Atoms, Molecules, and Ions

2.1 Atomic Theory of Matter

The concept of the atom as the fundamental building block of matter has evolved from ancient philosophy to a scientific theory. John Dalton's atomic theory, developed in the early 1800s, was based on several key experimental laws:

• Law of Constant Composition: Compounds always have the same relative number of atoms of each element, regardless of the sample source. Discovered by Joseph Proust.

• Law of Conservation of Mass: The total mass remains unchanged during a chemical reaction. Attributed to Antoine Lavoisier.

• Law of Multiple Proportions: When two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in small whole-number ratios. Discovered by John Dalton.

Dalton's atomic theory postulated:

• Elements are composed of tiny particles called atoms.

• All atoms of a given element are identical in mass and properties, but differ from those of other elements.

• Atoms are neither created nor destroyed in chemical reactions.

• Atoms of different elements combine in simple ratios to form compounds.

2.2 Discovery of Subatomic Particles

While Dalton considered atoms indivisible, later experiments revealed that atoms are composed of smaller subatomic particles:

• Electrons: Discovered by J. J. Thomson through cathode ray experiments, showing atoms contain negatively charged particles.

• Millikan Oil-Drop Experiment: Robert Millikan measured the charge of the electron, allowing calculation of its mass.

• Radioactivity: Henri Becquerel, and later Marie and Pierre Curie, discovered that atoms can emit high-energy radiation, indicating internal structure.

• Types of Radiation (Rutherford): Alpha (α, positive), beta (β, negative), and gamma (γ, neutral) rays.

• Plum Pudding Model: J. J. Thomson proposed atoms as spheres of positive charge with embedded electrons.

• Discovery of the Nucleus: Ernest Rutherford's gold foil experiment showed that atoms have a small, dense, positively charged nucleus.

2.3 Modern View of Atomic Structure

Atoms consist of a dense nucleus containing protons and neutrons, surrounded by electrons in mostly empty space. Key subatomic particles:

• Protons: Positive charge (+1), mass ≈ 1 amu, located in the nucleus.

• Neutrons: No charge, mass ≈ 1 amu, located in the nucleus.

• Electrons: Negative charge (–1), negligible mass, orbit the nucleus.

Atomic Number (Z): Number of protons in the nucleus; defines the element.

Mass Number (A): Total number of protons and neutrons in the nucleus.

Isotopes: Atoms of the same element with different numbers of neutrons (and thus different mass numbers).

2.4 Atomic Mass Unit (amu) and Atomic Weight

Atomic masses are extremely small, so chemists use the atomic mass unit (amu) as a standard. One amu is defined as one-twelfth the mass of a carbon-12 atom.

• Atomic Weight: The weighted average mass of all isotopes of an element, based on their natural abundance.

• Measured using a mass spectrometer.

Equation for average atomic weight:

2.5 The Periodic Table

The periodic table organizes elements by increasing atomic number. Elements in the same column (group) have similar chemical properties, while rows are called periods.

• Groups: Vertical columns; elements share chemical properties.

• Periods: Horizontal rows; properties repeat periodically.

• Metals: Left side; shiny, conductive, mostly solids.

• Nonmetals: Right side; can be solids, liquids, or gases.

• Metalloids: Border the stair-step line; have mixed properties.

2.6 Molecules and Molecular Compounds

Molecules are groups of atoms bonded together. Molecular compounds are composed of molecules and usually contain only nonmetals.

• Chemical Formula: Indicates the number and type of atoms in a molecule (e.g., H2O).

• Diatomic Molecules: Seven elements naturally occur as diatomic molecules: H2, N2, O2, F2, Cl2, Br2, I2.

• Empirical Formula: Simplest whole-number ratio of atoms.

• Molecular Formula: Actual number of atoms of each element.

• Structural Formula: Shows how atoms are connected.

2.7 Ions and Ionic Compounds

Ions are charged particles formed when atoms gain or lose electrons. Ionic compounds are formed from the electrostatic attraction between cations and anions.

• Cations: Positively charged ions (formed by loss of electrons, usually metals).

• Anions: Negatively charged ions (formed by gain of electrons, usually nonmetals).

• Ionic Compounds: Typically formed between metals and nonmetals; formula represents the lowest whole-number ratio of ions.

2.8 Naming Inorganic Compounds

Chemical nomenclature provides systematic names for compounds:

• Ionic Compounds: Name cation first, then anion. If cation has multiple charges, indicate with Roman numerals (e.g., iron(III) chloride).

• Polyatomic Ions: Groups of atoms with a charge (e.g., NH4+, SO42–).

• Oxyanions: Ions containing oxygen. Fewer oxygens: -ite; more oxygens: -ate (e.g., nitrite NO2–, nitrate NO3–).

• Acids: If anion ends in -ide, acid name begins with hydro- and ends in -ic acid (e.g., HCl: hydrochloric acid). If anion ends in -ate, acid ends in -ic acid; if -ite, ends in -ous acid.

• Binary Molecular Compounds: Use prefixes to indicate number of atoms (mono-, di-, tri-, etc.), and change ending of second element to -ide (e.g., CO2: carbon dioxide).

2.9 Some Simple Organic Compounds

Organic chemistry focuses on carbon-containing compounds. The simplest are alkanes (hydrocarbons with only single bonds). Alcohols are derived from alkanes by replacing a hydrogen with an –OH group, and their names end in -ol (e.g., methanol, ethanol).

• Isomers: Compounds with the same molecular formula but different structures (e.g., 1-propanol vs. 2-propanol).