Class Notes - Unit 4: Covalent Bonding and Molecule Properties
(Ch. 8 and some of 9)
Lewis Structures
Lewis structures are two dimensional representations based on electron dot structures to
show the arrangement of atoms when they are covalently bonded.
Nonmetals share electrons in covalent bonds in order to have a valence shell similar to
that of noble gases. The electrons shared between atoms belong to the outer shell of both
atoms. They are known as bonding electrons and are represented by a dash for ever pair.
They count towards the total valence number for both atoms involved in the bond.
Electrons that aren’t shared between atoms count only towards the atom that they are on.
They are called nonbonding or unshared electron pairs. They remain as dots.
There are a few simple steps for drawing Lewis structures that will work no matter
what molecule you are dealing with:
1. Total the number of valence electrons from all atoms involved in the molecule
2. Arrange the elements in the general way that they are positioned in relation to each
other. Compare how many bonds each element involved can potentially make. An
element can potentially make as many bonds as it has lone electrons in its electron dot
structure. For example, C can make 4, N can make 3, O can make 2, and F can make 1.
The central atom will be the atom that can make the most bonds. The other atoms will be
placed as symmetrically as possible around this central atom. If you have two atoms with
the same potential number of bonds, you either need to be told the central atom or the
element that there are fewer of is the central atom. For example in SO2, the S is the
central atom.
3. Distribute your total number of valence electrons. Start with distributing 2 electrons
into a single bond between each pair of atoms. Then distribute the remaining electrons
around the atoms starting with the most electronegative atom first. If there are multiple
atoms with the same electronegativity, the electrons get distributed 2 at a time evenly
between them.
4. Check to see if the octet rule is satisfied for each atom (remember exceptions to the
octet rule). If all atoms satisfy the octet rule, the Lewis structure is complete.
5. If one or more atoms do not satisfy the octet rule, double bonds need to be made. Any
pair of electrons that is not already being shared in a single bond can move from an
adjoining atom into a double bond. Remember to form double bonds symmetrically if
multiple ones need to be formed. Check to see if the octet rule is satisfied.
6. Form triple bonds. Keep in mind rules previously stated.
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A triple bond (made up of 6 electrons being shared between 2 atoms) is the maximum
number of electrons that can be shares in a covalent bond. A single bond is the most
common type of covalent bond, then a double bond (2 pairs of shared electrons), and
finally a triple bond.
The atmosphere is 80% nitrogen (N2 = two nitrogens triple bonded to each other) and
20% oxygen (O2 = two oxygens double bonded to each other). Very few compounds exist
with nitrogen however, compared to the variety that exist with oxygen. Why is this true if
there is so much more nitrogen present in the atmosphere? The triple bond in nitrogen is
much strong than the double bond in oxygen. The bonds in oxygen are much easier to
break (require less energy) and therefore potentially bond with other things.
The strength of bonds is: triple > double > single > ionic
The structure of molecules can be represented in a number of ways. A structural formula
shows the pairs of bonding electrons in the form of dashes and nonbonding electrons as
dots.
A molecular formula only displays the ratio of each element in the molecule and does not
contain any arrangement or structural information. Unlike in ionic compounds, which are
formula units (the lowest whole number ratio of elements in the compound), molecular
compounds are not in the lowest whole number ratios. For example, hydrogen peroxide is
H2O2 and the 2’s are not reduced.
Exceptions to the Octet Rule
There are four common exceptions to the octet rule:
1. H and He only need 2 electrons to be full/stable.
2. Be can be stable with only 4 electrons.
3. B can be stable with only 6 electrons.
4. Elements in the 3rd energy level and above can have more than 8 electrons (up
to 18) in their valence shell. These are known as “expanded octets.” This can
occur because once the 3rd energy level is reached, there are d orbital available
to hold electrons.
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VSEPR Theory and Molecular Geometry
Lewis dot structures are two dimensional representations of the structure of molecules. In
reality, molecules are three dimensional. The three-dimensionality of a molecule is
critically linked to its chemical and physical properties. A sound understanding of
molecular shape is necessary to fundamentally understand chemical phenomena.
Chemists routinely predict the structural characteristics of proposed molecules when
designing substances for specific applications. The structure of a molecule is inherently
linked to its function.
For main group compounds, valence-shell electron-pair repulsion (VSEPR) theory is a
systematic methodology for linking a two-dimensional Lewis structure to a molecule’s
three-dimensional shape. The application of VSEPR permits the reliable prediction of
molecular structure and bond angles. In VSEPR bonding theory, all bonds are viewed as
single bonds (even if in reality they are double or triple). This is possible because the four
electrons in a double bond or the six electrons in a triple bond have a very high
probability of being in the region between the two bonded atoms, since they are being
shared. Each group of electrons must be accommodated around the central atom. Because
electrons repel each other they will stay as far apart as possible. They can spread
anywhere in a three-dimensional sphere around the central atom to utilize the most space
possible.
