A student forgets that the N in ammonia, NH₃, has a lone pair as well as its three single bonds. After checking Table 6.2, he mistakenly draws the molecule—three bonding groups, no lone pairs— as having a trigonal planar shape. If the ammonia molecule really were trigonal planar, how would the intermolecular forces differ from what they actually are?

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Table 6.2 VSPER Molecular Shape Table Four electron groups (SN = 4) Molecular shape Electron-group arrangement 109.5° Four bonding groups Tetrahedral No lone pairs Tetrahedral Bond angles all 109.5⁰ Three bonding groups Pyramidal One lone pair Tetrahedral Bond angles compressed to ~107° by lone pair* -105° Two bonding groups Two lone pairs Tetrahedral Bent Bond angle compressed to ~105° by lone pair* Three electron groups (SN = 3) Molecular shape Electron-group arrangement 120° Three bonding groups No lone pairs Trigonal planar Trigonal planar Bond angles all 120° -118° Two bonding groups One lone pair Trigonal planar Bent Bond angle compressed to ~118° by lone pair* Two electron groups (SN = 2) Electron-group Molecular shape arrangement 180° Two bonding groups No lone pairs Linear Linear Bond angle 180° Key: SN = Steric Number = # of electron groups = # attached atoms Central atom around central atom and # of ion pairs Bonding electron group (single, double, or triple covalent bond) Lone pair of electrons