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Chapter 10

Chapter 10. The Shapes of Molecules. The Shapes of Molecules. 10.1 Depicting Molecules and Ions with Lewis Structures. 10.2 Using Lewis Structures and Bond Energies to Calculate Heats of Reaction. 10.3 Valence-Shell Electron-Pair Repulsion (VSEPR)

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Chapter 10

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  1. Chapter 10 The Shapes of Molecules

  2. The Shapes of Molecules 10.1 Depicting Molecules and Ions with Lewis Structures 10.2 Using Lewis Structures and Bond Energies to Calculate Heats of Reaction 10.3 Valence-Shell Electron-Pair Repulsion (VSEPR) Theory and Molecular Shape 10.4 Molecular Shape and Molecular Polarity

  3. Rules for Lewis Structures • 1. Attach atoms in reasonable fashion with single bonds • Consider ABn formulas, atom with lower group number goes in center with others attached to it. If same group number, place atom with highest period in center. • The atom with the lowest E.N. goes in middle • Consider # of free valence e-, this is often the number of bonds this atom will make. Atoms making more bonds will be central in molecules. • 2. Sum valence electrons (adding for neg charge and sub for pos.) • 3. Complete octets of peripheral atoms (not He like atoms) • 4. Place leftover e- on central atom • 5. If necessary use multiple bonds to fill center atom's octet • Last Slide

  4. Molecular formula For NF3 Atom placement : : N 5e- : F : : F : : Sum of valence e- F 7e- X 3 = 21e- N Total 26e- : F : : Remaining valence e- Lewis structure

  5. PROBLEM: Write a Lewis structure for CCl2F2, one of the compounds responsible for the depletion of stratospheric ozone. PLAN: Follow the steps outlined in Figure 10.1 . : : : Cl : : : Cl : C F : : : : F : SAMPLE PROBLEM 10.1 Writing Lewis Structures for Molecules with One Central Atom SOLUTION: Cl Step 1: Carbon has the lowest EN and is the central atom. The other atoms are placed around it. Cl C F F Steps 2-4: C has 4 valence e-, Cl and F each have 7. The sum is 4 + 4(7) = 32 valence e-. Make bonds and fill in remaining valence electrons placing 8e- around each atom.

  6. PROBLEM: Write the Lewis structure for methanol (molecular formula CH4O), an important industrial alcohol that is being used as a gasoline alternative in car engines. SAMPLE PROBLEM 10.2 Writing Lewis Structure for Molecules with More than One Central Atom SOLUTION: Hydrogen can have only one bond so C and O must be next to each other with H filling in the bonds. There are 4(1) + 4 + 6 = 14 valence e-. C has 4 bonds and O has 2. O has 2 pair of nonbonding e-. H : H C O H : H

  7. H H H H C C C C H H H H N : N : . . N : N : N : N : . . . . SAMPLE PROBLEM 10.3 Writing Lewis Structures for Molecules with Multiple Bonds. PROBLEM: Write Lewis structures for the following: (a) Ethylene (C2H4), the most important reactant in the manufacture of polymers (b) Nitrogen (N2), the most abundant atmospheric gas PLAN: For molecules with multiple bonds, there is a Step 5which follows the other steps in Lewis structure construction. If a central atom does not have 8e-, an octet, then e- can be moved in to form a multiple bond. SOLUTION: (a) There are 2(4) + 4(1) = 12 valence e-. H can have only one bond per atom. : (b) N2 has 2(5) = 10 valence e-. Therefore a triple bond is required to make the octet around each N.

  8. Resonance: Delocalized Electron-Pair Bonding O3 can be drawn in 2 ways - Neither structure is actually correct but can be drawn to represent a structure which is a hybrid of the two - a resonance structure. Resonance structures have the same relative atom placement but a difference in the locations of bonding and nonbonding electron pairs.

  9. SAMPLE PROBLEM 10.4 Writing Resonance Structures PROBLEM: Write resonance structures for the nitrate ion, NO3-. PLAN: After Steps 1-4, go to 5 and then see if other structures can be drawn in which the electrons can be delocalized over more than two atoms. SOLUTION: Nitrate has 1(5) + 3(6) + 1 = 24 valence e- N does not have an octet; a pair of e- will move in to form a double bond.

  10. Three criteria for choosing the more important resonance structure: Resonance (continued) Formal charge of atom = # valence e- - (# unshared electrons + 1/2 # shared electrons) Smaller formal charges (either positive or negative) are preferable to larger charges; Avoid like charges (+ + or - - ) on adjacent atoms; A more negative formal charge should exist on an atom with a larger EN value.

