MCTC Molecular Models Lewis Structure and VSEPR Theory Paper

10.2: REPRESENTING VALENCE ELECTRONS WITH DOTSLEARNING OBJECTIVE
Draw a Lewis electron dot diagram for an atom or a monatomic ion.
In almost all cases, chemical bonds are formed by interactions of valence electrons in atoms. To facilitate our understanding of how
valence electrons interact, a simple way of representing those valence electrons would be useful.
A Lewis electron dot diagram (or electron dot diagram or a Lewis diagram or a Lewis structure) is a representation of the valence
electrons of an atom that uses dots around the symbol of the element. The number of dots equals the number of valence electrons in the
atom. These dots are arranged to the right and left and above and below the symbol, with no more than two dots on a side. (It does not
matter what order the positions are used.) For example, the Lewis electron dot diagram for hydrogen is simply
H⋅
(10.2.1)
Because the side is not important, the Lewis electron dot diagram could also be drawn as follows:
˙
H
or ⋅H
or
H
(10.2.2)
.
The electron dot diagram for helium, with two valence electrons, is as follows:
He:
(10.2.3)
By putting the two electrons together on the same side, we emphasize the fact that these two electrons are both in the 1s subshell; this
is the common convention we will adopt, although there will be exceptions later. The next atom, lithium, has an electron configuration
of 1s22s1, so it has only one electron in its valence shell. Its electron dot diagram resembles that of hydrogen, except the symbol for
lithium is used:
Li⋅
(10.2.4)
Beryllium has two valence electrons in its 2s shell, so its electron dot diagram is like that of helium:
Be:
(10.2.5)
The next atom is boron. Its valence electron shell is 2s22p1, so it has three valence electrons. The third electron will go on another side
of the symbol:
˙
Be :
(10.2.6)
Again, it does not matter on which sides of the symbol the electron dots are positioned.
For carbon, there are four valence electrons, two in the 2s subshell and two in the 2p subshell. As usual, we will draw two dots together
on one side, to represent the 2s electrons. However, conventionally, we draw the dots for the two p electrons on different sides. As
such, the electron dot diagram for carbon is as follows:
˙
⋅ C:
(10.2.7)
With N, which has three p electrons, we put a single dot on each of the three remaining sides:
˙
⋅ N:
(10.2.8)
.
For oxygen, which has four p electrons, we now have to start doubling up on the dots on one other side of the symbol. When doubling
up electrons, make sure that a side has no more than two electrons.
¨
⋅ O:
(10.2.9)
.
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Fluorine and neon have seven and eight dots, respectively:
¨:
:F
(10.2.10)
.
¨
:Ne:
(10.2.11)
. .
With the next element, sodium, the process starts over with a single electron because sodium has a single electron in its highestnumbered shell, the n = 3 shell. By going through the periodic table, we see that the Lewis electron dot diagrams of atoms will never
have more than eight dots around the atomic symbol.
EXAMPLE 10.2.1: LEWIS DOT DIAGRAMS
What is the Lewis electron dot diagram for each element?
a. aluminum
b. selenium
SOLUTION
a. The valence electron configuration for aluminum is 3s23p1. So it would have three dots around the symbol for aluminum, two
of them paired to represent the 3s electrons:
˙
Al :
(10.2.12)
2. The valence electron configuration for selenium is 4s24p4. In the highest-numbered shell, the n = 4 shell, there are six
electrons. Its electron dot diagram is as follows:
˙
⋅ S e:
(10.2.13)
. .
EXERCISE 10.2.1
What is the Lewis electron dot diagram for each element?
a. phosphorus
b. argon
Answer a
Answer b
SUMMARY
Lewis electron dot diagrams use dots to represent valence electrons around an atomic symbol.
Lewis electron dot diagrams for ions have less (for cations) or more (for anions) dots than the corresponding atom.
CONTRIBUTORS
Marisa Alviar-Agnew (Sacramento City College)
Henry Agnew (UC Davis)
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10.3: LEWIS STRUCTURES OF IONIC COMPOUNDS:
ELECTRONS TRANSFERRED
LEARNING OBJECTIVES
State the octet rule.
Define ionic bond.
Draw Lewis structures for ionic compounds.
