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 foundational to chemistry. The atom is the smallest particle that retains the identity of an element. The term atom comes from the Greek atomos, meaning "indivisible." Modern atomic theory is based on several key assumptions.

• All matter is composed of atoms. Atoms are the building blocks of all substances.

• Atoms of a given element are unique. Each element has atoms that differ from those of other elements.

• Chemical compounds consist of atoms combined in specific ratios. Only whole atoms can combine to form compounds.

• Chemical reactions rearrange atoms. Reactions change how atoms are combined, but do not change the atoms themselves.

Subatomic Particles

Atoms are made up of three main subatomic particles: protons, neutrons, and electrons.

• Protons: Positively charged particles found in the nucleus.

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

• Electrons: Negatively charged particles with a much smaller mass than protons or neutrons, found in the space surrounding the nucleus.

Example: The hydrogen atom contains one proton and one electron, but no neutrons.

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; determines the element.

• Mass Number (A): Number of protons plus neutrons.

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

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

Isotopes and Atomic Weight

Isotopes

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

• Hydrogen Isotopes: Protium (A = 1), Deuterium (A = 2), Tritium (A = 3).

• Isotopes are represented as $^A_ZX$ where A is mass number, Z is atomic number, and 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: $\text{Atomic weight} = \sum [(\text{isotopic abundance}) \times (\text{isotopic mass})]$

• Example: Gallium has two isotopes: Ga-69 (60.4%, 68.9257 amu) and Ga-71 (39.6%, 70.9248 amu). 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 elements by increasing atomic number and groups elements with similar properties together.

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

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

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

Groups and Periodicity

Elements in the same vertical column (group) have similar chemical properties. The periodic table shows periodicity, a repeating pattern of properties.

• Group 1A (Alkali metals): Li, Na, K, Rb, Cs, Fr; soft, reactive metals.

• Group 2A (Alkaline earth metals): Be, Mg, Ca, Sr, Ba, Ra; less reactive than alkali metals.

• Group 7A (Halogens): F, Cl, Br, I, At; colorful, corrosive nonmetals.

• Group 8A (Noble gases): He, Ne, Ar, Kr, Xe, Rn; colorless, chemically inert gases.

Electronic Structure of Atoms

Quantum Mechanical Model

The arrangement of electrons in atoms determines their chemical properties. Electrons occupy energy levels called shells, which are divided into subshells (s, p, d, f) and further into orbitals.

• Shells: Numbered 1, 2, 3, ...; farther shells hold more electrons and have higher energy.

• Subshells: s (1 orbital), p (3 orbitals), d (5 orbitals), f (7 orbitals).

• Orbitals: Each holds up to 2 electrons with opposite spins.

Example: The second shell (n=2) contains one 2s orbital and three 2p orbitals, for a total of 8 electrons.

Electron Configurations

Rules for Electron Configuration

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

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

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

• Rule 3: Orbitals of equal energy are half-filled before any is completely filled (Hund's rule).

Notation: The number of electrons in each subshell is shown as a superscript, e.g., $1s^2 2s^2 2p^6$.

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

Electron Configurations and the 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 electron configurations.

• Valence shell: The outermost shell of electrons.

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

• Example: Group 6A elements have a general valence-shell configuration of $ns^2 np^4$.

Electron-Dot Symbols

Lewis Dot Symbols

Electron-dot (Lewis) symbols use dots around the atomic symbol to represent valence electrons.

• Each dot represents one valence electron.

• Dots are placed on four sides of the symbol, pairing as needed.

• Example: Nitrogen (Group 5A) has five valence electrons, shown as five dots around the symbol N.

Summary Table: Subatomic Particles

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

Summary Table: Classification of Elements

Type	Location on Periodic Table	Properties
Metals	Left side	Malleable, lustrous, good conductors
Nonmetals	Upper right	Poor conductors, brittle
Metalloids	Zigzag band	Intermediate properties

Additional info: These notes expand on the brief points in the slides, providing definitions, examples, and formulas for key concepts in atomic theory and the periodic table, suitable for GOB Chemistry students.