General Chemistry Lab Report

Instructions:For each structure provided:
1. Provide the preferred Lewis structure for the listed structure. (not typed)
2. Calculate the formal charge for each atom type in the structure.
You may include the formal charge calculation in the table or on a separate page.
All work should be neat and readable. Calculations do not have to be typed.
3. Give the number of electron domains.
4. Indicate the ABn formula for each Molecule.
5. Give the name of the electron Geometry.
6. Provide the number of lone pairs.
7. Give name of the molecular geometry.
8. Tell if the bonds are polar ( Yes-Y) or non-polar (No- N).
9. Tell if the molecule is polar (Yes-Y) or non-polar (No- N).
Bonus material: I have provided another file for bonus points. You are to complete the 3-D
VESPER structure for each of Lewis structures. It does not have to be typed, but it must be
correct and readable to be given credit.
General Chemistry 1
Laboratory 10: Lewis Dot Structures and Molecular Geometry
March 2013
Laboratory 10: Lewis Dot Structures and Molecular Geometry
10.1 Introduction
10.1.1 General
G. N. Lewis (in about 1916) observed that many elements are most stable when they contained
eight electrons in their valence shell. He suggested that atoms with fewer than eight valence
electrons bond together to share electrons and complete their valence shells. Bonds form
tending to stabilize a chemical system by releasing energy. The larger the amount of energy
released during the formation of a bond, the more stable the bond will be. If two atoms release
energy (Exothermic reaction) by forming a bond, then the atoms will be more stable by staying
together than they would be as individual atoms.
10.1.2 Valence Electrons
The electronic configuration of an atom is given by listing its subshells with the number of
electrons in each subshell, as shown in Table 1. According to the Aufbau principle, the
electrons of an atom occupy quantum levels or orbital’s starting from the lowest energy level,
and proceeding to the highest, with each orbital holding a maximum of two paired electrons
(opposite spins). Study the third column of complete electronic configurations carefully so you
understand how electrons are added to the subshell of lowest energy until it reaches its capacity;
then the subshell of the next energy level begins to be filled. The electrons on the highest
numbered subshells are the valence electrons, which comprise the valence shell of the atom and
participative in chemical bonding.
Table 1: Electron Configurations and Oxidation Numbers
Element Name
Hydrogen
Helium
Lithium
Beryllium
Boron
Carbon
Nitrogen
Oxygen
Fluorine
Atomic Number
1
2
3
4
5
6
7
8
9
Electron Configuration
1 s1
1 s2
1 s2 2 s1
1 s2 2 s2
1 s2 2 s2 2 p1
1 s2 2 s2 2 p2
1 s2 2 s2 2 p3
1 s2 2 s2 2 p4
1 s2 2 s2 2 p5
Valence Shell
1 s1
1 s2
2 s1
2 s2
2 s2 2 p1
2 s2 2 p2
2 s2 2p3
2 s2 2 p4
2 s2 2 p5
Many chemists record the electron configuration of an atom by listing the noble gas symbol in
brackets followed by the remaining electron configuration of an element. This process easily
provides the valance electrons. An example is potassium which would be abbreviated as [Ar]4s1
or carbon as [He]2s12p2.
10.1.3 Chemical Properties Associated with Valance Electrons
The chemical properties of the elements reflect their electron configurations and especially their
out shell electrons. For example, helium, neon and argon are exceptionally stable and un-reactive
mono-atomic gases. Helium is unique since its valence shell consists of a single s-orbital. The
other members of group 8 have a characteristic valence shell electron octet (ns2 + npx2 + npy2 +
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npz2). This group of inert (or noble) gases also includes Krypton (Kr: 4s2, 4p6), Xenon (Xe: 5s2,
5p6) and Radon (Rn: 6s2, 6p6) (Figure 1). In the periodic table above these elements are colored
beige.
