Atoms and the Periodic Table: Fundamentals for GOB Chemistry

Atomic Theory and the Structure of Atoms

Introduction to Atomic Theory

The concept of the atom is central to chemistry. An atom is the smallest particle of an element that retains its chemical identity. The term comes from the Greek atomos, meaning "indivisible." Atomic theory provides the foundational principles for understanding chemical behavior.

• Assumption 1: All matter is composed of atoms.

• Assumption 2: Atoms of a given element differ from those of other elements.

• Assumption 3: Chemical compounds consist of atoms combined in specific ratios; only whole atoms can combine.

• Assumption 4: Chemical reactions change only the way atoms are combined in compounds.

Subatomic particles make up atoms:

• Protons: Positively charged, located in the nucleus.

• Neutrons: Neutral, similar mass to protons, located in the nucleus.

• Electrons: Negatively charged, much lighter than protons/neutrons, move around the nucleus.

Atoms are mostly empty space, with a dense nucleus at the center.

Elements and Atomic Number

Defining Elements by Atomic Number

Each element is defined by its atomic number (Z), which is the number of protons in its nucleus. All atoms of a particular element have the same atomic number.

• Mass number (A): The sum of protons and neutrons in an atom.

• In a neutral atom, the number of electrons equals the number of protons.

Example: Carbon (Z = 6) has 6 protons and 6 electrons.

Isotopes and Atomic Weight

Isotopes and Calculating Atomic Weight

Isotopes are atoms of the same element (same Z) with different numbers of neutrons (different A). Most elements exist as mixtures of isotopes.

• Atomic weight: The weighted average mass of an element's atoms, based on the relative abundance and mass of each isotope.

Formula:

$\text{Atomic weight} = \sum [(\text{isotopic abundance}) \times (\text{isotopic mass})]$

Example: Gallium has two isotopes: Ga-69 (60.4%, 68.9257 amu) and Ga-71 (39.6%, 70.9248 amu). The atomic weight is:

$\text{Atomic weight} = (0.604 \times 68.9257) + (0.396 \times 70.9248) = 69.7$ amu

The Periodic Table

Organization and Classification

The periodic table arranges all known elements by increasing atomic number. Each box contains the element's symbol, name, and atomic mass.

• Metals: Malleable, lustrous, good conductors; found on the left side.

• Nonmetals: Poor conductors; found on the upper-right side.

• Metalloids: Intermediate properties; located in a zigzag band between metals and nonmetals.

Elements in the same vertical column (group) have similar chemical properties.

Some Characteristics of Different Groups

Main Element Groups

• Group 1A – Alkali Metals: Li, Na, K, Rb, Cs, Fr. Shiny, soft, low melting points, highly reactive, never found pure in nature.

• Group 2A – Alkaline Earth Metals: Be, Mg, Ca, Sr, Ba, Ra. Lustrous, silvery, less reactive than alkali metals, not found pure in nature.

• Group 7A – Halogens: F, Cl, Br, I, At. Colorful, corrosive nonmetals, found only in compounds.

• Group 8A – Noble Gases: He, Ne, Ar, Kr, Xe, Rn. Colorless, very low reactivity.

Electronic Structure of Atoms

Quantum Mechanical Model

The arrangement of electrons determines chemical properties. The quantum mechanical model describes electrons as occupying discrete energy levels (shells, subshells, orbitals).

• Shell: Main energy level, designated by principal quantum number (n = 1, 2, 3, ...).

• Subshell: s, p, d, f types, increasing in energy.

• Orbital: Region where electrons are most likely found; each orbital holds two electrons with opposite spins.

Electron capacity per shell:

• n = 1: 2 electrons

• n = 2: 8 electrons

• n = 3: 18 electrons

• n = 4: 32 electrons

Subshell orbital counts: s (1), p (3), d (5), f (7)

Electron Configurations

Rules for Electron Arrangement

Electron configuration describes how electrons are distributed among orbitals.

• Rule 1: Electrons occupy the lowest energy orbitals available.

• Rule 2: Each orbital holds two electrons of opposite spin.

• Rule 3: Orbitals of equal energy are half-filled before any is completely filled.

Notation: The number of electrons in each subshell is shown as a superscript (e.g., 1s2).

Example: Magnesium (Z = 12): $1s^2 2s^2 2p^6 3s^2$ or [Ne] $3s^2$

Electron Configurations and the Periodic Table

Periodic Table Blocks and Valence Electrons

The periodic table is divided into blocks (s, p, d, f) based on the subshell being filled. Elements in the same group have similar valence shell configurations.

• Valence shell: The outermost electron shell.

• Valence electrons: Electrons in the valence shell, important for chemical bonding.

Example: Sodium (Na): $1s^2 2s^2 2p^6 3s^1$; one valence electron in 3s.

Electron-Dot Symbols

Lewis Dot Symbols

Electron-dot (Lewis) symbols use dots around the atomic symbol to represent valence electrons. These are useful for visualizing bonding and electron sharing.

• Each dot represents one valence electron.

• Dots are placed on four sides of the symbol; pairs are formed as needed.

Example: Nitrogen (Group 5A): five dots around the symbol N.

Table: Comparison of Subatomic Particles

Particle	Symbol	Mass (amu)	Charge
Proton	p	1.007276	+1
Neutron	n	1.008665	0
Electron	e-	0.0005486	-1

Additional info:

• Periodic trends such as atomic radius, ionization energy, and electronegativity are also important but not fully covered in the provided material.

• Electron configuration shorthand uses noble gas symbols to simplify notation for larger atoms.