Chapter 9 The Shapes of Molecules 1 Cocaine
10.1 Depicting Molecules & Ions with Lewis Structures 2
Number of Covalent Bonds 3 The number of covalent bonds can be determined from the number of electrons needed to complete an octet.
Molecules with Single Bonds 4 Ammonia, NH 3 1. Decide on the central atom; never H. Central atom is atom of lowest affinity for electrons. Therefore, N is central 2. Count total valence electrons H = 1 and N = 5 Total = (3 x 1) + 5 = 8 electrons / 4 pairs
Drawing Lewis Structure 5 3. Form a single bond between the central atom and each surrounding atom H N H H 4. Remaining electrons form LONE PAIRS to complete octet as needed. 3 BOND PAIRS and 1 LONE PAIR. Note that N has a share in 3 pairs (6 electrons), while H shares 1 pair.
Sulfite ion, SO 3 2-6 Remaining pairs become lone pairs, first on outside atoms and then on central atom. Each atom is surrounded by an octet of electrons.
Molecules with Multiple Bonds Formaldehyde H 2 CO Step 4: Place lone pairs on outer atoms. 7 Step 5: To fulfill the octet on C, we form DOUBLE BONDS between C and O.
Double and even triple bonds are commonly observed for C, N, P, O, and S H 2 CO 8 SO 3 N 2 C 2 F 4 CO 2
RESONANCE: Delocalized Electron-Pair Bonding 9 These equivalent structures are called RESONANCE STRUCTURES. The true electronic structure is a HYBRID of the two.
Electron Delocalization 10 Lewis structures depict electrons as localized either on an individual atom (lone pairs) or in a bond between two atoms (shared pair). In a resonance hybrid, electrons are delocalized: their density is spread over a few adjacent atoms.
Fractional Bond Orders 11 Resonance hybrids often have fractional bond orders due to partial bonding. For O 3, bond order = 3 electron pairs 2 bonded-atom pairs = 1½
Nitrate ion, NO 3-12 Each atom is surrounded by an octet of electrons.
Formal Charge: Selecting the More Important Resonance Structure Formal charge is the charge an atom would have if all electrons were shared equally. Formal charge of atom = # of valence e - - (# of unshared valence e - + ½ # of shared valence e - ) 13 For O A in resonance form I, the formal charge is given by 6 valence e - - (4 unshared e - + ½(4 shared e - ) = 6 4 2 = 0
Determine Formal Charge - Formal charges must sum to the actual charge on the species for all resonance forms. O A [6 4 ½(4)] = 0 O B [6 2 ½(6)] = +1 O C [6 6 ½(2)] = -1 14 O A [6 6 ½(2)] = -1 O B [6 2 ½(6)] = +1 O C [6 4 ½(4)] = 0 For both these resonance forms the formal charges sum to zero, since O 3 is a neutral molecule. - Smaller formal charges (positive or negative) are preferable to larger ones. - The same nonzero formal charges on adjacent atoms are not preferred. Avoid like charges on adjacent atoms. - A more negative formal charge should reside on a more electronegative atom.
Choosing the More Important Resonance Form Example: NCO has 3 possible resonance forms: +2 0-1 -1 0 0 0 0-1 15 Resonance forms with smaller formal charges are preferred. Resonance form I is therefore not an important contributor. A negative formal charge should be placed on a more electronegative atoms. Resonance form III is preferred to resonance form II. The overall structure of the NCO - ion is still an average of all three forms, but resonance form III contributes most to the average.
Formal Charge Versus Oxidation Number For a formal charge, bonding electrons are shared equally by the atoms. The formal charge of an atom may change between resonance forms. 16 Formal charges +2 0-1 -1 0 0 0 0-1 -3 +4-2 -3 +4-2 -3 +4-2 Oxidation numbers For an oxidation number, bonding electrons are transferred to the more electronegative atom. The oxidation number of an atom is the same in all resonance forms.
Exceptions to the Octet Rule 17 Molecules with Electron-Deficient Atoms B and Be are commonly electron-deficient. Odd-Electron Species A molecule with an odd number of electrons is called a free radical.
Exceptions to the Octet Rule 18 Expanded Valence Shells An expanded valence shell is only possible for nonmetals from Period 3 or higher because these elements have available d orbitals.
10.2 The VSEPR Model 19 Valence Shell Electron Pair Repulsion theory. Most important factor in determining geometry is relative repulsion between electron pairs. In general, the relative repulsion strength are in the order: LP-LP > LP-BP > BP-BP Molecule adopts the shape that minimizes the electron pair repulsions.
