Chemistry Definitions: Atoms, Ions, and More

Basic Definitions

• Atom: The smallest unit of an element retaining its chemical properties, comprised of protons, neutrons, and electrons.
• Element: A pure substance made of only one type of atom (e.g., oxygen, carbon).
• Compound: A substance made of two or more different elements chemically bonded (e.g., water, H₂O).
• Molecule: Two or more atoms chemically bonded, elements (O₂) or compounds (H₂O).
• Ion: An atom or molecule that has gained or lost electrons, acquiring a charge (e.g., Na⁺, Cl⁻).
• Isotope: Atoms of the same element, same number of protons, different numbers of neutrons (e.g., carbon-12 and carbon-14).
• Cation: A positively charged ion (loses electrons, e.g., Na⁺).
• Anion: A negatively charged ion (gains electrons, e.g., Cl⁻).
• Atomic Number: The number of protons in an atom's nucleus, determining the element.
• Mass Number: The total number of protons and neutrons in an atom's nucleus.
• Molar Mass: The mass of one mole of a substance, measured in grams per mole (g/mol).
• Mole: A unit representing 6.022 × 10²³ particles of a substance (Avogadro's number).
• Valence Electrons: Electrons in the outermost energy level of an atom, important in bonding.
• Covalent Bond: A chemical bond where two atoms share electrons (e.g., H₂O, CO₂).
• Ionic Bond: A chemical bond where electrons are transferred, creating charged ions that attract (e.g., NaCl).
• Metallic Bond: A bond between metal atoms where electrons move freely, allowing conductivity.
• Electronegativity: An atom's ability to attract electrons in a bond.
• Polar Covalent Bond: A bond where electrons are shared unequally, creating partial charges (e.g., H₂O).
• Nonpolar Covalent Bond: A bond where electrons are shared equally (e.g., O₂, CH₄).
• Solute: The substance dissolved in a solution (e.g., salt in saltwater).
• Solvent: The substance that dissolves the solute (e.g., water in saltwater).
• Solution: A homogeneous mixture of solute and solvent (e.g., saltwater).
• Acid: A substance donating H⁺ ions in solution (pH < 7, e.g., HCl).
• Base: A substance accepting H⁺ ions or donating OH⁻ ions (pH > 7, e.g., NaOH).
• pH Scale: Measures acidity/basicity of a solution (0-14, with 7 as neutral).
• Buffer: A solution resisting pH changes when acids/bases are added.
• Oxidation: Loss of electrons in a reaction.
• Reduction: Gain of electrons in a reaction.
• Redox Reaction: A reaction involving both oxidation and reduction.
• Catalyst: A substance speeding up a reaction without being consumed.
• Endothermic Reaction: A reaction absorbing heat (e.g., melting ice).
• Exothermic Reaction: A reaction releasing heat (e.g., combustion).

Kinetic Molecular Theory of Matter

• Matter comprises atoms/particles.
• Particles are always moving (possessing kinetic energy).
• The movement dictates the matter's phase.
• Solid State: Matter maintains fixed volume/shape; particles vibrate in a fixed position.
• Liquid state: Matter maintains a fixed volume, taking the container's shape.
• Gas state: Matter expands to fill available volume.

State Dependence

• Kinetic Energy: Dictates particle movement level.
• Intermolecular Forces: Hold particles/molecules together in solid and liquid phases.
• Temperature: Higher temperatures mean faster particle movement, and lower temperatures mean slower particle movement.

Phase Changes

• Melting (Solid to Liquid): Solid absorbs heat, particles gain energy and move more freely (e.g., ice to water).
• Freezing (Liquid to Solid): Liquid loses heat, particles slow and form a rigid structure.
• Evaporation/Vaporization (Liquid to Gas): Boiling occurs throughout the liquid at a specific boiling point, while evaporation occurs at the liquid's surface below the boiling point (e.g., water boiling into steam).
• Condensation (Gas to Liquid): Gas loses heat; particles slow and form a liquid (e.g., water vapor to dew).
• Sublimation (Solid to Gas): Solid gains energy, transitioning directly into a gas (e.g., dry ice turning into carbon dioxide gas).
• Deposition (Gas to Solid): Gas loses energy, transitioning directly into a solid (e.g., frost forming on a cold surface).
• Ionization (Gas to Plasma): Gas gains energy, stripping electrons to form ions (e.g., lightning creating plasma).
• Demonization (Recombination) (Plasma to Gas): Plasma loses energy, allowing ions and electrons to recombine (e.g., plasma in fluorescent lights cooling).

