Atoms, Molecules, and Ions

Atomic Theory of Matter

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

• Law of Constant Composition: Compounds always contain the same proportion of elements by mass. Discovered by Joseph Proust, this law states that the relative number of atoms of each element in a compound is constant.

• Law of Conservation of Mass: The total mass of substances remains unchanged during a chemical reaction. Antoine Lavoisier formulated this law, which asserts that atoms are neither created nor destroyed in chemical reactions.

• 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. John Dalton discovered this law, highlighting the discrete nature of atoms.

Dalton’s Atomic Theory Postulates:

1. Each element is composed of extremely small particles called atoms.

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

3. Atoms are not changed into different elements by chemical reactions; they are neither created nor destroyed.

4. Atoms of different elements combine in simple whole-number ratios to form compounds.

Discovery of Subatomic Particles

Dalton’s model of the atom was later refined as scientists discovered that atoms are composed of smaller particles:

• Electrons and Cathode Rays: J. J. Thomson discovered electrons as negatively charged particles using cathode ray tubes (1897).

• Millikan Oil-Drop Experiment: Robert Millikan measured the charge of the electron (1909), allowing calculation of its mass.

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

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

• Plum Pudding Model: J. J. Thomson proposed that electrons were embedded in a positively charged sphere.

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

Modern View of Atomic Structure

Atoms consist of a dense nucleus containing protons and neutrons, surrounded by electrons in mostly empty space. The nucleus contains nearly all the atom’s mass.

• Protons: Positively charged particles in the nucleus (relative mass ≈ 1).

• Neutrons: Neutral particles in the nucleus (relative mass ≈ 1).

• Electrons: Negatively charged particles outside the nucleus (relative mass ≈ 0).

Table: Comparison of Subatomic Particles

Atomic Number, Mass Number, and Isotopes

• 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 masses).

Example: Carbon has isotopes such as 12C, 13C, and 14C, all with 6 protons but different numbers of neutrons.

Atomic Mass Unit (amu) and Atomic Weight

Atomic masses are measured in atomic mass units (amu), where 1 amu is defined as 1/12 the mass of a 12C atom. The atomic weight of an element is the weighted average of the masses of its naturally occurring isotopes.

• Mass Spectrometry: Used to measure atomic and molecular weights and determine isotope abundances.

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; intermediate properties.

Table: Names of Some Groups in the Periodic Table

Molecules and Molecular Compounds

Molecules are groups of atoms held together by chemical bonds. Molecular compounds typically contain only nonmetals.

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

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

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

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

• Structural Formula: Shows how atoms are connected.

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 (positive) and anions (negative).

• Cations: Formed by loss of electrons (usually metals).

• Anions: Formed by gain of electrons (usually nonmetals).

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

Naming Inorganic Compounds

Chemical nomenclature provides systematic rules for naming compounds:

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

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

• Oxyanions: For two ions of the same element, the one with fewer oxygens ends in -ite, the one with more in -ate (e.g., nitrite NO2−, nitrate NO3−).

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

• Binary Molecular Compounds: Use prefixes to indicate the number of atoms (e.g., CO2: carbon dioxide; N2O5: dinitrogen pentoxide).

Table: Prefixes Used in Naming Binary Compounds

Simple Organic Compounds

Organic chemistry focuses on compounds containing carbon. 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).