Periodic Properties of the Elements

Development of the Periodic Table

The periodic table is a fundamental tool in chemistry, organizing elements according to their properties. Dmitri Mendeleev and Lothar Meyer independently proposed grouping elements based on similarities in their properties. Mendeleev is especially credited for using chemical properties to organize the table and predicting the existence and properties of elements not yet discovered, such as germanium.

• Mendeleev’s Table: Organized by atomic mass, the most fundamental property known at the time.

• Discovery of Atomic Number: Henry Moseley established atomic number as the basis for periodicity, using X-ray experiments to count protons.

• Periodicity: Refers to the repetitive pattern of properties as a function of atomic number.

Effective Nuclear Charge

Many atomic properties depend on the attraction between valence electrons and the nucleus, which is influenced by both the nuclear charge and electron repulsion. The concept of effective nuclear charge (Zeff) is central to understanding periodic trends.

• Definition: The effective nuclear charge is the net positive charge experienced by an electron in a multi-electron atom.

• Formula: where Z is the atomic number and S is the screening constant (number of core electrons).

• Trend: Increases from left to right across a period.

Sizes of Atoms and Ions

The size of atoms and ions is determined by the space occupied by electrons, which is a probability distribution. Two main types of atomic radii are used:

• Nonbonding Atomic Radius (van der Waals radius): Half the shortest distance between nuclei during a collision.

• Bonding Atomic Radius (covalent radius): Half the distance between nuclei in a bond.

• Trends: Bonding atomic radius decreases across a period and increases down a group.

Sizes of Ions (Ionic Radii)

• Cations: Smaller than their parent atoms due to loss of electrons and reduced repulsion.

• Anions: Larger than their parent atoms due to gain of electrons and increased repulsion.

• Isoelectronic Series: Ions with the same number of electrons; size decreases with increasing nuclear charge.

Ionization Energy and Electron Affinity

Ionization energy and electron affinity are key properties that reflect how easily atoms lose or gain electrons.

• Ionization Energy (I): Minimum energy required to remove an electron from a gaseous atom or ion.

• First Ionization Energy: Energy to remove the first electron.

• Second Ionization Energy: Energy to remove the second electron.

• Trend: Increases across a period, decreases down a group.

• Irregularities: Occur when electrons enter new sublevels or pair in orbitals.

• Electron Affinity (EA): Energy change when an electron is added to a gaseous atom; typically exothermic (negative).

• Trend: Generally increases across a period, with exceptions in Groups 2A, 5A, and 8A.

Metals, Nonmetals, and Metalloids

Elements are classified as metals, nonmetals, or metalloids based on their physical and chemical properties.

• Metals: Shiny, conductive, malleable, ductile, mostly solids at room temperature, low ionization energies, form cations.

• Nonmetals: Found on the right side of the periodic table, can be solid, liquid, or gas, dull, brittle, poor conductors, large negative electron affinity, form anions.

• Metalloids: Exhibit properties of both metals and nonmetals; several are electrical semiconductors.

Comparison Table: Properties of Metals and Nonmetals

Group Trends

Elements in the same group share similar properties, with trends observed within each group.

• Group 1A (Alkali Metals): Soft, metallic solids, low densities and melting points, low ionization energies, react exothermically with water.

• Group 2A (Alkaline Earth Metals): Higher densities and melting points than alkali metals, form +2 cations, reactivity increases down the group.

• Group 6A: Oxygen, sulfur, selenium are nonmetals; tellurium is a metalloid; polonium is a metal. Metallic character increases down the group.

• Group 7A (Halogens): Typical nonmetals, highly negative electron affinities, react with metals to form halides.

• Group 8A (Noble Gases): Monatomic gases, very large ionization energies, positive electron affinities, generally unreactive.

Select Nonmetals: Hydrogen and Oxygen

• Hydrogen: Electron configuration suggests metallic character, but it behaves as a nonmetal. Occurs as a colorless diatomic gas, forms hydride anions with metals.

• Oxygen: Exists as O2 (colorless, odorless gas) and O3 (ozone, pale blue, poisonous gas). Forms oxides, peroxides, and superoxides.

Summary Table: Group Properties

Periodic Trends for Ionization Energy

The first ionization energy generally increases across a period and decreases down a group. The s- and p-block elements show a larger range of values, while the d-block increases slowly and the f-block shows small variations.

Additional info: The periodic table and its trends are foundational for understanding chemical reactivity, bonding, and the classification of elements. These notes provide a concise overview suitable for exam preparation.