In Section 4.7 we saw how ions are formed by losing electrons to make cations or by gaining electrons to form anions. The astute
reader may have noticed something: Many of the ions that form have eight electrons in their valence shell. Either atoms gain enough
electrons to have eight electrons in the valence shell and become the appropriately charged anion, or they lose the electrons in their
original valence shell; the lower shell, now the valence shell, has eight electrons in it, so the atom becomes positively charged. For
whatever reason, having eight electrons in a valence shell is a particularly energetically stable arrangement of electrons. The trend that
atoms like to have eight electrons in their valence shell is called the octet rule. When atoms form compounds, the octet rule is not
always satisfied for all atoms at all times, but it is a very good rule of thumb for understanding the kinds of bonding arrangements that
atoms can make.
It is not impossible to violate the octet rule. Consider sodium: in its elemental form, it has one valence electron and is stable. It is rather
reactive, however, and does not require a lot of energy to remove that electron to make the Na+ ion. We could remove another electron
by adding even more energy to the ion, to make the Na2+ ion. However, that requires much more energy than is normally available in
chemical reactions, so sodium stops at a 1+ charge after losing a single electron. It turns out that the Na+ ion has a complete octet in its
new valence shell, the n = 2 shell, which satisfies the octet rule. The octet rule is a result of trends in energies and is useful in
explaining why atoms form the ions that they do.
Now consider an Na atom in the presence of a Cl atom. The two atoms have these Lewis electron dot diagrams and electron
configurations:
¨
⋅ Cl :
Na ⋅
(10.3.1)
. .
1
2
[N e] 3s
5
[N e] 3s 3p
(10.3.2)
For the Na atom to obtain an octet, it must lose an electron; for the Cl atom to gain an octet, it must gain an electron. An electron
transfers from the Na atom to the Cl atom:
¨
Na ⋅ ↷ ⋅ Cl :
(10.3.3)
. .
resulting in two ions—the Na+ ion and the Cl− ion:
Na ⋅
+
−
¨
:Cl :
(10.3.4)
. .
[N e]
2
6
[N e] 3s 3p
(10.3.5)
Both species now have complete octets, and the electron shells are energetically stable. From basic physics, we know that opposite
charges attract. This is what happens to the Na+ and Cl− ions:
Na ⋅
+
−
+
−
¨
+ :Cl :
→ N a Cl
or
N aCl
(10.3.6)
. .
where we have written the final formula (the formula for sodium chloride) as per the convention for ionic compounds, without listing
the charges explicitly. The attraction between oppositely charged ions is called an ionic bond, and it is one of the main types of
chemical bonds in chemistry. Ionic bonds are caused by electrons transferring from one atom to another.
In electron transfer, the number of electrons lost must equal the number of electrons gained. We saw this in the formation of NaCl. A
similar process occurs between Mg atoms and O atoms, except in this case two electrons are transferred:
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The two ions each have octets as their valence shell, and the two oppositely charged particles attract, making an ionic bond:
2−
Mg
2+
¨
+ [:O :]
Mg
2+
2−
O
or M gO
(10.3.7)
. .
Remember, in the final formula for the ionic compound, we do not write the charges on the ions.
What about when an Na atom interacts with an O atom? The O atom needs two electrons to complete its valence octet, but the Na atom
supplies only one electron:
¨
Na ⋅ ↷ ⋅ O :
(10.3.8)
.
The O atom still does not have an octet of electrons. What we need is a second Na atom to donate a second electron to the O atom:
These three ions attract each other to give an overall neutral-charged ionic compound, which we write as Na2O. The need for the
number of electrons lost being equal to the number of electrons gained explains why ionic compounds have the ratio of cations to
anions that they do. This is required by the law of conservation of matter as well.
EXAMPLE 10.3.1: SYNTHESIS OF CALCIUM CHLORIDE FROM ELEMENTS
With arrows, illustrate the transfer of electrons to form calcium chloride from Ca atoms and Cl atoms.
SOLUTION
A Ca atom has two valence electrons, while a Cl atom has seven electrons. A Cl atom needs only one more to complete its octet,
while Ca atoms have two electrons to lose. Thus we need two Cl atoms to accept the two electrons from one Ca atom. The
transfer process looks as follows:
The oppositely charged ions attract each other to make CaCl2.
EXERCISE 10.3.1
With arrows, illustrate the transfer of electrons to form potassium sulfide from K atoms and S atoms.
