Periodic Properties of the Elements

Introduction to the Periodic Law

The periodic law, first proposed by Mendeleev in 1869, states that the chemical and physical properties of elements repeat periodically when the elements are arranged in order of increasing atomic number (Z). This foundational principle underlies the structure of the modern periodic table and explains recurring trends in atomic and molecular properties.

Trends in Atomic Properties

Atomic Size

Atomic size refers to the typical distance from the nucleus to the boundary of the surrounding cloud of electrons. Because atoms do not have sharply defined edges, atomic size is often described using different conventions, such as metallic radius, covalent radius, or van der Waals radius.

• Metallic radius: Half the distance between nuclei in a metallic crystal.

• Covalent radius: Half the distance between nuclei of identical atoms joined by a covalent bond.

To understand trends in atomic radii, it is essential to consider the effective nuclear charge (Zeff), which is the net positive charge experienced by valence electrons. It is calculated as:

where Z is the atomic number and S is the number of core electrons (shielding constant).

• Across a period: Atomic radius decreases due to increasing Zeff, which pulls electrons closer to the nucleus.

• Down a group: Atomic radius increases as additional electron shells are added, increasing the distance between the nucleus and the outermost electrons.

Example: Lithium (Li) is larger than Beryllium (Be) across a period, but Sodium (Na) is larger than Lithium (Li) down a group.

Ionization Energy (IE)

Ionization energy is the energy required to remove an electron from an isolated gaseous atom or ion. It is a measure of how strongly an atom holds onto its electrons.

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

• Down a group: IE decreases as atomic size increases and valence electrons are farther from the nucleus.

There are anomalies in these trends, such as between Be (IE = 899 kJ/mol) and B (IE = 800 kJ/mol), which can be explained by electron configurations and subshell stability.

Successive ionization energies can also be measured. Each successive IE is larger than the previous, with a significant jump when core electrons are removed.

Example: For a period 3 element with the following IEs (in kJ/mol): 1012, 1903, 2910, 4956, 6278, 22230, the large jump after the sixth IE indicates the removal of a core electron.

Electron Affinity (EA)

Electron affinity is the energy change that occurs when a gaseous atom or ion gains an electron. It reflects the tendency of an atom to accept an electron.

• Across a period: EA generally becomes more negative (increases in magnitude), indicating a greater tendency to gain electrons.

• Down a group: EA generally becomes less negative (decreases in magnitude), though there are exceptions.

Summary of Trends:

• Zeff: Increases across a period, slightly increases down a group.

• Atomic radius: Decreases across a period, increases down a group.

• Ionization energy: Increases across a period, decreases down a group.

• Electron affinity: Becomes more negative across a period, less negative down a group.

Trends in Physical Properties

Melting Point, Boiling Point, Conductivity, and Hardness

Physical properties such as melting point (mp), boiling point (bp), electrical conductivity, and hardness depend on the type of structural entities and interparticle forces present in a substance. These properties can be interpolated for elements such as bromine by examining trends among halogens.

Example: Bromine (Br2) is a liquid at room temperature, with mp and bp values between those of Cl2 and I2.

For period 3 elements, melting point, hardness, and conductivity vary due to changes in bonding and structure:

Trends in Chemical Properties

Metals vs. Nonmetals: Relation to IE and EA

Metals tend to have low ionization energies and low electron affinities, making them good reducing agents (they lose electrons easily). Nonmetals have high ionization energies and high electron affinities, making them good oxidizing agents (they gain electrons easily). Metallic character increases down a group and decreases across a period.

Redox Behavior: Tendency to Lose or Gain Electrons

In redox reactions, electrons are transferred between species, resulting in changes in oxidation states. The substance that loses electrons is oxidized (reducing agent), and the substance that gains electrons is reduced (oxidizing agent).

Example: In the reaction Na(s) + H2O(l) → NaOH(aq) + H2(g), sodium is oxidized and water is reduced. Alkali metals become more reactive down the group and are strong reducing agents.

Halogens are strong oxidizing agents, and their reactivity decreases down the group. In displacement reactions, the more reactive halogen will displace a less reactive halide ion from solution:

• Cl2(g) + 2 I–(aq) → 2 Cl–(aq) + I2(s)

• Br2(l) + 2 Cl–(aq) → No reaction (Cl2 is a stronger oxidizing agent than Br2)

Good reducing agents: Metals with low IE (e.g., Mg, CaH2) Good oxidizing agents: Nonmetals with high EA (e.g., O2, KMnO4)

Example: In the Space Shuttle booster reaction: N2O4(l) + 2 N2H4(l) → 3 N2(g) + 4 H2O(g), N2O4 is the oxidizing agent and N2H4 is the reducing agent.

Acid-Base Behavior of Oxides in Water

Metal oxides tend to be basic, reacting with water to form hydroxides, while nonmetal oxides are acidic, forming acids in water. These oxides are also called base or acid anhydrides.

• Basic oxide: Li2O(s) + H2O(l) → 2 LiOH(aq)

• Acidic oxide: SO3(g) + H2O(l) → H2SO4(aq)

There is a gradual trend in the periodic table from basic to acidic oxides as one moves from left to right across a period:

Additional info: Amphoteric oxides (e.g., Al2O3) can react as either acids or bases, depending on the reaction conditions.