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2 Polar Covalent Bonds, Shape, and Polarity

2 Polar Covalent Bonds, Shape, and Polarity

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4



A



n Introduction to

Organic Compounds



About this Chapter

Now that we have seen the basics of how atoms and compounds are put together, it

is time to set the stage for the organic and biochemistry topics that we will encounter

as we move through the text. This chapter will begin with an introduction to a range of

diverse topics, all of which relate to how bonding patterns, the distribution of electrons,

and the shape of a molecule influence the way that it interacts with others. In the last

part of the chapter we will focus on organic families of compounds. An organic compound

is placed into a particular family based on the types of atoms or bonds that it contains.



CHAPTER 4 OBJECTIVES

After completing this chapter, you should be able to:

1 Draw molecules and polyatomic ions using electron dot and Lewis structures.

2 Describe how condensed structural formulas and skeletal structures differ from electron dot and

Lewis structures.

3 Define electronegativity and explain its relationship to polar covalent bonds. Give a simple rule that

can be used to predict whether or not a covalent bond is polar.

4 List the five basic shapes about an atom in a molecule or ion and describe the rules used to predict

shape. Explain how shape plays a role in determining overall polarity.

5 Name the different types of noncovalent interactions and be able to predict the important noncovalent interactions that hold molecules to one another.

6 Given the structure of a molecule, be able to identify the family to which it belongs.



117



118   Chapter 4  An Introduction to Organic Compounds



4.1



S tr u ct u ral F or m u las

Before we can consider organic molecules, we must expand on the discussion of covalent

bonds and molecules that was initiated in Chapter 3. Section 3.5 introduced molecules,

the uncharged groups of nonmetal atoms that are connected to one another by covalent

bonds, and showed that the structure of a given molecule can be represented by an electron dot structure (all valence electrons are shown using dots) or a Lewis structure (each

pair of shared bonding electrons is represented by a line). In the structures of isopropyl

alcohol (rubbing alcohol) shown below, the carbon atoms and oxygen atom have formed

the number of covalent bonds required to reach an octet—four bonds for carbon atoms

and two bonds for oxygen atoms.

Electron dot structure



Lewis structure



H H H



H H H

HC C CH

H O H

H



ƒ



ƒ



ƒ



ƒ



ƒ



ƒ



H¬ C ¬ C ¬ C ¬ H

H O H

ƒ



Isopropyl alcohol



H



Either of these structural formulas provides more information about isopropyl alcohol than does the molecular formula (C3H8O), because molecular formulas tell nothing

about how atoms are attached to one another. Suppose, for example, that a toxicologist is

reporting on the health problems associated with exposure to high levels of the compound

C3H8O. Because this molecular formula does not tell us how the molecule is put together,

it is not clear whether she is talking about isopropyl alcohol or one of the other two molecules that have the same molecular formula. Representing the molecule using a structural

formula (Figure 4.1) clarifies the matter.

Given a molecular formula, how do we go about drawing a structural formula? For

small molecules, knowing the number of covalent bonds that an atom is expected to form

can be a good place to start. As we saw in Section 3.4, the number of covalent bonds that a

nonmetal atom forms is generally the same as the number of electrons that it needs to gain

in order to have an octet. An oxygen atom has six electrons, needs two more to gain an

octet, and can form two covalent bonds. A nitrogen atom has five valence electrons, needs

three more to gain an octet, and can form three covalent bonds. The expected bonding

patterns for hydrogen and the period two nonmetals are:

• A hydrogen atom can form 1 covalent bond.

• A carbon atom can form 4 covalent bonds.

• A nitrogen atom can form 3 covalent bonds.

• An oxygen atom can form 2 covalent bonds.

• A fluorine atom can form 1 covalent bond.

To draw the structural formula for HF, join H and F by a covalent bond to give each

its required number of bonds. For NH3, N becomes the central atom by forming three

■ ■ FIGURE 4.1

Structural formulas  Structural



formulas, like the Lewis formulas

shown here for the three molecules with the formula C3H8O,

indicate the relative position of

each atom within a molecule.



H H H

ƒ



ƒ



ƒ



ƒ



ƒ



ƒ



H¬C¬C¬C¬H

H O H

ƒ



H



H H H

ƒ



ƒ



ƒ



ƒ



ƒ



ƒ



H¬C¬ C ¬C¬ O ¬H

H H H



H

ƒ



H H

ƒ



ƒ



ƒ



ƒ



H ¬ C ¬O ¬ C ¬ C ¬ H

ƒ



H



H H



4.1  Structural Formulas   119



covalent bonds, one to each H atom. Note that in the structures of HF and NH3, the

F and N atoms each carry enough nonbonding electrons to have an octet. Figure 3.11

shows other structural formulas that can be drawn using this approach.

