chapter 5 chem

Linus Pauling and others (1930s) developed the

valance bond theory

of covalent bonding.

Electrons reside in

orbitals

.Covalent bonds form when electrons are

shared by the overlapping

of singly occupied orbitals.

Electrons in these overlapping orbitals must have

opposite spins.

to form a covalent bond, an atom must have

an unpaired electron.

Number of bonds formed by an atom is determined by the?

number of

unpaired electrons

This theory works well for explaining the

bonding in diatomic

molecules with only single bonds

VSPER theory explains

This theory also works well for explaining the lack of bonding

experienced by the noble gases, but it does not work well for other

elements or molecules.

valence bond theory must be modified to explain ?

covalent bonds

formed in other molecules

The basic principle of Valence Bond Theory is that?

a covalent bond is

formed when the atomic orbitals of two atoms overlap, and the

overlapping orbitals each contain one unpaired electron.

More overlap=

stronger bond

This sharing of electrons lowers the energy of the system, thus

stabilizing the

atoms when bonded

Lower energy=

happy molecules can exist

Overlapping or orbitals form covalent bonds. The shared space has two electrons with opposite spins, thus demonstrating what princple?

pauli exclusion

The overlap of two p orbitals is greatest when the orbitals are directed end to end. (b) Any

other arrangement results in

less overlap. The dots indicate the locations of the nucle

how many pi and sigmas

one sigma bond no pi bonds

how many sigma and pi bonds?

one sigma bond one pi bond

how many bonds and sigmas

one sigma bond two pi bonds

Single bonds are formed by the direct overlap of two hybrid orbitals,

p orbitals, or s orbitals.

• These are called

sigma (sigma) bonds.

The additional electrons shared in a multi bond are not a result of

directly overlapping hybrid orbitals, but are a result of

ide-by-side

overlap of two regular “p” atomic orbitals.

• These are called pi () bonds

Identify the number

of σ and π bonds contained in this molecule.

There are six σ C–H bonds and one σ C–C bond, for a total of seven

from the single bonds. There are two double bonds that each have a

π bond in addition to the σ bond. This gives a total nine σ and two π

bonds overall.

sigma bonds have head on overlap. This could include

s-orbitals overlapping s-orbitals

s-orbitals overlapping p-orbitals

p-orbitals overlapping p-orbitals if aligned correctly

sp, sp2, or sp3 hybrid orbitals overlapping each other or an s or p-orbital

pi-bonds have side by side overlap and is represented by the pi symbol. They form between two unhybridized parallel p orbitals

“side by side” pi overlap

look

dentify each illustration as depicting a σ or π bond:

(a) is a π bond with a node along the axis connecting the nuclei while (b) and (c) are σ

bonds that overlap along the axis.

hybridization

Mixing of the native atomic orbitals to form special orbitals for bonding.

Bonding in many molecules is best explained with

hybrid orbitals.

Hybrid orbitals only exist in covalent bonded atoms.

Keep in mind when working with hybrid orbitals.

• The number of hybrid orbitals formed always equals the?

number of atomic

orbitals that are combined.

H2O, The hypothetical overlap of two of the 2p orbitals on an oxygen atom (red) with the 1s orbitals of two?

hydrogen atoms (blue) would produce a bond angle of 90°. This is not consistent with experimental evidence. In reality it is 140°. (Note that orbitals may sometimes be drawn in an elongated “balloon” shape rather than in a more realistic “plump” shape in order to make the geometry easier to visualize.

A water molecule has four regions of electron density, so VSEPR theory predicts a?

tetrahedral arrangement of hybrid orbitals. Two of the hybrid orbitals on oxygen contain lone pairs, and the other two overlap with the 1s orbitals of hydrogen atoms to form the O–H bonds in H2O. This description is more consistent with the experimental structure.

two atomic orbitals can hybridize to produce two hybrid orbitals, what does that look like?

one s orbital + one p orbital  two sp hybrid orbitals

All orbitals that are in a set of hybrid orbitals will have the same energy

and shape. VESPR Theory will

predict the type of orbital formed based on its

electron-pair geometry.

hybridization of an s orbital (blue) and a p orbital (red) of the same atom produces two

sp hybrid orbitals (yellow). Each hybrid orbital is oriented primarily in just one direction. Note that each sp orbital contains one lobe that is significantly larger than the other. The set of two sp orbitals are oriented at 180°, which is consistent with the geometry for two domains.

sp Hybrid Orbitals

Occurs when a central atom is surrounded by two regions of electron

density.

• Has 2 areas of electron density

This orbital energy-level diagram shows the sp hybridized orbitals on Be in the linear BeCl2 molecule. Each of the two sp hybrid orbitals holds one electron and is thus?

half filled and

available for bonding via overlap with a Cl 3p orbital.

