Introduction to Chemistry: Matter and Properties

• Chemistry is the study of matter, its properties, and how it changes
• Matter is anything that possesses mass and occupies space

States of Matter

• Solid: Defined by a definite shape and volume.
• Liquid: Defined by a definite volume but adapts to the shape of its container.
• Gas: Lacks a definite shape or volume.
• Plasma: Described as an ionized gas existing at high temperatures, such as in stars.
• Phase changes are considered physical changes

Chemical vs. Physical Properties

• Physical properties are observed without altering a substance's composition, examples are color, density, and melting point.
• Chemical properties become evident when a substance undergoes a change in composition, examples are flammability and reactivity with acid
• Physical change: alters the form of a substance, not its composition, for example melting ice
• Chemical change: alters the composition of a substance, for example burning wood
• Indicators of chemical change include color change, precipitate formation, gas evolution, and heat absorption or release

Elements and Compounds

• Element: A substance that cannot be broken down into simpler substances through chemical means, like gold or oxygen.
• Compound: A substance containing two or more elements chemically combined in a fixed ratio, such as water or salt.
• Mixtures: Physical combinations of substances that can be separated through physical methods.
• Homogeneous mixture: Exhibits uniform composition throughout, for example saltwater.
• Heterogeneous mixture: Exhibits non-uniform composition, for example salad.

Atomic Structure

• Atom: An element's smallest unit, retaining its chemical properties.
• Protons: Positively charged particles located in the nucleus.
• Neutrons: Neutral particles located in the nucleus.
• Electrons: Negatively charged particles orbiting the nucleus.
• Atomic number (Z): Corresponds to the number of protons in the nucleus, identifying the element.
• Mass number (A): Represents the total number of protons and neutrons in the nucleus.
• Isotopes: Atoms of the same element distinguished by different numbers of neutrons.
• Ions: Atoms that have gained or lost electrons, leading to a net charge.
• Cations: Positively charged ions, formed through the loss of electrons.
• Anions: Negatively charged ions, formed through the gain of electrons.

The Mole Concept

• The mole (mol) serves as the SI unit for measuring the amount of a substance.
• One mole contains Avogadro's number (6.022 x 10^23) of particles such as atoms, molecules, or ions.
• Molar mass: Represents the mass of one mole of a substance in grams, numerically equivalent to the atomic or molecular weight in atomic mass units (amu).

Chemical Reactions and Equations

• Chemical equation: A symbolic representation of a chemical reaction.
• Reactants: Substances that undergo change in a chemical reaction, shown on the left side of the equation.
• Products: Substances resulting from a chemical reaction, shown on the right side of the equation.
• Balancing chemical equations: This process ensures the conservation of mass by ensuring the number of atoms for each element is identical on both sides of the equation.

Stoichiometry

• Stoichiometry is the quantitative relationship between reactants and products in a chemical reaction
• Mole ratio: The ratio of moles of reactants and products in a balanced chemical equation
• Limiting reactant: Reactant that is completely consumed in a reaction, determining the amount of product formed
• Percent yield: Actual yield / Theoretical yield x 100%

Solutions

• Solution: Defined as a homogeneous mixture of two or more substances.
• Solvent: The substance present in greater quantity within a solution.
• Solute: The substance(s) present in lesser quantity within a solution.
• Concentration: Denotes the quantity of solute within a specific amount of solvent or solution.
• Molarity (M): Defined as moles of solute per liter of solution.
• Dilution: The process of reducing a solution's concentration by adding more solvent, following the relationship M1V1 = M2V2.

Acids and Bases

• Acids: Proton (H+) donors in aqueous solutions (Arrhenius) or electron-pair acceptors (Lewis); pH < 7.
• Bases: Proton (H+) acceptors in aqueous solutions (Arrhenius) or electron-pair donors (Lewis); pH > 7.
• Strong acids/bases: Fully dissociate in water.
• Weak acids/bases: Partially dissociate in water.
• pH: Measures a solution's acidity or basicity; calculated as pH = -log[H+].
• Neutralization: The reaction between an acid and a base, forming a salt and water.

Chemical Bonding

• Ionic bond: The electrostatic attraction between ions, formed through the transfer of electrons, resulting in charged ions
• Covalent bond: Sharing of electrons between atoms
• Metallic bond: Sharing of electrons within a metal lattice
• Polar covalent bond: Unequal sharing of electrons, resulting in partial charges on atoms
• Electronegativity: Measures an atom's ability to attract electrons in a chemical bond

Organic Chemistry

• Organic chemistry is the study of carbon-containing compounds
• Hydrocarbons: Compounds containing only carbon and hydrogen
• Alkanes: Saturated hydrocarbons featuring single bonds (C-C)
• Alkenes: Unsaturated hydrocarbons containing at least one double bond (C=C)
• Alkynes: Unsaturated hydrocarbons containing at least one triple bond (C≡C)
• Functional groups: Specific groupings of atoms within molecules that have their own characteristic properties, regardless of the other atoms present in a molecule
• Isomers: Molecules possessing the same molecular formula but differing structural formulas

Thermodynamics

• Thermodynamics is the study of energy and its transformations
• Energy: The capacity to do work or transfer heat
• Kinetic energy: Energy of motion
• Potential energy: Energy of position or composition
• First law of thermodynamics: Energy is conserved
• Enthalpy (H): Measures heat content in a system at constant pressure; ΔH = ΔU + PΔV
• Exothermic process: Releases heat to the surroundings (ΔH < 0)
• Endothermic process: Absorbs heat from the surroundings (ΔH > 0)
• Entropy (S): Measures disorder or randomness in a system
• Second law of thermodynamics: The entropy of the universe increases in a spontaneous process
• Gibbs free energy (G): Measures a process's spontaneity; ΔG = ΔH - TΔS
• Spontaneous process: Occurs without external intervention (ΔG < 0)

This lesson introduces fundamental chemistry concepts, including the definition of matter, its states (solid, liquid, gas, plasma), and phase changes. It distinguishes between physical and chemical properties, providing examples and indicators of chemical change.