The halogens (F, Cl, Br etc.) are one electron short of a valence shell octet, and are among the
most reactive of the elements (they are colored red in this periodic table, Figure 1). In their
chemical reactions halogen atoms achieve a valence shell octet by capturing or borrowing the
eighth electron from another atom or molecule. The alkali metals Li, Na, K etc. (colored violet
above) are also exceptionally reactive, but for the opposite reason. These atoms have only one
electron in the valence shell, and on losing this electron arrive at the lower shell valence octet.
As a consequence of this electron loss, these elements are commonly encountered as cations
(positively charged atoms).
1A
2A
3A
4A
5A
6A
7A
1
H
1s1
8A
2
He
1s2
3
Li
1s2
2s1
4
Be
1s2
2s2
5
B
1s2
2s22p1
6
C
1s2
2s22p2
7
N
1s2
2s22p3
8
O
1s2
2s22p4
9
F
1s2
2s22p5
10
Ne
1s2
2s22p6
11
Na
[Ne]
3s1
12
Mg
[Ne]
3s2
13
Al
[Ne]
3s23p1
14
Si
[Ne]
3s23p2
15
P
[Ne]
3s23p3
16
S
[Ne]
3s23p4
17
Cl
[Ne]
3s23p5
18
Ar
[Ne]
3s23p6
Figure 1: Trends in reactivity associated with valance electron
The elements in groups 2 through 7 all exhibit characteristic reactivates and bonding patterns that
can in large part be rationalized by their electron configurations. It should be noted that hydrogen
is unique. Its location in the periodic table should not suggest a kinship to the chemistry of the
alkali metals, and its role in the structure and properties of organic compounds is unlike that of
any other element.
10.1.4 Chemical Bonding and Valance Electron
As noted earlier, the inert gas elements of Group 8A exist as monoatomic gases, and do not in
general react with other elements. In contrast, other gaseous elements exist as diatomic
molecules (H2, N2, O2, F2 & Cl2), and all but nitrogen are quite reactive. Some dramatic
examples of this reactivity are shown in the following equations (Table 2).
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Table 2: Selected reactions producing ionic or covalent compounds
2Na(s) + Cl2(g)
2NaCl(s)
2H2(g) + O2(g)
2H2O(l)
C(s) + O2(g)
CO2(g)
C(s) + 2F2(g)
CF4(g)
Why do the atoms of many elements interact with each other and with other elements to give
stable molecules?
10.1.5 Ionic Bonding
When sodium is burned in a chlorine atmosphere, it produces the compound sodium chloride.
This has a high melting point (~800 ºC) and dissolves in water to give a solution that conducts
electrical current. Sodium chloride is an ionic compound, and is a crystalline solid as shown in
Figure 2. When sodium and chlorine come into contact with each other, the valance electron
(3s1) of the Na is transferred to the chloride atom to complete 3S23P5 orbital (Cl-) which forms a
stable octet (argon electron configuration). The sodium cation (Na+) has the noble gas
configuration of Ne, due Na losing a single electron. Remember that atoms with low ionization
energies tend to form cations and those with high positive electron affinities tend to form anions.
The electrostatic attraction of the cations and the anions results in these oppositely charged ions
packing together in a crystal lattice and creating an electrically neutral compound. The attractive
force holding the ions in place is referred to as ionic bonds. Remember, that the strength of an
ionic bond is proportional to the product of the charges of the ions (cation charge X anion
charge) and inversely related to the radius squared (1/ r2).
Figure 2: Pictorial representation of the formation of an ionic compound (NaCl) and the crystal
formation by electrostatic forces.
For ions, the valence equals the electrical charge. The common oxidation numbers in the third
column of Table 3 are interpreted as the result of either losing the valence electron/s (leaving a
positive ion) or gaining enough electrons to fill that valence subshell. The charges on the
chlorine, potassium, and calcium ions result from a strong tendency of valence electrons to adopt
the stable configuration of the Nobel gases, with completely filled electronic shells. Notice that
in Table 3, the 3 ions have electronic configurations identical to that of argon.
