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Chemistry II-AP
Chemical Bonding Notes
Bonding of atoms - types
1) ionic (electrovalent)
2) covalent (including coordinate, or dative, bonding)
3) metallic
[Other types of intermolecular bonding will be discussed later.]
The above types of bonding are intramolecular.
1) ionic - transfer of electrons and subsequent union of an anion with a cation - formation of ionic bond. Ions
are held together by electrostatic forces - some of the strongest bonds that exist.
2) covalent - sharing of electrons to form molecules.
Electrons involved generally are those in the outer energy levels.
1
ex. for NaCl: 3s - loss of 1 electron for Na
5
3p - gain of 1 electron for Cl
Ions try to achieve a noble gas configuration.
-often called the octet structure - try to fill the s and p subshells.
Ionic bonding: occurs when atoms differ greatly in electronegativity; generally with a difference of 1.7 or
greater.
Typically, ionic bonding occurs between a metal of low ionization energy (Group IA, IIA, or transition metal)
and a highly electronegative nonmetal (Group VIA or VIIA).
Ionic compounds: characterized as solids with high M.P.; are good conductors of electricity; usually exhibit
cleavage.
Formation of pure ionic compounds from elements is always exothermic.
Bond strength:
- also known as the bond dissociation energy
- measured in terms of bond energy (amt. of energy required to make or break a bond)
- usually limited to gaseous compounds
- the more polar the molecule, the stronger the bond.
- for nonpolar molecules, the calculated value is close to the observed value, but for polar molecules, the
observed energy value is greater than the calculated value. (The difference is greater, too, as the polarity
increases.)
ex. for H-F Observed: 135 kcal/mole
104 + 37
Calc.: 1/2 D(H2) + 1/2 D(F2)
= ----------- = 70. kcal/mole
2
(found by splitting the bond energy of H2 and F2.
-2**
Linus Pauling made calculations to compensate for the extra strength due to polarity.
This was the reason for proposing the E.N scale!
2
Y = 23( E.N.)
*** BOND ENERGIES MAY BE USED TO CALCULATE THE H for a gaseous reaction if the Hf of the
substances are not known. THESE VALUES ARE ONLY APPROXIMATIONS!!
...............................................................
Bond distance:
1. For a nonpolar molecule, the distance is simply the sum of the two atomic radii.
2. For a polar molecule, the distance is slightly less, most probably due to the added strength of the ionic
character that shortens the bond.
o
o
ex. F-F = 2(0.64 A) = 1.28 A
o
o
o
but for H-F Calc. = 0.37 A + 0.64 A = 1.01 A
o
yet Observed = 0.92 A
3. With multiple bonds, the bond distance gets shorter for each additional bond
.................................................
Multiple Bonds:
In general, the bond distance and bond energy for an element is constant regardless of the molecule that is
formed. If the observed distance is much less, then a multiple bond probably exists.
examples:
formula
ethane
(C2H6)
propane
(C3H8)
butane
(C4H10)
ethene (ethylene) (C2H4)
ethyne (acetylene) (C2H2)
bond distance (in Angstroms)
1.54
1.54
1.54
1.33
1.20
for ethylene - a double bond exists
bond energy (in kcal/mole)
83
83
83
143
196
for acetylene - a triple bond exists
Multiple bonds show extra strength since the atoms are pulled closer together.
**(Draw structures of ethylene and acetylene)
Comparison of bond strengths for nitrogen:
nitrogen gas - N2 vs. hydrazine - N2H4
225 kcal/mole
38 kcal/mole
Hydrazine is used as a rocket fuel because upon combustion, it forms nitrogen gas which releases large
amounts of energy as it is formed (since it is such a stable molecule).
-3Drawing Lewis structures (Lewis Dot Structures or Lewis Diagrams)
- used to show bonding configuration for compounds
- shows the number of valence electrons and their arrangement
general order for an element,
3 6
4
7
E
1
2
"s" subshell electrons
5 8
Locations # 1 and 2 represent the "s" subshell electrons; the rest of the locations represent the six "p" subshell
electrons.
This single element configuration is not very useful since an element will rearrange its electrons to achieve an
octet configuration.
+1
for an ionic compound: ex NaCl
Na
for a molecular substance: ex: HCl
H
-1
Cl
Cl
Guidelines for drawing Lewis structures:
1.
Add up the total number of valence electrons for all elements. These must all be used.
If ions are involved, be sure to add extra electrons for anions, and reduce the number of valence electrons
for cations.