H F     H N



H



H

Not all structural formulas are this easy to come up with. Fortunately, for more challenging situations there is a systematic approach that can be used. This approach will be

introduced by showing how NCl3 is drawn.



Drawing Lewis Structures









1. Count the total number of valence electrons.  In a molecule, an atom’s valence

electrons end up either in covalent bonds or as nonbonding electrons. The structural formula of NCl3 should show 26 electrons:

1 N atom



+





5 valence electrons

+







5



3 Cl atoms

3¯˚˚˚˘˚˚˚˙

 7 valence electrons



+



21



= 26 valence electrons



2. Use single covalent bonds to connect the atoms to one another.  The one bit

of guesswork here will be deciding which atom to attach to another. Although

not always the case, central atoms often come first in the molecular formula. The

structure of NCl3 follows this pattern: N is the central atom. Note that although

hydrogen is sometimes listed first in molecular formulas, it can never be a central

atom because hydrogen forms just one covalent bond.

In NCl3, nitrogen is the central atom, so we will connect it to the chlorine

atoms using single bonds.

Cl



N Cl

Cl







3. Beginning with the atoms attached to the central atom, add the remaining

electrons to complete octets.  Six valence electrons have been used in the NCl3

structure above: two for each of the three covalent bonds. Of the original

26 electrons, 20 must still be added. Drawing enough electrons to give each

chlorine atom an octet gives

Cl



N Cl

Cl



This structure is not complete, because only 24 electrons are shown (3 bonds and

9 pairs of nonbonding electrons). The two remaining electrons are added to the

nitrogen atom.

Cl



N Cl

Cl







4. If the central atom does not have an octet, move pairs of nonbonding electrons from attached atoms to form multiple bonds with the central atom. Do

so until the central atom has an octet.  The NCl3 drawn in step 3 is completed

and need not be modified. The nitrogen atom has an octet (3 bonds and 1 pair

of nonbonding electrons), as do the chlorine atoms (1 bond and 3 pairs of nonbonding electrons each).



120   Chapter 4  An Introduction to Organic Compounds



Let us repeat this process for formaldehyde, which has the formula CH2O. In this molecule, C is the central atom and is attached to each of the other three atoms.

1. Count the total number of valence electrons.  CH2O has 12 valence electrons

available.





+



1 C atom





4 valence electrons +









1 O atom



2  1 valence electron +

¯˚˚˚˘˚˚˚˙



+



4



+



2 H atoms



6 valence electrons



+ 6 = 12 valence electrons



2



2. Use single covalent bonds to connect the atoms to one another.  As the hint

above explained, C is the central atom.

H



C



O



H





3. Beginning with the atoms attached to the central atom, add the remaining

electrons to complete octets.  The drawing above shows three single bonds, so 6

electrons are represented. That leaves 6 more to add. Hydrogen atoms form just

one covalent bond and can accept no more electrons, so the remaining electrons

must go to the oxygen atom.

H



C



O



H







Additional nonbonding electrons are not added to the carbon atom because all of

the 12 valence electrons have been used.

4. If the central atom does not have an octet, move pairs of nonbonding electrons from attached atoms to form multiple bonds with the central atom.

Do so until the central atom has an octet.  The carbon atom does not have an

octet, so a pair of nonbonding electrons is moved from the oxygen atom and

used to form a double bond with the carbon atom. This gives the carbon atom an

octet and the drawing is complete.

H



C



O



H



H



C



O



H



Carbon has 6 valence electrons   Carbon has an octet



SAMPLE PROBLEM   4.1



Drawing Lewis structures

Draw the Lewis structure of hydrogen cyanide, HCN. In this molecule, carbon is the central atom.

Strategy



After completing step 3 of the procedure just described, you will find that the carbon atom

does not have an octet. Move as many pairs of nonbonding electrons as needed to give

carbon an octet.

Solution



1.The molecular formula of HCN indicates that 10 valence electrons are available, 1 from

H, 4 from C, and 5 from N.

