Atoms & Elements

The Atom

The atom is the smallest unit of an element that retains the chemical properties of that element. It consists of subatomic particles: protons, neutrons, and electrons.

• Protons: Positively charged particles found in the nucleus.

• Neutrons: Neutral particles also located in the nucleus.

• Electrons: Negatively charged particles that orbit the nucleus.

• Nucleus: The dense center of the atom containing protons and neutrons.

Example: Protons and electrons have charges of the same magnitude but opposite signs. The number of protons equals the number of electrons in a neutral atom.

Subatomic Particles

Subatomic particles differ in mass and charge. Their properties are summarized below:

Example: Osmium has an actual mass of 190.23 grams. To find the number of atoms, divide by the atomic mass unit (amu).

Isotopes

Isotopes are atoms of the same element with different numbers of neutrons. They share the same atomic number but have different mass numbers.

• Atomic Number (Z): Number of protons; defines the element.

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

• Isotope Notation: AZX, where X is the element symbol.

Example: Calcium-43 (43Ca) has 20 protons, 23 neutrons, and 20 electrons.

Ion Formation

Ions are formed when atoms gain or lose electrons. Losing electrons forms cations (positive ions), while gaining electrons forms anions (negative ions).

• Cation: Atom loses electrons; positive charge.

• Anion: Atom gains electrons; negative charge.

• Isoelectronic: Species with the same number of electrons.

Example: The cation 27Al3+ has 13 protons, 14 neutrons, and 10 electrons.

Atomic Mass

Atomic mass is the weighted average mass of all isotopes of an element. It can be found using the periodic table or calculated from isotopic masses and abundances.

• Isotopic Mass: Mass of a specific isotope.

• Percent Abundance: Percentage of each isotope in nature.

• Fractional Abundance: Percent abundance divided by 100.

Atomic Mass Formula:

Example: If iron has two isotopes, Fe-54 and Fe-56, with known masses and atomic mass, you can solve for the percent abundance.

Periodic Table Classifications

The periodic table organizes elements by increasing atomic number and groups them by similar properties.

• Metals: Largest classification; good conductors, malleable, shiny.

• Nonmetals: Poor conductors, brittle, dull.

• Metalloids: Properties intermediate between metals and nonmetals.

Periodic Table Group Names

Elements are organized into periods (rows) and groups (columns). Groups share similar chemical properties.

• Representative Elements: Groups 1A-8A (s- and p-block).

• Transition Metals: Groups 3B-12B (d-block).

• Inert Gases: Group 8A (noble gases).

Element Symbols and History

Element symbols are based on Latin names and atomic numbers. The periodic law states that properties of elements recur periodically when arranged by atomic number.

• Example: Fluorine (F) is the halogen with the smallest atomic number.

Elemental Forms

Elements exist in different forms in nature:

• Monatomic Elements: Stable as single atoms (e.g., noble gases).

• Diatomic Elements: Stable as molecules of two atoms (e.g., H2, O2).

• Polyatomic Elements: Stable as molecules with more than two atoms (e.g., P4).

Homonuclear Compounds: Composed of atoms of the same element.

Heteronuclear Compounds: Composed of atoms of different elements.

Phases of Elements

At standard temperature and pressure, elements exist as solids, liquids, or gases.

• Solids: Definite shape and volume.

• Liquids: Definite volume, indefinite shape.

• Gases: Indefinite shape and volume.

Periodic Table Charges

Elements form ions to achieve stable electron configurations. Metals tend to lose electrons (form cations), while nonmetals gain electrons (form anions).

• Main Group Elements: Charge often equals group number for metals, or group number minus eight for nonmetals.

• Transition Metals: Can have multiple possible charges.

Calculating Molar Mass

Molar mass is the mass of one mole of a substance, expressed in grams per mole (g/mol).

Molar Mass Formula:

• Sum the atomic masses of all atoms in the formula.

• Example: For (NH4)2SO4, add the masses of N, H, S, and O.

Mole Concept

The mole is a counting unit in chemistry, representing 6.022 × 1023 particles (Avogadro's number).

• Formula Unit: Smallest unit of an ionic compound.

• Conversions: Moles ↔ Mass ↔ Particles

Example: To convert between moles and grams, use the molar mass as a conversion factor.

Law of Conservation of Mass

Mass is conserved in chemical reactions; the total mass of reactants equals the total mass of products.

Law of Definite Proportions

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

Example: CO2 always contains carbon and oxygen in a fixed mass ratio.

Atomic Theory

John Dalton's Atomic Theory (1803) laid the foundation for modern chemistry:

• All matter is composed of atoms.

• Atoms of the same element are identical.

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

• Atoms are rearranged in chemical reactions.

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.

Example: NO and NO2 show a 1:2 ratio of oxygen atoms per nitrogen atom.

Millikan Oil Drop Experiment

Robert Millikan measured the charge of the electron using oil droplets suspended in an electric field.

• Result: Fundamental charge of electron is C.

Rutherford Gold Foil Experiment

Ernest Rutherford discovered the nucleus by observing the scattering of alpha particles through gold foil.

• Most alpha particles passed through; some were deflected, indicating a dense, positively charged nucleus.

Rutherford's Postulates:

• Atoms have a small, dense nucleus.

• Electrons orbit the nucleus.

Additional info: These notes cover foundational concepts in atomic structure, periodic table organization, and basic chemical laws, suitable for introductory college-level General Chemistry.