using bond energies to estimate h° rxn we often use average bond energies to estimate the h rxn...

Using Bond Energies to Estimate H°rxn

• We often use average bond energies to estimate the Hrxn

– works best when all reactants and products in gas state

• Bond breaking is endothermic, H(breaking) = +• Bond making is exothermic, H(making) = −

Hrxn = ∑ (H(bonds broken)) + ∑ (H(bonds formed))

1

Tro: Chemistry: A Molecular Approach, 2/e

Example: Estimate the enthalpy of the following reaction

H C

H

H

H

Cl Cl H C

H

Cl

H

H Cl+ +

Bond breaking1 mole C─H +414 kJ1 mole Cl─Cl +243 kJ

total +657 kJ

Bond making1 mole C─Cl −339 kJ1 mole Cl─H −431 kJ

total −770 kJ

Hrxn = ∑ (H(bonds broken)) + ∑ (H(bonds made))

Hrxn = (+657 kJ) + (−770 kJ)Hrxn = −113 kJ

Tro: Chemistry: A Molecular Approach, 2/e 2

Bond Bond Energy(kJ/mol)

C-H 414Cl-Cl 243C-Cl 339H-Cl 431

Break1 mol C─H +414 kJ1 mol Cl─Cl +243 kJ

Make1 mol C─Cl −339 kJ1 mol H─Cl −431 kJ

Tro: Chemistry: A Molecular Approach, 2/e 3

Practice – Estimate the enthalpy of the following reaction

H H + O O H O O H

4

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 = 3.03.0 − 3.0 = 0

Pure Covalent

ENCl = 3.0ENH = 2.1

3.0 – 2.1 = 0.9Polar Covalent

ENCl = 3.0ENNa = 0.9

3.0 – 0.9 = 2.1Ionic

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)

• (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

• (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

(Without hybridization)

(With sp3 hybridization)

Why hybrid orbitals were conceived

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

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

s + p = two sp hybrid orbitals

From McMurry text

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

Copyright © Houghton Mifflin Company. All rights reserved.

9–22

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 Video

• http://www.youtube.com/watch?v=d1E18tBTlBg

• “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 -bonds (or lone pairs)

• The only way to get 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 -bonding

• Any single bond is a bond; any bond AFTER a single bond is a bond– A double bond is always 1 and ond

– A triple bond is always 1 and two

Example on Board

• H2CO

Str

ateg

y / A

ppro

ach

Tro notation: hybridization and bonding scheme

O C

H

H

Wedge and dash depiction

Copyright © Houghton Mifflin Company. All rights reserved.

9–31

Figure 9.11 The Sigma Bonds in Ethylene

Copyright © Houghton Mifflin Company. All rights reserved.

9–32

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

Copyright © Houghton Mifflin Company. All rights reserved.

9–33

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

Structure for Ethylene

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

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

Copyright © Houghton Mifflin Company. All rights reserved.

9–37

Figure 9.17 The Orbitals of an sp Hybridized Carbon Atom

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9–52

Figure 9.21 A Set of dsp3 Hybrid Orbitals on Phosphorus Atom

Copyright © Houghton Mifflin Company. All rights reserved.

9–53

Figure 9.22b The Orbitals Used to Form the Bonds in PCl5

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9–54

Figure 9.23 An Octahedral Set of d2sp3 Orbitals on Sulfur Atom

Copyright © Houghton Mifflin Company. All rights reserved.

9–55

Xenon