2.
Draw the symbols for the elements and connect them by single bonds. One line represents a single bond
pair. Compounds are often written in the order shown in the formula. If an element has a multiple
valence, it is usually the central atom. Some elements - such as hydrogen - only can form single bonds, so
they cannot serve as a central atom.
3.
If all the valence electrons are used and the octet configuration is not achieved, then multiple bonds must
exist. For every two electrons that are needed that do not exist, a multiple bond is needed. If four
electrons are needed, either two double bonds may be needed, or a triple bond may be needed.
4.
If the octet rule is achieved and extra valence electrons exist, the central atom violates the octet rule. The
extra electrons are put around the central atom.
-4Special cases when the octet rule is not achieved:
1. He configuration - 1st five elements
- often show tendency to form dative bonds.
2. Odd number of valence electrons (often form free radicals)
First two cases are rare; third case is more common.
3. Central atom holds more than 8 electrons.
Not found in second series elements because they don't have a "d" subshell capability.
Found when large central atom holds small, highly electronegative anion.
-reason that larger noble gases can form compounds.
ex. 1) BF3
2) NO2
3) SF6
RESONANCE:
Limitation of Lewis structure - some molecules cannot be represented by just one drawing because electron
pair(s) are delocalized.
- Need to draw all structures with double arrow between them.
ex. SO2
NO3-1 show benzene (aromatics) vs. aliphatics
Class problem: draw the Lewis structure for ozone
** WHAT WOULD YOU PREDICT ABOUT BOND STRENGTH AND BOND LENGTH??
..............................................................
FORMAL CHARGE:
- used to help determine the correct Lewis structure
Formal charge = # of valence electrons - (# of lone electrons) - # of bonded pairs
The best Lewis structure is the one that has the fewest number of elements with a formal charge other than
zero.
(Often compounds are written from the most electropositive element first with the most electronegative being
last.)
ex. carbon dioxide
for O (#1):
for O (#2):
for C:
<C=O=O>
1 2
or
6 - 0 - 4 = +2
6-4-2= 0
4 - 4 - 2 = -2
net:
<O=C=O> (remember 4 lone pair electrons)
1
2
6-4-2=0
6-4-2=0
4-4 -0=0
0
0 (preferred because of 3 zeros)
-
Class problem: do cyanate ion (CNO )
.....................................................................
[N=C=O is best structure]
Be sure to review sections on oxidation numbers; and redox reactions.
Review section on oxides - acidic, basic, and amphoteric.
Know the importance of ionic size for solubility and biological activity.
-5HYBRIDIZATION:
Rearrangement of electrons in available subshell orbitals to allow for maximum bonding.
atomic orbitals:
1. "s" subshell - spherical shape - result in direct overlap
"p" subshell - figure-8 configuration
- may result in direct overlap if orbitals are lying along the same axis; but may also produce
delocalized bonding.
direct overlap - sigma () bonding
delocalized bonding - pi () bonding
When a single bond is formed, a sigma bond is always involved.
For multiple bonds, the first bond is a sigma and the others are pi bonds.
Specifically:
single bond
double "
triple "
1 sigma 0 pi
1 sigma 1 pi
1 sigma 2 pi
Class problem: How many sigma and how many pi bonds exist in a molecule of benzene?
Hybridization: involves electrons that directly overlap
ex. show carbon as a lone element; then as a central atom
2 2
3
normal configuration s p
hybrid sp
Instead of 2 bond sites, there are now four bond sites and octet configuration can be achieved.
Class problem: do hybridization of the following:
of B in BF3
of S in SF6
of P in PCl5
of Be in BeH2
Types of Intermolecular forces (IMF) - discuss each
1. ionic
2. covalent network
3. metallic
4. dipole-dipole
5. hydrogen
6. ion-dipole
7. ion - induced dipole
8. dipole - induced dipole
9. van der Waals (or London dispersion froces)
-6Molecular Geometry
VSEPR theory (Valence Shell Electron Pair Repulsion theory)
- describes the 3-d (geometric) shape of the molecule
- based on the idea that electrons tend to repel each other as far as possible
- doesn't matter whether they are bonding or nonbonding electron pairs.
Classification of shapes:
# of e pairs
# of e pairs # of e pairs
shape
(total)
(bonding)
(nonbonding)
...........................................................................
bond angle
examples
1
1
0
linear
............................................................................
180
H-Cl
2
2
0
linear
.............................................................................
