Using Bond Energies to Estimate DH°rxn We often use average bond energies to estimate the DHrxn works best when all reactants and products in gas state Bond breaking is endothermic, DH(breaking) = + Bond making is exothermic, DH(making) = − DHrxn = ∑ (DH(bonds broken)) + ∑ (DH(bonds formed)) NOTE: Previous approach (with DHf’’s) used elements as intermediate state; this approach (with BE’s) uses “free atoms” as the intermediate state. Tro: Chemistry: A Molecular Approach, 2/e

Example: Estimate the enthalpy of the following reaction DHrxn = ∑ (DH(bonds broken)) + ∑ (DH(bonds made)) Endothermic! Bond breaking 1 mole C─H +414 kJ 1 mole Cl─Cl +243 kJ total +657 kJ DHrxn = (+657 kJ) + (−770 kJ) DHrxn = −113 kJ Bond Bond Energy (kJ/mol) C-H 414 Cl-Cl 243 C-Cl 339 H-Cl 431 Exothermic! Bond making 1 mole C─Cl −339 kJ 1 mole Cl─H −431 kJ total −770 kJ Tro: Chemistry: A Molecular Approach, 2/e 2

Break 1 mol C─H +414 kJ Make 1 mol Cl─Cl +243 kJ 1 mol C─Cl −339 kJ 1 mol H─Cl −431 kJ 3 Tro: Chemistry: A Molecular Approach, 2/e

Practice – Estimate the enthalpy of the following reaction equation + O DHrxn = (436 + 498 kJ) + [−142 kJ + 2(-464)] = (+934 kJ) + (−1070 kJ) = −136 kJ 4

Finishing up Covalent Bonding CHM 121 Fri, Wk 13F, Wk 10 Wk 13 (&14) Finishing up Covalent Bonding Electronegativity of Atoms; Polarity of Bonds (Section 9.6) (Note: Polarity of Molecules, Section 10.5, will be covered in Ppt 25 & PS12) Theories of Covalent Bonding Valence Bond Theory (VBT)_Section 10.6 Molecular Orbital Theory (Way better [but more complex] model! Not covered in this course [10.8]) VBT and Hybrid Orbitals (Section 10.7)

Electronegativity (of Atoms) and Polarity (of Bonds) See separate 1-page handout sheet See Section 9.6, but don’t worry about how electronegativity values were obtained by Pauling or about calculation of “% ionic character”. Be aware that Sec. 9.6 does not as clearly distinguish electronegativity (as a properties of an atom) from polarity (as a property of a bond) as I would like. See Figures 9.8 and 9.9 (in PowerPoint)

Electronegativity (EN) A number (no units!) which attempts to reflect how strongly an atom attracts bonding electrons. Bigger EN = Greater “pulling power” IS A PROPERTY OF ATOMS Trends as IE1 does (except for noble gases) Increases to the right (b/c Zeff increases) Decreases as you go down a column (r increases; farther away means less pull) F has greatest EN (4.0)

Figure 9.8 Pauling Electronegativities

Polarity The degree to which a covalent bond (or later, a whole molecule) has two differently charged ends. Greater polarity = Greater charge difference IS A PROPERTY OF BONDS (later, molecules) Results from a difference in EN values of two bonded atoms. “unequal sharing” No polarity = “nonpolar” (equal sharing)

Table 9.1,Fig. 9.9 Relationship Between Electro-negativity DIFFERENCE and Bond Type (i.e., Polarity) Nonpolar

Bond Polarity ENCl = 3.0 ENH = 2.1 3.0 – 2.1 = 0.9 Polar Covalent 3.0 − 3.0 = 0 Pure Covalent ENCl = 3.0 ENNa = 0.9 3.0 – 0.9 = 2.1 Ionic Tro: Chemistry: A Molecular Approach, 2/e

Valence Bond Theory Assumes that EACH ATOM making a bond has (uses) one of its ATOMIC orbitals to make the bond Each of these atomic orbitals contains one electron (“singly occupied orbital”) Bonding occurs only if the two orbitals “overlap” each other in the space between the atoms (better “orbital overlap” = stronger covalent bond)

Two Basic Kinds of Bonds in VBT (distinguished by the type of overlap) s (sigma) bond Overlap is along the bond axis (i.e., orbitals generally point “at” one another [except for s])  Rotation around the bond does not change overlap p (pi) bond Overlap is above and below the bond axis (i.e., “sideways” overlap)  Rotation around the bond does change overlap

See Board for O2 and N2 Youtube video showing this with “graphics”: https://www.youtube.com/watch?v=PRmYlUuDXVI

Why hybrid orbitals were conceived (With sp3 hybridization) (Without hybridization)

