Chapter 9: The Periodic Properties of the Elements

Development of the Periodic Table

Historical Background

The periodic table is a fundamental tool in chemistry, organizing elements according to recurring properties. Its development is credited to the Russian chemist Dmitri Mendeleev (1834–1907), who arranged elements by increasing atomic mass and noticed repeating patterns in their properties.

• Periodic Law: When elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically.

• Mendeleev grouped elements with similar properties in the same column, allowing him to predict the properties of undiscovered elements.

• When atomic mass order did not fit observed properties, Mendeleev reordered elements by other criteria (now known to be atomic number).

Additional info: The modern periodic table is arranged by increasing atomic number, not mass, which resolves earlier inconsistencies.

Predictive Power of the Periodic Table

Mendeleev's periodic law allows us to predict what the properties of an element will be based on its position. However, it does not explain why the pattern exists. Quantum mechanics later provided the theoretical explanation for these periodic trends.

Example: Mendeleev's Predictions

Mendeleev predicted the properties of elements that had not yet been discovered, such as eka-silicon (later identified as germanium). His predictions were remarkably accurate.

Element	Mendeleev's Predicted Properties	Actual Properties
Gallium	Atomic mass: ~68 amu Melting point: Low Density: 5.9 g/cm3 Formula of oxide: X2O3 Formula of chloride: XCl3	Atomic mass: 69.72 amu Melting point: 29.8°C Density: 5.90 g/cm3 Formula of oxide: Ga2O3 Formula of chloride: GaCl3
Germanium (eka-silicon)	Atomic mass: ~72 Density: 5.5 g/cm3 Formula of oxide: XO2 Formula of chloride: XCl4	Atomic mass: 72.6 Density: 5.47 g/cm3 Formula of oxide: GeO2 Formula of chloride: GeCl4

Quantum Mechanical Explanation of Periodicity

Electron Configuration

Quantum-mechanical theory describes the behavior of electrons in atoms. Electrons exist in regions called orbitals, and the arrangement of electrons in these orbitals is called the electron configuration.

• Orbital: A region in an atom where there is a high probability of finding electrons.

• Electron configuration: A listing of the orbitals occupied by electrons in an atom, with the number of electrons in each orbital indicated as a superscript.

Example: The electron configuration of hydrogen is written as H: 1s1, where '1s' is the orbital and the superscript '1' indicates one electron in that orbital.

Sublevels and Orbitals

Each principal energy level (n) contains one or more sublevels (s, p, d, f), each with a specific number of orbitals:

• s sublevel: 1 orbital (can hold 2 electrons)

• p sublevel: 3 orbitals (can hold 6 electrons)

• d sublevel: 5 orbitals (can hold 10 electrons)

• f sublevel: 7 orbitals (can hold 14 electrons)

Each orbital can hold a maximum of 2 electrons with opposite spins.

Quantum Numbers and Electron Spin

Electrons in atoms are described by four quantum numbers:

• Principal quantum number (n): Energy level

• Angular momentum quantum number (l): Sublevel (s, p, d, f)

• Magnetic quantum number (ml): Specific orbital within a sublevel

• Spin quantum number (ms): Electron spin (+1/2 or -1/2)

The Pauli Exclusion Principle states that no two electrons in an atom can have the same set of four quantum numbers, so each orbital can hold only two electrons with opposite spins.

Example: Electron Configuration Notation

For hydrogen (H):

• Electron configuration: 1s1

• n = 1 (first energy level), l = 0 (s orbital), ml = 0, ms = +1/2 or -1/2

Additional info: For helium (He), the configuration is 1s2, with two electrons of opposite spin in the same orbital.