180
CO2
120
BF3
< 120
SO2
3
3
0
trigonal
planar
3
2
1
4
4
0
tetrahedral
109.5
CH4
4
3
1
trigonal
pyramid
< 109.5
NH3
bent or
V-shaped
.............................................................................
(107 for NH3)
4
2
2
bent or
V-shaped
< 109.5
H2S
(104.5 for H2O – important!)
4
1
3
linear
180
H-Cl
.............................................................................
5
5
0
trigonal
bipyramid
90 (axial)
120 (equatorial)
PCl5
5
4
1
seesaw or
teetertotter
90 (a)
< 120 (eq)
SF4
5
3
2
T-shaped
90
ClF3
5
2
3
linear
180
XeF2
-7Classification of shapes (cont.):
-
# of e pairs
(total)
-
# of e pairs
(bonding)
-
# of e pairs
(nonbonding)
shape
bond angle
examples
6
6
0
octahedral or
square bipyramid
90
SF6
6
5
1
square-based
pyramid
90
BrF5
6
4
2
square planar
90
XeF4
?????????????????????????????????????????????????????????????????????????????
?6
?3
?3
? T-shaped
90
???
?6
?2
?4
? linear
180
???
Note #1: if multiple bonds exist, they are treated as single bonds concerning the shape.
Note #2: if lone pairs exist, these electrons are found to be the equatorial pairs, not the axial pairs. This is due
to the fact that repulsion is minimized with this arrangement. Consider a species with 1 lone pair: if
the lone pair were in an axial position, it would encounter electron pair repulsion with three pair of
o
electrons in a 90 angle; if the lone pair is an equatorial pair, it only encounters two electron pair for
repulsion (those closest to it). The other two pair on the equator are less significant that those axial
pair closest to it.
Note #3: The last two situations with six sets of electron pairs are doubtful because they would be violating the
octet rule for no reason - there is not special bonding that is taking place. The first three situations
occur because the majority of the electron pairs around the central atom are bonding electron pairs.
...........................................................................
Class problem: What about the shape of ions and complexes?
ex. ions: nitrate, triiodide, sulfate, sulfite, phosphate, carbonate, chlorate, perchlorate, hydronium, ammonium,
amide, hypochlorite
ex. complexes: hexachlorostannate (IV), tetrachloroiodo (III), tetrachloroplatinate(II),
(I), hexafluorosilicate (IV), hexaaquazinc(II)cation
.........................................................................
dichloroargentate
Polarity of bonds vs. polarity of molecules
1. You can have polar bonds within a molecule and still not have a polar molecule. Can you have a polar
molecule if polar bonds do not exist?
2. Polarity of the molecule deals with the symmetry of the molecule.
3. Polarity of the bond depends on the electronegativity difference.
-8Examples:
molecule
bond type
molecular shape
methane
CCl4
NH3
nonpolar (0.3)
polar (0.6)
polar (0.8)
tetrahedral
tetrahedral
trig. pyramid
polar or nonpolar?
nonpolar
nonpolar
polar
(axis of symmetry for ammonia runs from the lone pair, through the N, and through the middle of the H)
H2O
polar (1.3)
bent
polar
(axis of symmetry runs from between the lone pairs, through the O, and through the middle of the H)
HCl
HI
SO3
SF6
polar (1.0)
nonpolar (0.4)
polar (1.0)
polar (1.5)
linear
linear
trig. planar
octahedral
polar
nonpolar
nonpolar
nonpolar
Depending on the location of lone pair electrons, these electrons can produce highly polar molecules. The
strength of the polarity is directly related to the solubility of the molecule, especially in a highly polar solvent
like water.
Molecular Orbital (M.O.) Theory:
Atomic orbitals vs. Molecular orbitals
M.O. - describes the distribution of electrons throughout the molecule - not just around the individual atoms
- describes the energy of the different electrons
Energy - based on quantum mechanics
- based on the shape of the region in which the electrons move "orbital"
- electron may move around the whole molecule - "molecular orbital"
Molecular orbital - still holds a maximum of two electrons (w/ opposite spin)
Types of orbitals - sigma and pi
actually 3 types - sigma-s, sigma-p, and pi-p
Bonding vs. anti-bonding orbitals
Bonding orbitals: electrons are generally found between the nuclei - they are attracted by both nuclei which
results in a lower energy level (more stable) than when they were isolated.
-serve to stabilize the molecule.
Anti-bonding
*
*
*
sigma-s , sigma-p , and pi-p