Atoms and the Periodic Table: Structured Study Notes

Atoms and the Periodic Table

2.1 Atomic Theory and the Structure of Atoms

The atomic theory forms the foundation of modern chemistry, describing the nature and behavior of atoms, the smallest identifiable units of elements.

• Atomic Theory Assumptions:

  • All matter is composed of atoms.

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

  • Chemical compounds consist of atoms combined in specific ratios; only whole atoms combine.

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

• Subatomic Particles:

  • Proton: Positive charge, mass ≈ 1 amu.

  • Neutron: No charge, mass ≈ 1 amu.

  • Electron: Negative charge, mass ≈ 1/1836 amu.

• Atomic Mass Unit (amu): Defined relative to the mass of a carbon-12 atom.

• Nucleus: Dense core containing protons and neutrons; electrons move rapidly in the surrounding space.

• Electrostatic Forces: Opposite charges attract; like charges repel.

Particle	Charge	Relative Mass (amu)
Proton	+1	1
Neutron	0	1
Electron	-1	1/1836

2.2 Elements and Atomic Number

Each element is defined by its atomic number, which is the number of protons in its nucleus. The mass number is the sum of protons and neutrons.

• Atomic Number (Z): Number of protons; unique to each element.

• Mass Number (A): Sum of protons and neutrons.

• Neutral Atoms: Number of electrons equals number of protons.

Example: Phosphorus (Z = 15, A = 31) has 15 protons, 15 electrons, and 16 neutrons (A - Z = 31 - 15 = 16).

2.3 Isotopes and Atomic Weight

Isotopes are atoms of the same element (same Z) with different mass numbers (A). Atomic weight is the weighted average mass of an element's isotopes.

• Isotope Notation: $^A_Z \text{Element}$ (e.g., $^3_1 \text{H}$ for tritium).

• Atomic Weight Calculation: Weighted average based on isotope masses and abundances.

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

Formula:

$\text{Atomic Weight} = \sum (\text{Isotope Mass} \times \text{Fractional Abundance})$

For gallium: $(68.9257 \times 0.604) + (70.9248 \times 0.396) = 69.723$ amu

2.4 The Periodic Table

The periodic table organizes elements by increasing atomic number and groups elements with similar properties together.

• Classification:

  • Metals: Left side; malleable, lustrous, good conductors.

  • Nonmetals: Upper-right; poor conductors.

  • Metalloids: Zigzag band; intermediate properties.

• Groups:

  • Main group elements

  • Transition metals

  • Inner transition metals

2.5 Characteristics of Different Groups

Elements in the same group share similar chemical properties due to periodicity.

• Group 1A (Alkali Metals): Shiny, soft, highly reactive, never found pure in nature.

• Group 2A (Alkaline Earth Metals): Lustrous, silvery, less reactive than 1A, never found pure.

• Group 7A (Halogens): Colorful, corrosive nonmetals, found only in compounds.

• Group 8A (Noble Gases): Colorless, chemically inert gases.

2.6 Electronic Structure of Atoms

The arrangement of electrons in shells, subshells, and orbitals determines the chemical properties of elements.

• Quantum Mechanical Model: Electrons have quantized energy levels described by wave functions.

• Shells: Groupings of electrons by energy; farther shells hold more electrons and have higher energy.

• Subshells: Types: s, p, d, f (increasing energy).

• Orbitals: Regions where electrons are likely found; s (1), p (3), d (5), f (7) orbitals per subshell.

• Each orbital holds two electrons with opposite spins.

Subshell	Number of Orbitals	Maximum Electrons
s	1	2
p	3	6
d	5	10
f	7	14

2.7 Electron Configurations

Electron configuration describes the distribution of electrons among orbitals in an atom.

• Rule 1: Electrons fill lowest energy orbitals first.

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

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

• Notation: Subshell with superscript (e.g., 1s2).

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

2.8 Electron Configurations and the Periodic Table

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: Outermost shell; contains valence electrons.

• Valence Electrons: Electrons in the valence shell; determine chemical reactivity.

Group	Valence-Shell Configuration
1A	ns1
2A	ns2
7A	ns2np5
8A	ns2np6

2.9 Electron-Dot Symbols (Lewis Symbols)

Electron-dot symbols visually represent the number of valence electrons around an element's symbol.

• Lewis Dot Symbol: Dots placed around the atomic symbol indicate valence electrons.

• Distribution: First four electrons are placed singly; additional electrons pair up.

Element	Group	Valence Electrons	Lewis Symbol
Na	1A	1	Na•
Cl	7A	7	Cl•••••••
O	6A	6	O••••••
N	5A	5	N•••••

Example: Group 5A element (e.g., nitrogen): Five dots, with four singly placed and one paired.

Additional info:

• Periodic trends such as atomic radius, ionization energy, and electronegativity are related to electron configuration and group membership.

• Shorthand electron configuration uses noble gas symbols to represent filled inner shells (e.g., [Ne]3s2).