Given the following molecules: – Draw the lewis structure (don’t forget to show the number of valance electrons) – Determine the polarity (polar vs. non-polar). Show directions of dipole moments/arrows to support your answers. – List all intermolecular forces. Indicate strongest force. a) BH3 b) SO2 c) H2CO3 d) XeF2 e) C2H4
The Correct Answer and Explanation is:
Let’s break down the Lewis structures, polarity, and intermolecular forces for each molecule:
a) BH3 (Borane)
- Lewis Structure: Boron (B) is in the center with three single bonds to three hydrogen atoms. Boron only has six valence electrons (3 from H atoms and 3 from its own electrons), so it is electron-deficient.
- Valence Electrons: Boron (3) + Hydrogen (3 × 1) = 6 electrons.
- Polarity: Non-polar. The molecule has a trigonal planar shape due to the sp2 hybridization of boron. Since all the bonds are identical and there is no lone pair on boron, the dipoles cancel out.
- Intermolecular Forces: London Dispersion Forces (LDF) are the only intermolecular force present here, as there are no significant dipoles or hydrogen bonds.
b) SO2 (Sulfur Dioxide)
- Lewis Structure: Sulfur (S) is the central atom with two double bonds to two oxygen (O) atoms. Sulfur has a lone pair of electrons.
- Valence Electrons: Sulfur (6) + Oxygen (2 × 6) = 18 electrons.
- Polarity: Polar. The molecule has a bent shape due to the lone pair on sulfur. The two oxygen atoms pull electron density away from sulfur, creating a net dipole pointing toward the more electronegative oxygens.
- Intermolecular Forces: Dipole-Dipole Interactions are the strongest intermolecular force due to the permanent dipole of the molecule. LDF also plays a role.
c) H2CO3 (Carbonic Acid)
- Lewis Structure: Carbon (C) is at the center with two single bonds to oxygen atoms and one double bond to another oxygen. One of the oxygen atoms is attached to a hydrogen atom (OH group).
- Valence Electrons: Carbon (4) + Oxygen (3 × 6) + Hydrogen (2 × 1) = 18 electrons.
- Polarity: Polar. The molecule has a trigonal planar shape with polar bonds. The dipole moments do not cancel, resulting in a net dipole.
- Intermolecular Forces: Hydrogen Bonding is the strongest force, in addition to dipole-dipole interactions and LDF.
d) XeF2 (Xenon Difluoride)
- Lewis Structure: Xenon (Xe) is at the center with two single bonds to fluorine (F) atoms. Xenon also has three lone pairs of electrons.
- Valence Electrons: Xenon (8) + Fluorine (2 × 7) = 22 electrons.
- Polarity: Non-polar. Even though the bonds are polar, the molecule has a linear shape, and the dipoles cancel out because of symmetry.
- Intermolecular Forces: London Dispersion Forces are the primary intermolecular forces.
e) C2H4 (Ethene or Ethylene)
- Lewis Structure: Two carbon atoms double-bonded to each other, each bonded to two hydrogen atoms.
- Valence Electrons: Carbon (2 × 4) + Hydrogen (4 × 1) = 12 electrons.
- Polarity: Non-polar. The molecule is symmetrical, with equal sharing of electrons between the carbon atoms and hydrogens. The dipoles cancel out.
- Intermolecular Forces: London Dispersion Forces are the strongest force due to the relatively simple structure.
Summary:
- BH3: Non-polar, LDF.
- SO2: Polar, Dipole-Dipole.
- H2CO3: Polar, Hydrogen Bonding.
- XeF2: Non-polar, LDF.
- C2H4: Non-polar, LDF.
The strongest intermolecular forces vary, but hydrogen bonding in H2CO3 is the strongest among these examples.