Intermolecular Forces

• Intramolecular forces hold atoms together within a molecule, including covalent, ionic, and metallic bonds.
• Intermolecular forces exist between molecules, including dipole-dipole interactions, hydrogen bonds, and London dispersion forces.
• London Dispersion Forces: Weakest intermolecular forces, exist in all substances via instantaneous dipoles; strength increases with more electrons (e.g., O₂, CH₃-CH₃⁺, halogens).
• Dipole-Dipole Interactions: Occur in polar molecules with permanent dipole moments; arise from electronegativity differences causing uneven electron sharing.
• The difference in electronegativity results in uneven sharing of electron that one results in partially positive and another partially negative.
• Polarity: Polar liquids mix with polar liquids; water (polar) does not mix with oil (non polar).
• Hydrogen Bonds: Strongest dipole-dipole interactions between hydrogen and highly electronegative atoms (F, O, N) in inorganic (e.g., water) and organic molecules (e.g., DNA, proteins).

Properties of Liquids

• Surface Tension: Result of cohesive forces between surface molecules, allowing resistance to external forces and droplet formation; influenced by intermolecular forces and temperature.
• Capillary Action: Liquid's ability to flow in narrow spaces without external force, driven by adhesion and cohesion.
• Viscosity: Liquid's resistance to flow, depends on intermolecular forces and temperature, and high-viscosity liquids (e.g., honey) flow slower.
• Vapor Pressure: Pressure exerted by a liquid's vapor when in equilibrium with its liquid phase; higher vapor pressure indicates greater volatility.
• Boiling Point: Temperature at which a liquid's vapor pressure equals atmospheric pressure, leading to phase transition.
• Molar Heat of Vaporization: Energy required to convert one mole of liquid into vapor, influenced by intermolecular forces.

Additional Examples Related to Liquid Properties

• Incompressibility: Solids and liquids cannot be squeezed into smaller volumes due to tightly packed particles. When pressing water in a bottle, its volume remains constant.
• Diffusion: Particles move from high to low concentration, spreading evenly over time. For example, perfume sprayed in a room slowly spreads.
• Evaporation: Slow change of a liquid into a gas at the surface without boiling, like a puddle disappearing on a sunny day.
• Cooling Effect of Evaporation: Liquid absorbs heat, cooling the surface, such as when sweat evaporates.

Properties of Water

• Water as a Solvent: Water molecules are attracted to other polar molecules and ions.
• Hydrophilic (water-loving) substances dissolve well in water (e.g., salt NaCl, sugar), while hydrophobic (water-fearing) substances do not (e.g., oil, fats).
• Cohesion: Property of water molecules sticking together.
• Adhesion: Tendency of water molecules to be attracted to other substances.
• Density: Mass per unit volume. Liquid water's density is 1 g/cm³, while ice is 0.9168 g/cm³.
• Specific Heat Capacity: Heat needed to raise one gram of substance by one degree Celsius.
• Heat of Vaporization: Heat needed to turn 1g of liquid into vapor without temperature increase; requires much heat to vaporize water due to breaking hydrogen bonds.

Types of Solids

• Solids are classified into amorphous and crystalline based on the arrangement of their components.
• Crystalline Solids: Organized with components arranged in a highly ordered and repeated manner, usually periodic with a high melting point; example of jewelry
• Amorphous solids components don't have the regular organization that crystalline molecules do; examples are glass and rubber.

Types of Solutions

• Solutions are homogenous of two or more substances distributed through a single phase with uniform composition and properties.
• A solution has components of solute and solvent.

Classifications

• Physical State: Solutions classified based on their physical attributes after solute dissolved components.
• Solid Solution: compacted solution with jewelry, coins, and bronze brass.
• Liquid Solution are juices, soft drinks, and coffee.
• Gas solutions are smoke, air.
• Concentration: Solutions that get classified according to their ration of solute to solvent
• Dilute: Contains relatively small amount of solvent in a given solution
• Concentrated solutions contains a large amount of solute in a given solvent.
• Saturation: Solutions that are defined to the solubility limit
• Solubility Limit: maximum about of solvent.
• Unsaturated solutions contains amount that is less that the solubility limit.
• Saturated solution contains the maximum about of solution that a solvent can dissolve.
• Super saturated solutions contains solution that has more of a solution that can dissolve.

Energy and Solutions

• Energy of soultions is formed by spontaneous process
• Entropy: The measure of disorder of uncertainty in the system
• Exothermic energy is released
• Endothermic energy is absorbed
• Enthapy is the total heat contained in the system

Expressing Solutions

• Concentration: Ratio of solutions and solvents
• General term to express the quantity of soultions and solvents
• Percent by mass: mass of soultions x 100 / Mass of solution.
• Percent by volume is expressed that amount of solutants can be more easily dissolved in solvent by volume
• Percent by volume : Volume of solute x 100 / volume of solution.
• Percent by volume: the mass volume is the mass
• Grams of solute is 100 ml of solute