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Table 3: Comparison of three ions and a neutral atom to form a complete octet
(Noble Gas Configuration). Note: the resulting ions below are isoelectronic.
Chemical Element Valence Shell Electron Transfer Resulting Ion
Cl
3 s23 p5
gain 1
Cl−
2
6
Ar
3s 3p
None
Ar0
K
4 s1
lose 1
K+
Ca
4 s2
lose 2
Ca2+
Ion Configuration
1 s2 2 s2 2 p6 3 s2 3 p6
1 s2 2 s2 2 p6 3 s2 3 p6
1 s2 2 s2 2 p6 3 s2 3 p6
1 s2 2 s2 2 p6 3 s2 3 p6
10.1.6 Covalent Bonding
The other three reactions shown in Table 2 give products that are very different from sodium
chloride. Water is a liquid at room temperature; carbon dioxide and carbon tetrafluoride are
gases. None of these compounds are composed of ions. A different attractive interaction between
atoms, called covalent bonding, is involved here. Covalent bonding occurs by a sharing of
valence electrons, rather than an electron transfer to form ions. Similarities in physical properties
suggest that the diatomic elements H2, N2, O2, F2, and Cl2 also have covalent bonds, since they
are all gases at room temperature.
Examples of covalent bonding shown below include hydrogen, fluorine, carbon dioxide and
carbon tetrafluoride. These illustrations use a simple Bohr notation, with valence electrons
designated by colored dots. Each hydrogen atom has one valance electron. The hydrogen atoms
achieve a helium electron configuration by sharing a pair of 1s-electrons and form a single
covalent bond. In the other examples carbon, oxygen and fluorine achieve neon valence octets
by a similar sharing of electron pairs. Please notice that carbon dioxide uses two pairs of
electrons (four in all) in forming two covalent bonds with each oxygen atom. Both oxygen and
the carbon atoms have completed the octet (neon electron configuration) by forming a double
bond (4 electrons) between the carbon and each of the oxygen atoms.
Figure 4: Diagrammatic presentation of the sharing of electrons to form covalent bonds in
accordance with the octet rule (Except hydrogen: it holds only two electrons)
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10.1.7 Polar Covalent Bonds
Not all covalent bonds are created equal. Because of their differing nuclear charges, and as a
result of shielding by inner electron shells, the different atoms of the periodic table have different
affinities for nearby electrons. The ability of an element to attract an electron in a bond is called
electronegativity and is related to electron affinity and electron ionization. An atom like
fluorine which has a high electron affinity and a high ionization energy, will have a high
electronegativity, Sodium on the other hand has a low electron affinity and a low ionization
energy and has a low electronegativity. A rough quantitative scale of electronegativity values
was established by Linus Pauling, and some of these are given in the Table 4. A larger number
on this scale signifies a greater affinity for electrons. Again, fluorine has the greatest
electronegativity of all the elements, and the heavier alkali metals such as potassium, rubidium
and cesium have the lowest electronegativities. It should be noted that carbon is about in the
middle of the electronegativity range, and is slightly more electronegative than hydrogen.
Table 4: Electronegativity values for some selected elements
H
2.1
Li
0.98
Be
1.57
B
2.04
C
2.5
N
3.04
O
3.5
F
3.98
Na
0.90
Mg
1.31
Al
1.61
Si
1.90
P
2.19
S
2.58
Cl
3.16
K
0.82
Ca
1.00
Ga
1.81
Ge
2.01
As
2.18
Se
2.55
Br
2.96
When two different atoms are bonded covalently, the shared electrons are attracted to the more
electronegative atom of the bond, resulting in a shift of electron density toward the more
electronegative atom. This type of covalent bond is called polar, and will have a dipole (one end
is partially positive and the other end partially negative).
A
B
C
Figure 5: Pictorial representation of the effect of increasing electronegativity. Note the location
of the pair of electrons as the electronegativity increases from right to left resulting in
a greater dipole moment in these covalent bonds.