Electron Group Arrangements & Molecular Shapes 20
Molecular shape with TWO Electron Groups (Linear Arrangement) Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. 21 AX 2 Examples: CS 2, HCN, BeF 2 This key refers to Figures 10.3 through 10.8. Figure 10.3
Molecular Shapes with THREE Electron Groups (Trigonal Planar Arrangement) Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. 22 AX 3 Examples: SO 2, O 3, PbCl 2, SnBr 2 AX 2 E Examples: SO 3, BF 3, NO 3, CO 3 2 Figure 10.4
Factors Affecting Bond Angles 23 Nonbonding (Lone) Pairs A lone pair repels bonding pairs more strongly than bonding pairs repel each other. This decreases the angle between the bonding pairs. Double Bonds A double bond has greater electron density than a single bond, and repels the single bond electrons more than they repel each other.
Molecular Shapes with FOUR Electron Groups (Tetrahedral Arrangement) Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. 24 AX 4 Examples: CH 4, SiCl 4, SO 4 2, ClO 4 AX 3 E Examples: NH 3, PF 3 ClO 3, H 3 O + AX 2 E 2 Examples: H 2 O, OF 2, SCl 2 Figure 10.5
Lewis structures do not indicate molecular shape. 25 twist to the right twist to the left Figure 10.6
Molecular Shapes with FIVE Electron Groups (Trigonal Bipyramidal Arrangement) Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. AX 5 Examples: PF 5, AsF 5, SOF 4 AX 4 E Examples: SF 4, XeO 2 F 2 IF 4+, IO 2 F 2 26 AX 3 E 2 Examples: ClF 3, BrF 3 AX 2 E 3 Examples: XeF 2, I 3, IF 2 Figure 10.7
Axial and Equatorial Positions 27 A five electron-group system has two different positions for electron groups, and two ideal bond angles. Equatorial-equatorial repulsions are weaker than axial-equatorial repulsions. Where possible, lone pairs in a five electron-group system occupy equatorial positions.
Molecular Shapes with SIX Electron Groups (Octahedral Arrangement) 28 AX 6 Examples: SF 6, IOF 5 AX 5 E AX 4 E 2 Examples: BrF 5, TeF 5, XeOF 4 Examples: XeF 4, ICl 4 Figure 10.8
Molecular shapes for central atoms in Period 2 and in higher periods 29 Figure 10.9
The four steps in converting a molecular formula to a molecular shape 30 Molecular Formula Step 1 Draw Lewis structure. Lewis structure Step 2 Count all e - groups around central atom (A). Electrongroup arrangement Step 3 Bond angles Note positions of any lone pairs and double bonds. Step 4 Count bonding and nonbonding e - groups separately. Figure 10.11 Molecular shape (AX m E n )
31 Structure Determination by VSEPR Ammonia, NH 3 1. Draw electron dot structure 2. Count BP s and LP s = 4 3. The 4 electron pairs are at the corners of a tetrahedron. H N H lone pair of electrons in tetrahedral position H
Structure Determination by VSEPR Formaldehyde, H 2 CO 1. Draw electron dot structure 2. Count BP s and LP s at C H O 3. There are 3 electron lumps around C at the corners of a planar triangle. C H 32 O C The electron pair geometry is TRIGONAL PLANAR with 120 o bond angles. H H The molecular geometry is also trigonal planar.
Molecular Shapes with More than One Central Atom Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. 33 ethane CH 3 CH 3 ethanol CH 3 CH 2 OH Figure 10.12
Molecular Shapes with More than One Central Atom 34 Methanol, CH 3 OH Determine H-C-H and C-O-H bond angles H-C-H = 109 o C-O-H = 109 o In both cases the atom is surrounded by 4 electron groups. H H C O H H
Molecular Shapes with More than One Central Atom 35 Acetonitrile, CH 3 CN Determine bond angles H H C C N H-C-H = 109 o C-C-N = 180 o 109 H 180 One C is surrounded by 4 electron groups and the other C by 2 electron groups
10.3 Molecular Shape 36 & Molecular Polarity Overall molecular polarity depends on both shape and bond polarity. The polarity of a molecule is measured by its dipole moment (μ), which is given in the unit debye (D). A molecule is polar if - it contains one or more polar bonds and - the individual bond dipoles do not cancel.
The orientation of polar molecules in an electric field. 37 Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. Molecules are randomly oriented. Figure 10.13 Molecules become oriented when the field is turned on.
Bond Polarity, Bond Angle, and Dipole Moment 38 Compare CO 2 and H 2 O. Which one is polar?
Bond Polarity, Bond Angle, and Dipole Moment Molecules with the same shape may have different polarities. CH 4 CCl 4 Polar or Not? 39 Only CH 4 and CCl 4 are NOT polar. They are the two molecules that are symmetrical.
40 Polar or Nonpolar? Consider AB 3 molecules: BF 3, Cl 2 CO, and NH 3.
Effect of Molecular Polarity on Behavior 41 Example: cis and trans isomers of 1,2-dichloroethylene, C 2 H 2 Cl 2 The cis isomer is polar while the trans isomer is not. The boiling point of the cis isomer boils is 13 C higher than that of the trans isomer.