2.Carbon is the central atom, so the initial structure is



H C



N



4.1  Structural Formulas   121



3.Having two covalent bonds, the structure above uses only 4 of the available 10 valence

electrons. The hydrogen atom cannot accept any more electrons, so the remaining 6 are

added to nitrogen.



H



C



N



4.While the nitrogen atom has an octet, carbon does not. Pairs of nonbonding electrons

from N are moved to form multiple bonds with C, until C has an octet.



H



C



H



N



Carbon has 4

valence electrons



C



H



N



Carbon has 6

valence electrons



C



N



Carbon has an octet



Practice Problem   4.1



Draw the Lewis structure of carbon disulfide, CS2. In this molecule, C is the central atom.



Structural formulas of polyatomic ions can be drawn using the same set of four guidelines that we used to draw molecules. Let us use those guidelines to draw nitrite ion

(NO2-).

1. Count the total number of valence electrons.  NO2- has 18 electrons available.

In the case of polyatomic anions, one electron must be added for each negative

charge. For polyatomic cations, one electron must be subtracted for each positive

charge.





1 N atom



+





5 valence electrons +









5



2 O atoms



+



 6 valence electrons

2¯˚˚˚˘˚˚˚˙



+



12



+



+



1 negative charge

1 electron

1  = 18 electrons



2. Use single covalent bonds to connect the atoms to one another.  Nitrogen,

which is listed first in the formula, is the central atom.

O N O







3. Beginning with the atoms attached to the central atom, add the remaining electrons to complete octets.  In the drawing above, 4 of the 18 available

electrons have been used. The remaining 14 will be added, beginning with the O

atoms and then moving to the N atom.

O







N



O



4. If the central atom does not have an octet, move pairs of nonbonding electrons from attached atoms to form multiple bonds with the central atom. Do

so until the central atom has an octet.  The nitrogen atom does not have an

octet, so a pair of nonbonding electrons is moved from one of the oxygen atoms

(it does not matter which one). This gives the nitrogen atom an octet.

O N



O N



O



O



Nitrogen has 6 valence electrons   Nitrogen has an octet



To show the charge on nitrite ion, square brackets are placed around the structural formula, and the charge is written as a superscript on the right side. The final correct structure is



[O



N







O]



122   Chapter 4  An Introduction to Organic Compounds

SAMPLE PROBLEM   4.2



Drawing Lewis structures of polyatomic ions

Draw the Lewis structure of NO2+.

Strategy



As in NO2-, nitrogen is the central atom. When counting available electrons in polyatomic

cations, subtract one electron for each positive charge.

Solution



1.A nitrogen atom has 5 valence electrons and each oxygen atom has 6. This gives an initial total of 17 available electrons. Because of the positive charge on the ion, 1 electron

is subtracted, giving a total of 16 available electrons for NO2+.

2.Nitrogen is the central atom, so the initial structure is



O N O

3.Having two covalent bonds, the structure above uses only 4 of the available 16 valence

electrons. The remaining 12 are added to the oxygen atoms.



O N



O



4.To give the nitrogen atom an octet, one pair of nonbonding electrons from each oxygen

atom is used to make a multiple bond.



O N



O N



O



Nitrogen has 4

valence electrons



O N



O



Nitrogen has 6

valence electrons



[O



N



O



Nitrogen has an octet



+



O]



PRACTICE PROBLEM   4.2



Draw the Lewis structure of PO33-.



Many compounds, especially those encountered in organic and biochemistry, are quite

large, and drawing their electron dot or Lewis structural formulas can be a time-consuming

task. For this reason, chemists have devised more abbreviated methods for representing

structure. A condensed structural formula describes the attachment of atoms to one

another, without showing all of the bonds. For example, a carbon atom with three attached

hydrogen atoms can be written CH3 and one with two attached hydrogen atoms can be

written CH2. Several examples of condensed structural formulas are shown in Figure 4.2.

In skeletal structures (Figure 4.2), covalent bonds are represented by lines, carbon

atoms are not shown, and hydrogen atoms are drawn only when attached to atoms other

than carbon. To read a skeletal structure, you assume that a carbon atom appears where

lines (bonds) meet and at the end of each line. To simplify matters, nonbonding electrons

are sometimes omitted from skeletal and other structural formulas.

SAMPLE PROBLEM   4.3



Drawing condensed and skeletal structures

Draw condensed and skeletal structures of diethyl ether, a compound once used as a general anesthetic.



