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 are unique and differ from those of other elements.

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

  • Chemical reactions rearrange atoms but do not change the atoms themselves.

• Subatomic Particles:

  • Protons: Positively charged particles found in the nucleus.

  • Neutrons: Electrically neutral particles with a mass similar to protons, also in the nucleus.

  • Electrons: Negatively charged particles with a much smaller mass, moving rapidly around the nucleus.

• Atomic Mass Unit (amu): A unit for atomic mass, defined by the mass of a carbon-12 atom.

• Nucleus: Dense core containing protons and neutrons; electrons occupy the surrounding space.

• Electrostatic Forces: Opposite charges attract (electrons to nucleus), like charges repel (electrons from each other).

2.2 Elements and Atomic Number

Each element is defined by its atomic number, which determines its chemical identity and properties.

• Atomic Number (Z): Number of protons in the nucleus; unique for each element.

• Mass Number (A): Total number of protons and neutrons in an atom.

• Neutral Atoms: Number of electrons equals number of protons, resulting in no net charge.

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

• Calculating Neutrons: $\text{Number of neutrons} = A - Z$

2.3 Isotopes and Atomic Weight

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

• Isotopes: Same atomic number (Z), different mass number (A).

• Notation: $^A_Z\text{X}$, where X is the element symbol.

• Examples: Hydrogen has three isotopes: protium, deuterium, and tritium.

• Atomic Weight: Weighted average of the masses of all naturally occurring isotopes of an element.

• Calculation: $\text{Atomic weight} = \sum (\text{isotope mass} \times \text{fractional abundance})$

• Example: Gallium has two isotopes: 60.4% Ga-69 (68.9257 amu) and 39.6% Ga-71 (70.9248 amu). The atomic weight is calculated as: $\text{Atomic weight} = (0.604 \times 68.9257) + (0.396 \times 70.9248)$

2.4 The Periodic Table

The periodic table organizes elements by increasing atomic number and recurring chemical properties.

• Structure: Each element is represented by its symbol, atomic number, and atomic mass.

• Classification by Physical Properties:

  • 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.

• Classification by Chemical Behavior:

  • Main group elements (Groups 1A–8A)

  • Transition metals (Groups 3B–2B)

  • Inner transition metals (lanthanides and actinides)

2.5 Some Characteristics of Different Groups

Elements in the same group (vertical column) share similar chemical properties due to their valence electron configurations.

• Periodicity: Properties of elements show repeating patterns (periodicity) across periods.

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

• Group 2A – Alkaline Earth Metals: Lustrous, silvery, less reactive than alkali metals, not found pure in nature.

• Group 7A – Halogens: Colorful, corrosive nonmetals, found in compounds (e.g., NaCl).

• Group 8A – Noble Gases: Colorless, inert gases; very low chemical reactivity.

2.6 Electronic Structure of Atoms

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

• Quantum Mechanical Model: Electrons have both particle and wave properties; their behavior is described by wave functions.

• Quantization: Electrons occupy discrete energy levels (shells), not continuous ranges.

• Shells: Groupings of electrons by energy level (n = 1, 2, 3, ...).

• Subshells: Within shells, electrons are further grouped into s, p, d, and f subshells (increasing energy).

• Orbitals: Regions within subshells where electrons are most likely found; each orbital holds up to two 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)

2.7 Electron Configurations

Electron configuration describes the arrangement of electrons in an atom’s shells and subshells.

• Rules for Electron Configuration:

  1. Electrons fill the lowest energy orbitals first (Aufbau principle).

  2. Each orbital holds a maximum of two electrons with opposite spins (Pauli exclusion principle).

  3. Orbitals of equal energy are each half-filled before any is fully filled (Hund’s rule).

• Notation: Subshells are labeled with the number of electrons as a superscript (e.g., 1s2 2s2 2p6).

• Shorthand Notation: Uses the previous noble gas in brackets to represent core electrons (e.g., [Ne] 3s2).

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

2.8 Electron Configurations and the Periodic Table

The periodic table’s structure reflects the filling of electron subshells and helps predict valence electron configurations.

• Blocks: s-block, p-block, d-block, and f-block correspond to the type of subshell being filled.

• Valence Shell: The outermost shell containing electrons; determines chemical reactivity.

• Valence Electrons: Electrons in the valence shell; elements in the same group have similar valence electron configurations.

• Example Table:

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

2.9 Electron-Dot Symbols (Lewis Symbols)

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

• Construction: Place one dot for each valence electron around the element symbol, pairing only after each side has one dot.

• Example Table:

Element	Group	Valence Electrons	Electron-Dot Symbol
Na	1A	1	Na•
Mg	2A	2	Mg••
Cl	7A	7	Cl•••••••
Ne	8A	8	Ne••••••••

• General Rule: The number of valence electrons equals the group number for main group elements (A groups).

Worked Examples Summary

• Calculating Subatomic Particles: Use atomic number for protons/electrons; subtract from mass number for neutrons.

• Isotope Notation: $^A_Z\text{X}$ (e.g., $^{12}_6\text{C}$ for carbon-12).

• Average Atomic Weight: Multiply each isotope’s mass by its fractional abundance and sum the results.

• Electron Configuration: Fill orbitals in order of increasing energy, following the three rules.

• Valence Electrons: Identify from electron configuration or group number.

• Electron-Dot Symbols: Place dots around the symbol to represent valence electrons.

Additional info: For more detailed electron configurations, refer to the periodic table and the order of orbital filling (1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc.). The periodic table’s structure is a direct result of the quantum mechanical arrangement of electrons in atoms.