covalent bonds and molecules --you don’t mind a trick question, do you?

Covalent Bonds and Molecules

• --you don’t mind a trick question, do you?

• Show the formation of a covalent bond between a carbon and oxygen atom by sharing electrons

OC


• Carbon monoxide—colorless, odorless gas. Toxic. Careful, you can’t afford any more brain damage.

OC


OC


??!?!!

OC


The oxygen has a full octet, while the carbon is not yet full. Does this mean carbon

monoxide can’t form?

OC


No. This is a demonstration of why we chemists don’t circle electrons anymore.

OC


The correct bonding diagram.

OC


We notice that the atoms have (4 + 6=) 10 valence electrons.

OC


Two octets require 16 electrons.

OC


The 6 electrons short represent three bonds.

OC


Each atom needs an unshared pair of electrons for an octet

OC

• This is called a Lewis diagram (after Gilbert Norton Lewis)

• The bars represent covalent bond

( 2 shared electrons)

• Unshared pairs fill out the octets.

• Double and triple bars represent double and triple bonds.

Lewis diagramsStep 1: Count the total valence electrons available

--use the columns of the periodic chart--negative ions have extra electrons, --positive ions are missing electrons

Step 2: Count the total valence electrons needed--duet rule for hydrogen, or the --octet rule for everything else

Step 3: Number of bonds = (electrons needed-electrons available) / 2 electrons per bond

Lewis diagrams

Step 4: Choose the central atom (almost always the odd one), surround it with the others.

Step 5: Connect with one bond to each outer atom. (PS Recheck your formula!)

Step 6: Fill in enough multiple bonds to satisfy step 3

Step 7: Draw in unshared pairs to fill valence levels.

Don’t…

…try to figure out whose electrons are whose. Electrons are identical.

…string the atoms along. Put one atom in the center, unless you have 6 or more atoms.

…EVER put two bonds or an unshared pair on H.

Lewis diagrams

• Draw a Lewis diagram of hydrogen cyanide, HCN

Lewis diagrams


• Step 1: Total valence electrons available =1 (from H) + 4 (from C) + 5 (from N)= 10

Lewis diagrams


• Step 2: Total valence electrons needed = 2 (for H) + 8 (for C) + 8 (for N) = 18

Lewis diagrams


• Step 3: Number of bonds = (18-10) / 2 =8/2= 4 bonds

Lewis diagrams


H C N• Step 4: The central atom is the carbon,

since it’s written that way; HCN

Lewis diagrams


H C N• Step 5:

Lewis diagrams


H C N• Step 6: Since hydrogen can't make more

than one bond, the 3rd and 4th bonds have to be between the C and N

Lewis diagrams


H C N• Step 7

Unshared pair!

Try an ion.

• Draw a Lewis diagram for the nitrite, NO2- ,

ion

Anions have extra electrons!

Lewis diagrams

• Draw a Lewis diagram for the nitrite, NO2- ,

ion

[ O N O ]-

Use [brackets] around an ion

Resonance

• Which one is preferable?

[ O N O ]-

[ O N O ]-

Resonance

• Each is valid. The multiple bond exists in both locations. This is called resonance.

[ O N O ]-

[ O N O ]-

(the double-headed arrow signifies resonance)

Resonance

• Draw three resonance structures for carbon dioxide.

Coordinate covalent bonds

• How did this

become this?

OC

OC


• How did this

become this?

Carbon monoxide really does have the third bond. The oxygen donates both electrons to share. This is a coordinate covalent bond

OC

OC


• Draw a Lewis diagram of the ozone (O3) molecule.

Exceptions to the octet rule

• Draw a Lewis diagram for the triiodide ion, I3

-



-

• …Hey, Wait a minute!



-

• When you try to find the number of bonds, (24-22)/2=1 bond. That’s not enough to tie the ion together.

• When that happens—go old school. Circle your electrons



[ I I I ]-



[ I I I ]-

Two single bonds will satisfy the outer two iodine atoms, the inner one breaks the octet rule (with 10 electrons).


• Draw a Lewis diagram for XeF4



• (Don’t start with me. I know.)



XeF

F

F

F

…now that you’re big kids…

• We know electronegativity!

• The type of bond is better described using electronegativity differences instead of a simple metal/nonmetal distinction

• Subtract the two electronegativities (table on page 405)

• We usually claim that an electronegativity difference of

0-.4 is a nonpolar covalent bond

.5-1.8 is a polar covalent bond

1.9 and above is an ionic bond

• What kind of bond forms between chlorine and phosphorus atoms?


