Atomic Structure, Periodic Table, and Periodic Trends: Study Notes for General Chemistry

4.2 The Periodic Table

Introduction to the Periodic Table

The periodic table is a systematic arrangement of all known chemical elements, organized by increasing atomic number. It is a fundamental tool in chemistry for understanding element properties and relationships.

• Groups (columns) contain elements with similar chemical properties.

• Periods (rows) represent elements with increasing atomic number.

• Elements in Group 1A are called the alkali metals; Group 2A, the alkaline earth metals; Group 7A, the halogens; and Group 8A, the noble gases.

• Metals are located on the left and center, nonmetals on the right, and metalloids border the stair-step line.

Example: Sodium (Na) is an alkali metal in Group 1A and Period 3.

4.3 The Atom

Structure and Subatomic Particles

An atom is the smallest unit of an element that retains its chemical properties. Atoms are composed of three main subatomic particles:

• Protons (p+): Positively charged particles located in the nucleus.

• Neutrons (n0): Neutral particles also found in the nucleus.

• Electrons (e-): Negatively charged particles that move around the nucleus in electron clouds.

Example: A carbon atom has 6 protons, 6 neutrons, and 6 electrons.

4.4 Atomic Number and Mass Number

Defining Atomic Identity

The atomic number and mass number are essential for identifying and distinguishing atoms of different elements.

• Atomic Number (Z): The number of protons in the nucleus of an atom. It defines the element.

• Mass Number (A): The total number of protons and neutrons in the nucleus.

$A = Z + N$

• Where A is the mass number, Z is the atomic number, and N is the number of neutrons.

Example: For 23Na, Z = 11, N = 12, so A = 23.

4.5 Isotopes and Atomic Mass

Isotopes and Average Atomic Mass

Isotopes are atoms of the same element with different numbers of neutrons, resulting in different mass numbers.

• Isotopes have identical chemical properties but different physical properties due to mass differences.

• The atomic mass of an element is the weighted average of the masses of its naturally occurring isotopes.

$\text{Atomic mass} = \sum (\text{isotope mass} \times \text{fractional abundance})$

Example: Magnesium has three naturally occurring isotopes: 24Mg, 25Mg, and 26Mg.

4.6 Electron Energy Levels

Energy Levels, Sublevels, and Orbitals

Electrons in atoms occupy specific energy levels, sublevels, and orbitals, which determine the atom's chemical behavior.

• Energy levels (n = 1, 2, 3, ...) are quantized and increase in energy as n increases.

• Each energy level contains one or more sublevels (s, p, d, f).

• Each sublevel contains a specific number of orbitals, each holding up to two electrons.

• Electromagnetic radiation, such as light, is related to the energy changes of electrons moving between levels.

Example: The second energy level (n = 2) contains s and p sublevels.

4.7 Electron Configurations

Writing Electron Configurations and Orbital Diagrams

Electron configuration describes the arrangement of electrons in an atom's orbitals. The Aufbau principle, Pauli exclusion principle, and Hund's rule guide the filling order.

• Electrons fill orbitals in order of increasing energy: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc.

• Each orbital holds a maximum of two electrons with opposite spins.

• Within a sublevel, electrons occupy separate orbitals before pairing (Hund's rule).

Example: The electron configuration of oxygen (Z = 8) is 1s2 2s2 2p4.

4.8 Trends in Periodic Properties

Periodic Trends Explained by Electron Configuration

The periodic table reveals trends in atomic and chemical properties, explained by electron configurations and atomic structure.

• Atomic radius increases down a group and decreases across a period.

• Ionization energy decreases down a group and increases across a period.

• Metallic character increases down a group and decreases across a period.

• Elements in the same group have similar valence electron configurations, leading to similar chemical properties.

Example: Lithium (Li), sodium (Na), and potassium (K) are all alkali metals with similar reactivity due to their single valence electron.

Summary Table: Periodic Trends

Property	Down a Group	Across a Period (Left to Right)
Atomic Radius	Increases	Decreases
Ionization Energy	Decreases	Increases
Metallic Character	Increases	Decreases