H H

H H

ƒ

ƒ

ƒ

ƒ

Hi C i C iO

i C i C iH











H H

H H



4.2 Polar Covalent Bonds, Shape, and Polarity   123

Strategy



To write the condensed formula, begin by thinking of the molecule as a chain of five atoms

(C i C i O i C i C) and then add in the atoms attached to these five (CH3, etc.). The

skeletal structure of diethyl ether is drawn by leaving out all of the carbon and hydrogen

atoms, showing only the oxygen atom and the C i C and C i O bonds.

Solution



O



CH3CH2OCH2CH3



Condensed structure   Skeletal structure



PRACTICE PROBLEM   4.3



Research suggests that drinking green tea may help boost the immune system. Ethylamine

(below), produced when one of the compounds in green tea is broken down in the liver,

may be responsible for this immune response. Draw condensed and skeletal structures for

the molecule.



H H

ƒ



ƒ



ƒ



ƒ



H¬ C ¬ C ¬ N¬ H

ƒ



H H H

Lewis

structure



Condensed

formula



■ ■ FIGURE 4.2



Skeletal

structure



Condensed and skeletal

structural formulas  Condensed



H H H

H



C



C



formulas show the attachment of

atoms without always showing

covalent bonds and nonbonding

electrons. In skeletal structures

carbon atoms are not shown and

hydrogen atoms appear only

when attached to atoms other

than carbon.



CH3CH2CH3



C H



H H H

Propane



H H H H

H



C



C



C



H H F



C H



CH3CH2CHCH3

F



H



F



2-Fluorobutane



O

H

H



C

C



O H



C

C



C

C



C



O H



O C



C H

H



H



H



O

HC

HC



C

C

C

H



OH

C



O



O CCH3



CH



O



OH

O

O



Aspirin



4.2



P O L A R C O V A L E N T B O N D S , S H A P E , A N D P O L A R I T Y



We have described a covalent bond as consisting of a shared pair of valence electrons.

This is not the complete story, however, because the electrons in a covalent bond are not

always shared equally between the two atoms. An unequal sharing of electrons is attributed to differences in electronegativity, the ability of an atom to attract bonding electrons.



124   Chapter 4  An Introduction to Organic Compounds

1

H

2.1

3

Li

1.0

11

Na

0.9



5

B

2.0



4

Be

1.5

12

Mg

1.2

22

Ti

1.4



23

V

1.6



19

K

0.8



20

Ca

1.0



21

Sc

1.3



37

Rb

0.8



38

Sr

1.0



39

Y

1.2



40

Zr

1.4



41

Nb

1.6



55

Cs

0.7



56

Ba

0.9



57–71

La–Lu

1.1–1.2



72

Hf

1.3



73

Ta

1.5



87

Fr

0.7



88

Ra

0.9



89–103

Ac–Lr

1.1–1.7



42

Mo

1.8



43

Tc

1.9



26

Fe

1.8

44

Ru

2.2



74

W

1.7



75

Re

1.9



76

Os

2.2



24

Cr

1.6



25

Mn

1.5



27

Co

1.8



28

Ni

1.8



45

Rh

2.2



46

Pd

2.2



77

Ir

2.2



78

Pt

2.2



29

Cu

1.9

47

Ag

1.9

79

Au

2.4



30

Zn

1.6

48

Cd

1.7

80

Hg

1.9



6

C

2.5



13

Al

1.5



14

Si

1.8



31

Ga

1.6



32

Ge

1.8



7

N

3.0

15

P

2.1

33

As

2.0



9

F

4.0



8

O

3.5

16

S

2.5

34

Se

2.4



49

In

1.7



50

Sn

1.8



51

Sb

1.9



52

Te

2.1



81

Tl

1.8



82

Pb

1.8



83

Bi

1.9



84

Po

2.0



17

Cl

3.0

35

Br

2.8

53

I

2.5

85

At

2.2



He



Ne



Ar



Kr



Xe



Rn



■ ■ FIGURE 4.3

Electronegativity  Electronegativity reflects the ability of an atom to attract bonding electrons. Fluorine

(F) is the most electronegative atom, and cesium (Cs) and francium (Fr) are the least electronegative. Inert

gases do not attract bonding electrons and are not assigned an electronegativity value.

�−



�+



H ¬F



�−



�+



�+



H¬O ¬H

H



ƒ �+ �−



H ¬ C ¬ Cl

ƒ



H

H



ƒ �+ �− �+



H¬ C ¬ N ¬H

ƒ



ƒ



H H �+

■ ■ FIGURE 4.4

Polar covalent bonds 

Electrons are shared unequally

in polar covalent bonds, which

sets up partial positive (d+) and

partial negative (d-) charges.