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The degree of polarity and the magnitude of the bond dipole will be proportional to the
difference in electronegativity of the bonded atoms. So the bonds O–H and C-H are both polar,
but O-H has the greatest dipole moment. If we calculate the differences in the electronegativities
of the two bonds, the electronegativity of the O-H (1.4) is greater C-H (0.4). Both oxygen and
carbon have a greater electronegativity when compared to hydrogen, attracting the shared
elections creating a dipole moment, Figure 6 (A). Likewise, C–Cl and C–Li bonds are both polar
bonds. Chlorine has a greater electronegativity than carbon, so the chlorine will attract the
elections, creating a dipole moment, leaving the carbon end of the molecule more positive. For
C–Li bond, carbon has a greater electronegativity when compared to lithium; the electrons will
spend most of their time with carbon leaving the lithium with a positive partial charge, Figure 6
(B). The dipolar nature of these bonds is often indicated by a partial charge notation (δ+/–) or by
an arrow pointing to the negative end of the bond.
A
B
Figure 6: Dipoles of selected bonds
So a diatomic molecule has two atoms bonded to each other by a covalent bond. In such a
molecule, the dipole moment of the bond gives the dipole moment of the molecule. Thus, a
diatomic molecule is polar if the bond formed between the atoms is polar, Figure 7. The greater
the electronegativity difference between the atoms, the greater the dipole moment. In
polyatomic molecules the dipole moment not only depends upon the individual dipole moments
of the bonds but also on the spatial arrangement of the various bonds in the molecule. In such
molecules the dipole moment of the molecule is the vector sum of the dipole moments of various
bonds.
Figure 7: Diatomic molecule with a polar bond resulting in the molecule having a
dipole moment.
A good example of a polar covalent molecule is water.
Figure 8: Dipole moment of water as the result of the molecular geometry of the polar bonds
(O-H)
This molecule has one oxygen atom covalently bonded to two hydrogen atoms. The oxygen
atom shares two valence electrons and each hydrogen atom shares one. The oxygen atom has a
greater electronegativity than the hydrogen. Each O-H bond is polar and the resulting molecule
of water has an overall dipole as shown in Figure 8.
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10.1.8 Lewis Dot Structure
Although the octet rule and Lewis structures alone do not present a complete picture of covalent
bonding, they do help us account for some of the properties of molecules. In addition, Lewis
structures provide a starting point for other bonding theories. Therefore it is critical for you to be
able to correctly draw Lewis structures for molecules and polyatomic ions.
10.1.8.1 Rules for Lewis Structure
1. Count the total number of valence electrons for each atom.
Remember the group number of the atom gives the number of valance electrons.
Remember to adjust for charge if an ion is present:
Add one electron for each negative charge
Subtract one electron for each positive charge.
2. From the molecular formula, draw the skeletal structure of the compound.
For simple compounds, this step is easy.
Often there will be a central atom surrounded by other identical atoms.
Figure 9: Central atom (Carbon) surround by identical atoms (Hydrogen).
A two dimensional picture
In general the least electronegative atom will be the central atom.
Hydrogen is never a central atom.
Hydrogen and the halogens are usually outside atoms.
Place a single dash or two dots (pair of electrons) between the central atom and the
surroundings atoms as shown in Figure 9.
3. For each bond in the skeletal structure, subtract two electrons from the total valence
electrons to determine the remaining number of electrons.
4. Complete the octet (8 electrons) of the terminal atoms (except H, which needs only 2
electrons) with the remaining electrons by placing pairs of electrons around each atom.
If there is more than one type of terminal atom, complete the octet of the most
electronegative atoms first.
Note: Boron will have only six outer shell electrons (incomplete octet).
5. Place extra electrons, if any, on the central atom in pairs.
6. If the octet on the central atom is not complete, move one or more pairs from the terminal
atoms to form multiple bonds between the central atom and the terminal atoms.