Cl P

3.0 2.1

Look up the electronegativities


Cl P

3.0 - 2.1=.9

• Subtract


Cl P

3.0 - 2.1=.9

• This is a polar covalent bond


Cl P

3.0 - 2.1=.9

• This is a polar covalent bond

• (use an absolute value for the difference)

Polar bonds

• We use a symbol to show a polar covalent bond.

• The arrow points toward the more electronegative atom, the (+) end is less electronegative

H

O

HH

Polar bonds

• Or, mark the molecules (+) and (-) parts

• The is the small Greek delta, indicates a small change. In this case, a partial charge

H

O

HH

Polar bonds

• The less electronegative end of a polar

bond: + -

H Cl• --is more positive

• --cannot attract the electrons as well

• --is farther from the shared pair of electrons

What kind of bond forms between these atoms?

H and H H and C

C and F F and F

Na and O Mg and N

Mg and Mg N and O

Molecular Shapes

• Most molecules have a central atom that satisfies the octet rule. This allows the following shapes.– Tetrahedral– Trigonal pyramid (trigonal=having three– Bent corners)– Linear and– Trigonal planar

Molecular Shapes

• Four bonds in four directions makes a tetrahedral shape

Molecular Shapes

• Three bonds and one lone pair in four directions makes a trigonal pyramid shape

Molecular Shapes

• Two bonds and two lone pairs in four directions makes a bent shape

Molecular Shapes

• A double bond holds two electron pairs in the same direction. With no lone pairs, this makes a trigonal planar molecule

Molecular Shapes

• One lone pair, with a single and a double bond gives a bent molecule.

Molecular Shapes

• Two double bonds, or a single and a triple makes a linear molecule

Molecular Shapes

• Two atoms are always in a straight line, a linear molecule.

If A=central atom, B=atoms bonded to it, E=e- pairs:

• AB4 — tetrahedral

• AB3E-trigonal pyramid

• AB2E2 — bent

• ABE3 — linear

• AB3 — trigonal planar

• AB2E — bent

• ABE2 — linear

• AB2 — linear

• ABE — linear

Look for double bonds and unshared pairs

• AB4—tetrahedral• (no double bonds)• AB3E-trigonal pyramid• (no double bonds)• AB2E2 –bent• (no double bonds)• ABE3 –linear• (no double bonds)

• AB3—trigonal planar• (one double bond)• AB2E –bent• (one double bond)• ABE2—linear• (one double bond)• AB2—linear• (2 doubles or 1 triple)• ABE—linear (“)

• Determine the shape of each molecule and ion on the lab that has a single central atom.

Polarity of molecules

• When polar bonds are not cancelled by symmetry, you get a polar molecule. A polar molecule has (+) and (-) parts.

• POLARITY is the first property to look for when analyzing a molecule !

Polarity

• CH4 has no polar bonds. It is symmetric

• PH3 has no polar bonds It is not symmetric

• CO2 has polar bonds. It is symmetric

• H2O has polar bonds. It is not symmetric•

Polarity

• CH4 has no polar bonds. It is symmetric• Not polar!

• PH3 has no polar bonds It is not symmetric• Not polar!

• CO2 has polar bonds. It is symmetric• Not polar!

• H2O has polar bonds. It is not symmetric• Polar!

• Mark each molecule on the lab that is polar.

• For those that are not polar—why not?

• (PS—don’t even look at the ions. If it has a whole charge, ignore the partial charges)

Hybridization

• Consider CH4:

C 1s22s22p2

H 1s1 H 1s1 H 1s1 H 1s1

Who shares where?

Hybridization

C 1s22s22p2

H 1s1 H 1s1 H 1s1 H 1s1

If the first two hydrogen pair electrons with the 2p electrons of C and the last two enter an empty orbital, you would see…

Hybridization

• Consider CH4

C

H H H H

If the first two hydrogens pair electrons with the 2p electrons of C and the last two enter an empty orbital, you would see…

Hydridization

• Consider CH4

C

H H H H

If the first two hydrogens pair electrons with the 2p electrons of C and the last two enter an empty orbital, you would see…

• …a difference in the bonds.

You don’t.

• All four bonds are identical.

