Chemical and Physical Properties and Changes

Definitions and Examples

• Chemical Properties: Characteristics that describe a substance's ability to undergo a specific chemical change (e.g., flammability, reactivity with acid).

• Physical Properties: Characteristics that can be observed or measured without changing the substance's identity (e.g., color, melting point, density).

• Chemical Change: A process where one or more substances are converted into new substances (e.g., rusting of iron).

• Physical Change: A change that does not alter the chemical composition (e.g., melting ice).

• Example: Boiling water is a physical change; burning wood is a chemical change.

Naming Compounds

Ionic and Molecular Compounds

• Ionic Compounds: Name the cation (metal) first, then the anion (nonmetal) with '-ide' ending (e.g., NaCl: sodium chloride).

• Molecular Compounds: Use prefixes to indicate the number of atoms (e.g., CO2: carbon dioxide).

• Acids: Binary acids use 'hydro-' and '-ic' (e.g., HCl: hydrochloric acid); oxyacids use '-ic' or '-ous' based on the polyatomic ion.

Isotopes and Atomic Notation

Understanding Isotopes

• Isotopes: Atoms of the same element with different numbers of neutrons.

• Atomic Notation: where X is the element symbol, A is mass number, Z is atomic number.

• Example: is a carbon isotope with 8 neutrons.

Periodic Law and Families of Elements

Organization of the Periodic Table

• Periodic Law: Properties of elements repeat periodically when arranged by increasing atomic number.

• Families/Groups: Vertical columns with similar properties (e.g., alkali metals, halogens, noble gases).

• Example: Group 1: Alkali metals; Group 17: Halogens.

Phases of Matter and Changes of State

Solid, Liquid, Gas, and Phase Changes

• Phases: Solid (definite shape/volume), liquid (definite volume, variable shape), gas (variable shape/volume).

• Changes of State: Melting, freezing, vaporization, condensation, sublimation, deposition.

• Example: Ice melting to water is fusion; water boiling is vaporization.

Trends in Electronegativity

Periodic Trends

• Electronegativity: The ability of an atom to attract electrons in a bond.

• Trend: Increases across a period (left to right), decreases down a group.

• Example: Fluorine is the most electronegative element.

Balancing Reactions

Law of Conservation of Mass

• Ensure the same number of each atom on both sides of the equation.

• Example: becomes

Single and Double Replacement Reactions

Types of Chemical Reactions

• Single Replacement:

• Double Replacement:

• Example: (single replacement)

Net Ionic Equations and Spectator Ions

Identifying Key Participants

• Net Ionic Equation: Shows only the species that change during the reaction.

• Spectator Ions: Ions that do not participate in the actual chemical change.

• Example:

Polarity: Bonds and Molecules

Electronegativity and Molecular Shape

• Polar Bond: Unequal sharing of electrons due to difference in electronegativity.

• Polar Molecule: Molecule with an uneven distribution of charge (e.g., H2O).

• Nonpolar: Equal sharing of electrons or symmetrical shape (e.g., O2).

Nuclear Reactions

Types and Notation

• Alpha Decay:

• Beta Decay:

• Gamma Emission:

Acids, Bases, and pH

Properties and Calculations

• Acids: Taste sour, turn litmus red, donate H+ ions.

• Bases: Taste bitter, turn litmus blue, accept H+ or donate OH-.

• pH:

• Example: A solution with M has pH 3.

Valence Electrons and Electron Configuration

Determining Electron Arrangement

• Valence Electrons: Electrons in the outermost shell, determine chemical reactivity.

• Electron Configuration: Order of filling: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, etc.

• Example: Sodium (Na): 1s2 2s2 2p6 3s1

Lewis Dot Structures and Molecular Shapes

Visualizing Molecules

• Lewis Dot Structures: Show valence electrons as dots around element symbols.

• VSEPR Theory: Predicts molecular shapes (e.g., linear, bent, trigonal planar, tetrahedral).

• Example: H2O is bent; CH4 is tetrahedral.

Hydrogen Bonding

Special Intermolecular Force

• Occurs when H is bonded to N, O, or F.

• Responsible for high boiling point of water.

Solubility: Like Dissolves Like, Temperature, and Pressure Effects

Factors Affecting Solubility

• Like Dissolves Like: Polar solvents dissolve polar solutes; nonpolar dissolves nonpolar.

• Temperature: Solubility of solids increases with temperature; gases decrease.

• Pressure: Solubility of gases increases with pressure (Henry's Law).

Ideal Gases

Gas Laws and Behavior

• Ideal Gas Law:

• P = pressure, V = volume, n = moles, R = gas constant, T = temperature (K).

• Assumes no intermolecular forces and negligible volume of particles.

Naming Hydrocarbons and Organic Functional Groups

Basics of Organic Chemistry

• Hydrocarbons: Alkanes (single bonds, -ane), alkenes (double bonds, -ene), alkynes (triple bonds, -yne).

• Functional Groups: Alcohol (-OH), carboxylic acid (-COOH), amine (-NH2), etc.

• Example: CH3CH2OH is ethanol (an alcohol).

Isomers

Structural Variations

• Isomers: Compounds with the same molecular formula but different structures.

• Example: C4H10: butane and isobutane.

Calculations in Chemistry

Key Quantitative Skills

• Metric Conversions: Use conversion factors (e.g., 1 kg = 1000 g).

• Significant Figures: Reflect precision in measurements and calculations.

• Density:

• Mole Calculations:

• Nuclear Half-Life:

• Mass Percent:

• Molarity:

• Dilution:

• Ideal Gas Law:

• Combined Gas Law:

• Percent Composition:

• Empirical Formula: Simplest whole-number ratio of atoms.

• Molecular Formula: Actual number of atoms; multiple of empirical formula.

• Stoichiometry: Use balanced equations to relate moles, mass, and volume.

• Limiting Reactant: Reactant that runs out first, limits product formed.

• Percent Yield:

Summary Table: Key Chemical Families

Summary Table: Common Functional Groups

Additional info:

• For empirical and molecular formula calculations, determine the simplest ratio first, then compare to molar mass for the molecular formula.

• For net ionic equations, only include ions and molecules directly involved in the reaction.

• For isomers, structural isomers differ in connectivity; geometric isomers differ in spatial arrangement.