  11. For OC For OA # valence e- = 6 # valence e- = 6 # nonbonding e- = 6 # nonbonding e- = 4 # bonding e- = 2 X 1/2 = 1 # bonding e- = 4 X 1/2 = 2 For OB Formal charge = -1 Formal charge = 0 # valence e- = 6 # nonbonding e- = 2 # bonding e- = 6 X 1/2 = 3 Formal charge = +1 Formal Charge: Selecting the Best Resonance Structure An atom “owns” all of its nonbonding electrons and half of its bonding electrons. Formal charge of atom = # valence e- - (# unshared electrons + 1/2 # shared electrons)

  12. A B C Resonance (continued) EXAMPLE: NCO- has 3 possible resonance forms - formal charges -2 0 +1 -1 0 0 0 0 -1 Forms B and C have negative formal charges on N and O; this makes them more important than form A. Form C has a negative charge on O which is the more electronegative element, therefore C contributes the most to the resonance hybrid.

  13. Exceptions to the Octet Rule • Electron deficient – have fewer than eight • ex BeCl2, BF3 • may attain an octet by coordinate covalent bond • Odd number of electrons – aka free radicals • ex NO2 • May attain an octet by pairing with another free radical • Expanded Octets – only on period 3 and higher • Expanded octets form when an atom can decrease (or maintain at 0) it’s formal charge • ex SF6, PCl5, H2SO4, SO2, SO3, SO4

  14. PROBLEM: Write Lewis structures for (a) H3PO4 and (b) BFCl2. In (a) decide on the most likely structure. SAMPLE PROBLEM 10.5 Writing Lewis Structures for Exceptions to the Octet Rule. PLAN: Draw the Lewis structures for the molecule and determine if there is an element which can be an exception to the octet rule. Note that (a) contains P which is a Period-3 element and can have an expanded valence shell. SOLUTION: (a) H3PO4 has two resonance forms and formal charges indicate the more important form. -1 0 +1 0 (b) BFCl2 will have only 1 Lewis structure. 0 0 0 0 0 0 0 0 0 0 more stable 0 0 lower formal charges

  15. Bond Energies and Hrxn • In chemical reactions, reactant bonds are broken and new product bonds are formed. The overall energy change of a reaction is the energy change of this process, plus the energy associated with changes of state. • In gaseous reactions (no state changes) • Hºrxn = Hºbonds broken + Hºbonds formed

  16. Figure 10.2 Using bond energies to calculate DH0rxn DH0rxn = DH0reactant bonds broken + DH0product bonds formed Enthalpy, H DH01 = + sum of BE DH02 = - sum of BE DH0rxn

  17. 2[-BE(C O)]= -1598kJ DH0rxn= -818kJ Figure 10.3 Using bond energies to calculate DH0rxn of methane BOND BREAKAGE 4BE(C-H)= +1652kJ 2BE(O2)= + 996kJ DH0(bond breaking) = +2648kJ BOND FORMATION 4[-BE(O-H)]= -1868kJ Enthalpy,H DH0(bond forming) = -3466kJ

  18. PROBLEM: Use Table 9.2 (button at right) to calculate DH0rxn for the following reaction: CH4(g) + 3Cl2(g) CHCl3(g) + 3HCl(g) PLAN: Write the Lewis structures for all reactants and products and calculate the number of bonds broken and formed. SAMPLE PROBLEM 10.6 Calculating Enthalpy Changes from Bond Energies SOLUTION: bonds broken bonds formed

  19. SAMPLE PROBLEM 10.6 Calculating Enthalpy Changes from Bond Energies continued bonds broken bonds formed 4 C-H = 4 mol(413 kJ/mol) = 1652 kJ 3 C-Cl = 3 mol(-339 kJ/mol) = -1017 kJ 3 Cl-Cl = 3 mol(243 kJ/mol) = 729 kJ 1 C-H = 1 mol(-413 kJ/mol) = -413 kJ 3 H-Cl = 3 mol(-427 kJ/mol) = -1281 kJ DH0bonds broken = 2381 kJ DH0bonds formed = -2711 kJ DH0reaction = DH0bonds broken + DH0bonds formed = 2381 kJ + (-2711 kJ) = - 330 kJ

  20. Molecular Shapes Lewis structures show which atoms are connected where, and by how many bonds, but they don't properly show 3-D shapes of molecules. To find the actual shape of a molecule, first draw the Lewis structure, and then use VSEPR Theory.

  21. Valence Shell Electron-Pair Repulsion Theory or VSEPR • molecular shape is determined by the repulsions of electron pairs • Electron pairs around the central atom stay as far apart as possible. • electron pair geometry - based on number of regions of electron density • Consider non-bonding (lone pairs) as well as bonding electrons. • Electron pairs in single, double and triple bonds are treated as single electron clouds. • molecular geometry - based on the electron pair geometry, this is the shape of the molecule

  22. Figure 10.4 A balloon analogy for the mutual repulsion of electron groups.

  23. Figure 10.5 Electron-group repulsions and the five basic molecular shapes.

  24. Figure 10.6 The single molecular shape of the linear electron-group arrangement. Examples: CS2, HCN, BeF2

  25. Class Shape Figure 10.7 The two molecular shapes of the trigonal planar electron-group arrangement. Examples: SO2, O3, PbCl2, SnBr2 Examples: SO3, BF3, NO3-, CO32-