Note: halides and hydrogen do not form double bonds. If more than one possible Lewis dot
structure is possible, choose the one that yields the lowest formal charge on the atoms of the
molecule. Remember that the most electronegative atom should have the negative formal charge
if there is a choice. Use these guides lines to help you through the following examples.
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10.1.8.2 Sample Lewis Dot Problems
Lewis dot for NH3
Step 1
Determine the Total Number of Valiance Electrons
Element
Valence Electrons
Number of Atoms
Total Valance Electrons
Nitrogen
5
1
1
3
5
3
8
Hydrogen
Total electrons for
bonding
Step
2
From the molecular formula, draw the skeletal structure of the compound.
Often there will be a central atom surrounded by other identical atoms.
In general the least electronegative atom will be the central atom.
Hydrogen is never a central atom.
Step
3
For each bond in the skeletal structure, subtract two electrons from the total valence electrons to
determine the remaining number of electrons.
8 Valance Electrons– 6 Bonding Electrons = 2 Electrons Remain
Step
4
Complete the octet around the outer atoms first.
Hydrogen can have only two electrons, so the hydrogen atoms are complete. We
have two extra electrons which will be placed on the central atom (Nitrogen) (Step
5)
Step
5
Place extra electrons, if any, on the central atom in pairs, i.e. 1 electron pair.
Step
6
If the octet on the central atom is not complete, move one or more pairs from the terminal atoms to
form multiple bonds between the central atom and the terminal atoms
Nitrogen has a complete octet and each hydrogen has 2 electrons. Finished
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Construct the Lewis dot for CH4
Step
1
Determine the Total Number of Valiance Electrons
Element
Valence
Electrons
Number
of Atoms
Total Valance Electrons
4
1
1
4
4
4
Carbon
Hydrogen
Total
electrons for
bonding
8
From the molecular formula, draw the skeletal structure of the compound.
Step Often there will be a central atom surrounded by other identical atoms. In
general the least electronegative atom will be the central atom. Hydrogen is
2
never a central atom.
Step For each bond in the skeletal structure, subtract two electrons from the total valence electrons to
determine the remaining number of electrons.
3
8 Valance Electrons – 8 Bonding Electrons = Zero Electrons Left
Step Complete the octet around the outer atoms first, once this has been completed, place extra electrons, if
any, on the central atom in pairs. Hydrogen can have only two electrons, so the hydrogen atoms are
4
complete
Carbon has a complete octet a and hydrogen has 2 electrons
Step Place extra electrons, if any, on the central atom in pairs.
5
We have no extra electrons
Step If the octet on the central atom is not complete, move one or more pairs from the terminal atoms to
form multiple bonds between the central atom and the terminal atoms.
6
All Atoms have the Octet with the exception of hydrogen which will hold only 2 electrons. Finished.
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Lewis dot for CH2O
Step
1
Determine the Total Number of Valiance Electrons
Element
Carbon
Hydrogen
Oxygen
Step
2
Valence
Electrons
Number
of Atoms
4
1
6
1
2
1
Total Valance Electrons
4
2
6
12
Total electrons for
bonding
From the molecular formula, draw the skeletal structure of the compound.
Often there will be a central atom surrounded by other identical atoms. In
general the least electronegative atom will be the central atom. Hydrogen is
never a central atom.
Step For each bond in the skeletal structure, subtract two electrons from the total valence electrons to
determine the remaining number of electrons.
3
12 Valance Electrons– 6 Bonding Electrons = 6 Electrons Left
Step
Complete the octet around the outer atoms first. Hydrogen can have only two electrons, so the
4
hydrogen atoms are complete
We will complete the octet for oxygen.
Step Place extra electrons, if any, on the central atom in pairs.
5
We have no extra electrons.
Step If the octet on the central atom is not complete, move one or more pairs from the terminal atoms to
form multiple bonds between the central atom and the terminal atoms.
6
We need to move two electrons (circled in red) from
the oxygen atom and form a double bond between the
carbon and oxygen atom to complete the octet for
oxygen and carbon, as seen in the figure to the right.
