configuration of Ions

When forming ions, electrons are removed from the highest principal quantum number (n) orbitals first.

• Example a: Na+ (sodium ion)

Neutral Na: 1s2 2s2 2p6 3s1 Na+: 1s2 2s2 2p6

• Example b: Zn2+ (zinc ion)

Neutral Zn: [Ar] 4s2 3d10 Zn2+: [Ar] 3d10

• Example c: S2− (sulfide ion)

Neutral S: 1s2 2s2 2p6 3s2 3p4 S2−: 1s2 2s2 2p6 3s2 3p6

Electron Affinity

Electron affinity is the energy change when an electron is added to a gaseous atom, forming an anion.

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

• Negative electron affinity: Energy is released (exothermic) when an electron is added.

• Most nonmetals (except group 8A): Have negative electron affinities; they tend to gain electrons easily.

Trends in Electron Affinity

• Across a period: Electron affinity generally becomes more negative (more energy released) from left to right.

• Down a group: Electron affinity becomes less negative (less energy released) as atomic size increases.

Note: Group 8A (noble gases) have positive or near-zero electron affinities because their electron shells are full.