Atoms, Elements, and Compounds

What is Chemistry?

Chemistry is the branch of science dedicated to the study of matter, its transformations, and the associated energy changes. It is a central science that connects and overlaps with fields such as physics, biology, medicine, engineering, and earth sciences.

• Matter: Anything that occupies space and has mass.

• Transformation of Matter: Chemical and physical changes that alter the composition or state of substances.

States of Matter

Matter exists in different states, each with distinct physical properties:

• Solid: Definite shape and volume; particles are closely packed.

• Liquid: Definite volume but takes the shape of its container; particles are less tightly packed than in solids.

• Gas: No definite shape or volume; particles move freely and occupy the entire volume of their container.

Composition of Matter

• Atoms: The simplest unit of matter that retains the properties of a chemical element.

• Molecules: Substances composed of two or more atoms held together by chemical bonds (often covalent).

• Compounds: Substances formed from two or more different elements chemically bonded together.

• Mixtures: Physical combinations of two or more substances where each retains its identity; can be homogeneous (uniform) or heterogeneous (non-uniform).

Atomic Theory and Structure

Historical Development of Atomic Theory

• Democritus (5th century B.C.): Proposed that matter is composed of indivisible units called "atomos."

• Dalton (1803): Revived the atomic theory, suggesting that atoms are the fundamental building blocks of matter.

• Thomson (1897): Discovered the electron using cathode ray experiments; proposed the "plum pudding" model of the atom.

• Rutherford (1911): Gold foil experiment revealed a dense, positively charged nucleus with electrons orbiting around it.

• Bohr (1913): Electrons occupy specific energy levels or orbits around the nucleus.

• Schrödinger (1923): Introduced the quantum mechanical model, describing electrons as wave-like and existing in probability clouds.

• Chadwick (1932): Discovered the neutron, explaining atomic mass discrepancies.

Structure of the Atom

• Nucleus: Contains protons (p+) and neutrons (n).

• Electrons (e-): Occupy energy levels (shells) around the nucleus.

Additional info: 1 atomic mass unit (u) = 1/12 the mass of a carbon-12 atom = 1.66057×10-27 kg.

Elements, Molecules, Compounds, and Mixtures

• Element: Pure substance consisting of only one type of atom; cannot be broken down by chemical means.

• Molecule: Two or more atoms bonded together (e.g., O2, H2O).

• Compound: Substance formed from two or more elements chemically bonded (e.g., NaCl, CO2).

• Mixture: Combination of substances not chemically bonded; can be separated by physical means.

The Periodic Table, Atomic Mass, and the Mole

The Periodic Table

The periodic table organizes elements by increasing atomic number and similar chemical properties. Groups (columns) share similar valence electron configurations, influencing their reactivity.

Atomic Number, Mass Number, and Isotopes

• 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: Atoms of the same element with different numbers of neutrons (e.g., protium, deuterium, tritium for hydrogen).

Ions

• Anion: Atom that gains electrons, becoming negatively charged (e.g., Cl-).

• Cation: Atom that loses electrons, becoming positively charged (e.g., Na+).

The Mole and Molar Mass

• Mole (mol): The amount of substance containing 6.02×1023 entities (Avogadro's number).

• Molar Mass: Mass of one mole of a substance, expressed in g/mol.

Examples:

• 1 mol H2 = 2 g

• 1 mol Au = 196.97 g

• 1 mol CO2 = 44 g

Electronic Structure and Quantum Numbers

Quantum Numbers

Quantum numbers describe the properties of atomic orbitals and the electrons in them:

Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers.

Electron Configuration

Electrons fill orbitals in order of increasing energy, following the Aufbau principle, Hund's rule, and the Pauli exclusion principle.

• Aufbau Principle: Electrons occupy the lowest energy orbitals first.

• Hund's Rule: Electrons fill degenerate orbitals singly before pairing.

Example: Sodium (Na, atomic number 11): 1s2 2s2 2p6 3s1

Chemical Bonding

Types of Chemical Bonds

• Ionic Bond: Formed by electron transfer from a metal to a nonmetal; large difference in electronegativity.

