Chemical Bonding IV Valence Bond Theory Builds Directly Upon What We Have Learned About the Electron Configuration of Atoms. V

www.apchemsolutions.com Valence Bond Theory builds directly upon Lecture 9 what we have learned about the electron configuration of atoms.

Chemical Bonding IV Valence Bond Theory is very good at predicting what will bond with what in order Valence Bond Theory to complete octets, but it doesn’t explain the Hybrid Orbital Theory bond angles we observe in VSEPR Theory. Multiple Bonds Polar and Non-Polar Molecules Hybrid Orbital Theory builds on VSEPR Theory to help explain the angles we observe.

Valence Bond Theory

• Combines Lewis’ theory of filling octets by sharing pairs of electrons, with the According to Valence Bond Theory, electronic configuration of atomic orbitals. electrons can be shared when atomic orbitals from different atoms overlap. • This theory states that bonding occurs when atomic orbital overlap.

Remember that Boron produces an incomplete octet with only 6 valence electrons.

Building BF3 with Right now is has 3 valence electrons, so it Valence Bond Theory needs 3 more.

B: Where do they go? 1s 2s 2p F: It is clear that Fluorine atoms can overlap 1s 2s 2p their 2pz orbitals with half filled orbitals of another atoms, but Boron only has one half filled orbital.

It would appear as though Boron can only accept one electron from one Fluorine atom in order to make one bond.

Building BF3 with Valence Bond Theory Now we have three orbitals that can accept electrons. B: 1s 2s 2p The 2pz orbital of three different fluorine atoms can now overlap with one of these Boron enters an excited state, where an three half filled orbitals in order to share electron from the 2s orbital is promoted to electrons and create bonds. the 2py orbital.

Valence Bond Theory has The central atom is the one that determines the shape, so we don’t need to draw all of the problems with the shape s and p orbitals in Fluorine.

2py F In this structure two of the fluorine atoms are o B at 90 to one another and the other is floating 2px F 2s around on the s orbital.

Valence Bond Theory o Views bonding as an Gets the bond VSEPR Theory predicts bond angles of 120 overlapping of atomic angle wrong F (Trigonal Planar – 3 charge clouds, 3 bonds, orbitals. 0 lone pairs) Problems with Valence Bond Theory

and (BF3) Problem (the bond angle is wrong) • Valence Bond Theory says 90o and a random angle for the fluorine bonded to the s-orbital. • VSEPR theory says 120o Solution (Hybrid Orbitals)

• When an electron is promoted; the 2s, 2px, and 2py orbitals of boron morph into three separate sp2 hybrid orbitals that are identical in shape and size. The size of the sp2 orbitals are exaggerated.

sp2 Hybrid Orbitals They are called sp2 hybrid orbitals as they are made out of one s and two p orbitals. 2 F Each sp is one part s and two parts p.

sp2 You can also think of the name in terms of charge clouds around the central atom. sp2 sp2 There are 3 charge clouds and s+p+p=three F F charge clouds. The 2s and two 2p orbitals morph into three identical sp2 hybrid orbitals The bond angel is now 120o, as VSEPR Theory predicts.

© 2008 AP Chem Solutions. All rights reserved. 2 Tutorials to assist you with this material are available online at www.apchemsolutions.com www.apchemsolutions.com If you combine one s orbital with 2 p orbitals you get 3 identical sp2 hybrid orbitals. sp2 Hybrid Orbitals The large lobe of the sp2 orbital is the part sp2 that overlaps with the bonding orbital of the terminal atom. s ++p = p sp2 As the shared electrons spend 95% of their time in this orbital the shape is reasonable.

sp2 Both shared electrons will spend the majority of their time between the two nuclei, and only a short time in the small lobe on the other side of the central atom’s nucleus.

3 sp Hybrid Orbitals (e.g. CH4) Carbon only appears to have two orbitals H: with single unpaired electrons. 1s C: Thus one would think that it can only bond 1s 2s 2p with a maximum of 2 Hydrogen atoms.

Carbon enters an excited state as it is getting ready to bond with the Hydrogens. sp3 Hybrid Orbitals (e.g. CH ) 4 An electron is promoted from the 2s to the 2pz orbital. H: 1s This gives four bonding sites. C: 1s 2s 2p Hybridization of Carbon These four sites morph into four orbitals that C: are exactly the same. 1s sp3 They are all one part s and three parts p, so we call them sp3 hybrid orbitals.

sp3 hybrid orbitals have a tetrahedral geometry.

In this case it is a perfect tetrahedral with bond angles of 109.5o (4 charge clouds, 4 bonds, zero lone pairs). sp3 Hybrid Orbitals (e.g. CH ) 4 Water also has a tetrahedral geometry of charge clouds but the bond angle is 104.5o. H Bond Angle sp = 109.5o 3 All of the sp3 hybrid orbitals are exactly the

sp3 sp3 same. H sp3 H H Each is one part s-orbital and three parts p- orbital.

Thus they are called sp3 hybrid orbitals.

