States of Matter and Classification of Substances

Physical States of Matter

The three primary physical states of matter are solid, liquid, and gas. Each state is characterized by distinct particle arrangements and properties.

• Solids: Definite shape and volume; particles are closely packed in a fixed arrangement.

• Liquids: Definite volume but no definite shape; particles are close but can move past one another.

• Gases: No definite shape or volume; particles are far apart and move freely.

State changes include processes such as melting (solid to liquid), freezing (liquid to solid), condensation (gas to liquid), and vaporization (liquid to gas).

Physical and Chemical Properties

Physical properties can be observed without changing the substance's identity (e.g., melting point, density). Chemical properties describe a substance's ability to undergo chemical changes (e.g., flammability, reactivity).

Classification of Matter

• Elements: Pure substances consisting of only one type of atom (e.g., O2).

• Compounds: Pure substances composed of two or more elements chemically combined (e.g., H2O).

• Mixtures: Physical combinations of two or more substances.

  • Homogeneous mixtures (solutions): Uniform composition throughout (e.g., saltwater).

  • Heterogeneous mixtures: Non-uniform composition (e.g., salad, sand in water).

Molecules are two or more atoms bonded together. All compounds are molecules, but not all molecules are compounds (e.g., O2 is a molecule but not a compound).

Reading chemical formulas: The subscript indicates the number of atoms of each element (e.g., H2O has 2 hydrogen and 1 oxygen atom).

Measurement, Significant Figures, and Calculations

Significant Figures (Sig Figs)

• Definition: Digits in a measurement that are known with certainty plus one estimated digit.

• Exact numbers: Have an infinite number of significant figures (e.g., counting objects, defined conversions).

• Rules for determining sig figs:

  • Nonzero digits are always significant.

  • Zeros between nonzero digits are significant.

  • Leading zeros are not significant.

  • Trailing zeros are significant only if there is a decimal point.

• Rounding: Round to the correct number of sig figs based on the operation performed.

Dimensional Analysis and Unit Conversions

• Dimensional analysis: A method to convert units using conversion factors so that unwanted units cancel.

• SI prefixes: nano (n, 10-9), micro (μ, 10-6), milli (m, 10-3), kilo (k, 103).

• English-metric conversions: Use provided conversion factors (e.g., 1 km = 1094 yards).

Density Calculations

• Density (d): The mass per unit volume of a substance.

• Given any two of density, mass, and volume, you can solve for the third.

Atomic Structure and the Periodic Table

Atomic Number, Mass Number, and Isotopes

• Atomic number (Z): Number of protons in the nucleus; defines the element.

• Mass number (A): Total number of protons and neutrons.

• Isotopes: Atoms of the same element (same Z) with different numbers of neutrons (different A); chemically similar because they have the same number of electrons.

Nuclide symbol: AZX, where X is the element symbol.

Subatomic Particles

• Protons: Positive charge, located in the nucleus.

• Neutrons: No charge, located in the nucleus.

• Electrons: Negative charge, located outside the nucleus.

For ions, electrons = protons - charge.

Average Atomic Mass

The average atomic mass is calculated from the masses and relative abundances of isotopes:

The Periodic Table

• Know names and symbols for elements 1–36, Rb, Cs, Sr, Ba, and I.

• Metals: Left and center of the table; conduct electricity, malleable.

• Nonmetals: Right side; poor conductors, brittle.

• Common ion charges:

  • Group 1A: 1+ cations

  • Group 2A: 2+ cations

  • Group 3A: 3+ cations

  • Group 5A: 3- anions

  • Group 6A: 2- anions

  • Group 7A: 1- anions

• Diatomic elements: H2, N2, O2, F2, Cl2, Br2, I2

• Alkali metals: Group 1A

• Alkaline earth metals: Group 2A

• Halogens: Group 7A

• Noble gases: Group 8A

Molar Mass and Avogadro's Number

• Molar mass: Mass of one mole of a substance (g/mol).

• Avogadro's number: particles/mol.

• Conversions:

  • Grams ↔ moles ↔ molecules ↔ atoms

Light, Atomic Structure, and Quantum Theory

Electromagnetic Spectrum and Light

• Wavelength (λ), frequency (ν), and energy (E): Related by the equations:

  where c = speed of light, h = Planck's constant.

• Visible spectrum: 400–750 nm; violet (~400 nm), red (~750 nm), with colors in order ROYGBIV.

• Relationship: Wavelength and frequency are inversely related; energy is directly proportional to frequency.

Absorption and Emission

• Absorption: Electron absorbs energy and moves to a higher energy level.

• Emission: Electron releases energy and falls to a lower energy level.

• Spectra: Absorption and emission spectra are unique to each element.

Bohr Model and Electron Transitions

• Bohr model: Electrons orbit the nucleus in quantized energy levels.

• Ground state: Lowest energy level.

• Excited state: Any higher energy level.

• Rydberg equation: Used to calculate energy changes for electron transitions in hydrogen.

where is the Rydberg constant, and are initial and final energy levels.

Quantum Numbers and Orbitals

• Principal quantum number (n): Energy level (n = 1, 2, 3, ...)

• Angular momentum quantum number (l): Subshell (l = 0 to n-1)

• Magnetic quantum number (ml): Orientation (-l to +l)

• Spin quantum number (ms): +1/2 or -1/2

• Number of orbitals in a shell: n2

• Number of electrons in a shell: 2n2

• Shapes: s orbitals are spherical; p orbitals are dumbbell-shaped.

Electron Configurations and the Periodic Table

• Aufbau principle: Electrons fill lowest energy orbitals first.

• Pauli exclusion principle: No two electrons in an atom can have the same set of four quantum numbers.

• Hund's rule: Electrons fill degenerate orbitals singly before pairing.

• Complete and condensed configurations: Use noble gas shorthand for condensed form.

• Core electrons: Inner electrons; valence electrons: outermost electrons involved in bonding.

• Octet rule: Atoms tend to gain, lose, or share electrons to achieve eight valence electrons (duet for H).

• Isoelectronic: Atoms/ions with the same electron configuration.

• Metals form cations (lose electrons); nonmetals form anions (gain electrons).

• Ground state: Lowest energy configuration; excited state: at least one electron in a higher orbital.

Periodic Trends

• Effective nuclear charge (Zeff): Net positive charge experienced by valence electrons; increases across a period.

• Atomic radius: Decreases across a period, increases down a group.

• Ionic radius: Cations are smaller, anions are larger than their parent atoms.

• Ionization energy: Energy required to remove an electron; increases across a period, decreases down a group.

• Successive ionization energies: Each subsequent electron removed requires more energy; large jump after valence electrons are removed.

Example Table: Periodic Trends

Example: Sodium (Na) has a lower ionization energy than chlorine (Cl) because it is further left on the periodic table and has a larger atomic radius.

Additional info: These notes synthesize and expand upon the study guide points, providing definitions, explanations, and examples for each topic relevant to a first-semester General Chemistry course.