Honors Chemistry I
Chemical Bonding Notes
Since most elements do not have a filled outer shell, they attempt to achieve an octet arrangement by
combining with one or more other elements (or with other atoms of the same element). This drive to fill
the outer shell (the “s” and the “p” subshells) is called the “Octet Rule”.
Bonding of atoms - types
1) ionic (electrovalent) – atoms actually transfer electrons (gain or lose electrons)
2) covalent (including coordinate, or dative, bonding) – atoms “share” electrons
3) metallic – “core” of positively-charged nuclei surrounded by a “sea” of electrons
- noted for its loose association of electrons with nuclei
[Other types of intermolecular bonding (IMF) 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. Forms a
crystal lattice – no distinct piece of matter (no molecules) – only “unit cells” – the basic
repeating pattern
2) covalent - sharing of electrons to form molecules.
Coordinate covalent bonding occurs when one element contributes both electrons of the electron pair
when the chemical bond forms.
Ex. NH3 + H+  NH4+
Electrons involved in bonding generally are those in the outer energy levels.
ex. for NaCl: 3s1 - loss of 1 electron for Na
3p5 - 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
(EN) of 1.7 or greater. (Learn this number!)
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 in
molten form or aqueous solution; nonconductors as solids; usually exhibit
cleavage; have a definite crystalline pattern.
Metallic Bonds – may occur for a single element (such as a piece of copper wire) or as an alloy, like
brass (a homogeneous solution of copper and zinc)
-2Bond 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.
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 are 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
bond energy (in kcal/mole)
83
83
83
143
196
for ethylene - a double bond exists; 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).
Drawing 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, E
36
7
1
"s" subshell electrons
4
2
58
E
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.
for an ionic compound: ex NaCl Na+ Cl
for a molecular substance: ex: HCl H
Cl
-3Guidelines 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 - can only 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 (is lacking electrons),
then one or more 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.
Special cases (3) when the octet rule is not achieved:
1. He configuration - 1st five elements (realistically, only H, Be, and B)
- 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 period 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- show benzene (aromatics) vs. aliphatics
Class problem: draw the Lewis structure for ozone
** WHAT WOULD YOU PREDICT ABOUT BOND STRENGTH AND BOND LENGTH FOR A
SPECIES THAT EXHIBITS RESONANCE??
-4FORMAL CHARGE:
Used to help determine the correct Lewis structure when more than one structure is possible.
Formal charge = # of valence electrons - (# of lone electrons) - # of bonded pair
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
<C=O=O> or <O=C=O> (remember 4 lone pair electrons)
1 2
1
2
for O (#1):
for O (#2):
for C:
6 - 0 - 4 = +2
6-4-2= 0
4 - 4 - 2 = -2
net:
6-4-2=0
6-4-2=0
4-4 -0=0
0
Class problem: do cyanate ion (CNO-1-)
0 (preferred because of 3 zeros)
[ <N=C=O> is best structure]
.....................................................................
Be sure to review sections on oxidation numbers; and redox reactions.
HYBRIDIZATION:
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 (s) bonding, show s-s, s-p, and p-p overlaps
delocalized bonding - pi (p) 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
1 sigma
1 sigma
0 pi
1 pi
2 pi
Class problem: how many sigma and how many pi bonds exist in a molecule of benzene?
-5Hybridization: involves electrons that directly overlap
ex. show carbon as a lone element; then as a central atom
normal configuration s2p2
hybrid sp3
Instead of 2 bond sites, there are now four bond sites and octet configuration can be achieved.
Discuss the type of hybridization of diamond vs. the hybridization of graphite
Class problem: do hybridization of the following:
of B in BF3 of S in SF6 of P in PCl5 of Be in BeH2
Polarity of bonds vs. Polarity of molecules
For most molecules, even though the bonded electron pair is said to be “shared” between two atoms,
often one atom will have a stronger attraction for the electron pair. This imbalance in sharing produces
some polarity between the atoms.
If the difference in electronegativity (EN) is greater than or equal to 0.5, but less than 1.7, the bond is
said to be polar covalent. If the difference is less than 0.5, then the bond is called nonpolar covalent.
Since the diatomic molecules share their electron pair equally, they are called pure covalent.
For binary compounds such as hydrogen chloride gas, the bonded electron pair migrates slightly closer to
the chlorine atom, thus producing a polar molecule, called a “dipole”.