There are arrangements for electron pairs surrounding a particular atom that are most
favorable. These arrangements are found using simple geometric constructions. This
involves placing the central atom at the center of a sphere and then placing the electron
pairs on the surface of the sphere so that they are as far apart as possible.
Number of Electron Pairs
2
3
4
5
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Basic Geometry
linear
trigonal planar
tetrahedral
trigonal bipyramid
octahedral
Bond Angle
180°
120°
109.5°
90° and 120°
90°
Unshared electron pairs play an important in predicting the shapes of molecules. Because
no other atom is competing for these electrons they are held closer to the atom. These
unshared pairs therefore, strongly repel the bonding pairs pushing them closer together
and take up more space than the bonded pair(s) of electrons. If there are unshared
electron pairs involved in the molecule, the bond angles between the atoms involved will
be slightly smaller for the same general shape as if there were only bonded electron pairs
surrounding the atom.
Basic Geometry
Tetrahedral
Tetrahedral
Trigonal Bipyramid
Trigonal Bipyramid
Octahedral
# of Unshared e- pairs
1
2
1
2
2
New Geometry
Trigonal Pyramid
Bent
Seesaw
T-shaped
Square Planar
Bond Angle
107°
105°
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Polarity
Electronegativity is the ability of an atom to attract electrons when that atom is in a
compound.
In ionic compounds, there is a big difference in electronegativity between the atoms
involved (≥1.7). The anion has a high electronegativity value and really wants/attracts
electrons. The cation has a low electronegativity value and willingly gives up electrons.
In molecular compounds, the atoms involved are closer in electronegativity (with a
difference < 1.7 between the two atoms). Molecular compounds involve covalent bonds
where the electrons are shared between the atoms in the bond. However, all electrons in
covalent bonds are not shared equally and how they’re shared depends on how
similar or different the electronegativities between the atoms are.
Nonpolar covalent bonds
- the electrons in the covalent bond are shared equally
- occurs in diatomic molecule (the atoms are the same and therefore the
electronegativity difference is 0)
- occurs when the atoms involved have very similar electronegativity values (a
difference of <0.4)
Polar covalent bonds
- the electrons in the covalent bond are shared unequally (one atom attracts the
electrons significantly more than the other)
- the electron distribution is not symmetrical
- moderately polar covalent bonds are between atoms whose electronegativity
difference is between 0.4 and 1.0
- very polar covalent bonds are between atoms whose electronegativity difference
is between 1.0 and 1.7
- the more electronegative atom attracts the electrons with more force and is
“partially negative” (δ-)
- the less electronegative atom is more willing to have the electrons not around as
often and is “partially positive” (δ+)
Electronegativity Difference
Between Atoms
0 - 0.4
0.4 - 1.7
≥ 1.7
Most Probable Type of Bond
Nonpolar Covalent
Polar Covalent
Ionic
Ex: Cl2 electronegativity of Cl = 3.0
electronegativity difference between Cl and Cl = 0
nonpolar covalent
Ex: HCl electronegativity of H = 2.1 and of Cl = 3.0
electronegativity difference between H and Cl = 0.9
polar covalent
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Demo:
Polar rubber bands (one strong and one weak) with 2 paper clips in between - slowly pull
rubber bands apart. The paper clips stay closer to the thicker rubber band
Nonpolar rubber bands (two of the same strength) with 2 paper lips in between - slowly
pull apart and paper clips stay in the middle of the rubber bands
Covalent bonds are a sharing of electrons that can also be thought of as a tug of war for
electrons. The atom with the higher electronegativity value gets the electron more of the
time. Electrons are negatively charged. The atom that has electrons around it more of the
time (has negative charges around it more of the time) has a partial negative charge
Ex: H2O electronegativity of H = 2.1 and of O = 3.5
electronegativity difference between H and O = 1.4
polar covalent bonds
Bond polarity helps determine molecular properties (such as melting point, boiling point,
etc.).
A polar molecule is a molecule with two “poles” (also known as a dipole molecule). A
polar molecule can happen in two different ways:
(1) The molecule has one end with a partial positive charge and the other end with
a partial negative charge.
(2) The molecule has one end with all polar bonds and the other end with all
nonpolar bonds.
You must draw the molecule in the correct geometry in order to see if it is polar or not.