• Covalent Bond: Atoms share electrons to achieve stable configurations (octet rule).

Electronegativity Difference:

• 0: Non-polar covalent

• <2: Polar covalent

• >2: Ionic

Lewis Structures

Lewis structures represent valence electrons as dots around element symbols, showing bonding and lone pairs. The octet rule guides the arrangement, but exceptions exist (e.g., molecules with odd electrons, expanded octets).

Formal Charge

Used to determine the most stable Lewis structure:

$\text{Formal Charge} = \text{Valence electrons} - \text{Non-bonded electrons} + \frac{1}{2} \text{Bonded electrons}$

Resonance Structures

Some molecules can be represented by two or more valid Lewis structures; the actual structure is a resonance hybrid.

Molecular Geometry and VSEPR Theory

VSEPR Theory

The Valence Shell Electron Pair Repulsion (VSEPR) theory predicts the three-dimensional arrangement of atoms in a molecule based on the repulsion between electron pairs around a central atom.

• Electron Geometry: Arrangement of all electron pairs (bonding and lone pairs).

• Molecular Geometry: Arrangement of only the atoms (bonding pairs).

Chemical Reactions and Stoichiometry

Chemical Equations

Chemical reactions are represented by equations showing reactants and products, their physical states, and energy changes.

• Balancing Equations: Ensures the law of conservation of mass is obeyed; can be done by trial and error or algebraic methods.

Types of Chemical Reactions

• Single Displacement: A + BC → B + AC

• Double Displacement: AB + CD → AD + CB

• Combination (Synthesis): A + B → AB

• Decomposition: AB → A + B

• Combustion: Hydrocarbon + O2 → CO2 + H2O

Stoichiometry

• Limiting Reagent: The reactant that is completely consumed first, limiting the amount of product formed.

• Theoretical Yield: Maximum amount of product that can be formed from given reactants.

• Percent Yield: Ratio of actual yield to theoretical yield, expressed as a percentage.

$\%\text{Yield} = \frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100$

Intermolecular Forces and States of Matter

Types of Intermolecular Forces

• Ion-Ion: Between charged ions.

• Ion-Dipole: Between an ion and a polar molecule.

• Dipole-Dipole: Between polar molecules.

• Hydrogen Bonding: Special case of dipole-dipole, involving H bonded to N, O, or F.

• Dipole-Induced Dipole: Between a polar and a nonpolar molecule.

• Induced Dipole-Induced Dipole (London Dispersion): Present in all molecules, dominant in nonpolar substances.

Properties of Liquids

• Surface Tension: Energy required to increase the surface area of a liquid; decreases with temperature.

• Viscosity: Resistance to flow; decreases with temperature and increases with stronger intermolecular forces.

• Vapor Pressure: Pressure exerted by a vapor in equilibrium with its liquid; increases with temperature.

Phase Diagrams

Phase diagrams show the state of a substance at various temperatures and pressures, including critical points and lines for melting, vaporization, and sublimation.

Gases and Gas Laws

Kinetic Molecular Theory of Gases

• Gases consist of particles in constant, random motion.

• Collisions between particles and with container walls are elastic.

• Particles are far apart relative to their size; volume of particles is negligible.

• No attractive or repulsive forces between particles.

• Average kinetic energy is proportional to temperature (in Kelvin).

Ideal Gas Law

The behavior of ideal gases is described by:

$PV = nRT$

• P = pressure (atm or Pa)

• V = volume (L or m3)

• n = moles of gas

• R = gas constant (0.0821 L·atm/mol·K or 8.314 J/mol·K)

• T = temperature (K)

Real gases deviate from ideal behavior at high pressures and low temperatures; corrections are made using the Van der Waals equation.

$\left(P + \frac{a n^2}{V^2}\right) (V - n b) = nRT$

Additional info: 'a' corrects for intermolecular attractions, 'b' corrects for molecular volume.