Atomic Theory of Matter

Historical Development of Atomic Theory

The concept of the atom originated with ancient Greek philosophers, but it was John Dalton in the early 1800s who formulated a scientific atomic theory based on experimental evidence. Dalton's theory was built upon three fundamental laws:

• Law of Constant Composition: Compounds always contain the same proportion of elements by mass, as discovered by Joseph Proust.

• Law of Conservation of Mass: The total mass remains unchanged during a chemical reaction, as established by 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, as discovered by John Dalton.

Dalton’s Atomic Theory

Dalton’s atomic theory consists of several postulates that form the foundation of modern chemistry:

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

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

• Atoms are neither created nor destroyed in chemical reactions; they are simply rearranged.

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

Discovery of Subatomic Particles

The Electron and Cathode Rays

J. J. Thomson discovered the electron in 1897 through experiments with cathode ray tubes, demonstrating that atoms contain negatively charged particles.

• Thomson measured the charge-to-mass ratio of the electron as C/g.

Millikan Oil-Drop Experiment

Robert Millikan determined the charge of the electron in 1909 using the oil-drop experiment, allowing calculation of the electron's mass.

Radioactivity

Radioactivity is the spontaneous emission of radiation from an atom, first observed by Henri Becquerel and further studied by Marie and Pierre Curie. Ernest Rutherford identified three types of radiation:

• Alpha (α) particles: Positively charged

• Beta (β) particles: Negatively charged (like electrons)

• Gamma (γ) rays: Uncharged

Discovery of the Nucleus

Rutherford’s gold foil experiment demonstrated that atoms have a small, dense, positively charged nucleus, disproving the "plum pudding" model.

Structure of the Atom

Atoms consist of a nucleus (containing protons and neutrons) surrounded by electrons. Most of the atom's volume is empty space occupied by electrons.

Subatomic Particles

The three main subatomic particles are:

Atomic Number, Mass Number, and Isotopes

Atomic Number and Mass Number

The atomic number (Z) is the number of protons in the nucleus and defines the element. The mass number (A) is the total number of protons and neutrons.

• Atoms of the same element have the same atomic number but may have different mass numbers (isotopes).

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons. They have identical chemical properties but different masses.

Atomic Mass Unit and Atomic Weight

The atomic mass unit (amu) is defined as . The atomic weight of an element is the weighted average of the masses of its isotopes:

Mass Spectrometry

Mass spectrometry is used to measure atomic and molecular weights and determine isotopic abundances.

The Periodic Table

Organization and Structure

The periodic table arranges elements in order of increasing atomic number. Rows are called periods, and columns are called groups. Elements in the same group have similar chemical properties.

Reading the Periodic Table

Each element's box lists the atomic number (above the symbol) and atomic weight (below the symbol).

Classification of Elements

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

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

• Metalloids: Border the steplike line; properties intermediate between metals and nonmetals.

Groups with Special Names

Chemical Formulas and Molecular Structure

Chemical Formulas

Chemical formulas indicate the types and numbers of atoms in a compound. Subscripts show the number of each atom present.

• Molecular compounds: Composed of molecules, usually only nonmetals.

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

Types of Formulas

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

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

• Structural formula: Shows how atoms are bonded.

Ions and Ionic Compounds

Formation of Ions

Ions are formed when atoms gain or lose electrons:

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

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

Common Cations and Anions

Polyatomic Ions

Polyatomic ions are groups of atoms that carry a charge, such as ammonium (NH4+) and sulfate (SO42–).

Ionic Compounds

Ionic compounds are formed from the electrostatic attraction between cations and anions, typically between metals and nonmetals. The formula is written to reflect the lowest whole-number ratio of ions.

Chemical Nomenclature

Naming Ionic Compounds

• Name the cation first; if it can have more than one charge, indicate the charge with Roman numerals.

• If the anion is an element, change its ending to -ide; if polyatomic, use its name.

Oxyanion Nomenclature Patterns

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

• Prefixes: hypo- (fewest), per- (most).

Naming Acids

• Anion ends in -ide: hydro-...-ic acid (e.g., HCl: hydrochloric acid).

• Anion ends in -ate: ...-ic acid (e.g., HClO3: chloric acid).

• Anion ends in -ite: ...-ous acid (e.g., HClO2: chlorous acid).

Naming Binary Molecular Compounds

• The element farther left or lower in the group is named first.

• Prefixes indicate the number of atoms (mono-, di-, tri-, etc.).

• The second element ends in -ide.

Nomenclature of Organic Compounds

Organic compounds are primarily composed of carbon and hydrogen. The simplest are alkanes (single bonds only), named with the suffix -ane. When a hydrogen is replaced by an –OH group, the compound is an alcohol (suffix -ol).

Isomers

Isomers are compounds with the same molecular formula but different structures. For example, 1-propanol and 2-propanol both have the formula C3H8O but differ in the position of the –OH group.