  26. 1220 Effect of Double Bonds 1160 real Effect of Nonbonding Pairs Factors Affecting Actual Bond Angles Bond angles are consistent with theoretical angles when the atoms attached to the central atom are the same and when all electrons are bonding electrons of the same order. 1200 larger EN 1200 ideal greater electron density lone pairs repel bonding pairs more strongly than bonding pairs repel each other 950

  27. Figure 10.8 The three molecular shapes of the tetrahedral electron-group arrangement. Examples: CH4, SiCl4, SO42-, ClO4- NH3 PF3 ClO3 H3O+ H2O OF2 SCl2

  28. Figure 10.9 Lewis structures and molecular shapes

  29. Figure 10.10 The four molecular shapes of the trigonal bipyramidal electron-group arrangement. PF5 AsF5 SOF4 SF4 XeO2F2 IF4+ IO2F2- XeF2 I3- IF2- ClF3 BrF3

  30. Figure 10.11 The three molecular shapes of the octahedral electron-group arrangement. SF6 IOF5 BrF5 TeF5- XeOF4 XeF4 ICl4-

  31. See Figure 10.1 Figure 10.12 The steps in determining a molecular shape. Molecular formula Step 1 Lewis structure Count all e- groups around central atom (A) Step 2 Electron-group arrangement Note lone pairs and double bonds Step 3 Count bonding and nonbonding e- groups separately. Bond angles Step 4 Molecular shape (AXmEn)

  32. PROBLEM: Draw the molecular shape and predict the bond angles (relative to the ideal bond angles) of (a) PF3 and (b)COCl2. SOLUTION: (a) For PF3 - there are 26 valence electrons, 1 nonbonding pair SAMPLE PROBLEM 10.7 Predicting Molecular Shapes with Two, Three, or Four Electron Groups The shape is based upon the tetrahedral arrangement. The F-P-F bond angles should be <109.50 due to the repulsion of the nonbonding electron pair. The final shape is trigonal pyramidal. <109.50 The type of shape is AX3E

  33. 124.50 1110 SAMPLE PROBLEM 10.7 Predicting Molecular Shapes with Two, Three, or Four Electron Groups continued (b) For COCl2, C has the lowest EN and will be the center atom. There are 24 valence e-, 3 atoms attached to the center atom. C does not have an octet; a pair of nonbonding electrons will move in from the O to make a double bond. Type AX3 The shape for an atom with three atom attachments and no nonbonding pairs on the central atom is trigonal planar. The Cl-C-Cl bond angle will be less than 1200 due to the electron density of the C=O.

  34. PROBLEM: Determine the molecular shape and predict the bond angles (relative to the ideal bond angles) of (a) SbF5 and (b) BrF5. SOLUTION: (a) SbF5 - 40 valence e-; all electrons around central atom will be in bonding pairs; shape is AX5 - trigonal bipyramidal. SAMPLE PROBLEM 10.8 Predicting Molecular Shapes with Five or Six Electron Groups (b) BrF5 - 42 valence e-; 5 bonding pairs and 1 nonbonding pair on central atom. Shape is AX5E, square pyramidal.

  35. Figure 10.13 The tetrahedral centers of ethane.

  36. Figure 10.13 The tetrahedral centers of ethanol.

  37. PROBLEM: Determine the shape around each of the central atoms in acetone, (CH3)2C=O. PLAN: Find the shape of one atom at a time after writing the Lewis structure. tetrahedral tetrahedral trigonal planar >1200 <1200 SAMPLE PROBLEM 10.9 Predicting Molecular Shapes with More Than One Central Atom SOLUTION:

  38. Electric field OFF Electric field ON Figure 10.14 The orientation of polar molecules in an electric field.

  39. PROBLEM: From electronegativity (EN) values (button) and their periodic trends, predict whether each of the following molecules is polar and show the direction of bond dipoles and the overall molecular dipole when applicable: PLAN: Draw the shape, find the EN values and combine the concepts to determine the polarity. SOLUTION: (a) NH3 SAMPLE PROBLEM 10.10 Predicting the Polarity of Molecules (a) Ammonia, NH3 (b) Boron trifluoride, BF3 (c) Carbonyl sulfide, COS (atom sequence SCO) The dipoles reinforce each other, so the overall molecule is definitely polar. ENN = 3.0 ENH = 2.1 molecular dipole bond dipoles

  40. SAMPLE PROBLEM 10.10 Predicting the Polarity of Molecules continued (b) BF3 has 24 valence e- and all electrons around the B will be involved in bonds. The shape is AX3, trigonal planar. F (EN 4.0) is more electronegative than B (EN 2.0) and all of the dipoles will be directed from B to F. Because all are at the same angle and of the same magnitude, the molecule is nonpolar. 1200 (c) COS is linear. C and S have the same EN (2.0) but the C=O bond is quite polar(DEN) so the molecule is polar overall.

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