Atoms and the Periodic Table: Structure, Properties, and Electron Configurations

Study Guide - Smart Notes

Tailored notes based on your materials, expanded with key definitions, examples, and context.

Atomic Theory and the Structure of Atoms

Foundations of Atomic Theory

The atomic theory forms the basis of modern chemistry, describing the nature and behavior of matter at the atomic level.

All matter is composed of atoms, the smallest identifiable units of elements.
Atoms of a given element are unique and differ from those of other elements.
Chemical compounds consist of atoms combined in specific, whole-number ratios.
Chemical reactions rearrange how atoms are combined, but do not change the atoms themselves.

Subatomic Particles

Protons: Positively charged particles found in the nucleus; mass ≈ 1 atomic mass unit (amu).
Neutrons: Electrically neutral particles in the nucleus; mass ≈ 1 amu (slightly more than a proton).
Electrons: Negatively charged particles moving around the nucleus; mass ≈ 1/1836 that of a proton.

The nucleus is a dense core containing protons and neutrons, while electrons occupy the surrounding space. Opposite charges attract, so electrons are held near the nucleus, but like charges repel, causing electrons to spread out.

Comparison of Subatomic Particles

Particle	Symbol	Charge	Relative Mass (amu)
Proton	p+	+1	1
Neutron	n0	0	1
Electron	e-	-1	1/1836

Elements and Atomic Number

Atomic Number and Mass Number

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

Atomic number (Z): Number of protons (and, in a neutral atom, electrons).
Mass number (A): Number of protons + number of neutrons.

Atoms are electrically neutral, so the number of electrons equals the number of protons.

Example Calculations

Phosphorus (Z = 15, A = 31): 15 protons, 15 electrons, 16 neutrons (31 - 15 = 16).
Nickel (Z = 28, A = 60): 28 protons, 28 electrons, 32 neutrons (60 - 28 = 32).

Isotopes and Atomic Weight

Isotopes

Isotopes are atoms of the same element (same Z) with different numbers of neutrons (different A).

Example: Hydrogen has three isotopes—protium (1H), deuterium (2H), and tritium (3H).
Isotopes are represented as AZX, where X is the element symbol.

Atomic Weight

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

Calculated using the formula:

Example: Gallium has two isotopes, Ga-69 (68.9257 amu, 60.4%) and Ga-71 (70.9248 amu, 39.6%).
Atomic weight = (0.604 × 68.9257) + (0.396 × 70.9248) = 69.72 amu

The Periodic Table

Structure and Classification

The periodic table organizes elements by increasing atomic number and groups elements with similar properties into columns (groups or families).

Metals: Left side; shiny, malleable, good conductors.
Nonmetals: Upper right; poor conductors, often gases or brittle solids.
Metalloids: Zigzag band; properties intermediate between metals and nonmetals.

Element Groups

Group	Name	Properties
1A	Alkali metals	Shiny, soft, highly reactive, not found pure in nature
2A	Alkaline earth metals	Lustrous, silvery, less reactive than 1A, not found pure in nature
7A	Halogens	Colorful, corrosive nonmetals, form salts, found in compounds
8A	Noble gases	Colorless, inert gases, very low reactivity

Electronic Structure of Atoms

Quantum Mechanical Model

Electrons occupy specific energy levels (shells), subshells, and orbitals around the nucleus. Their arrangement determines chemical properties.

Shells: Main energy levels (n = 1, 2, 3, ...).
Subshells: s, p, d, f (increasing energy).
Orbitals: Regions within subshells where electrons are likely found; each orbital holds up to 2 electrons with opposite spins.

Number of orbitals per subshell:

s: 1 orbital (2 electrons)
p: 3 orbitals (6 electrons)
d: 5 orbitals (10 electrons)
f: 7 orbitals (14 electrons)

Example

An atom with filled first and second shells and four electrons in the third shell has 2 + 8 + 4 = 14 electrons, corresponding to silicon (Si).

Electron Configurations

Rules for Electron Configuration

Aufbau Principle: Electrons fill lowest energy orbitals first.
Pauli Exclusion Principle: Each orbital holds two electrons with opposite spins.
Hund's Rule: Orbitals of equal energy are singly occupied before any is doubly occupied.

Electron configurations are written by listing the occupied orbitals and the number of electrons in each as a superscript.

Examples

Magnesium (Mg, Z = 12): 1s2 2s2 2p6 3s2
Phosphorus (P, Z = 15): 1s2 2s2 2p6 3s2 3p3

Shorthand notation uses the previous noble gas in brackets, e.g., [Ne] 3s2 3p3 for phosphorus.

Electron Configurations and the Periodic Table

Periodic Table Blocks and Valence Electrons

The periodic table is divided into s, p, d, and f blocks, corresponding to the subshell being filled. Elements in the same group have similar valence electron configurations, which determine their chemical properties.

Valence shell: Outermost shell containing electrons.
Valence electrons: Electrons in the valence shell; involved in chemical bonding.

Valence-Shell Configurations (Selected Groups)

Group	General Valence-Shell Configuration
1A	ns1
2A	ns2
7A	ns2 np5
8A	ns2 np6

Example

Tin (Sn, Z = 50): Shells 1–5 have 2, 8, 18, 18, and 4 electrons, respectively. Tin has four valence electrons in the 5s and 5p subshells.

Electron-Dot (Lewis) Symbols

Definition and Construction

An electron-dot symbol (Lewis symbol) represents the valence electrons of an atom as dots around the element symbol.

Dots are placed singly on each side before pairing.
The number of dots equals the number of valence electrons (group number for main group elements).

Examples

Element	Group	Valence Electrons	Lewis Symbol
Na	1A	1	Na·
Cl	7A	7	Cl·······
O	6A	6	O······

For group 5A elements, five dots are placed: one on each side, then pairing begins.

Additional info: The above notes expand on the provided outline and examples, filling in missing details for clarity and completeness, and reconstructing tables and diagrams as described in the source material.