You can remember that they are sp3 hybrid orbitals as there is 4 charge clouds and 1s+1p+1p+1p= 4sp3. Valence Bond Theory and Lone Pairs

e.g. Building H O with valence bond theory 2 Valence bond theory works perfectly for H: H: predicting that 2 hydrogen atoms combine 1s 1s with 1 oxygen atom to produce 1 water O: 1s 2s 2p molecule. Oxygen could accept one electron from one

Hydrogen in its py orbital and another from the other Hydrogen in its pz orbital.

px has been drawn so that it is coming out of the board (straight at you), so you cannot see Valence Bond Theory and Lone Pairs the back half of it.

• Valence Bond Theory views bonding as an Again, there are three identical p orbitals overlapping of atomic orbitals. here. They only differ in orientation. One on the x axis (turned to be pointing straight at H py py you in this case), one on the y axis, and one

p px z p p on the z axis. z x Valence H Bond Theory According to the previous slide we had 2 get the angle wrong again! electrons in px, one in py and one in pz before bonding.

© 2008 AP Chem Solutions. All rights reserved. 4 Tutorials to assist you with this material are available online at www.apchemsolutions.com www.apchemsolutions.com After bonding the orbitals are full but the angel is wrong.

Valence Bond theory gives an angle of 90o and VSEPR theory gives an angle of 104.5o.

In the picture on the right hand side, the py and pz orbitals are not the right shape for sharing a pair of electrons with hydrogen.

An orbital is a picture that represents where an electron spends 95% of its time. In this picture at least one of the shared electrons is spending 95% of its time on the other side of the oxygen nucleus from where the hydrogen is sitting. In Valence Bond Theory, shared electrons will move into and out of the lobe of the p- orbital that is on the opposite side of Problems With Valence Bond oxygen’s nucleus from where the bonding is taking place. Theory in the H O Example 2 • The bond angle is wrong. If these orbitals did not change shape during • Valence Bond theory says 90o bonding, the bonding electrons would be • VSEPR Theory says 104.5o spending half of their time in a region where they do not attract the nucleus of hydrogen • The orbital shape must be wrong. atom. • Shared electrons are not spending enough time with the Hydrogen. The orbitals must change shape so that the bonding electrons will spend most of their time between the nuclei of oxygen and hydrogen.

Hybrid Orbital Theory and Lone Pairs

Since there are 4 charge clouds, four sp3 H: H: hybrid orbitals will be created. 1s 1s O: Two of them have lone pairs and two of them 1s sp3 acquire bonds with Hydrogen. Electrons are not promoted here, but every orbital in the n=2 energy level becomes hybridized.

© 2008 AP Chem Solutions. All rights reserved. 5 Tutorials to assist you with this material are available online at www.apchemsolutions.com www.apchemsolutions.com In 3D this structure has an imperfect tetrahedral geometry. There are 4 sp3 hybrid orbitals here (as there are 4 charge clouds).

3 sp Hybrid Orbitals (H2O) They are not identical. The lone pairs will spend most of their time closer to oxygen’s nucleus than will the bonding pairs.

3 sp 3 sp3 Thus, the sp hybrid orbitals containing the sp3 H lone pairs will be smaller than the sp3 hybrid Four sp3 orbitals that are involved in bond. hybrid orbitals sp3 Bond Angle are formed = 104.5o This pushes the charge clouds that bond with H the hydrogen’s a little closer together, thereby reducing the bond angle from 109.5o to 104.5o.

5 Charge Clouds! (e.g. PCl5)

.. .. :Cl: :Cl .. .. P Cl.. : :Cl.. : :Cl.. :

The extra 3d orbitals don’t disappear as they never actually existed. Orbitals, as just a sp3d Hybrid Orbitals concept. They are the area where an electron Atomic Orbitals can be found 95% of the time. They are not P: there unless the electrons are there. 3s 3p 3d Promotion These hybrid orbitals are each made of one P: part s, three parts p, and one part d. 3s 3p 3d Hybridization Thus, they are called sp3d hybrid orbitals. P: 3 sp d They occur when there are 5 charge clouds around the central atom.

6 Charge Clouds!!! (e.g. SF6)

.. .. :F: .. Sulfur hexafluoride :F.. ..F: S :F: .. :F.. : :..F:

sp3d2 Hybrid Orbitals Atomic Orbitals S: In this case two electrons are promoted. 3s 3p 3d Promotion The end result is 6 different but identical S: 3 2 3s 3p 3d sp d hybrid orbitals. Hybridization S: sp3d2

Summary

Charge Clouds Hybridization If you add up the number of orbitals in the 2spname of the hybridization, it will equal the 3sp2 number of charge clouds. 4sp3 5sp3d 1 s + 3 p + 2 d = 6 charge clouds 6sp3d2

Double Bonds Both carbon atoms are considered to be central atoms here. C H 2 4 H H Each is central with another carbon and two C C hydrogens as terminal atoms. H H Each carbon has 3 charge clouds They each have three charge clouds around sp2 hybrid orbitals them.