Why Hybrid Orbitals? (“pure” [unhybridized] vs. hybrid orbitals) s, p, d, (and f) orbitals are “pure” atomic orbitals. If these were the only orbitals used for bonding, only 180 and 90 bond angles could be explained! (see board example) p orbitals are perpendicular to one another Concept of “hybrid orbitals” was conceived From mathematical combination of 2 or more “pure” orbitals on same atom Have names that “show” the number and kinds of pure orbitals used to create them sp, sp2, sp3, sp3d, etc. Explains the angles predicted by VSEPR (and observed)… 120 and 109

“Conservation of orbitals” Idea, & Shapes of hybrid orbitals When you hybridize atomic orbitals, you always get the SAME NUMBER of orbitals that you started out with. two pure atomic orbitals => two hybrid orbitals (in a “set”) three pure orbitals => a set of three hybrid orbitals Every hybrid orbital has two “lobes”, but one is bigger than the other, so the smaller one is typically ignored (not drawn)

Shapes (Continued) Orbitals in a “set” of hybrids will all have the same shape, but point in different directions similar idea to that of a set of three p orbitals: px, py, pz Short, simple summary of hybrid orbital sets (very little “theory”, but okay for basic “result”): https://www.youtube.com/watch?v=vDREeV9nLEo

Tro, Figure 10.7 s + p + p + p yields four sp3 orbitals

s + p = two sp hybrid orbitals CHM 121 Wk 13 (&14) s + p = two sp hybrid orbitals From McMurry text

s + p + p = three sp2 hybrid orbitals (like Fig. 10.8 in Tro)

Figure 9.10, Zumdahl When An s and Two p Orbitals Are Hybridized to Form a Set of Three sp2 Orbitals, One p Orbital Remains Unchanged and is Perpendicular to the Plane of the Hybrid Orbitals

~10-minute YouTube Videos http://www.youtube.com/watch?v=d1E18tBTlBg **Was not accessible on 11/22/16** ??? https://www.youtube.com/watch?v=cPDptc0wUY https://www.youtube.com/watch?v=vHXViZTxLXo “molecular shape and orbital hybridization” This will help you visualize what is happening better than I can show on the board or in a static picture (as in PowerPoint). It will not help you determine which hybrid orbitals are used by a given atom—that is most easily done using an LDS [as I will describe]

Consistent with VSEPR Predictions It turns out that the “directions” in which the hybrid orbitals in a set “point” match precisely with the VSEPR geometries!!!

Final Generalizations (“It turns out that….”) The # hybrid orbitals needed always = the # of e- -clouds (from LDS/VSEPR model) Hybrid orbitals are ONLY used for s-bonds (or lone pairs) The only way to get p bonds for simple molecules (i.e., in this class) is to have/use “leftover” (pure) p orbitals i.e., it only happens when sp or sp2 hybrids are used for the s-bonding Any single bond is a s bond; any bond AFTER a single bond is a p bond A double bond is always 1 s and ond p A triple bond is always 1 s and two p

Example on Board H2CO

Strategy / Approach

Tro notation: hybridization and bonding scheme Wedge and dash depiction

Figure 9.11 The Sigma Bonds in Ethylene

Figure 9.12 A Carbon-Carbon Double Bond Consists of a s and a p Bond

Figure 9.13 (Zumdahl) (a)The Orbitals Used to Form the Bonds in Ethylene (b) The Lewis Structure for Ethylene

Tro’s VBT sketch of Ethylene Chapter 10, Unnumbered Figure 1, Page 432

Tro shows dichloroethene (bonding analogous to ethene except for Cl’s)

sp hybrids used for either two double bonds or a triple bond

Figure 9.17 The Orbitals of an sp Hybridized Carbon Atom

Chapter 10, Unnumbered Figure, Page 428

Chapter 10, Figure 10.11 sp 3d Hybridization

Chapter 10, Figure 10.12 sp 3d 2 Hybridization

Chapter 10, Unnumbered Figure, Page 429

Click to add notes

Chapter 10, Unnumbered Figure 2, Page 431

Chapter 10, Unnumbered Figure 4, Page 431

Chapter 10, Unnumbered Figure 3, Page 431

Figure 9.21 A Set of dsp3 Hybrid Orbitals on Phosphorus Atom

Figure 9.22b The Orbitals Used to Form the Bonds in PCl5 NOTE: This text author (UNLIKE TRO) uses hybrid orbitals rather than p orbitals on terminal atoms. I will use Tro’s convention.

Figure 9.23 An Octahedral Set of d2sp3 Orbitals on Sulfur Atom

Xenon

Chapter 10, Unnumbered Figure 4, Page 450