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Organic Chemistry Exam 1 Review

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Termini in questo insieme (26)

  • Basic structure of atoms

    Atoms consist of protons (positive charge, in nucleus), neutrons (neutral, in nucleus), and electrons (negative charge, orbiting nucleus).

  • Electron configuration

    Describes the distribution of electrons in atomic orbitals, reflecting the wavelike properties of electrons.

  • Valence electrons

    Electrons in the outermost shell of an atom; determine bonding behavior and reactivity.

  • Lewis dot structures

    Diagrams showing valence electrons as dots around element symbols to represent covalent bonding.

  • Octet rule

    Atoms tend to gain, lose, or share electrons to achieve 8 valence electrons, except H (2), sometimes Be and B.

  • Formal charge formula

    Formal charge = (valence electrons in free atom) – (lone pair electrons) – ½(bonding electrons).

  • Common bonding patterns

    Elements like C, H, O, N, halogens have typical numbers of bonds and lone pairs; deviations affect formal charge.

  • Multiple bonds in Lewis structures

    Use double or triple bonds when single bonds and lone pairs don't satisfy octet rule; triple > double > single in strength and bond order.

  • Resonance structures

    Different valid Lewis structures for the same molecule; actual structure is a hybrid; resonance stabilizes ions.

  • Sigma (σ) vs Pi (π) bonds

    σ bonds form by end-to-end orbital overlap; π bonds form by side-to-side overlap with a nodal plane between nuclei.

  • Bond rotation

    Single bonds (σ) can rotate freely; double/triple bonds cannot because rotation breaks π bonds.

  • Double and triple bond composition

    A double bond = 1 σ + 1 π bond; a triple bond = 1 σ + 2 π bonds.

  • Electronegativity

    Measure of an atom's ability to attract electrons in a bond; trends increase across a period and up a group.

  • Bond polarity classification

    Nonpolar covalent: ΔEN < 0.5; Polar covalent: 0.5 ≤ ΔEN ≤ 2.0; Ionic: ΔEN > 2.0.

  • Electron pair geometry

    Shape based on electron pairs around central atom: linear (2), trigonal planar (3), tetrahedral (4).

  • Molecular shape

    Shape considering bonding and lone pairs: linear, trigonal planar, bent, tetrahedral, trigonal pyramidal.

  • Molecular polarity

    Determined by bond polarity and shape; polar bonds may cancel in symmetrical molecules, making them nonpolar.

  • Lone pair effects on shape

    Lone pairs repel more than bonding pairs, causing bond angles to adjust and shapes to distort.

  • Isomers

    Compounds with the same molecular formula but different structures or shapes.

  • Hybridization types

    Atomic orbitals mix to form hybrids: sp (linear, 180°), sp2 (trigonal planar, ~120°), sp3 (tetrahedral, ~109°).

  • Drawing organic structures

    Convert among Lewis, Kekulé, condensed, and skeletal structures; interpret number of atoms, bonds, and lone pairs.

  • Intermolecular forces

    Include London dispersion, dipole-dipole, and hydrogen bonding; strength order: London < dipole-dipole < hydrogen bonds << covalent bonds.

  • Boiling and melting points

    Increase with stronger intermolecular forces and larger molecular size; branching lowers boiling points.

  • Solubility principles

    "Like dissolves like"; polar or ionic compounds dissolve in polar solvents (hydrophilic), nonpolar dissolve in nonpolar solvents (hydrophobic).

  • Arrhenius acids and bases

    Acids produce \(\mathrm{H_3O^+}\) in water; bases produce \(\mathrm{OH^-}\).

  • Brønsted acids

    Proton donors that release \(\mathrm{H^+}\) ions.