Answer:
SUMMARY
The tendency to form species that have eight electrons in the valence shell is called the octet rule.
The attraction of oppositely charged ions caused by electron transfer is called an ionic bond.
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The strength of ionic bonding depends on the magnitude of the charges and the sizes of the ions.
CONTRIBUTORS
Marisa Alviar-Agnew (Sacramento City College)
Henry Agnew (UC Davis)
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10.4: COVALENT LEWIS STRUCTURES: ELECTRONS
SHARED
LEARNING OBJECTIVES
Define covalent bond.
Illustrate covalent bond formation with Lewis electron dot diagrams.
Ionic bonding typically occurs when it is easy for one atom to lose one or more electrons and another atom to gain one or more
electrons. However, some atoms won’t give up or gain electrons easily. Yet they still participate in compound formation. How? There
is another mechanism for obtaining a complete valence shell: sharing electrons. When electrons are shared between two atoms, they
make a bond called a covalent bond.
Let us illustrate a covalent bond by using H atoms, with the understanding that H atoms need only two electrons to fill the 1s subshell.
Each H atom starts with a single electron in its valence shell:
H⋅
⋅ H
(10.4.1)
The two H atoms can share their electrons:
H: H
(10.4.2)
We can use circles to show that each H atom has two electrons around the nucleus, completely filling each atom’s valence shell:
Because each H atom has a filled valence shell, this bond is stable, and we have made a diatomic hydrogen molecule. (This explains
why hydrogen is one of the diatomic elements.) For simplicity’s sake, it is not unusual to represent the covalent bond with a dash,
instead of with two dots:
H–H
Because two atoms are sharing one pair of electrons, this covalent bond is called a single bond. As another example, consider fluorine.
F atoms have seven electrons in their valence shell:
These two atoms can do the same thing that the H atoms did; they share their unpaired electrons to make a covalent bond.
Note that each F atom has a complete octet around it now:
We can also write this using a dash to represent the shared electron pair:
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There are two different types of electrons in the fluorine diatomic molecule. The bonding electron pair makes the covalent bond.
Each F atom has three other pairs of electrons that do not participate in the bonding; they are called lone pair electrons. Each F atom
has one bonding pair and three lone pairs of electrons.
Covalent bonds can be made between different elements as well. One example is HF. Each atom starts out with an odd number of
electrons in its valence shell:
The two atoms can share their unpaired electrons to make a covalent bond:
We note that the H atom has a full valence shell with two electrons, while the F atom has a complete octet of electrons.
EXAMPLE 10.4.1:
Use Lewis electron dot diagrams to illustrate the covalent bond formation in HBr.
SOLUTION
HBr is very similar to HF, except that it has Br instead of F. The atoms are as follows:
The two atoms can share their unpaired electron:
EXERCISE 10.4.1
Use Lewis electron dot diagrams to illustrate the covalent bond formation in Cl2.
Answer:
When working with covalent structures, it sometimes looks like you have leftover electrons. You apply the rules you learned so far and
there are still some electrons hanging out there unattached. You can’t just leave them there. So where do you put them?
MULTIPLE COVALENT BONDS
Some molecules are not able to satisfy the octet rule by making only single covalent bonds between the atoms. Consider the compound
ethene, which has a molecular formula of C H . The carbon atoms are bonded together, with each carbon also being bonded to two
hydrogen atoms.
2
4
two C atoms = 2 × 4 = 8 valence electrons
four H atoms = 4 × 1 = 4 valence electrons
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total of 12 valence electrons in the molecule
If the Lewis electron dot structure was drawn with a single bond between the carbon atoms and with the octet rule followed, it would
look like this:
Figure 10.4.1 : Incorrect dot structure of ethene.
This Lewis structure is incorrect because it contains a total of 14 electrons. However, the Lewis structure can be changed by
eliminating the lone pairs on the carbon atoms and having to share two pairs instead of only one pair.
Figure 10.4.2 : Correct dot structure for ethene.
A double covalent bond is a covalent bond formed by atoms that share two pairs of electrons. The double covalent bond that occurs
between the two carbon atoms in ethane can also be represented by a structural formula and with a molecular model as shown in the
figure below.
Figure 10.4.3 : (A) The structural model for C H consists of a double covalent bond between the two carbon atoms and single
bonds to the hydrogen atoms. (B) Molecular model of C H .
