AL

Chapter 1–7: Intermolecular Forces, Dipoles, and Bonding (Vocabulary Flashcards)

  • IMF and polarity recap

    • IMF stands for intermolecular forces, which are the attractive forces that exist between molecules. These are distinct from intramolecular forces (covalent or ionic bonds) which hold atoms within a molecule together. Major categories discussed include:

      • Dipole-dipole forces: These exist exclusively between polar molecules. They arise from the electrostatic attraction between the permanent partial positive end of one molecule and the permanent partial negative end of an adjacent polar molecule. The strength of these interactions depends directly on the magnitude of the dipole moment.

      • London dispersion forces (LDFs): Also known as induced dipole-induced dipole interactions, these forces are present in all molecules, both polar and nonpolar. They originate from instantaneous, temporary dipoles that form due to the momentary uneven distribution of electron density around a nucleus. These instantaneous dipoles then induce temporary dipoles in neighboring molecules, leading to weak, temporary attractions.

      • Ion-dipole interactions: These occur between an ion (a fully charged species) and a polar molecule (which has a permanent dipole moment). For example, when an ionic compound dissolves in water, the charged ions are attracted to the polar water molecules.

    • Permanent dipoles vs. induced dipoles: Permanent dipoles are an intrinsic property of polar molecules, arising from persistent differences in electronegativity between bonded atoms, leading to a net uneven distribution of electron charge across the molecule. Induced dipoles, by contrast, are transient and arise when the electron cloud of an atom or molecule is temporarily distorted (polarized) by the presence of a nearby electric field (e.g., from another instantaneous dipole or an ion).

    • A higher dipole moment generally signifies a greater separation of partial positive and negative charges within a molecule, which leads to stronger electrostatic attractions between molecules in polar species. Consequently, this results in stronger intermolecular attractions.

  • Polarity and VSEPR basics (with a practical workflow)

    • It is crucial to understand that the presence of polar bonds within a molecule does not automatically make the entire molecule polar. The overall molecular polarity is a vector sum of all individual bond dipoles, and it significantly depends on the molecule's three-dimensional geometry, as determined by Valence Shell Electron Pair Repulsion (VSEPR) theory. If the bond dipole vectors cancel out due to symmetry, the molecule will be nonpolar; otherwise, it will be polar.

    • Step-by-step approach to determine molecular polarity:
      1) Identify electronegative atoms and estimate bond polarities: Use the concept of electronegativity (EN) differences between bonded atoms. A bond dipole moment is created when there is an uneven sharing of electrons. The difference in electronegativity, represented as \Delta EN = EN{\text{X}} - EN{\text{Y}}, helps quantify this. A larger difference typically yields a larger bond dipole. For general classification, bonds with (approximately) \Delta EN < 0.5 are considered nonpolar, between 0.5 and 1.7 are polar covalent, and > 1.7 are ionic.
      2) Determine electron-domain geometry (VSEPR): Identify the central atom(s) and count the number of electron domains (bonding pairs and lone pairs). Use VSEPR theory to predict the arrangement of these electron domains, which dictates the electron-domain geometry (e.g., linear, trigonal planar, tetrahedral). This, in turn, influences the molecular geometry.
      3) Draw bond dipole vectors: For each polar bond, draw a vector arrow pointing from the less electronegative (partially positive, $\delta^+$) atom to the more electronegative (partially negative, $\delta^-$) atom. The length of the arrow can qualitatively represent the magnitude of the bond dipole.
      4) Assess for net dipole: Sum the bond dipole vectors. If the vectors are symmetrical and cancel each other out, the net dipole moment is zero, and the molecule is nonpolar. If the vectors do not cancel, resulting in a net nonzero dipole moment, the molecule is polar. This vector sum considers both the magnitude and direction of each bond dipole.

    • Example note: Bonds between carbon and hydrogen (CH bonds) are generally considered nonpolar due to their small electronegativity difference (approx. 0.35). However, when heteroatoms like chlorine (Cl), fluorine (F), or oxygen (O) are present, the bond dipoles can be significantly larger due to their higher electronegativity. The resulting net molecular polarity depends on the overall vector sum of these individual bond dipoles and the molecular geometry, not just the presence or absence of individual polar bonds.

    • Geometry example: For a central atom with three electron domains and no lone pairs, the electron-domain geometry and molecular geometry are trigonal planar, yielding ideal bond angles of approximately 120^{\circ}. In a perfectly symmetrical trigonal planar molecule (e.g., BF_3), even if individual B-F bonds are polar, the symmetry causes the bond dipoles to cancel, making the molecule nonpolar.

  • Dipole moment, geometry, and boiling points

    • A larger net dipole moment indicates a greater overall separation of charge within a molecule, which subsequently leads to stronger intermolecular attractions, particularly stronger dipole-dipole forces between adjacent molecules. These stronger attractions require more energy to overcome during phase transitions.

    • Consequently, molecules with larger net dipole moments (i.e., polar molecules) tend to have higher boiling points and melting points compared to nonpolar molecules of comparable size and molecular mass. For example, in a typical pair of molecules with similar molecular formulas and masses (e.g., propene vs. cyclopropane), the polar molecule often exhibits a significantly higher boiling point than its nonpolar analog due to these stronger dipole-dipole interactions enhancing the overall IMF strength.