Atoms and the Periodic Table: Fundamentals for GOB Chemistry

Atoms and the Periodic Table

Atomic Theory and the Structure of Atoms

The study of chemistry is grounded in the understanding of atoms, the smallest units of matter that retain the identity of an element. The modern atomic theory provides a framework for understanding the composition and behavior of matter.

• Atom: The smallest particle of an element that retains its chemical identity. The term comes from the Greek atomos, meaning "indivisible."

• Atomic Theory: Chemistry is based on four main assumptions:

  1. All matter is composed of atoms.

  2. The atoms of a given element differ from those of all other elements.

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

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

• Subatomic Particles: Atoms are composed of protons, neutrons, and electrons.

  • Proton: Positively charged particle found in the nucleus.

  • Neutron: Electrically neutral particle with a mass similar to a proton, also in the nucleus.

  • Electron: Negatively charged particle with a much smaller mass, found outside the nucleus.

Example Table: Comparison of Subatomic Particles

Particle	Symbol	Mass (amu)	Charge
Proton	p	1.007	+1
Neutron	n	1.008	0
Electron	e-	0.00055	-1

Atomic Mass Unit (amu): A unit for describing the mass of atoms, defined as one-twelfth the mass of a carbon-12 atom.

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

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

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

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

Isotopes and Atomic Weight

Isotopes are atoms of the same element with different numbers of neutrons, resulting in different mass numbers. The atomic weight of an element is the weighted average of the masses of its naturally occurring isotopes.

• Isotope: Atoms with the same atomic number but different mass numbers.

• Notation: Isotopes are represented as $^A_ZX$, where A is the mass number, Z is the atomic number, and X is the element symbol.

• Atomic Weight Calculation: $\text{Atomic weight} = \sum [(\text{isotopic abundance}) \times (\text{isotopic mass})]$

Example: Gallium has two isotopes: 60.4% Ga-69 (68.9257 amu) and 39.6% Ga-71 (70.9248 amu). $\text{Atomic weight} = (0.604 \times 68.9257) + (0.396 \times 70.9248) = 69.7 \text{ amu}$

The Periodic Table

The periodic table organizes elements by increasing atomic number and similar chemical properties. Elements are classified as metals, nonmetals, or metalloids based on their physical and chemical properties.

• Metal: Malleable, lustrous, good conductor of heat and electricity; found on the left side of the table.

• Nonmetal: Poor conductor, often gases or brittle solids; found on the upper-right side.

• Metalloid: Properties intermediate between metals and nonmetals; found in a zigzag band between metals and nonmetals.

Groups: Vertical columns with similar chemical properties. Main groups include:

• Group 1A (Alkali metals): Very reactive, soft, low melting points.

• Group 2A (Alkaline earth metals): Less reactive than alkali metals, silvery.

• Group 7A (Halogens): Colorful, corrosive nonmetals.

• Group 8A (Noble gases): Colorless, unreactive gases.

Electronic Structure of Atoms

The arrangement of electrons in an atom determines its chemical properties. Electrons occupy shells, subshells, and orbitals according to specific rules.

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

• Subshell: Subdivision of shells (s, p, d, f) with increasing energy.

• Orbital: Region of space where electrons are most likely to be found; each orbital holds two electrons with opposite spins.

Shell Capacities:

Shell (n)	Electron Capacity
1	2
2	8
3	18
4	32

Subshell Capacities:

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

Electron Configurations

Electron configuration describes the arrangement of electrons in an atom's shells and subshells. Electrons fill the lowest energy orbitals first (Aufbau principle), each orbital holds two electrons (Pauli exclusion principle), and electrons fill degenerate orbitals singly before pairing (Hund's rule).

• Notation: The number of electrons in each subshell is indicated by a superscript (e.g., 1s2 2s2 2p6).

• Shorthand: Use the previous noble gas in brackets to represent core electrons (e.g., [Ne] 3s1 for sodium).

Example: Magnesium (Z = 12): Full: 1s2 2s2 2p6 3s2 Shorthand: [Ne] 3s2

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 electron configurations, which determine their chemical properties.

• Valence Shell: The outermost electron shell of an atom.

• Valence Electrons: Electrons in the valence shell; important for chemical bonding.

Example: Group 6A elements have a general valence-shell configuration of ns2np4.

Electron-Dot (Lewis) Symbols

Electron-dot symbols (Lewis symbols) use dots around the atomic symbol to represent valence electrons. These symbols help visualize the number of electrons available for bonding.

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

Example: Nitrogen (Group 5A) has five valence electrons, so its Lewis symbol is N with five dots arranged around it.