Atoms, Molecules, and Ions

Atomic Theory of Matter

The concept of the atom as the fundamental building block of matter originated with ancient Greek philosophers, but was formalized by John Dalton in the early 1800s. Dalton's atomic theory was based on several key laws:

• Law of Constant Composition: Compounds have a definite composition; the relative number of atoms of each element in a compound is the same in any sample. Discovered by Joseph Proust.

• Law of Conservation of Mass: The total mass of substances present before and after a chemical process remains unchanged. Discovered by Antoine Lavoisier.

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

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 atoms of other elements.

3. Atoms are not changed into atoms of another element by chemical reactions; atoms are neither created nor destroyed in chemical reactions.

4. Atoms of different elements combine to form compounds; a given compound always has the same relative number and kind of atoms.

Discovery of Subatomic Particles

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

• Electrons: Discovered by J. J. Thomson through cathode ray experiments (1897). Electrons are negatively charged particles.

• Millikan Oil-Drop Experiment: Robert Millikan determined the charge of the electron in 1909, allowing calculation of its mass.

• Radioactivity: Spontaneous emission of high-energy radiation, first observed by Henri Becquerel and studied by Marie and Pierre Curie. Revealed the existence of more subatomic particles.

• Types of Radiation: Ernest Rutherford identified three types: alpha (α, positively charged), beta (β, negatively charged), and gamma (γ, uncharged).

• Plum Pudding Model: J. J. Thomson proposed a model with electrons embedded in a positive sphere.

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

Modern View of Atomic Structure

Atoms consist of a small, dense nucleus containing protons and neutrons, surrounded by electrons. Most of the atom is empty space.

• Protons: Positively charged, found in the nucleus.

• Neutrons: Neutral, found in the nucleus.

• Electrons: Negatively charged, orbit the nucleus.

• Protons and neutrons have nearly equal mass (relative mass 1); electron mass is negligible (relative mass 0).

Comparison of Subatomic Particles

Atomic Number, Mass Number, and Isotopes

• Atomic Number (Z): Number of protons in the nucleus; determines 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).

Example: Carbon Isotopes

Atomic Mass Unit (amu) and Atomic Weight

Atoms have extremely small masses, so a relative scale is used:

• Atomic Mass Unit (amu): Defined as 1/12 the mass of a carbon-12 atom.

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

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

The Periodic Table

The periodic table organizes elements by atomic number. Elements in the same group (column) have similar chemical properties.

• Periods: Rows in the periodic table.

• Groups: Columns in the periodic table; elements in a group share chemical properties.

• Metals: Left side; shiny, conduct heat/electricity, mostly solids.

• Nonmetals: Right side; can be solid, liquid, or gas.

• Metalloids: Border between metals and nonmetals; properties intermediate between both.

Example: Group Names

Molecules and Molecular Compounds

Molecular compounds are composed of molecules, typically containing only nonmetals. Chemical formulas indicate the number of atoms of each element in a molecule.

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

• Empirical Formula: Shows the simplest whole-number ratio of atoms.

• Molecular Formula: Shows the exact number of atoms of each element.

• Structural Formula: Shows how atoms are connected; does not indicate 3D shape.

Ions and Ionic Compounds

Atoms or groups of atoms can gain or lose electrons to form ions:

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

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

• Ionic Compounds: Formed between metals and nonmetals; electrons are transferred, and the resulting ions attract each other.

Naming Inorganic Compounds

Chemical nomenclature is the system for naming compounds:

• Ionic Compounds: Name the cation first, then the anion. If the cation has multiple charges, indicate with Roman numerals.

• Acids: Naming depends on the anion; -ide becomes hydro-...-ic acid, -ite becomes ...-ous acid, -ate becomes ...-ic acid.

• Binary Molecular Compounds: Name the element farther left or lower in the periodic table first; use prefixes to indicate number of atoms.

Prefixes for Binary Compounds

Some Simple Organic Compounds

Organic chemistry studies carbon-containing compounds. The simplest hydrocarbons are alkanes, named based on the number of carbon atoms (meth-, eth-, prop-, etc.) and ending in -ane.

• Alcohols: When a hydrogen in an alkane is replaced by a functional group (–OH), the compound is an alcohol, named with the suffix -ol.

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

Example: 1-Propanol and 2-propanol have the same molecular formula (C3H8O) but different structural formulas.

Additional info: The textbook cover image visually reinforces the study of molecular structure and atomic theory, which are central to this chapter.