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 (discovered by Dalton).

Dalton’s Atomic Theory: Postulates

• 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. This discovery revealed the presence of subatomic particles and atomic energy.

• Alpha (α) particles: Positively charged

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

• Gamma (γ) rays: Uncharged

Models of the Atom

Thomson’s Plum Pudding Model

Thomson proposed that atoms consist of a positively charged sphere with embedded electrons, resembling a "plum pudding." This model was later disproven.

Rutherford’s Nuclear Model

Ernest Rutherford’s gold foil experiment demonstrated that atoms have a small, dense, positively charged nucleus, with electrons surrounding it. Most of the atom is empty space.

Subatomic Particles

• Protons: Positive charge (+1), mass ≈ 1 amu

• Neutrons: No charge, mass ≈ 1 amu

• Electrons: Negative charge (–1), mass ≈ 0 amu (very small)

Atomic Number, Mass Number, and Isotopes

Atomic Number and Mass Number

• 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

Isotopes are atoms of the same element with different numbers of neutrons, resulting in different mass numbers.

Atomic Mass Unit (amu) and Atomic Weight

• 1 amu = g

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

• Calculated as:

Measuring Atomic and Molecular Mass

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

The Periodic Table

Structure and Organization

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).

Groups and Periodicity

Certain groups have specific names and characteristic properties:

Classification of Elements

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

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

• Metalloids: Border the steplike line; have properties of both metals and nonmetals.

Chemical Formulas and Molecular Structure

Chemical Formulas

• Subscripts indicate the number of atoms of each element in a molecule.

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

Diatomic Molecules

Seven elements exist naturally 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 in a molecule.

• Structural formula: Shows how atoms are bonded.

Ions and Ionic Compounds

Formation of Ions

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

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

Common Cations and Anions

Polyatomic Ions

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

Ionic Compounds

Ionic compounds are formed from metals and nonmetals, with electrons transferred from metal to nonmetal. The formula is written in the lowest whole-number ratio.

Writing Formulas for Ionic Compounds

• The charge on the cation becomes the subscript for the anion and vice versa.

• Reduce subscripts to the lowest whole-number ratio.

Chemical Nomenclature

Naming Ionic Compounds

• Name the cation first; if it can have more than one charge, indicate 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– and –ic acid (e.g., HCl: hydrochloric acid).

• Anion ends in –ite: –ous acid (e.g., HClO: hypochlorous acid).

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

Introduction to Organic Compounds

Alkanes

Alkanes are the simplest hydrocarbons, containing only single bonds between carbon atoms. Their names are based on the number of carbon atoms (meth–, eth–, prop–, etc.) followed by –ane.

Alcohols and Isomerism

Alcohols are derived from alkanes by replacing a hydrogen with an –OH group, and their names end in –ol. Isomers are compounds with the same formula but different structures.