2
4
2
4
A triple covalent bond is a covalent bond formed by atoms that share three pairs of electrons. The element nitrogen is a gas that
composes the majority of Earth’s atmosphere. A nitrogen atom has five valence electrons, which can be shown as one pair and three
single electrons. When combining with another nitrogen atom to form a diatomic molecule, the three single electrons on each atom
combine to form three shared pairs of electrons.
Figure 10.4.4 : Triple bond in N .
2
Each nitrogen atom follows the octet rule with one lone pair of electrons and six electrons that are shared between the atoms.
SUMMARY
Covalent bonds are formed when atoms share electrons.
Lewis electron dot diagrams can be drawn to illustrate covalent bond formation.
Double bonds or triple bonds between atoms may be necessary to properly illustrate the bonding in some molecules.
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CONTRIBUTORS
CK-12 Foundation by Sharon Bewick, Richard Parsons, Therese Forsythe, Shonna Robinson, and Jean Dupon.
Anonymous
Marisa Alviar-Agnew (Sacramento City College)
Henry Agnew (UC Davis)
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10.5: WRITING LEWIS STRUCTURES FOR COVALENT
COMPOUNDS
SKILLS TO DEVELOP
Draw Lewis structures for covalent compounds.
The following procedure can be used to construct Lewis electron structures for more complex molecules and ions:
HOW-TO: CONSTRUCTING LEWIS ELECTRON STRUCTURES
1. Determine the total number of valence electrons in the molecule or ion.
Add together the valence electrons from each atom. (Recall that the number of valence electrons is indicated by the position of
the element in the periodic table.)
If the species is a polyatomic ion, remember to add or subtract the number of electrons necessary to give the total charge on the
ion.
For CO32−, for example, we add two electrons to the total because of the −2 charge.
2. Arrange the atoms to show specific connections.
When there is a central atom, it is usually the least electronegative element in the compound. Chemists usually list this central
atom first in the chemical formula (as in CCl4 and CO32−, which both have C as the central atom), which is another clue to the
compound’s structure.
Hydrogen and the halogens are almost always connected to only one other atom, so they are usually terminal rather than
central.
3. Place a bonding pair of electrons between each pair of adjacent atoms to give a single bond.
In H2O, for example, there is a bonding pair of electrons between oxygen and each hydrogen.
4. Beginning with the terminal atoms, add enough electrons to each atom to give each atom an octet (two for hydrogen).
These electrons will usually be lone pairs.
5. If any electrons are left over, place them on the central atom.
We will explain later that some atoms are able to accommodate more than eight electrons.
6. If the central atom has fewer electrons than an octet, use lone pairs from terminal atoms to form multiple (double or
triple) bonds to the central atom to achieve an octet.
This will not change the number of electrons on the terminal atoms.
7. Final check
Always make sure all valence electrons are accounted for and each atom has an octet of electrons except for hydrogen (with
two electrons).
The central atom is usually the least electronegative element in the molecule or ion; hydrogen and the halogens are usually
terminal.
Now let’s apply this procedure to some particular compounds, beginning with one we have already discussed.
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EXAMPLE 10.5.1: WATER
Write the Lewis Structure for H2O.
SOLUTION
Example 10.5.1
Steps for Writing Lewis Structures
1. Determine the total number of valence electrons in the molecule
or ion.
Each H atom (group 1) has 1 valence electron, and the O atom
(group 16) has 6 valence electrons, for a total of 8 valence
electrons.
H O H
Because H atoms are almost always terminal, the arrangement
2. Arrange the atoms to show specific connections.
within the molecule must be
3. Place a bonding pair of electrons between each pair of adjacent
atoms to give a single bond.
4. Beginning with the terminal atoms, add enough electrons to each
atom to give each atom an octet (two for hydrogen).
HOH.
Placing one bonding pair of electrons between the O atom and each
H atom gives
H -O- H
with 4 electrons left over.
Each H atom has a full valence shell of 2 electrons.
Adding the remaining 4 electrons to the oxygen (as two lone pairs)
gives the following structure:
5. If any electrons are left over, place them on the central atom.
6. If the central atom has fewer electrons than an octet, use lone
pairs from terminal atoms to form multiple (double or triple) bonds
to the central atom to achieve an octet.