As shown in Figure 4.3, electronegativity displays a periodic trend—moving to the right

across a period or up a group, electronegativity generally increases, with fluorine atoms

being the most electronegative.

When the electrons in covalent bonds are shared by atoms with different electronegativities, a polar covalent bond results. In a polar covalent bond, the unequal sharing of

electrons gives the bond a partially positive end and a partially negative end. Fluorine is much

more electronegative than hydrogen, so when a fluorine atom and a hydrogen atom combine to form a covalent bond, the fluorine atom pulls the shared pair of bonding electrons

toward itself. As a result, the fluorine atom carries a partial negative charge, represented by

d- (delta minus), the hydrogen atom carries a partial positive charge (d+), and the bond

is polar (Figure 4.4).

The range of bond types that we have seen in this and the previous chapter can be

explained in terms of the electronegativity differences between the atoms involved (Table

4.1). Combining nonmetal atoms of identical electronegativity produces nonpolar covalent bonds in which electrons are shared equally. When the electronegativity difference

between a pair of atoms involved in a covalent bond is small (less than 0.5), the level of

polarity is so small that the bond behaves as if it were nonpolar. A greater difference in

electronegativity (0.5–1.9) produces a polar covalent bond. Electronegativity differences

of greater than 1.9 usually indicate an ionic bond, in which electrons have been completely transferred from the less electronegative atom (a metal) to a more electronegative

one (a nonmetal).



Table | 4.1   B OND TYPES





Bond



Nonpolar Covalent



Polar Covalent



Ionic





Characteristics





Equal sharing of

bonding electrons



Unequal sharing of

bonding electrons



Attraction of opposite

charges







Example



C ¬ C







Electronegativity difference



Less than 0.5



�+



�−



C¬O



0.5–1.9



Na+



Cl-



Greater than 1.9



4.2  Polar Covalent Bonds, Shape, and Polarity   125



You may not always have access to electronegativity values. For the compounds that

we are most likely to encounter in organic and biochemistry, the important polar covalent

bonds are those in which either hydrogen or carbon atoms are covalently attached to nitrogen,

oxygen, fluorine, or chlorine atoms.

SAMPLE PROBLEM   4.4



Identifying polar covalent bonds

The compound below has been used to kill any insect larvae present in cereal and dried

fruit. Label the polar covalent bonds, using the symbols d+ and d-.

≠O

H





ƒ

H i Ci N

¶ i C iH

ƒ

ƒ

H H



Strategy



A bond is polar covalent if the electronegativity difference between the two atoms forming

the bond ranges between 0.5 and 1.9. Alternatively, we have defined polar covalent bonds

as those between H or C atoms and N, O, F, or Cl atoms. In a polar covalent bond, the d+

belongs with the less electronegative atom (H or C) and the d- with the more electronegative one (N, O, F, or Cl).

Solution

�−



O



H



‘ �+ �− �+ ƒ



H¬ C ¬ N ¬ C ¬ H

�+



ƒ



H



ƒ



H



PRACTICE PROBLEM   4.4



Label the polar covalent bonds in aspirin (see Figure 4.2), using the symbols d+ and d-.



Electron dot, Lewis, condensed, and skeletal structures are designed to show how the

atoms in a compound are attached to one another but are not intended to necessarily

represent its actual shape. Shape can be predicted, however, based on these structural

formulas.

Consider the way that four groups of electrons arrange themselves around an atom.

The carbon atom in methane (Table 4.2) has four single bonds, one to each of the four

hydrogen atoms. The best way for these four groups of electrons to be as far apart from

one another as possible (electrons all have the same 1 - charge and similar charges repel)

is for them to point to the four corners of a tetrahedron. This results in a tetrahedral

shape, in which the hydrogen atoms of methane are placed at each of the four corners of

the same tetrahedron.

In ammonia, the nitrogen atom also has four groups of electrons surrounding it—three

single bonds and one pair of nonbonding electrons. As with methane, the four groups

of electrons arrange themselves to point to the four corners of a tetrahedron (Table 4.2).

With the three hydrogen atoms of ammonia sitting at three corners of the tetrahedron,

the molecule has a pyramidal shape. Using a similar argument, the oxygen atom in water

has four groups of electrons (two single bonds and two nonbonding pairs) and the molecule has a bent shape.