For molecules that contain more than two atoms, whether or not a molecule is polar will depend on the
symmetry of the molecule.
Some molecules can have polar bonds, yet the molecule itself is nonpolar, ex. CO2. Carbon monoxide,
on the other hand has a polar bond, making the molecule highly polar.
Compounds such as methane, CH4, contain nonpolar bonds, and the molecule is nonpolar.
Compounds such as CF4 contain polar bonds, but the molecule is nonpolar.
Still other compounds, such as CH3F, contain both polar and nonpolar bonds, and because of the
asymmetrical shape, the molecule is polar.
What would you predict about the polarity of the nitrate ion? Sulfur dioxide molecule?
-6Types of Intermolecular Forces (IMF)
#1 ionic – no distinction between inter- and intra- due to uniform electrostatic attraction between ions
#2 covalent network – again, no distinction because of uniformity between atoms;
ex. Diamond – sp3 hybrid in a tetrahedral shape (cause of its extreme hardness)
graphite, silicon dioxide (sand)
#3 metallic – again, no distinction because of loose association of electrons to their nuclei
#4 dipole – dipole, first of the true IMF
when molecules that exhibit polarity are brought together, the opposite ends attract each other.
Ex. hydrogen chloride gas - the slightly positive hydrogen end of one HCl molecule is
attracted to the partially negative chloride end of another HCl molecule. Dipole-dipole
attraction can be the main IMF for a single substance, or a combination of polar substances.
#5 hydrogen
special type of dipole – dipole attraction; only occurs between molecules that contain hydrogen
and either N, O, or F. These three elements have unusually high electronegativities, and because they are
on the far right side in the second period, they are unusually small is size. If more than one compound is
in the mixture, both compound must each contain hydrogen and one of the three elements listed. ex.
water, ammonia, HF, or ammonia in water.
#6 ion – dipole
when an ionic compound, such as NaCl, is placed in a polar substance, such as water, the positive
sodium ions are bonded to the negative (oxygen) end of the water molecule. Likewise, the chloride ions
are attracted to the hydrogen end of the water molecule.
#7 ion – induced dipole
some substances, even if they are classified as nonpolar, have a small degree of polarity, for
example, SeBr2. If these “nonpolar” substances are placed in the presence of a highly polar substance
(ion vapor, such as NaCl), the polarity of the ionic substance causes the nonpolar compound to become
polarized.
#8 dipole – induced dipole
this type of IMF forms much like the ion – induced dipole, except that a strongly polar molecule,
like water, causes a nonpolar substance to become polarized.
Ex. SeBr2 in water.
#9 van der Waals (or London dispersion forces)
even for substances, such as He, which would have no polarity, there is evidence of
intermolecular forces. For example, since helium has a melting point and a boiling point, this indicates
that there is some type of attraction between atoms. This is due to a temporary relocation of the
electrons, producing a “temporary dipole”. The strength of the force increases with:
1)
increased number of electron (ex. M.Pt increases as you move from He to Ne to Ar, etc.
2)
increased points of contact between molecules; n-pentane has a higher M.Pt than neopentane
(show structural models).
Honors Chemistry I
Molecular 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.
Shape is described by the arrangement of the atoms, but this arrangement is affected by the
lone pair electrons as well as the bonded electrons.
** make sure hybridization is understood!!
Classification of shapes:
# of e- pairs
# of e- pairs # of e- pairs
shape
(total)
(bonding)
(nonbonding)
...........................................................................
1
1
0
linear
............................................................................
2
2
0
linear
.............................................................................
3
3
0
trigonal
planar
bent or
V-shaped
.............................................................................
bond angle
examples
180
H-Cl
180
CO2
120
BF3
< 120
SO2
109.5
CH4
3
2
1
4
4
0
tetrahedral
4
3
1
trigonal
pyramid
4
2
2
bent or
V-shaped
< 109.5
(104.5 for H2O)
H2S
linear
180
H-Cl
4
1
3
...................................................................
< 109.5
(107.5 for NH3)
NH3
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
-7# 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 with respect to 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 one
lone pair: if the lone pair were in an axial position, it would encounter electron pair repulsion
with three pair of electrons in a 90o 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 situation 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?
ex. ions: nitrate, triiodide, sulfate, sulfite, phosphate, carbonate, chlorate, perchlorate, hydronium,
ammonium, amide, hypochlorite