Double Bonds (e.g. C H ) 2 4 There is a 2p orbital containing an electron 2 Atomic Orbitals that is left over after the three sp hybrid C: orbitals have been formed. 2s 2p Promotion Remember that 3 charge clouds result in sp2 C: hybridization. 2s 2p Hybridization In this case there is one un-morphed p-orbital C: left over on each carbon atom. sp2 2p A σ Bond results from the direct head-to- head overlap of orbitals. The bonds with Hydrogen atoms are also sigma bonds.

A π Bond results from the side to side attractive forces between the unmorphed p- orbitals. Double Bond e.g. C2H4 As each carbon has one unmorphed p-orbital, the attraction between the two p-orbitals is p H H considered to be one bond. sp2 sp2 sp2 sp2 sp2 sp2 All double bonds result from one σ Bond and H σ Bond H p one π Bond.

π Bond All single bonds are σ Bonds.

Atoms can spin on a single bond, but Double bonds are ridged, meaning that the =CH2 one one side of the double bond cannot flip upside-down while the one on the other side remains as it is.

They must move together as a sold structure.

Triple Bonds

C2H2

H C C H

Each carbon has two charge clouds sp hybrid orbitals

Triple Bonds (e.g. C2H2) Atomic Orbitals C: 2s 2p Each carbon has 2 p-orbitals with unpaired Promotion electrons left over after the sp hybridization C: has occurred. 2s 2p Hybridization C: sp 2p

Triple Bond (e.g. C H ) The other half of one of the p-orbitals from 2 2 each carbon cannot be seen as they are behind the molecule. p σ Bond p There are two π Bonds here as there are sp sp sp sp p p bonds from the side by side interactions of two different p-orbitals.

Two π Bonds This molecule is also rigid.

Bond Polarity • Shared electrons spend more time around the most electronegative element in the chemical bond.

• This gives the more electronegative element a slightly negative charge, and the less electronegative element a slightly positive charge.

Hydrogen has an electronegativity value of 2.1 and Chlorine has an electronegativity value of 3.0. The elecronegativity difference is 0.9 making it a polar bond.

As chlorine is more electronegative, the shared electrons are pulled closer to its nucleus.

The excess electron density around Cl gives it a slightly negative charge.

The reduced electron density around H gives it a slightly positive charge.

Polar Bonds The fact that one side is slightly positive and e.g. HCl the other is slightly negative can be demonstrated in one of the following ways. H Cl The arrow above points toward the more δ+ δ− electronegative element, and the cross over H Cl the H means that hydrogen carries a partially positive charge.

The δ+ symbol means slightly positive and the δ− symbol means slightly negative.

As long as the bond between the elements remains in tack, the electron density will be shared.

Thus, the chlorine never fully gains an electron to obtain a full –1 charge and the hydrogen never completely loses its electron to obtain a +1 charge.

Polar Molecules

To know if a molecule is polar: • You must know if the bonds are polar. • You must know the overall shape.

Both ends of this linear molecule are slightly negative, while the center is slightly positive. e.g. CO2 It is non-polar, even though it has polar .. .. bonds. O C O .. .. To be a polar molecule the overall structure must have a positive and a negative end. Non-Polar Bond Dipoles Cancel You need to draw the Lewis structure and figure out it shape in order to determine if the overall structure is polar. The oxygen acquires a slightly negative charge, as it is more electronegative.

e.g. H2O The lone pairs that sit close to oxygen’s nucleus make the oxygen end of the structure 2δ− .. even more negative.

:O H δ+ Polar The two hydrogens become slightly positive.

Bond dipoles H do not cancel This molecule is considered to by polar δ+ because it is bent.

The Oxygen end is negative and the hydrogen ends are positive.

e.g. BF3

F Non-Polar This structure is non-polar as all the ends are Bond Dipoles negative and the center is positive. Cancel B

F F

e.g. CCl4

Cl Non-Polar This structure is non-polar as all the ends are Bond Dipoles negative and the center is positive. Cancel C Cl Cl Cl

This molecule is the same shape as the last one; however, here we have replaced three of the chlorines with hydrogens.

Hydrogen is less electronegative than carbon, and thus, the electrons will spend more of their time around the Carbon – making the e.g. CH Cl hydrogens a little bit positive. 3 The Chlorine is slightly negative as it is more Cl electronegative than carbon.

Polar Thus, the top has an overall negative charge Bond C Dipoles H H and the bottom has an overall positive charge Do Not H Cancel (The electronegativity difference between C and H is only 0.4 – normally considered to be non-polar. However, when there is any electronegativity difference at all there is some polarity. Even if we consider these bonds to be non-polar the overall molecule is still polar, as chlorine is slightly negative and carbon is slightly positive.)

e.g. NH3

The top is slightly negative and the three Polar 3δ− hydrogen atoms are all slightly positive. Bond N Dipoles The overall structure is polar. Do Not H H Cancel δ+ H δ+ δ+