If you were to build a model of a molecule that had a tetrahedral shape, the angle

between the bonds would all be the same at 109.5° (Figure 4.5). This is the bond angle

in methane (Table 4.2). Bond angles will not always be 109.5° in molecules whose shape

is based on a tetrahedron. This is because nonbonding pairs of electrons repel electrons



n



Bonds between carbon or

hydrogen atoms and nitrogen,

oxygen, fluorine, or chlorine

atoms are polar covalent.



126   Chapter 4  An Introduction to Organic Compounds

Table | 4.2   C ommon Molecular Shapes











Number of Groups

of Electrons a

Shape b

Example



Lewis

Structure



Three-dimensional

View c



Tetrahedral



4



4



H





Methane



3



H¬ C ¬H

ƒ



H



1









2



A tetrahedron









H



ƒ



H



C



H

H 109.5°







Pyramidal

(nonbonding electron pair)

4



3





2











1



Ammonia



H ¬ N ¬ N

ƒ



H







H



N



H

H 107.5°



Bent

(nonbonding electron pairs)

4



3







Water



H¬O¬H



1



H



H



104.5°



2









O



Trigonal planar



3



3



O



O







A triangle

2







Formaldehyde







H¬C¬ H



H



C

118°



H



1



Bent

(nonbonding electron pair)

3









2



Nitrite ion

1



[O



-



N¬ O ]



O



N

115°



-



O



4.2  Polar Covalent Bonds, Shape, and Polarity   127







Number of Groups

of Electrons a

Shape b

Example



Lewis

Structure



2

Linear

O C

——————



Carbon

A line

dioxide

1

2











Three-dimensional

View c



O



O C O

180°



a



 he number of groups of electrons around an atom is the sum of the attached atoms (the single, double, or triple bond to an atom counts as

T

one group of electrons) and attached pairs of nonbonding electrons.

b

Molecular shape is based on the relative placement of atoms. Nonbonding electron pairs are ignored.

c

Wedge and dashed line notation can be used to show three-dimensional shapes. A solid wedge points out of the plane of the paper and a dashed

wedge points in to the paper.



Did You

Know



■ ■ FIGURE 4.5



?



14.11



From a historical point of

view, our understanding

of organic chemistry is

not all that old. Electron

dot structures were introduced in 1916 and the

theory used to predict

the shapes of molecules

was first described

in 1957.



Ken Karp for John Wiley & Sons, Inc.



Shapes  (a) For a tetrahedral arrangement groups around an atom, the angle between any two

bonds is 109.5°. (b) A trigonal planar arrangement has bond angles of 120° and (c) a linear

arrangement has bond angles of 180°.



in single bonds and tend to compress bond angles. Instead of the predicted 109.5° bond

angles, those of ammonia and water are 107.5° (H i N i H) and 104.5° (H i O i H),

respectively.

When three groups of electrons surround an atom, the two common shapes observed

are trigonal planar and bent. The carbon atom in formaldehyde has three groups of electrons (two single bonds and one double bond) and is trigonal planar. The nitrogen atom

in nitrite ion has three groups of electrons (one single bond, one double bond, and one

pair of nonbonding electrons) and is bent. The model of a molecule with a trigonal planar

shape has bond angles of 120° (Figure 4.5), and most real molecules with three attached

groups of electrons have bond angles near this value. In formaldehyde (Table 4.2) the

H i C i H bond angle is reduced to 118° by the nonbonding electrons on the oxygen

atom, and the O i N i O bond angle in nitrite is reduced to 115° due to the effects

of the pair of nonbonding electrons on the N atom. Two groups of electrons assume a

linear shape, as seen for carbon dioxide, and have a bond angle of 180°. When many

atoms are present in a molecule it can assume a very complex shape (Figure 4.6).



Polarity

Now that we have seen how to identify polar covalent bonds and to determine the shape

about an atom, we can see how these two characteristics determine if a molecule, as a

whole, is polar. In a polar molecule, one side has a partial positive charge and the other has a

partial negative charge. Knowing whether or not a molecule is polar helps us to understand

how it interacts with other compounds, as we will see in the next section.

If a molecule has no polar covalent bonds, it is nonpolar. The molecule F2, for example,

is nonpolar because it contains only one covalent bond—a nonpolar one (Figure 4.8a).



n



In a polar molecule, one side is

partially positive and the other is

partially negative.



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2 Polar Covalent Bonds, Shape, and Polarity

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