Atoms, Elements, Compounds, and the Mole: Foundational Concepts in Chemistry

Atoms and Elements

Dalton’s Atomic Theory

Dalton’s Atomic Theory laid the groundwork for modern chemistry by describing the nature of atoms and their role in chemical reactions.

• Indivisible Atoms: All matter is composed of indivisible atoms.

• Elemental Identity: Atoms of the same element are identical in mass and properties.

• Simple Ratios: Atoms combine in simple whole-number ratios to form compounds.

• Rearrangement: Chemical reactions involve the rearrangement of atoms.

Discovery of Subatomic Particles

Key Experiments and Models

The structure of the atom was revealed through several landmark experiments.

• Electrons: Discovered by J.J. Thomson using cathode ray tubes, which revealed negatively charged particles.

• Plum Pudding Model: Proposed that atoms are positive 'puddings' with electrons scattered throughout.

• Rutherford’s Gold Foil Experiment: Showed that atoms are mostly empty space with a dense, positively charged nucleus at the center.

Subatomic Particles and Properties

Atoms are composed of three main subatomic particles, each with distinct properties.

Particle	Charge	Symbol	Mass (amu)	Location
Proton	+1	p+	1	Nucleus
Neutron	0	n0	1	Nucleus
Electron	-1	e-	1/1836	Outside nucleus

The Periodic Table

Classification of Elements

The periodic table organizes elements based on their properties and atomic structure.

• Metals: Located on the left and center; shiny, conductors, malleable, form cations.

• Nonmetals: Located on the right; brittle, poor conductors, form anions.

• Metalloids: Border the staircase (e.g., B, Si, As); have mixed properties.

Groups and Naming

• Main Group: Groups 1A-8A.

• Transition Elements: Groups 1B-8B.

• Naming:

  • Group 1A: Alkali metals

  • Group 2A: Alkaline earth metals

  • Group 7A: Halogens

  • Group 8A: Noble gases

Ions

• Main Group Elements: Charge predicted from group number (e.g., 1A = +1, 2A = +2, 6A = -2, 7A = -1).

• Transition Elements: Have variable charges, indicated by Roman numerals (e.g., Fe2+, Fe3+).

Isotopes and Atomic Mass

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons.

• Example: Carbon-12 and Carbon-14 are isotopes of carbon.

Weighted Average Atomic Mass

The atomic mass listed on the periodic table is a weighted average of all naturally occurring isotopes.

• Formula:

$\text{Average Mass} = \sum (\text{fractional abundance}) \times (\text{isotopic mass})$

Molecules and Compounds

Law of Constant Composition

A given compound always contains the same proportion of elements by mass.

• Example: Water (H2O) always contains hydrogen and oxygen in a 2:1 ratio by atoms.

Chemical Symbols and Formulas

• Chemical Formula: Indicates elements and their proportions (e.g., H2O).

• Polyatomic Ions: Ion groups acting as a single unit (e.g., NO3-, SO42-).

Classification of Substances

Type	Composition	Example
Atomic element	Single atoms	Na
Molecular element	Diatomic or polyatomic	O2, P4
Ionic compound	Metal + nonmetal	NaCl
Molecular compound	Nonmetal + nonmetal	CO2

Writing Ionic Compound Formulas

• Write cation (metal) and anion (nonmetal).

• Balance charges to form a neutral compound.

Naming Compounds

• Binary Ionic: Name metal + nonmetal (end with -ide). E.g., NaCl = Sodium chloride.

• Transition Metals: Use Roman numerals. E.g., FeCl3 = Iron(III) chloride.

• Polyatomic Ions: Name cation first, then polyatomic anion. E.g., CaSO4 = Calcium sulfate.

• Molecular Compounds: Use prefixes (mono-, di-, tri-). E.g., CO2 = Carbon dioxide.

• Binary Acids: Hydro + base name + ic acid. E.g., HCl = Hydrochloric acid.

• Oxyacids: Based on polyatomic ions: -ate → -ic acid; -ite → -ous acid.

Formula Mass

The formula mass is the sum of atomic masses of all atoms in the molecule or formula unit.

• Formula:

$\text{Formula Mass} = \sum \text{(atomic masses of all atoms in the formula)}$

The Mole Concept

Definition and Avogadro’s Number

The mole is a counting unit in chemistry, representing a specific number of particles.

• 1 mole = $6.022 \times 10^{23}$ particles (Avogadro’s number).

Molar Mass and Atomic Mass

• Molar Mass: Mass (in g/mol) equal to atomic or molecular mass (in amu).

• Formula:

$\text{Moles} = \frac{\text{mass}}{\text{molar mass}}$

Conversions

• mass → moles → molecules

Examples

• From mass to moles: $\text{moles} = \frac{\text{mass}}{\text{molar mass}}$

• From moles to molecules: Multiply by Avogadro’s number.

Mass Percent

Mass percent expresses the proportion of an element in a compound.

• Formula:

$\text{Mass % of element} = \frac{\text{mass of element in compound}}{\text{total mass of compound}} \times 100\%$

Empirical and Molecular Formulas

Empirical formulas show the simplest whole-number ratio of atoms in a compound, while molecular formulas show the actual number of atoms.

• Use experimental data or mass percent to find moles of each element.

• Divide by the smallest mole number.

• Adjust to whole numbers.

• To find molecular formula: divide molar mass by empirical formula mass.

Example: If a compound has an empirical formula CH2O and a molar mass of 180 g/mol, the molecular formula is C6H12O6.