Atoms and the Periodic Table: Study Notes for GOB Chemistry

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 atoms, the smallest identifiable units of elements.

• Atomic Theory Assumptions:

  • 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, not 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 mass approximately 1/1836 that of a proton, 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 are held near the nucleus), like charges repel (electrons spread apart).

Example: The nucleus is to the atom as a pea is to a stadium in terms of relative size.

2.2 Elements and Atomic Number

Each element is defined by its atomic number, which is the number of protons in its atoms. The mass number is the sum of protons and neutrons.

• 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: Neutral carbon has 6 protons and 6 electrons.

Worked Example: For phosphorus (Z = 15, A = 31): 15 protons, 15 electrons, 16 neutrons (A - Z = 31 - 15 = 16).

2.3 Isotopes and Atomic Weight

Isotopes are atoms of the same element with different numbers of neutrons, resulting in different mass numbers. Atomic weight is the weighted average of all naturally occurring isotopes.

• Isotopes: Same atomic number, different mass numbers (e.g., hydrogen: protium, deuterium, tritium).

• Isotope Notation: $^A_ZX$ (A = mass number, Z = atomic number, X = element symbol).

• Atomic Weight Calculation: Weighted average based on isotope masses and their natural abundances.

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

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

• Classification by Physical Properties:

  • Metals: Left side; malleable, lustrous, good conductors.

  • Nonmetals: Upper-right; poor conductors.

  • Metalloids: Zigzag band; intermediate properties.

• 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) have similar chemical properties due to similar valence electron configurations.

• Periodicity: Properties repeat in a predictable pattern across periods.

• Group 1A – Alkali Metals: Li, Na, K, Rb, Cs, Fr; soft, shiny, highly reactive, not found pure in nature.

• Group 2A – Alkaline Earth Metals: Be, Mg, Ca, Sr, Ba, Ra; lustrous, less reactive than 1A, not found pure in nature.

• 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, inert gases, very low reactivity.

2.6 Electronic Structure of Atoms

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

• Quantum Mechanical Model: Electrons have quantized energy levels and are described by wave functions.

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

• Subshells: s, p, d, f (increasing energy); each shell contains one or more subshells.

• Orbitals: Regions within subshells where electrons are likely found; s (1 orbital), p (3), d (5), f (7).

• Electron Spin: Each orbital holds two electrons with opposite spins.

Example: The first shell (1s) holds 2 electrons; the second shell (2s, 2p) holds 8 electrons.

2.7 Electron Configurations

Electron configuration describes the arrangement of electrons in an atom’s orbitals.

• Rules for Electron Configuration:

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

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

  3. Orbitals of equal energy are half-filled before any is completely 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 filled inner shells (e.g., [Ne] 3s2).

Example: Magnesium (Mg, Z = 12): 1s2 2s2 2p6 3s2

2.8 Electron Configurations and the Periodic Table

The periodic table reflects the filling order of electron subshells and helps identify valence electrons.

• Blocks: s-block, p-block, d-block, 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 the same number of valence electrons.

Example: Sodium (Na): 1 valence electron in 3s; Chlorine (Cl): 7 valence electrons (3s2 3p5).

2.9 Electron-Dot Symbols (Lewis Symbols)

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

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

• Group Number: Main group number (A) indicates the number of valence electrons (e.g., Group 5A elements have 5 dots).

Example: Nitrogen (N, Group 5A): N with five dots arranged singly before pairing.

Table: Comparison of Subatomic Particles

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

Table: Electron Distribution in Shells (Example: Tin, Sn)

Shell Number	1	2	3	4	5
Number of Electrons	2	8	18	18	4

Table: Electron-Dot Symbols for Main Group Elements (Selected)

Element	Group	Valence Electrons	Electron-Dot Symbol
Li	1A	1	Li•
Be	2A	2	Be••
B	3A	3	B•••
C	4A	4	C••••
N	5A	5	N•••••
O	6A	6	O••••••
F	7A	7	F•••••••
Ne	8A	8	Ne••••••••

Key Equations

• Number of Neutrons:

  $\text{Number of neutrons} = A - Z$

• Atomic Weight (Weighted Average):

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

Summary Table: Valence-Shell Electron Configurations for Main Groups

Group	General Valence-Shell Configuration
1A	ns1
2A	ns2
6A	ns2 np4
7A	ns2 np5
8A	ns2 np6

Additional info: Where the original content was incomplete or referenced figures/tables, standard academic context and logical inferences were added to ensure completeness and clarity for GOB Chemistry students.