10.1.9 Exceptions to the Octet Rule and Resonance Structures
10.1.9.1 Expanded Octet
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or
Many atoms expand their octet. Only atoms with d orbitals can expand their octet. This requires
that the atom have a principal quantum number of 3 or more. Therefore these atoms will be in
the third or higher period of the periodic table and have an atomic number of 12 or more.
Note: Although these atoms can expand their octet, they do not always do so. Only the central
atom will expand its octet. After drawing a structure in the normal way, if the formal charges on
the molecule are decreased by creating a double bond, the double bond will form.
Lewis dot for XeF4
Step 1
Determine the Total Number of Valiance Electrons
Element
Xenon
Fluorine
Valence
Electrons
Number
of Atoms
8
7
1
4
Total Valance Electrons
Total electrons for
bonding
8
28
36
Step 2
From the molecular formula, draw the skeletal structure of the compound.
Often there will be a central atom surrounded by other identical atoms.
In general the least electronegative atom will be the central atom.
Halides as a general rule form single bonds.
Step 3
For each bond in the skeletal structure, subtract two electrons from the total valence electrons to
determine the remaining number of electrons.
36 Valance Electrons– 8 Electrons (bonds) = 24 Electrons Left
Step 4
Complete the octet around the outer atoms first.
We will complete the octet for the fluorine atoms.
Step
5
Place extra electrons, if any, on the central atom in pairs.
We have four extra electrons label in red to add to the central atom.
Step 6
If the octet on the central atom is not complete, move one or more pairs from the terminal atoms to
form multiple bonds between the central atom and the terminal atoms. All atoms have the octet.
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10.1.9.2 Resonance Structures
It’s not uncommon to have a molecule where a Lewis structure does not perfectly describe the
molecule. For example, consider the sulfur dioxide molecule, SO2:
Figure 10: Lewis structure for SO2
The above structure would indicate that there are two different types of bonds in the molecule:
one single S-O bond and one double S=O bond. However, experimental data shows that this is
not correct: there is only one S-O bond length, equal to about a bond and a half. This is an
example where the basic Lewis diagram breaks down: to indicate the fact that there is only one
type of bond, we draw a resonance structure
Figure 11: Pictorial representation of the resonance structures of SO2
Here, we draw the two possible structures and a double headed arrow between them to indicate
that the structure is an average of these two structures. Both of these resonance Lewis dot
structures are equally valid.
It cannot be overemphasized that the various structures in a resonance structure do NOT really
exist. SO2 does not exist as two structures flipping back and forth, but only as the average of the
two structures. Think of it as a mutt: a dog that is 1/2 doberman and 1/2 retriever is an average of
the two breeds; it’s not a doberman 1/2 the time and a retriever the other half.
It’s possible to have more than two resonance forms: for example, the SO3 molecule has three:
Figure 12: Pictorial representation of the resonance structures of SO3.
Resonance structures are typically needed when you can draw a molecule multiple different
ways, and all structures are equally valid.
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Example:
Draw the resonance Lewis structure for the formate ion COOH-.
Solution: If we simply write a Lewis structure for the ion, we end up with two possible Lewis
structures of the formate ion:
Figure 13: Pictorial representation of two Lewis structures for COOH-.
If we use the formal charge calculation:
Equation 1: FC = [# of valence e- in free atom] – ½[# of bonding e-] _ [# of non-bonding e-]
[# of valence e-]
Atom
Carbon
Hydrogen
Oxygen Double Bond
Oxygen Single Bond
4
1
6
6
½[# of bonding e-]
½*8
½*2
½*4
½*2
[# of non-bonding e-]
Formal Charge
0
0
4
6
0
0
0
-1
We see that both oxygen’s will have -1and 0, carbon and H both have 0 for their formal charges.
We can’t decide which is correct based on formal charge, but both are correct.
In fact, this is a resonance structure, Figure 14.
