Atomic Structure, Chemical Compounds, and Nomenclature: Study Notes

Atomic Structure and Subatomic Particles

Introduction

Understanding the structure of the atom and its constituent particles is fundamental to general chemistry. This section covers the basic components of atoms, their properties, and how they relate to the periodic table.

• Atom: The smallest unit of an element that retains the chemical properties of that element.

• Subatomic Particles: Atoms are composed of three main subatomic particles:

  • Protons (p+): Positively charged particles found in the nucleus. The number of protons defines the atomic number (Z) and the identity of the element.

  • Neutrons (n0): Neutral particles also located in the nucleus. Neutrons contribute to the mass of the atom but not its charge.

  • Electrons (e-): Negatively charged particles that orbit the nucleus in electron clouds or shells.

• Atomic Number (Z): The number of protons in the nucleus of an atom.

• Mass Number (A): The total number of protons and neutrons in the nucleus.

Example: Carbon has 6 protons, 6 neutrons, and 6 electrons in its most common isotope.

Isotope Notation

Introduction

Isotopes are atoms of the same element with different numbers of neutrons. Isotope notation is used to represent specific isotopes of an element.

• Isotope: Atoms of the same element (same Z) with different mass numbers (A) due to varying numbers of neutrons.

• Isotope Notation: Written as $^A_Z\text{X}$, where X is the element symbol, A is the mass number, and Z is the atomic number.

Example: $^{14}_6\text{C}$ represents carbon-14, with 6 protons and 8 neutrons.

Ions

Introduction

Ions are charged species formed when atoms gain or lose electrons. They play a crucial role in chemical reactions and compound formation.

• Cation: A positively charged ion formed when an atom loses one or more electrons.

• Anion: A negatively charged ion formed when an atom gains one or more electrons.

• Ion Notation: The charge is indicated as a superscript (e.g., Na+, Cl-).

Example: Sodium (Na) loses one electron to form Na+; chlorine (Cl) gains one electron to form Cl-.

The Periodic Table

Introduction

The periodic table organizes elements based on their atomic number and recurring chemical properties. It is a key tool for understanding element relationships and predicting chemical behavior.

• Groups: Vertical columns; elements in the same group have similar chemical properties.

• Periods: Horizontal rows; properties change progressively across a period.

• Main Group Elements: Groups 1, 2, and 13-18; also called representative elements.

• Transition Metals: Groups 3-12; elements with partially filled d subshells.

Example: Group 1 elements (alkali metals) are highly reactive and form +1 cations.

Ionic and Molecular Compounds

Introduction

Chemical compounds are classified based on the types of elements involved and the nature of their bonding. The two main types are ionic and molecular (covalent) compounds.

• Ionic Compounds: Formed from the electrostatic attraction between cations (usually metals) and anions (usually nonmetals).

• Molecular (Covalent) Compounds: Formed when two or more nonmetals share electrons.

Example: NaCl (sodium chloride) is an ionic compound; H2O (water) is a molecular compound.

Empirical and Molecular Formulas

Introduction

Formulas represent the composition of compounds. The empirical formula shows the simplest whole-number ratio of atoms, while the molecular formula shows the actual number of atoms in a molecule.

• Empirical Formula: The simplest ratio of elements in a compound.

• Molecular Formula: The actual number of each type of atom in a molecule.

• Relationship: The molecular formula is a whole-number multiple of the empirical formula.

Example: The empirical formula of hydrogen peroxide is HO; the molecular formula is H2O2.

Inorganic Nomenclature

Introduction

Naming inorganic compounds follows systematic rules to ensure clarity and consistency. The nomenclature depends on the type of compound (ionic or molecular) and the elements involved.

• Ionic Compounds:

  • Name the cation first, then the anion.

  • For metals with variable charges, indicate the charge with Roman numerals (e.g., FeCl2 is iron(II) chloride).

  • For polyatomic ions, use the standard ion name (e.g., NaNO3 is sodium nitrate).

• Molecular Compounds:

  • Use prefixes to indicate the number of each atom (mono-, di-, tri-, etc.).

  • The more electronegative element is named last and ends with “-ide.”

  • Example: CO2 is carbon dioxide; N2O4 is dinitrogen tetroxide.

Example Table: Common Prefixes for Molecular Compounds

Number	Prefix
1	mono-
2	di-
3	tri-
4	tetra-
5	penta-
6	hexa-
7	hepta-
8	octa-
9	nona-
10	deca-