Not necessary
7. Final check
The Lewis structure gives oxygen an octet and each hydrogen two
electrons,
EXAMPLE 10.5.2
Write the Lewis structure for the C H O molecule
2
Solution
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Example 10.5.2
Steps for Writing Lewis Structures
1. Determine the total number of valence electrons in the molecule
or ion.
Each hydrogen atom (group 1) has one valence electron, carbon
(group 14) has 4 valence electrons, and oxygen (group 16) has 6
valence electrons, for a total of [(2)(1) + 4 + 6] = 12 valence
electrons.
2. Arrange the atoms to show specific connections.
Because carbon is less electronegative than oxygen and hydrogen
is normally terminal, C must be the central atom.
Placing a bonding pair of electrons between each pair of bonded
atoms gives the following:
3. Place a bonding pair of electrons between each pair of adjacent
atoms to give a single bond.
Six electrons are used, and 6 are left over.
Adding all 6 remaining electrons to oxygen (as three lone pairs)
gives the following:
4. Beginning with the terminal atoms, add enough electrons to each
atom to give each atom an octet (two for hydrogen).
Although oxygen now has an octet and each hydrogen has 2
electrons, carbon has only 6 electrons.
Not necessary
5. If any electrons are left over, place them on the central atom.
There are no electrons left to place on the central atom.
6. If the central atom has fewer electrons than an octet, use lone
pairs from terminal atoms to form multiple (double or triple) bonds
to the central atom to achieve an octet.
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To give carbon an octet of electrons, we use one of the lone pairs of
electrons on oxygen to form a carbon–oxygen double bond:
10.5.3
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Example 10.5.2
Steps for Writing Lewis Structures
Both the oxygen and the carbon now have an octet of electrons, so
this is an acceptable Lewis electron structure. The O has two
bonding pairs and two lone pairs, and C has four bonding pairs.
This is the structure of formaldehyde, which is used in embalming
fluid.
7. Final check
EXERCISE 10.5.1
Write Lewis electron structures for CO2 and SCl2, a vile-smelling, unstable red liquid that is used in the manufacture of rubber.
Answer CO2
Answer SCl2
The United States Supreme Court has the unenviable task of deciding what the law is. This responsibility can be a major challenge
when there is no clear principle involved or where there is a new situation not encountered before. Chemistry faces the same challenge
in extending basic concepts to fit a new situation. Drawing of Lewis structures for polyatomic ions uses the same approach, but tweaks
the process a little to fit a somewhat different set of circumstances.
WRITING LEWIS STRUCTURES FOR POLYATOMIC IONS
Recall that a polyatomic ion is a group of atoms that are covalently bonded together and which carry an overall electrical charge. The
ammonium ion, NH , is formed when a hydrogen ion (H ) attaches to the lone pair of an ammonia (NH ) molecule in a coordinate
covalent bond.
+
+
4
3
Figure 10.5.3 : The ammonium ion.
When drawing the Lewis structure of a polyatomic ion, the charge of the ion is reflected in the number of total valence electrons in the
structure. In the case of the ammonium ion:
1N
4H
atom = 5 valence electrons
atoms = 4 × 1 = 4 valence electrons
subtract 1 electron for the 1+ charge of the ion
total of 8 valence electrons in the ion
It is customary to put the Lewis structure of a polyatomic ion into a large set of brackets, with the charge of the ion as a superscript
outside the brackets.
EXERCISE 10.5.2
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10.5.4
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Draw the Lewis electron dot structure for the sulfate ion.
Answer
EXCEPTIONS TO THE OCTET RULE
As important and useful as the octet rule is in chemical bonding, there are some well-known violations. This does not mean that the
octet rule is useless—quite the contrary. As with many rules, there are exceptions, or violations.
There are three violations to the octet rule. Odd-electron molecules represent the first violation to the octet rule. Although they are few,
some stable compounds have an odd number of electrons in their valence shells. With an odd number of electrons, at least one atom in
the molecule will have to violate the octet rule. Examples of stable odd-electron molecules are NO, NO2, and ClO2. The Lewis
electron dot diagram for NO is as follows:
Although the O atom has an octet of electrons, the N atom has only seven electrons in its valence shell. Although NO is a stable
compound, it is very chemically reactive, as are most other odd-electron compounds.
