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UNIT 6. Theories of Covalent Bonding and Intro to Organic Chemistry VSEPR Theory, Molecular Shapes, and Valence Bond Theory. Electron Domains . Electron domains are the regions in the molecules where it is most likely to find electrons.

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slide1

UNIT 6

Theories of Covalent Bonding and Intro to Organic Chemistry

VSEPR Theory, Molecular Shapes, and Valence Bond Theory

slide2

Electron Domains

  • Electron domains are the regions in the molecules where it is most likely to find electrons.
    • For a bond (single, double, or triple), the electron domain is between the two atoms in the bond and consists of all the electrons involved in the bond.
    • For nonbonding pairs of electrons, the domain is the nonbonding pair and is centered on a single atom.
slide3

Valence Electrons

  • Valence electrons are the electrons in the outermost unfilled shell of the atom or ion. These are usually s and p electrons, but can be d electrons.
  • For the main group elements, the number of valence electrons is the last digit of the group number.
  • Knowing the number of valence electrons allows us to draw Lewis symbols for the elements and Lewis structures for compounds.
slide4

Lewis Structures

  • Lewis structures are very helpful in studying bonding because they show only the valence electrons of the atoms or ions.

The Octet Rule

  • Atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons (the s and p subshells are full). We use this rule to draw Lewis structures for compounds.
slide5

Lewis Structures of Covalent CompoundsFollow these steps in order.

1. Decide which atoms are bonded.

2. Count all valence electrons.

3. Put two electrons in each bond.

4. Complete the octets of the atoms attached to the central atom except H, which takes a duet.

5. Put any remaining electrons on the central atom.

6. If the central atom has less than an octet, form double or triple bonds.

slide6

Putting Formal Charges on Lewis Structures

The formal charge of any atom in a compound or ion may be calculated using the following:

FC = # of valence electrons – number of bonds – number of nonbonding electrons

FC of O = 6-2-4 = 0

FC of O = 6-1-6 = -1

FC of O = 6-3-2 = +1

slide7

Resonance Structures

Three completely equivalent Lewis structures can be drawn for the nitrate ion, NO3-.

Reality is a blend of the three. There are no double bonds in the nitrate ion, but each bond is more stable than just a single bond. All three structures are resonance structures.

slide8

From Lewis Structure to Electron Domain Geometry via VSEPR

  • The Lewis structure shows the covalent bonds (solid lines) and the nonbonding electrons (dots) that are present in a compound.
  • This allows the identification of the electron domains of the molecule. Domains are the regions in the molecules where it is most likely to find electrons.
  • Electron domain geometry gives bond angle and hybridization.
slide9

VSEPR Theory

  • Valence Shell Electron Pair Repulsion – qualitative explanation of molecular shapes
  • Electrons in a domain are subject to electrostatic repulsion from the electrons in the other domains. The domains will orient themselves so as to minimize this repulsion.
  • The orientation of these domains is a function of the number of domains around the central atom and is one of several simple geometric figures.
slide10

Bond Angles when a Nonbonding Electron Pair or Multiple Bond is Present

  • Nonbonding electron pairs take up more space than bonding pairs and have the effect of squeezing (decreasing) the bond angles among the atoms.
  • Multiple bonds exert more repulsion than single bonds, and have the same effect on bond angle as the nonbonding electron pair.
slide11

Bond Angles for Atoms in Organic Molecules: Effect of Nonbonding Electron Pairs

Example: O in an ether

4 electron domains

tetrahedral geometry

Bond angles on O are <109.5° because the two e- pairs take up extra space.

Example: C in an alkane

4 bonding domains

tetrahedral geometry

Bond angles 109.5°

Example: N in an amine

4 electron domains

tetrahedral geometry

Bond angles on N are <109.5° because the e- pair takes up extra space.

slide13

Valence Bond Theory

Lewis Structures – Explain bonding as a sharing of electron pairs and geometry through VSEPR.

VB Theory - A more quantitative approach to explaining bonding. Here bonds are explained by the overlap of orbitals on the two atoms in the bond.

slide14

Orbital Energy Diagram

4p

4s

3d

3s

3p

ENERGY

2p

2s

1s

slide15

The periodic table may also be used to determine the electron configuration of the elements.

slide16

Valence Bond Theory

Bonds occur from the overlap of atomic orbitals.

Cl: 1s22s22p63s23p5

H: 1s1

The bond in H-Cl is formed by the overlap of the H 1s orbital with the Cl 3p orbital.

How would you describe the Cl-Cl bond?

The bond in H-H is formed by the overlap of the two H 1s orbitals.

Bonds formed by end-to-end overlap are called sigma (σ)bonds.

slide17

Valence Bond Theory – Hybrid Orbitals

Describing bonds as the overlap of s and p orbitals explains some geometries, but certainly not the tetrahedral geometries (e.g. H2O).

