Introduction to Chemistry: Matter and Composition

Chemistry Overview

• Chemistry examines matter, its characteristics, and changes it undergoes.
• Matter possesses mass and occupies space, existing as solids, liquids, or gases.
• Physical propreties include color and density.
• Chemical properties include reactivity and flammability.

Basic Concepts

• An element is a pure substance with only one type of atom, identified by its atomic number, or number of protons.
• A compound forms when two or more elements chemically bond in a fixed ratio.
• A molecule represents the smallest unit of a compound with the compound's properties.
• Mixtures combine substances physically, not chemically, and can be homogeneous (uniform) or heterogeneous (non-uniform).

Atoms, Molecules, and Ions

• Atoms consist of protons, neutrons, and electrons.
• Protons, with a positive charge, reside in the nucleus.
• Neutrons, with no charge, also reside in the nucleus.
• Electrons, negatively charged, orbit the nucleus in energy levels or shells.
• The number of protons determines an element; isotopes arise from varying neutron numbers.
• Ions form when atoms gain or lose electrons, acquiring a net charge.
• Cations are positive ions formed by electron loss.
• Anions are negative ions formed by electron gain.

Chemical Bonds

• A chemical bond is an attraction between atoms, enabling the formation of substances with multiple atoms.
• Ionic bonds arise from electrostatic forces between oppositely charged ions, common between metals and nonmetals.
• Covalent bonds involve electron sharing between atoms, typically nonmetals.
• Metallic bonds involve electron sharing within a "sea" of electrons, found in metals.

Chemical Reactions

• A chemical reaction rearranges atoms and molecules to create new substances.
• Reactants are substances that change during a reaction.
• Products are substances resulting from the reaction.
• Balanced reactions maintain equal atom numbers on both sides, obeying mass conservation.
• Common reaction types include synthesis, decomposition, single replacement, double replacement, and combustion.

Stoichiometry

• Stoichiometry is the quantitative study of reactant and product relationships in chemical reactions.
• The mole is the SI unit for substance amount, containing as many entities as atoms in 12 grams of carbon-12.
• Avogadro's number (approximately 6.022 x 10^23) defines the entities in one mole.
• Molar mass is the mass of one mole of a substance, in grams per mole (g/mol).
• Stoichiometric calculations use balanced equations to find reactant or product amounts.
• Limiting reactant is fully consumed, dictating maximum product yield.
• Percent yield is the actual product yield relative to the theoretical yield, expressed as a percentage.

Acids and Bases

• Acids donate protons (H+) or accept electrons.
• Bases accept protons or donate electrons.
• The pH scale measures acidity or basicity from 0 to 14.
• A pH below 7 indicates acidic solutions.
• A pH above 7 indicates basic or alkaline solutions.
• A pH of 7 indicates a neutral solution.
• Strong acids and bases fully dissociate in water, while weak ones partially dissociate.
• Acid-base reactions transfer protons from acid to base, forming salt and water (neutralization).

Organic Chemistry

• Organic chemistry studies carbon-containing compounds.
• Carbon uniquely forms stable bonds with itself and other elements, creating diverse molecules.
• Hydrocarbons contain only carbon and hydrogen, and can be saturated (single bonds only) or unsaturated (multiple bonds).
• Functional groups are atoms or groups in molecules that dictate chemical reactions.
• Common functional groups: alcohols (-OH), ethers (-O-), aldehydes (-CHO), ketones (-CO-), carboxylic acids (-COOH), amines (-NH2), and amides (-CONH2).
• Isomers share molecular formulas but differ in structure and properties.

Thermodynamics

• Thermodynamics studies energy and its transformations.
• The first law of thermodynamics: energy is conserved, only converted in form.
• Enthalpy (H) measures a system's heat content.
• Exothermic reactions release heat (ΔH < 0).
• Endothermic reactions absorb heat (ΔH > 0).
• Entropy (S) measures a system's disorder or randomness.
• The second law of thermodynamics: entropy in an isolated system increases over time.
• Gibbs free energy (G) combines enthalpy and entropy to determine reaction spontaneity (ΔG = ΔH - TΔS).
• Spontaneous reactions have negative Gibbs free energy change (ΔG < 0).

Kinetics

• Kinetics studies reaction rates and influencing factors.
• Reaction rate is the change in reactant or product concentration per time unit.
• Factors affecting reaction rates: reactant concentration, temperature, surface area, and catalysts.
• A catalyst accelerates a reaction without being consumed.
• The rate law relates reaction rate to reactant concentrations.
• Reaction order is the exponent of a reactant's concentration in the rate law.
• Activation energy is the minimum energy needed for a reaction.

This lesson introduces fundamental chemistry concepts, including the definition of matter, its properties, and its various states. It covers elements, compounds, molecules, and mixtures, differentiating between homogeneous and heterogeneous compositions. The lesson also explores the structure of atoms, including protons, neutrons, and electrons.