Electron-deficient molecules represent the second violation to the octet rule. These stable compounds have less than eight electrons
around an atom in the molecule. The most common examples are the covalent compounds of beryllium and boron. For example,
beryllium can form two covalent bonds, resulting in only four electrons in its valence shell:
Boron commonly makes only three covalent bonds, resulting in only six valence electrons around the B atom. A well-known example
is BF3:
The third violation to the octet rule is found in those compounds with more than eight electrons assigned to their valence shell. These
are called expanded valence shell molecules. Such compounds are formed only by central atoms in the third row of the periodic table
or beyond that have empty d orbitals in their valence shells that can participate in covalent bonding. One such compound is PF5. The
only reasonable Lewis electron dot diagram for this compound has the P atom making five covalent bonds:
Formally, the P atom has 10 electrons in its valence shell.
EXAMPLE 10.5.3: OCTET VIOLATIONS
Identify each violation to the octet rule by drawing a Lewis electron dot diagram.
a. ClO
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b. SF6
Solution
a. With one Cl atom and one O atom, this molecule has 6 + 7 = 13 valence electrons, so it is an odd-electron molecule. A Lewis
electron dot diagram for this molecule is as follows:
b. In SF6, the central S atom makes six covalent bonds to the six surrounding F atoms, so it is an expanded valence shell molecule.
Its Lewis electron dot diagram is as follows:
EXERCISE 10.5.3: XENON DIFLUORIDE
Identify the violation to the octet rule in XeF2 by drawing a Lewis electron dot diagram.
Answer:
SUMMARY
Lewis dot symbols provide a simple rationalization of why elements form compounds with the observed stoichiometries. A plot of the
overall energy of a covalent bond as a function of internuclear distance is identical to a plot of an ionic pair because both result from
attractive and repulsive forces between charged entities. In Lewis electron structures, we encounter bonding pairs, which are shared
by two atoms, and lone pairs, which are not shared between atoms. Lewis structures for polyatomic ions follow the same rules as those
for other covalent compounds. There are three violations to the octet rule: odd-electron molecules, electron-deficient molecules, and
expanded valence shell molecules
CONTRIBUTORS
Modified by Joshua Halpern (Howard University), Scott Sinex, and Scott Johnson (PGCC)
CK-12 Foundation by Sharon Bewick, Richard Parsons, Therese Forsythe, Shonna Robinson, and Jean Dupon.
Marisa Alviar-Agnew (Sacramento City College)
Henry Agnew (UC Davis)
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10.5.6
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10.6: RESONANCE – EQUIVALENT LEWIS STRUCTURES FOR
THE SAME MOLECULE
SKILLS TO DEVELOP
Explain the following laws within the Ideal Gas Law
RESONANCE
There are some cases in which more than one viable Lewis structure can be drawn for a molecule. An example is the ozone (O )
molecule in the figure below. There are a total of 18 electrons in the structure and so the following two structures are possible.
3
Figure 10.6.2 : Resonance forms of ozone.
The structure on the left (see figure above) can be converted to the structure on the right by a shifting of electrons without altering the
positions of the atoms.
It was once thought that the structure of a molecule such as O consisted of one single bond and one double bond which then shifted
back and forth as shown above. However, further studies showed that the two bonds are identical. Any double covalent bond between
two given atoms is typically shorter than a single covalent bond. Studies of the O and other similar molecules showed that the bonds
were identical in length. Interestingly, the length of the bond is in between the lengths expected for an O−O single bond and a double
bond.
3
3
Resonance is the use of two or more Lewis structures to represent the covalent bonding in a molecule. One of the valid structures is
referred to as a resonance structure. It is now understood that the true structure of a molecule which displays resonance is that of an
average or a hybrid of all the resonance structures. In the case of the O molecule, each of the covalent bonds between O atoms is best
thought of as being “one and a half” bonds, as opposed to either a pure single bond or a pure double bond. This “half-bond” can be
shown as a dotted line in both the Lewis structure and the molecular model (see figure below).
3
Figure 10.6.3 : “Half-bond” model of ozone molecule.
Many polyatomic ions also display resonance. In some cases, the true structure may be an average of three valid resonance structures,
as in the case of the nitrate ion, NO (see figure below).
−
3
Figure 10.6.4 : Resonance structure of nitrate anion.