What orbitals are overlapping here?

slide18

Valence Bond Theory – Hybrid Orbitals

Since the orbitals (s, p, d, f, etc.) are all mathematical solutions to the Schrödinger wave equation, it is true that linear combinations of these orbitals are also solutions to the Schrödinger wave equation.

In other words, we may mix s, p, and d orbitals to make new, hybrid orbitals that are also valid.

slide19

sp Hybrid Orbitals

Consider HC≡CH. VSEPR says the molecule is linear.

How can hybrid orbitals explain the geometry?

2p

2s

C ground state (the orbitals of C are what determine the geometry of HC≡CH.)

Orbital energy

Two e- are available for bonding, but the geometry is wrong.

1s

slide20

sp Hybrid Orbitals

Consider HC≡CH. VSEPR says the molecule is linear.

How can hybrid orbitals explain the geometry?

Energy is used to promote one 2s e- to a 2p orbital.

2p

2s

Orbital energy

A 2s orbital and a 2p orbital mix to make two new orbitals. The other two 2p orbitals are unchanged.

1s

slide21

sp Hybrid Orbitals

Consider HC≡CH. VSEPR says the molecule is linear.

How can hybrid orbitals explain the geometry?

2p

sp

The 2s orbital and one 2p orbital mix to form two sp hybrid orbitals. The energy of mixing is more than paid back when the C-H and C≡C bonds are formed.

Orbital energy

1s

slide22

sp Hybrid Orbitals

C

C

C

C

C

Hybrid orbitals have a small lobe and a large one. The large lobe allows more overlap and, therefore, the formation stronger bonds. The energy needed to make the hybrid orbital is paid back, with interest, in the formation of stronger bonds.

slide23

sp2 Hybrid Orbitals

Consider H2C=CH2. VSEPR says the molecule is linear.

How can hybrid orbitals explain the geometry?

One 2p orbital remains unhybridized.

2p

sp2

Orbital energy

Mixing the three orbitals gives three sp2 hybrid orbitals with the same energy.

1s

What orbitals overlap to form each of the C-H bonds?

slide25

sp3 Hybrid Orbitals

Consider CH4. How can hybrid orbitals explain this tetrahedral geometry?

sp3

Orbital energy

Mixing the four orbitals gives four sp3 hybrid orbitals with the same energy.

1s

The VSEPR geometry tells you the hybridization: Tetrahedral VSEPR geometry  sp3 hybrid orbitals.

slide26

sp3 Hybrid Orbitals

What orbitals overlap to form each of the C-H bonds in methane?

slide27

sp3d Hybrid Orbitals

PCl5 has the shape of a trigonal bipyramid.

P ground state:

3s

3p

3d

Energy is used to promote 3s e- to 3d:

Hybridize:

sp3d

3d

Trigonal bipyramid VSEPR geometry  sp3d hybrid orbitals.

slide28

sp3d2 Hybrid Orbitals

SF6 has the shape of an octahedron.

S ground state:

3s

3p

3d

Energy is used to promote 3s and a 3p e- to 3d:

Hybridize:

sp3d2

3d

Octahedral VSEPR geometry  sp3d2 hybrid orbitals.

slide29

Use Lewis structures to get the electron domain geometry, and that will lead to the bond angles and hybridization.

# of e- e- domain bond angle hybridization

domains geometry

2 linear 180° sp

3 trigonal 120° sp2

planar

4 tetrahedral 109.5° sp3

5 trigonal 90°,120°,180° sp3d

bipyramidal

6 octahedral 90°,180° sp3d2

slide30

Hybridization and Bond Angles in Larger Molecules

Just identify the geometry/hybridization around each atom in succession.

sp2, trigonal planar geometry, bond angle is >120° (due to double bond)

H : O :

| ||

H — C — C — O — H acetic acid

| ¨

H

..

sp3, bent geometry, bond angle is <109.5°

What orbitals overlap to form each of the bonds in acetic acid?

sp3, tetrahedral geometry, bond angle is 109.5°

slide31

Hybridization and Bond Angles in Larger Molecules

H

|

H — C — C ≡ C — H propyne

|

H

sp, linear geometry, 180°

sp3, tetrahedral geometry, 109.5°

slide32

Valence Bond Descriptions

H

|

H — C— C ≡ C — H

|

H

What orbitals overlap to form each of the bonds in propyne?

  • The three C-H bonds shown at the left are each formed by the overlap of a H 1s orbital with one of the four C sp3 hybrid orbitals.
  • The C-C bond is formed by the overlap of the remaining sp3 orbital on the first C with one of the sp orbitals on the second C.
  • The C≡C triple bond is formed by: 1) the overlap of C sp orbitals (sigma bond), 2) the overlap of C 2pyorbitals (pi bond), and 3) the overlap of C 2pzorbitals (pi bond).
  • The final C-H bond is formed by the overlap of the C sp orbital with a H 1s orbital.