Introduction to Chemistry

Chemistry Fundamentals

• Chemistry is the study of matter, its properties, and how it changes.
• Matter is anything possessing mass and occupying space.
• Chemistry is the central science, linking physics, geology, and biology.

Branches of Chemistry

• Analytical Chemistry focuses on the composition of substances.
• Organic Chemistry studies carbon-containing compounds, excluding some like carbon oxides.
• Inorganic Chemistry studies substances lacking carbon.
• Physical Chemistry examines the physics underlying chemical systems.
• Biochemistry explores chemical processes within living organisms.

States of Matter

• Solids have definite shapes and volumes.
• Liquids have definite volumes and assume the shape of their containers.
• Gases lack definite shape or volume, expanding to fill available space.
• Plasma is ionized gas, often at high temperatures.

Properties of Matter

• Physical properties are observable without changing a substance's composition (e.g., color, density, melting point).
• Chemical properties describe how a substance changes into a new one (e.g., flammability, reactivity).
• Intensive properties do not depend on the amount of matter (e.g., density, boiling point).
• Extensive properties depend on the amount of matter (e.g., mass, volume).

Changes of Matter

• Physical changes alter a substance's form without changing its chemical identity (e.g., melting, boiling, cutting).
• Chemical changes rearrange atoms to form new substances (e.g., burning, rusting, cooking).

Mixtures

• Mixtures combine two or more physically combined substances.
• Homogeneous mixtures have uniform composition (e.g., saltwater, air).
• Heterogeneous mixtures have non-uniform composition (e.g., salad, sand and water).

Separating Mixtures

• Filtration separates solids from liquids.
• Distillation separates liquids based on boiling points.
• Chromatography separates substances by affinity for a stationary phase.

Elements and Compounds

• Elements are pure substances that cannot be broken down chemically; represented by symbols.
• Compounds are two or more elements chemically combined in fixed proportions.

Chemical Formulas

• Molecular formulas show the exact number of atoms in a molecule.
• Empirical formulas show the simplest whole-number ratio of atoms in a compound.

Chemical Reactions

• Chemical equations represent reactions, showing reactants and products.
• Reactants are substances undergoing change.
• Products are substances formed in a reaction.
• Balancing equations ensures the same number of atoms on both sides, obeying mass conservation.

The Mole

• The mole is the SI unit for the amount of substance.
• One mole contains as many elementary entities as atoms in 12 grams of carbon-12.
• Avogadro's number is approximately 6.022 x 10^23 entities per mole.
• Molar mass is the mass of one mole in grams per mole (g/mol).

Stoichiometry

• Stoichiometry is the quantitative relationship between reactants and products.
• The limiting reactant is completely consumed, determining product amount.
• The excess reactant remains after the limiting reactant is used up.
• Theoretical yield is the maximum product amount from given reactants.
• Actual yield is the product amount obtained from a reaction.
• Percent yield is (Actual Yield / Theoretical Yield) x 100%.

Atomic Structure

• Atoms are the basic unit of matter, with protons, neutrons, and electrons.
• Protons are positively charged particles in the nucleus.
• Neutrons are neutral particles in the nucleus.
• Electrons are negatively charged particles orbiting the nucleus.
• Atomic number is the number of protons, defining the element.
• Mass number is the total number of protons and neutrons.
• Isotopes are atoms of the same element with different neutron numbers.

Electron Configuration

• Electron configuration describes electron arrangement within an atom.
• Orbitals (s, p, d, f) are regions where electrons are likely to be found.
• The Aufbau principle dictates electrons fill orbitals in increasing energy order.
• Hund's Rule: Electrons fill orbitals individually before pairing up in the same orbital.
• Pauli Exclusion Principle states no two electrons share the same four quantum numbers.

The Periodic Table

• The periodic table arranges elements by atomic number and groups with similar properties.
• Groups (columns) contain elements with similar chemical properties.
• Periods (rows) contain elements with the same number of electron shells.
• Metals are typically shiny, conduct electricity and heat, and are malleable and ductile.
• Nonmetals are typically dull, poor conductors.
• Metalloids have properties of both metals and nonmetals.

Chemical Bonding

• Chemical bonds are attractive forces holding atoms together.
• Ionic bonds involve electron transfer between metals and nonmetals.
• Covalent bonds involve electron sharing between nonmetals.
• Metallic bonds involve electron sharing among metal atoms, creating an electron "sea".

Types of Covalent Bonds

• Single bonds share one electron pair.
• Double bonds share two electron pairs.
• Triple bonds share three electron pairs.
• Polar covalent bonds involve unequal electron sharing, resulting in partial charges.
• Nonpolar covalent bonds involve equal electron sharing.

Molecular Geometry

• VSEPR (Valence Shell Electron Pair Repulsion) theory predicts molecular shapes based on electron pair repulsion.
• Common geometries include linear, trigonal planar, tetrahedral, bent, and trigonal pyramidal.

Intermolecular Forces

• Intermolecular forces are attractive forces between molecules.
• Dipole-dipole forces occur between polar molecules.
• Hydrogen bonds are strong dipole-dipole forces with H bonded to N, O, or F.
• London dispersion forces are weak forces in all molecules due to temporary electron distribution fluctuations.

Gases

• Ideal Gas Law: PV = nRT (Pressure x Volume = moles x Ideal Gas Constant x Temperature)
• Boyle's Law: At constant temperature, volume is inversely proportional to pressure (P1V1 = P2V2).
• Charles's Law: At constant pressure, volume is directly proportional to absolute temperature (V1/T1 = V2/T2).
• Avogadro's Law: At constant temperature and pressure, volume is directly proportional to the number of moles (V1/n1 = V2/n2).

Solutions

• Solutions are homogeneous mixtures of two or more substances.
• The solute is the substance being dissolved.
• The solvent is the substance that dissolves the solute.
• Concentration is the amount of solute in a solvent or solution (e.g., molarity, molality, percent concentration).
• Molarity (M) is moles of solute per liter of solution.
• Solubility is the maximum solute that can dissolve in a solvent at a specific temperature.

Acids and Bases

• Acids donate protons (H+) or accept electrons.
• Bases accept protons (H+) or donate electrons.
• pH measures acidity or alkalinity; pH = -log[H+].
• Acid-base reactions are neutralization reactions that form water and a salt.

Thermodynamics

• Thermodynamics studies energy and its transformations.
• Enthalpy (H) measures a system's heat content.
• Entropy (S) measures a system's disorder or randomness.
• Gibbs Free Energy (G) measures reaction spontaneity; G = H - TS.

Chemical Kinetics

• Chemical kinetics studies reaction rates and mechanisms.
• The rate law relates reaction rate to reactant concentrations.
• Activation energy is the minimum energy for a reaction to occur.
• A catalyst speeds up a reaction without being consumed.

Equilibrium

• Chemical equilibrium is when forward and reverse reaction rates are equal.
• The equilibrium constant (K) measures relative reactant and product amounts at equilibrium.
• Le Chatelier's Principle states a system at equilibrium will counteract changes to restore equilibrium.