Ex: H2O
electronegativity of H = 2.1 and of O = 3.5
electronegativity difference between H and O = 1.4 = polar covalent bonds
= polar molecule (due to bent geometry; one positive side, one negative side)
Ex: CO2
electronegativity of C = 2.5 and of O = 3.5
electronegativity difference between C and O = 1.0 = polar covalent bonds
= nonpolar molecule (due to linear geometry; the bond polarities cancel out due to
linearity and the charge distribution being on the same axis but in opposite directions; no
two distinct sides)
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Ex: CH3CH2COOH (or C3H6O2)
electronegativity of C = 2.5, O = 3.5, H = 2.1
electronegativity difference between C and H = 0.4 = nonpolar cov. bonds
electronegativity difference between C and O = 1.0 = polar covalent bonds
electronegativity difference between O and H = 1.4 = polar covalent bonds
= polar molecule (due to one polar side, one nonpolar side)
Demo: A burette filled with water (polar). Rub a balloon with hair to charge it negatively.
Bring the balloon near the open end of the water quickly. Because the water is polar is
has positive and negative ends to its molecule. The polar water is attracted to the
negatively charged balloon. The polar water molecules align themselves with the external
charge and the water stream is pulled towards the balloon.
Like Dissolves Like
One of the reasons we learn about polarity is because the polarity of molecules is one of
the key factors that determine how they will interact with other molecules.
Like dissolves like
- polar molecules dissolve and mix with other polar molecules (forming a
homogeneous mixture)
- nonpolar molecules dissolve and mix with other nonpolar molecules (forming
a homogeneous mixture)
-
polar molecules DO NOT dissolve and mix with nonpolar molecules (forming
a heterogeneous mixture)
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Intermolecular Forces
Ionic compounds are 3D arrangements of positive and negative charges. It is the ionic
bonding that holds these ions together (give and take of electrons). Ionic compounds and
sample of ionic compounds are solids due to these electrostatic forces between ions.
Molecular compounds are a group of atoms held together by covalent bonds (a sharing of
electrons). These covalent bonds that bind the atoms together are very strong. Between
one molecular and another molecule however, there are much weaker, sometimes almost
nonexistent forces holding them together. For this reason, a sample of a molecular
compound at room temperature is a gas or liquid.
If there were absolutely no other forces between the molecules, all covalently bonded
molecules would be gases at any temperature. There would be no attractive forces present
to pull them together to make a liquid or solid.
Between molecules there are three kinds of weak forces, known as “intermolecular
forces.”
There are Van der Waals forces which include: dipole-dipole interactions and dispersion
(or London) forces. Dipole-dipole interactions are the attractions between the oppositely
charged poles of polar molecules. It is similar to ionic bonding in that the molecules tend
to align due to the electrostatic attraction between oppositely charged things. However, it
is not nearly as strong as ionic bonding because the + and - charges are only partial
charges.
Dispersion forces are the weakest of all intermolecular forces. They arise from the
random motion of electrons around an atom that creates instantaneous dipoles. These
forces are present in all compounds (ionic, polar molecular, and nonpolar molecular). The
strength of the dispersion force increases as the number of electrons increase. This occurs
because the more electrons you have orbiting the atoms, the more occurrences of
instantaneous dipoles.
The second class of intermolecular forces is hydrogen bonding. Hydrogen bonding is a
specific type of dipole interaction where H is the partial positive charge. The occurs when
H is bonded to a very electronegative atom (such as F, O, or N). In this strongly polar
bond, the hydrogen is electron deficient therefore it tries to share one of its neighbors
nonbonding electron pairs to compensate for its deficiency. A hydrogen bond is the
strongest of the intermolecular forces and is about 5% the strength of the average
covalent bond.
The intermolecular forces present in a molecule have a direct impact on the molecule’s
properties. For example, the stronger the intermolecular forces, the higher the melting
and boiling point of the molecule; the longer it takes the molecule to evaporate; the more
viscous the molecule; and the greater the surface tension of the molecule.
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When you boil water, you are not actually breaking up the covalent bonds between the
hydrogen and oxygen atoms within the molecule to product the gas. You are breaking up
the hydrogen bonding between one water molecule and another water molecule to allow
the water molecule as a whole to change into the gas state.
Demo:
Write names and structures for acetone (CH3COCH3), ethyl alcohol (CH3CH2OH), and
water (HOH) on the board. Squirt several mL of the liquid on the board under its
structure. Observe the rate of spreading on the surface and the rate of evaporation.
Evaporation is the change from a liquid to a gas. If two liquids evaporate at the same
temperature, the difference in evaporation rate is directly dependent on the intermolecular
forces holding the liquid together.
The evaporation rate for this demo is acetone > ethyl alcohol > water. The slow rate of
evaporation of water molecules reflects the strong hydrogen bonding in water. Only the
water molecules at the liquid surface that have enough energy to break the hydrogen
bonds will escape into the gas phase. Hydrogen bonding is also present in ethyl alcohol,
although to a lesser extent than water. Acetone is a polar compound. The molecules are
held together in the liquid phase by dipole-dipole interactions, which are weaker than
hydrogen bonds.
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