The bond lengths between the central N atom and each O atom are identical and the bonds can be approximated as being equal to one
and one-third bonds.
SUMMARY
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Resonance structures are averages of different Lewis structure possibilities.
Bond lengths are intermediate between covalent bonds and covalent double bonds.
CONTRIBUTORS
CK-12 Foundation by Sharon Bewick, Richard Parsons, Therese Forsythe, Shonna Robinson, and Jean Dupon.
Marisa Alviar-Agnew (Sacramento City College)
Henry Agnew (UC Davis)
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10.6.2
Updated 8/29/2018
10.7: PREDICTING THE SHAPES OF MOLECULES
LEARNING OBJECTIVE
Determine the shape of simple molecules.
Molecules have shapes. There is an abundance of experimental evidence to that effect—from their physical properties to their chemical
reactivity. Small molecules—molecules with a single central atom—have shapes that can be easily predicted. The basic idea in
molecular shapes is called valence shell electron pair repulsion (VSEPR). It basically says that electron pairs, being composed of
negatively charged particles, repel each other to get as far away from each other as possible. VSEPR makes a distinction between
electron group geometry, which expresses how electron groups (bonds and nonbonding electron pairs) are arranged, and molecular
geometry, which expresses how the atoms in a molecule are arranged. However, the two geometries are related.
There are two types of electron groups: any type of bond—single, double, or triple—and lone electron pairs. When applying VSEPR
to simple molecules, the first thing to do is to count the number of electron groups around the central atom. Remember that a multiple
bond counts as only one electron group.
Any molecule with only two atoms is linear. A molecule whose central atom contains only two electron groups orients those two
groups as far apart from each other as possible—180° apart. When the two electron groups are 180° apart, the atoms attached to those
electron groups are also 180° apart, so the overall molecular shape is linear. Examples include BeH2 and CO2:
Figure 10.7.1 : Beryllium hydride and carbon dioxide bonding.
The two molecules, shown in the figure below in a “ball and stick” model.
Figure 10.7.2 : Beryllium hydride and carbon dioxide models.
A molecule with three electron groups orients the three groups as far apart as possible. They adopt the positions of an equilateral
triangle—120° apart and in a plane. The shape of such molecules is trigonal planar. An example is BF3:
Figure 10.7.3 : Boron trifluoride bonding.
Some substances have a trigonal planar electron group distribution but have atoms bonded to only two of the three electron groups. An
example is GeF2:
Figure 10.7.4 : Germanium difluoride bonding.
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From an electron group geometry perspective, GeF2 has a trigonal planar shape, but its real shape is dictated by the positions of the
atoms. This shape is called bent or angular.
A molecule with four electron groups about the central atom orients the four groups in the direction of a tetrahedron, as shown in
Figure 10.7.1 Tetrahedral Geometry. If there are four atoms attached to these electron groups, then the molecular shape is also
tetrahedral. Methane (CH4) is an example.
Figure 10.7.5 : Tetrahedral structure of methane.
This diagram of CH4 illustrates the standard convention of displaying a three-dimensional molecule on a two-dimensional surface. The
straight lines are in the plane of the page, the solid wedged line is coming out of the plane toward the reader, and the dashed wedged
line is going out of the plane away from the reader.
Figure 10.7.3 : Methane bonding.
NH3 is an example of a molecule whose central atom has four electron groups but only three of them are bonded to surrounding atoms.
Figure 10.7.3 : Ammonia bonding.
Although the electron groups are oriented in the shape of a tetrahedron, from a molecular geometry perspective, the shape of NH3 is
trigonal pyramidal.
H2O is an example of a molecule whose central atom has four electron groups but only two of them are bonded to surrounding atoms.
Figure 10.7.3 : Water bonding.
Although the electron groups are oriented in the shape of a tetrahedron, the shape of the molecule is bent or angular. A molecule with
four electron groups about the central atom but only one electron group bonded to another atom is linear because there are only two
atoms in the molecule.
Double or triple bonds count as a single electron group. CH2O has the following Lewis electron dot diagram.
The central C atom has three electron groups around it because the double bond counts as one electron group. The three electron
groups repel each other to adopt a trigonal planar shape:
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(The lone electron pairs on the O atom are omitted for clarity.) The molecule will not be a perfect equilateral triangle because the C–O
double bond is different from the two C–H bonds, but both planar and triangular describe the appropriate approximate shape of this
molecule.