Three atomic orbitals can hybridize to produce three hybrid orbitals.

one s orbital + two p orbitals  three sp2 hybrid orbitals

Occurs when a central atom is surrounded by three regions of electron density.

The hybridization of an s orbital (blue) and two p orbitals (red) produces three equivalent sp2 hybridized orbitals (yellow) oriented at 120° with

respect to each other. The remaining unhybridized p orbital is not shown here, but is located along the z axis.

sp2 Hybridization

Combination of one s and two p orbitals.

• Gives a trigonal planar arrangement of atomic

orbitals.

• One p orbital is not used.

–Oriented perpendicular to the plane of the sp2

orbitals.

–Has 3 areas of electron density

This alternate way of drawing the trigonal planar sp2 hybrid orbitals is sometimes used in

more crowded figures.

look

H3 is an electron-deficient molecule with a trigonal planar structure. draw it and mark the angle

look

In an isolated B atom, there are one 2s and three 2p valence orbitals. When boron is in a molecule with three regions of electron density, three of the orbitals hybridize and create a

set of three sp2 orbitals and one unhybridized 2p orbital. The three half-filled hybrid orbitals each overlap with an orbital from a hydrogen atom to form three σ bonds in BH3

look

These all show sp2 hybridization because they have 3 regions of electron density. The central atom(s) in each of the structures shown contain three regions of electron density and are sp2 hybridized. As we know from the discussion of VSEPR theory, a region of electron density contains all of the electrons that point in one direction. A lone pair, an unpaired electron, a single bond, or a multiple bond would each count as one region of electron density.

Four atomic orbitals can hybridize to produce four hybrid orbitals.

one s orbital + three p orbitals  four sp3 hybrid orbitals

• Occurs when a central atom is surrounded by four regions of electron

  density.

the hybridization of an s orbital (blue) and three p orbitals (red) produces

four equivalent sp3 hybridized orbitals (yellow) oriented at 109.5° with respect to each other.

in the ethane molecule, C2H6, each carbon has?

four sp3 orbitals. (b) These four orbitals overlap to form seven σ bonds.

Five atomic orbitals can hybridize to produce five hybrid orbitals

one s orbital + three p orbitals + one d orbital  five sp3d hybrid

orbitals

Occurs when a central atom is surrounded by five regions of electron

density.

sp3d Hybridization

Combination of one d, one s, and three p

orbitals.

•Gives a trigonal bipyramidal arrangement of

five equivalent hybrid orbitals.

•Has 4 areas of electron density

Draw the Lewis structure for PCl5

What is the shape of a phosphorus pentachloride

molecule?

What are the bond angles?

trigonal bipyramidal

90° and 120

These have 5 areas of electron density so they are sp3d. The three compounds pictured exhibit

sp3d hybridization in the central atom and a trigonal bipyramidal form. SF4 and ClF4+ have one lone pair of electrons on the central atom, and ClF3 has two lone pairs giving it the T-shape shown.

The five regions of electron density around phosphorus in PCl5 require?

five hybrid sp3d orbitals. (b) These orbitals combine to form a trigonal bipyramidal structure with each large lobe of the hybrid orbital pointing at a vertex. As before, there are also small lobes pointing in the opposite direction for each orbital (not shown for clarity).

sp3d2 Hybrid Orbitals

Six atomic orbitals can hybridize to produce six hybrid orbitals.

one s orbital + three p orbitals + two d orbitals  six sp3d2 hybrid

orbitals

• Occurs when a central atom is surrounded by six regions of electron

density.

sp3d2 Hybridization

Combination of two d, one s, and three p

orbitals.

•Gives an octahedral arrangement of six

equivalent hybrid orbitals.

•Has 6 areas of electron density

Draw the Lewis structure for XeF4.’

What is the shape of a xenon tetrafluoride molecule?

What are the bond angles?

octahedral

90° and 180°

look

Sulfur hexafluoride, SF6, has an octahedral structure that requires sp3d2 hybridization. (b) The six sp3d2 orbitals form an octahedral structure around sulfur. Again, the minor lobe of each orbital is not shown for clarity.

summary of electron geometry

VSEPR theory predicts the shapes of molecules, and hybrid orbital theory provides an explanation for how

those shapes are formed.

To find hybridization of the central atom:

1. Determine the Lewis structure of the molecule.

2. Determine the number of regions of electron density around an atom using VSEPR theory, in which

single bonds, multiple bonds, radicals, and lone pairs each count as one region.

3. Assign the set of hybridized orbitals from the table that corresponds to this geometry.