What really happens is…

Hybridization

• The first step is a hybridization of the valence level

C

H H H H

forms…

Hybridization


C

H H H H

The s and p orbitals hybridize to form sp3 orbitals. The sp3 designation shows one s orbital and 3 p orbitals make the new ones

Hybridization


C

H H H H

The number of orbitals is preserved

(4 in 4 out)

Hybridization

C

H H H H

All four bonds are identical. Methane is a symmetrical molecule.

sp2 Hybridization

• When one p orbital is left out of the hybridization, it is used to make a double bond

• C

• H H O

…forms….

sp2 Hybridization


• C

• H H O

sp2 orbitals

Unused p orbitals

sp2 Hybridization


• C

• H H O

Makes the double bond!

sp2 Hybridization

H

C O::

H

Carbon shares electrons in sp2 orbitals

Carbon & oxygen share electrons in unused p orbitals

sp Hybridization

• When two p orbitals are left out of the hybridization, it is used to make two double bonds, or a triple bond

• C

• O O

…forms….

sp Hybridization


• C

• O O

sp orbitals

sp2 orbitals

Unused p orbitals

sp Hybridization


• C

• O O

sp Hybridization

::O C O::

Carbon shares electrons in sp orbitals

Carbon & oxygen share electrons in unused p orbitals

What is the hybridization of the carbon atoms in…

• CCl4

• H2CO

• C2H6

• C2H4

• C2H2

• CO

• CH3OH

• HCOOH

Molecular orbitals

• Overlapping s orbitals, or hybridized orbitals makes a (sigma) bond

• The electron density is on the SAME line as the nuclei

s s

Molecular orbitals

• Overlapping p orbitals, makes a (pi) bond

• The electron density is on a PARALLEL line to the line of the nuclei

p p + +

Molecular orbitals

• A single bond is a bond

• A double bond is a bond, and a bond above and below the

• A triple bond is a bond, with two bonds– above/below and front/back

For every bonding molecular orbital ( or ) an antibonding orbital is formed (* or *)

A bond is formed when there are more bonding than antibonding electrons

• Count the and bonds in the following molecule. Label each bond as or

C C CC

O

C

H

H

H H

H

H

• Determine the hybridization of the carbons and the oxygen atom

C C CC

O

C

H

H

H H

H

H

The molecular aufbau order

1s2*1s

22s2*2s

22px22py,z

42py,z42px

2….

• For example:

• O2 has 16 electrons. Its electron configuration is:

O2 1s2*1s

22s2*2s

22px22py,z

42py,z2

The molecular aufbau order

What is the electron configuration of…

N2

NO

Ne2

Remember: we couldn’t do a Lewis diagram!

The molecular aufbau diagram

VSEPR

• Valence Shell Electron Pair Repulsion

Theory

VSEPR

• Valence Shell Electron Pair Repulsion

Theory

--pronounced “Vesper”

Electron pairs repel each other. Just as it says.

• VSEPR is used to predict bond angles. The pairs will space themselves out as far as possible.

• A lone pair will take as much room as a bond AND MORE!

• Consider sp3 hybridization

• AB4—like methane. Tetrahedral 109.5o

• AB3E—like ammonia. Pyramidal 107o

• AB2E2—like water. Bent 104.5o angles

• --the unshared pairs force the bonds closer together

With sp2 hybridization:

• AB3—like carbonate. Trigonal planar: 120o

• AB2E—like nitrite. Bent: less than 120o

• ABE2—like O2(2 atoms, has to be linear)

With sp hybridization:

• AB2—like carbon dioxide. Linear: 180o

• ABE—like carbon monoxide. Linear: 180o

• --but that’s just if you always follow the rules– like the octet rule.

With dsp3 hybridization:

• AB5—trigonal bipyramid

• AB4E—seesaw

• AB3E2—t-shaped

• AB2E3—linear

• ABE4—linear

With d2sp3 hybridization:

• AB6— octahedral

• AB5E—square pyramid

• AB4E2—square planar

• AB3E3—t-shaped

• AB2E4—linear

• ABE5—linear

What is the shape of…

• All of the molecules and ions on the lab?

• I3-, SF6, XeF4, PCl5, IF5

-?

Bond Energies

The energy it takes to break a bond is the amount of energy released as the bond is formed.

• --measured in kJ/mol

• --can be used to estimate Hrxn

• --can be absorbed or emitted as light.

What is the Hf of NH3?


• Write the reaction

N2 + 3H2 2NH3


• Count the bonds made and broken

N2 + 3H2 2NH3

1 NN triple bond, 3 HH single bonds broken

6 NH single bonds made


• Look up bond energies, and find a total

N2 + 3H2 2NH3

1 molx941kJ/mol+3 molx436kJ/mol= 2249kJ used

6 molx393 kJ/mol=2358 kJ released


• Find the difference, express as kJ/mol

N2 + 3H2 2NH3

2358 kJ-2249kJ= 109 kJ more is released, as 2 mol NH3 is produced,Hf=109kJ/2mol=-55kJ/mol

• It’s an estimate.

• My book claims -46 kJ/mole.

What is the heat of reaction for

2H2O2 2H2O +O2

covalent bonds and molecules --you don’t mind a trick question, do you?

Documents

oxygen atom

covalent bonds

central atom

electronseach atom

outer atom

extra electrons

electrons short

electronsthe oxygen