Figure 14: Pictorial representation of the resonance structure for COOH-.
10.2 Molecular Geometry
Molecular geometry is a three dimensional representation of the molecule. We can use VSEPR
modeling and the Lewis structure to help predict the geometry of the molecule.
10.2.1 VSEPR Model
The VSEPR (Valence-Shell Electron-Pair Repulsion) model is one way to generally represent
the geometric shape individual molecules will take. It is not perfectly accurate but instead gives a
general impression of how atoms and electrons will be arranged. The AX method is commonly
used in formatting molecules to fit the VSEPR model. The A represents the central atom and is
always (implied) subscript one. The X represents the electron domain; the electron domain can
be represented by a lone pair of electrons or a bond. Regardless how X is attached to the central
atom (single, double or triple bond), it is still counted as a single electron domain.
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Figure 15: The basic structures used in the VSEPR geometries.
To draw the VSEPR geometry, follow these steps:
1. Draw the Lewis structure.
2. Determine the number of electron domains on the central atom and its electron-domain
geometry. Please, note that a single, double or triple bond counts as one electron domain.
3. Determine the number of lone pairs and their locations,
4. Predict the molecular geometry.
Example Problem NF3
1. 4 electron domains on the central atom.
Electron-domain geometry: tetrahedral
2. One lone pair on the central atom.
3. Molecular geometry: trigonal pyramidal
Table 5: VSEPR Geometries
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10.2.2 Polarity and Non Polarity
In Figure 15, the first three geometries have each geometry has its own bond angles. In the
trigonal bipyramidal and the octahedral geometries, there are two separate bond angle types: 1)
axial position (up and down) and 2) equatorial at 90 degrees to the axial position, Figure 16.
Figure 16: Location of axial and equatorial positions in the trigonal bipyramidal geometry.
With the first three geometries (linear, trigonal planer and the tetrahedral, seen in Figure 15, if all
the outer atoms are identical with no unpaired electrons, then the molecule will be non-polar.
However, if we replace an atom with an atom having a different electronegativity or with a lone
pair of electrons, the molecule will be polar.
When drawing the geometry for the trigonal bipyramidal, electron pairs are located at the
equatorial position. If all five outer atoms are identical in electronegativity then the molecule is
non-polar. The molecule can be non-polar if all three outer atoms in the equatorial position are
identical (-H) and the axial outer atoms are identical (-Cl) this includes three electrons pairs in
the equatorial position. Any other arrangement is polar.
With the octahedral geometry, if all the outer atoms are identical or outer atoms are identical
located directly opposite one another, the molecule is non-polar
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Examples of Free Radical Molecules
Recall that the Lewis structure of a molecule must depict the total number of valence electrons from
all the atoms which are bonded together.
Nitric oxide has the formula NO. The total number of valence electrons is 5 (N) +6 (O)=11.
Therefore, no matter how electrons are shared between the nitrogen and oxygen atoms, there is no
way for nitrogen to have an octet. It will have seven electrons, assuming the oxygen
atom does satisfy the octet.
Nitric oxide is a by-product of combustion reactions that occur in engines, like those in automobile
engines and fossil fuel power plants. It is also produced naturally during the electrical discharge of
lightning during thunderstorms.
Nitrogen dioxide is the chemical compound with the formula NO2. Again, nitrogen dioxide does not
follow the octet rule for one of its atoms, namely nitrogen. The total number of valence electrons is
5+2(6)=17. There is persistent radical character on nitrogen because it has an unpaired electron.
The two oxygen atoms in this molecule follow the octet rule.
Nitrogen dioxide: Nitrogen dioxide is another stable molecule that disobeys the octet rule. Note
the seven electrons around nitrogen. Formal charges and the molecule’s resonance structures are
indicated.
Nitrogen dioxide is an intermediate in the industrial synthesis of nitric acid, millions of tons of
which is produced each year. This reddish-brown toxic gas has a characteristic sharp, biting odor
and is a prominent air pollutant.