Atoms, Molecules, and Ions

Atomic Theory of Matter

The concept of the atom as the fundamental building block of matter dates back to 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 several foundational laws:

• Law of Constant Composition: Compounds always contain the same proportion of elements by mass. For example, water (H2O) always contains two hydrogen atoms for every oxygen atom, regardless of the sample size. (Discovered by Joseph Proust)

• Law of Conservation of Mass: The total mass of substances remains unchanged during a chemical reaction. (Discovered 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 ratios of small whole numbers. (Discovered by John Dalton)

Dalton's atomic theory postulated that:

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 neither created nor destroyed in chemical reactions.

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

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, electrons are 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 spontaneously, indicating internal structure.

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

• 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. Key points:

• Protons: Positively charged, mass ≈ 1 amu.

• Neutrons: Neutral, mass ≈ 1 amu.

• Electrons: Negatively charged, mass ≈ 0 amu (very small compared to protons/neutrons).

• Protons and neutrons are in the nucleus; electrons move around 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

Isotopes are atoms of the same element with different numbers of neutrons (and thus different masses). For example, carbon has isotopes such as 12C, 13C, and 14C.

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.

Atomic and molecular weights can be measured using a mass spectrometer.

The Periodic Table

The periodic table organizes elements by increasing atomic number. Key features:

• Periods: Horizontal rows.

• Groups: Vertical columns; elements in the same group have similar chemical properties.

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

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

• Metalloids: Border the stair-step line; properties intermediate between metals and nonmetals.

Molecules and Molecular Compounds

Molecules are groups of atoms bonded together. Molecular compounds typically contain only nonmetals. Chemical formulas indicate the number and type of atoms present:

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

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

• Structural formula: Shows how atoms are connected.

Diatomic molecules: Seven elements naturally exist as molecules with two atoms (H2, N2, O2, F2, Cl2, Br2, I2).

Ions and Ionic Compounds

Ions are charged particles formed when atoms gain or lose electrons:

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

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

Ionic compounds are formed from metals and nonmetals; electrons are transferred from metal to nonmetal, and the resulting ions are held together by electrostatic attraction. Only empirical formulas are used for ionic compounds.

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.

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

• Binary molecular compounds: Use prefixes to indicate the number of each atom (e.g., CO2: carbon dioxide).

Oxyanions: For ions with different numbers of oxygen atoms:

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

• Prefixes hypo- (less) and per- (more) are used for series with more than two oxyanions.

Some Simple Organic Compounds

Organic chemistry studies carbon-containing compounds. The simplest are alkanes (hydrocarbons with only single bonds). Naming is based on the number of carbon atoms (meth-, eth-, prop-, etc.) and the suffix -ane.

Alcohols are derived from alkanes by replacing a hydrogen with an -OH group; their names end in -ol (e.g., methanol, ethanol).

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