Atoms, Molecules, and Ions

Atomic Theory of Matter

The concept of the atom as the fundamental building block of matter has evolved from ancient philosophy to modern science. Key experimental laws led to the development of Dalton’s Atomic Theory, which forms the basis of our understanding of chemical substances and reactions.

• Law of Constant Composition: Compounds always contain the same proportion of elements by mass. Discovered by Joseph Proust, this law states that any sample of a compound will have the same ratio of constituent atoms.

• Law of Conservation of Mass: The total mass of substances remains unchanged during a chemical reaction. Antoine Lavoisier established that mass is neither created nor destroyed in chemical processes.

• 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 formulated this law, highlighting the discrete nature of atoms.

Dalton’s Atomic Theory (1808):

1. Elements are composed of extremely small particles called atoms.

2. All atoms of a given element are identical in mass and properties; atoms of different elements differ in these respects.

3. Atoms are not created or destroyed in chemical reactions; they are simply rearranged.

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 themselves 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) particles: positively charged

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

  • Gamma (γ) rays: uncharged

• 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, with electrons surrounding it.

Modern View of Atomic Structure

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

• 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 and Mass Number

The identity of an element is determined by its atomic number (number of protons). The mass number is the sum of protons and neutrons in the nucleus.

• Atomic Number (Z): Number of protons in the nucleus; also equals the number of electrons in a neutral atom.

• Mass Number (A): Total number of protons and neutrons.

• Isotopes: Atoms of the same element with different numbers of neutrons (and thus different mass numbers).

Example: Carbon-12 (12C) has 6 protons and 6 neutrons; Carbon-14 (14C) has 6 protons and 8 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 carbon-12 atom. The atomic weight of an element is the weighted average of the masses of its naturally occurring isotopes.

• Atomic Weight Calculation:

  • Weighted average:

• 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 and groups elements with similar chemical properties together.

• Periods: Horizontal rows

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

• Metals: Left side; shiny, conductive, mostly solids

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

• Metalloids: Border the stair-step line; have properties of both metals and nonmetals

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 are composed of molecules and 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 (e.g., CH2O for glucose).

• Molecular Formula: Actual number of atoms (e.g., C6H12O6 for glucose).

• 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 and anions.

• Cation: Positively charged ion (formed by loss of electrons; usually metals).

• Anion: Negatively charged ion (formed by gain of electrons; usually nonmetals).

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

Naming Inorganic Compounds

Chemical nomenclature provides systematic rules for naming compounds and ions.

• 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).

• Monatomic Anions: Change the ending to -ide (e.g., chloride, oxide).

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

• Oxyanions: For two ions of the same element, -ite (fewer O), -ate (more O); prefixes hypo- (less) and per- (more) for series with more than two.

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

• Binary Molecular Compounds: Name the more metallic element first; use prefixes to indicate number of atoms (mono-, di-, tri-, etc.). Second element ends in -ide.

Table: Prefixes for Binary Molecular Compounds

Simple Organic Compounds

Organic chemistry focuses on compounds containing carbon. The simplest organic compounds are hydrocarbons (alkanes), named according to the number of carbon atoms.

• Alkanes: Saturated hydrocarbons with only single bonds; names end in -ane (e.g., methane, ethane, propane).

• Alcohols: Derived from alkanes by replacing a hydrogen with an -OH group; 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).