Atoms and the Periodic Table: Structured Study Notes

Atoms and the Periodic Table

Atomic Theory and the Structure of Atoms

The foundation of chemistry is atomic theory, which explains the nature of matter and the structure of atoms. Atoms are the smallest units of elements that retain their identity.

• 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: Neutral, 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

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; also equals the number of electrons in a neutral atom.

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

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

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

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: $^A_Z \text{Element}$ (A = mass number, Z = atomic number)

• Examples: Hydrogen isotopes: protium ($^1_1\text{H}$), deuterium ($^2_1\text{H}$), tritium ($^3_1\text{H}$)

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

Formula:

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

Example: Gallium: 60.4% Ga-69 (68.9257 amu), 39.6% Ga-71 (70.9248 amu). $\text{Atomic Weight} = (0.604 \times 68.9257) + (0.396 \times 70.9248)$

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: Vertical columns; elements in the same group have similar chemical properties.

• Types of Groups:

  • Main group elements

  • Transition metals

  • Inner transition metals

Characteristics of Different Groups

Groups in the periodic table exhibit periodicity, a repeating pattern of properties.

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

• Group 2A (Alkaline Earth Metals): Lustrous, silvery, less reactive than 1A; not found pure in nature.

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

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

Electronic Structure of Atoms

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

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

• Shell: Grouping of electrons by energy level; 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 spin.

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

Electron Configurations

Electron configuration describes the distribution of electrons in an atom's orbitals. Three rules govern electron filling:

• Rule 1: Electrons occupy the 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: Subshells are written with a superscript indicating the number of electrons (e.g., 1s2).

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

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

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	Lewis Dot Symbol
Na	Na•
Cl	Cl•••••••
O	O••••••
N	N•••••
C	C••••

Example: Group 5A element (e.g., nitrogen): Five dots around the symbol, with the first four singly placed and the fifth paired.

Summary Table: Electron Distribution in Tin (Sn)

Shell Number	Number of Electrons
1	2
2	8
3	18
4	18
5	4

Valence electrons: 4 (in 5s and 5p subshells)

General Valence-Shell Configuration for Group 6A: $ns^2\ np^4$

Additional info: These notes expand on brief points by providing definitions, examples, and formulas for key concepts in atomic structure and periodicity, suitable for GOB Chemistry students.