Atoms and the Periodic Table: Foundations of GOB Chemistry

Chapter 2: Atoms and the Periodic Table

2.1 Atomic Theory and the Structure of Atoms

The atomic theory forms the foundation of chemistry, describing the nature and structure of matter at the atomic level. Atoms are the smallest identifiable units of elements, derived from the Greek word atomos, meaning "indivisible." Chemistry is based on several key assumptions about atoms and their behavior.

• 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.

Atoms are made up of three main 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 mass about 1/1836 that of a proton, moving rapidly around the nucleus.

The nucleus is extremely small compared to the overall size of the atom, analogous to a pea in a stadium. Opposite charges attract (protons and electrons), while like charges repel (proton-proton or electron-electron).

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

Atomic mass unit (amu): The unit for atomic mass, based on the mass of a carbon-12 atom.

2.2 Elements and Atomic 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.

• In a neutral atom, the number of electrons equals the number of protons.

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

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

Worked Example: An atom with 28 protons and A = 60 is Nickel (Ni), with 28 electrons and 32 neutrons (60 - 28 = 32).

2.3 Isotopes and Atomic Weight

Isotopes are atoms of the same element (same Z) but different mass numbers (A). For example, hydrogen has three isotopes: protium, deuterium, and tritium.

• Isotopes are represented as AZX, where X is the element symbol.

• Most elements are mixtures of isotopes.

Atomic weight is the weighted average mass of an element's atoms, calculated as:

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

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

Worked Example: 19478X is platinum (Pt), with 78 protons, 78 electrons, and 116 neutrons (194 - 78 = 116).

2.4 The Periodic Table

The periodic table organizes elements by increasing atomic number and similar chemical properties. Each box lists the atomic number, symbol, and atomic mass.

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

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

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

Elements in the same group (vertical column) have similar chemical properties and are classified as main group elements, transition metals, or inner transition metals.

2.5 Some Characteristics of Different Groups

Groups in the periodic table exhibit periodicity—repeating patterns in properties.

• Group 1A (Alkali metals): Li, Na, K, Rb, Cs, Fr; shiny, soft, highly reactive, never found pure in nature.

• Group 2A (Alkaline earth metals): Be, Mg, Ca, Sr, Ba, Ra; lustrous, less reactive than 1A, never found pure.

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

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

2.6 Electronic Structure of Atoms

The arrangement of electrons in atoms determines their properties. The quantum mechanical model describes electrons as occupying quantized energy levels (shells, subshells, orbitals).

• Shells: Principal energy levels (n = 1, 2, 3, ...), farther shells hold more electrons and have higher energy.

• Subshells: s, p, d, f (increasing energy).

• Orbitals: Regions where electrons are likely found; s (1 orbital), p (3), d (5), f (7); each orbital holds 2 electrons with opposite spins.

Shell Number	Subshells	Number of Orbitals	Electron Capacity
1	s	1	2
2	s, p	1, 3	8
3	s, p, d	1, 3, 5	18
4	s, p, d, f	1, 3, 5, 7	32

2.7 Electron Configurations

The electron configuration describes the arrangement of electrons in an atom's shells and subshells. Three rules guide electron configurations:

1. Electrons occupy the lowest energy orbitals available (Aufbau principle).

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

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

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

Phosphorus (Z = 15): $1s^2 2s^2 2p^6 3s^2 3p^3$

2.8 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, which determine their chemical properties.

• Valence shell: Outermost electron shell.

• Valence electrons: Electrons in the valence shell, involved in chemical bonding.

Example: Sodium (Na, Z = 11): $1s^2 2s^2 2p^6 3s^1$ or [Ne] $3s^1$ (1 valence electron)

Chlorine (Cl, Z = 17): $1s^2 2s^2 2p^6 3s^2 3p^5$ or [Ne] $3s^2 3p^5$ (7 valence electrons)

2.9 Electron-Dot Symbols

Electron-dot (Lewis) symbols represent the valence electrons of an atom as dots around the element symbol. These are useful for visualizing bonding and chemical reactivity.

Group	Valence Electrons	Lewis Symbol Example
1A	1	Na •
2A	2	Mg ••
5A	5	N •••••
8A	8	Ne ••••••••

Example: For group 5A elements, the Lewis symbol has five dots, with the first four distributed singly around the symbol and the fifth paired.