Table 10.7.1 summarizes the shapes of molecules based on their number of electron groups and surrounding atoms.
Table 10.7.1: Summary of Molecular Shapes
Number of Electron
Groups on Central
Atom
Number of Bonding
Pairs
Number of Lone Pairs
Electron Geometry
Molecular Shape
2
2
0
linear
linear
3
3
0
trigonal planar
trigonal planar
3
2
1
trigonal planar
bent
4
4
0
tetrahedral
tetrahedral
4
3
1
tetrahedral
trigonal pyramidal
4
2
2
tetrahedral
bent
EXAMPLE 10.7.1:
What is the approximate shape of each molecule?
a. PCl3
b. NOF
SOLUTION
The first step is to draw the Lewis structure of the molecule.
1. For PCl3, the electron dot diagram is as follows:
The lone electron pairs on the Cl atoms are omitted for clarity. The P atom has four electron groups with three of them bonded
to surrounding atoms, so the molecular shape is trigonal pyramidal.
2. The electron dot diagram for NOF is as follows:
The N atom has three electron groups on it, two of which are bonded to other atoms. The molecular shape is bent.
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EXERCISE 10.7.1
What is the approximate molecular shape of CH2Cl2?
Answer
EXERCISE 10.7.2
Ethylene (C2H4) has two central atoms. Determine the geometry around each central atom and the shape of the overall molecule.
Hint, hydrogen is a terminal atom.
Answer
SUMMARY
The approximate shape of a molecule can be predicted from the number of electron groups and the number of surrounding atoms.
CONTRIBUTORS
Marisa Alviar-Agnew (Sacramento City College)
Henry Agnew (UC Davis)
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10.8: ELECTRONEGATIVITY AND POLARITY: WHY OIL AND
WATER DON’T MIX
SKILLS TO DEVELOP
Explain how polar compounds differ from nonpolar compounds.
Determine if a molecule is polar or nonpolar.
Identify whether or not a molecule can exhibit hydrogen bonding.
List important phenomena which are a result of hydrogen bonding.
Given a pair of compounds, predict which would have a higher melting or boiling point.
BOND POLARITY
The ability of an atom in a molecule to attract shared electrons is called electronegativity. When two atoms combine, the difference
between their electronegativities is an indication of the type of bond that will form. If the difference between the electronegativities of
the two atoms is small, neither atom can take the shared electrons completely away from the other atom and the bond will be covalent.
If the difference between the electronegativities is large, the more electronegative atom will take the bonding electrons completely
away from the other atom (electron transfer will occur) and the bond will be ionic. This is why metals (low electronegativities) bonded
with nonmetals (high electronegativities) typically produce ionic compounds.
A bond may be so polar that an electron actually transfers from one atom to another, forming a true ionic bond. How do we judge the
degree of polarity? Scientists have devised a scale called electronegativity, a scale for judging how much atoms of any element attract
electrons. Electronegativity is a unitless number; the higher the number, the more an atom attracts electrons. A common scale for
electronegativity is shown in Figure 10.8.1 .
Figure 10.8.1 : Electronegativities of the Elements. Electronegativities are used to determine the polarity of covalent bonds.
The polarity of a covalent bond can be judged by determining the difference of the electronegativities of the two atoms involved in the
covalent bond, as summarized in the following table:
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Electronegativity Difference
Bond Type
0
nonpolar covalent
0–0.4
slightly polar covalent
0.5–2.0
definitely polar covalent
>2.0
likely ionic
NONPOLAR COVALENT BONDS
A bond in which the electronegativity difference is less than 1.9 is considered to be mostly covalent in character. However, at this point
we need to distinguish between two general types of covalent bonds. A nonpolar covalent bond is a covalent bond in which the
bonding electrons are shared equally between the two atoms. In a nonpolar covalent bond, the distribution of electrical charge is
balanced between the two atoms.
Figure 10.8.2 : A nonpolar covalent bond is one in which the distribution of electron density between the two atoms is equal.
The two chlorine atoms share the pair of electrons in the single covalent bond equally, and the electron density surrounding the Cl
molecule is symmetrical. Also note that molecules in which the electronegativity difference is very small (