Kaplan MCAT General Chemistry Review полностью

NH₃ has three hydrogen atoms bonded to the central nitrogen and one lone pair on the central nitrogen. These four groups—three atoms, one lone pair—lead NH₃ to be sp³ hybridized. By hybridizing all three p-orbitals and the one s-orbital, four groups are arranged about the central atom, maximizing the distances between the groups to minimize the energy of the configuration. NH₃’s hybridization leads to its tetrahedral electronic geometry yet trigonal pyramidal molecular geometry. In contrast, BF₃ has three atoms but no lone pairs, resulting in sp² hybridization. Its shape is called trigonal planar. (A) is incorrect; although BF₃ does has three bonded atoms and no lone pairs, its geometry is not trigonal pyramidal. (C) is tempting because NH₃ has a lone pair, but this answer choice is not complete enough to explain the differences in molecular geometry. Finally, although (D) is also true, the polarity of the molecules does not explain their geometry; rather, the molecules’ different geometries contribute to the overall polarity of the molecules.

16. B

Most atoms require eight valence electrons to follow the octet rule. However, some atoms, like beryllium, can have fewer than eight valence electrons (suboctet). As a result, when bonding with chloride, beryllium is likely to form only two bonds, using its own two outer valence electrons and one from each chloride to form BeCl₂, as drawn in (B). (A) is incorrect because Be cannot both bond to chloride and have lone pairs. (C) assumes that all three atoms need their octets filled, and (D) places 10 electrons on each chlorine.

17. B

This answer requires an understanding of the trends that cause higher or lower bond energies. Bonds of high energy are those that are difficult to break. These bonds tend to have more shared pairs of electrons and, thus, cause a stronger attraction between the two atoms in the bonds. This stronger attraction also means that the bond length of a high-energy, high-order bond (i.e., a triple bond) is shorter than that of its lower-energy counterparts (i.e., single or double bonds). Thus, (A) is incorrect; bond energy increases with decreasing bond length. Single bonds are longer than double bonds. Thus, (C) is incorrect; single bonds are easier to break than double bonds. Finally, as previously discussed, bond energy is inversely related to bond length, making (D) incorrect.

18. A

Polarity is precisely described by the first answer choice. Dipole moments describe the relationship of shared electrons between two bonded atoms. If there is a dipole moment between two atoms, then the electrons in their shared bond are preferentially centered on one atom, typically the more electronegative of the two. The polarity of a molecule is then defined as the vector sum of these dipole moments in the molecule’s three-dimensional configuration. For example, although individual carbon–chloride bonds have dipole moments with a partial negative charge on chloride, the molecule CCl₄ has a tetrahedral configuration. Thus, the four polarized bonds cancel each other out for no net dipole moment, creating a nonpolar molecule. Based on this description, it is clear that (B) is incorrect, because molecular geometry is an essential component of determining molecular polarity. For the same reasons described previously, it is possible for a nonpolar molecule to contain one or more polar bonds, making (C) incorrect. Finally, it is possible for a molecule to contain at least one nonpolar bond and yet be a polar molecule. For example, CH₃CH₂Cl has a nonpolar bond (C–C), yet a molecular dipole exists. In its tetrahedral arrangement, the one polar bond between C–Cl will have a dipole moment, creating a net polarity of the molecule in the direction of the chloride atom.

Chapter 4: Compounds and Stoichiometry

Oh—what is that smell? You and your mother are taking an afternoon walk through the public rose garden, a pleasant activity for a pleasant summer day, and everything was so pleasant until you smelled . . . it, whatever it is. Your mother doesn’t have to wonder long until her own olfactory system is screeching the alarm, too. Oh, it’s just horrible. It smells like rancid almonds. Although your better instincts tell you to walk quickly away, you pause and, holding your breath, crouch down to get a better look. You are a premed, after all, and you did just synthesize cinnamaldehyde in your orgo lab, so your curiosity is understandable.