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# 1.3: Basics of bonding

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## Home1.3.1. How and why atoms bond

Atoms can join together by forming a chemical bond, which is a very strong attraction between two atoms. Chemical bonds are formed when electrons in different atoms interact with each other to make an arrangement that is more stable than when the atoms are apart.

What causes atoms to make a chemical bond with other atoms, rather than remaining as individual atoms? A clue comes by considering the noble gas elements, the rightmost column of the periodic table. These elements—helium, neon, argon, krypton, xenon, and radon—do not form compounds very easily, which suggests that they are especially stable as lone atoms. What else do the noble gas elements have in common? Except for helium, they all have eight valence electrons. Chemists have concluded that atoms are especially stable if they have eight electrons in their outermost shell. This useful rule of thumb is called the octet rule, and it is a key to understanding why compounds form.

There are two ways for an atom that does not have an octet of valence electrons to obtain an octet in its outer shell. One way is the transfer of electrons between two atoms until all atoms have octets. Because some atoms will lose electrons and some atoms will gain electrons, there is no overall change in the number of electrons, but individual atoms acquire a nonzero electric charge. Those that lose electrons become positively charged, and those that gain electrons become negatively charged. Charged atoms are called ions. Because opposite charges attract (while like charges repel), these oppositely charged ions attract each other, forming ionic bonds. The resulting compounds are called ionic compounds and are the primary subject of this chapter.

The second way for an atom to obtain an octet of electrons is by sharing electrons with another atom. These shared electrons simultaneously occupy the outermost shell of more than one atom. The bond made by electron sharing is called a covalent bond. Covalent bonding and covalent compounds will be discussed in detail below.

NOTE: Despite our focus on the octet rule, we must remember that for small atoms, such as hydrogen, helium, and lithium, the first shell is, or becomes, the outermost shell and hold only two electrons. Therefore, these atoms satisfy a “duet rule” rather than the octet rule.

### Learning Objective

1. Compare covalent bonds in terms of bond length and bond polarity.

Covalent bonds have certain characteristics that depend on the identities of the atoms participating in the bond. Two characteristics are bond length and bond polarity.

### Bond length

The covalent bond in the hydrogen molecule (H2) has a certain length (about 7.4 × 10−11 m). Other covalent bonds also have known bond lengths, which are dependent on both the identities of the atoms in the bond and whether the bonds are single, double, or triple bonds. Table 1.1. lists the approximate bond lengths for some single covalent bonds. The exact bond length may vary depending on the identity of the molecule but will be close to the value given in the table.

#### Table 1.1. Approximate bond lengths of some single bonds

Bond Length (× 10−12 m)
H–H 74
H–C 110
H–N 100
H–O 97
H–I 161
C–C 154
C–N 147
C–O 143
N–N 145
O–O 145

Table 1.2. compares the lengths of single covalent bonds with those of double and triple bonds between the same atoms. Without exception, as the number of covalent bonds between two atoms increases, the bond length decreases. With more electrons between the two nuclei, the nuclei can get closer together before the internuclear repulsion is strong enough to balance the attraction.

#### Table 1.2. Comparison of bond lengths for single and multiple bonds

Bond Length (× 10−12 m)
C–C 154
C=C 134
C≡C 120
C–N 147
C=N 128
C≡N 116
C–O 143
C=O 120
C≡O 113
N–N 145
N=N 123
N≡N 110
O–O 145
O=O 121

### Electronegativity and bond polarity

Figure 1.6. Polar versus nonpolar covalent bonds. (a) The electrons in the covalent bond are equally shared by both hydrogen atoms. This is a nonpolar covalent bond. (b) The fluorine atom attracts the electrons in the bond more than the hydrogen atom does, leading to an imbalance in the electron distribution. This is a polar covalent bond.

Although we defined covalent bonding as electron sharing, the electrons in a covalent bond are not always shared equally by the two bonded atoms. Unless the bond connects two atoms of the same element, there will always be one atom that attracts the electrons in the bond more strongly than the other atom does, as shown in Figure 1.6. “Polar versus Nonpolar Covalent Bonds”. When such an imbalance occurs, there is a resulting buildup of some negative charge (called a partial negative charge and designated δ−) on one side of the bond and some positive charge (designated δ+) on the other side of the bond. A covalent bond that has an unequal sharing of electrons, as in part (b) of Figure 1.6., is called a polar covalent bond. A covalent bond that has an equal sharing of electrons (part (a) is called a nonpolar covalent bond.

Any covalent bond between atoms of different elements is a polar bond, but the degree of polarity varies widely. Some bonds between different elements are only minimally polar, while others are strongly polar. Ionic bonds can be considered the ultimate in polarity, with electrons being transferred rather than shared. To judge the relative polarity of a covalent bond, chemists use electronegativity, which is a relative measure of how strongly an atom attracts electrons when it forms a covalent bond. There are various numerical scales for rating electronegativity. Figure 1.7. “Electronegativities of Various Elements” shows one of the most popular—the Pauling scale. The polarity of a covalent bond can be judged by determining the difference in the electronegativities of the two atoms making the bond. The greater the difference in electronegativities, the greater the imbalance of electron sharing in the bond. Although there are no hard and fast rules, the general rule is if the difference in electronegativities is less than about 0.4, the bond is considered nonpolar; if the difference is greater than 0.4, the bond is considered polar. If the difference in electronegativities is large enough (generally greater than about 1.8), the resulting compound is considered ionic rather than covalent. An electronegativity difference of zero, of course, indicates a nonpolar covalent bond.

When a molecule’s bonds are polar, the substance is usually found to be polar. The polarity of water has an enormous impact on its physical and chemical properties. (For example, the boiling point of water [100°C] is high for such a small molecule and is due to the fact that polar molecules attract each other strongly.)  This aspect will be discussed in more detail in section 1.8. on intermolecular forces.

Figure 1.7. Electronegativities of Various Elements. A popular scale for electronegativities has the value for fluorine atoms set at 4.0, the highest value.

### Looking Closer: Linus Pauling

Arguably the most influential chemist of the 20th century, Linus Pauling (1901–94) is the only person to have won two individual (that is, unshared) Nobel Prizes. In the 1930s, Pauling used new mathematical theories to enunciate some fundamental principles of the chemical bond. His 1939 book The Nature of the Chemical Bond is one of the most significant books ever published in chemistry.

By 1935, Pauling’s interest turned to biological molecules, and he was awarded the 1954 Nobel Prize in Chemistry for his work on protein structure. (He was very close to discovering the double helix structure of DNA when James Watson and James Crick announced their own discovery of its structure in 1953.) He was later awarded the 1962 Nobel Peace Prize for his efforts to ban the testing of nuclear weapons.

In his later years, Pauling became convinced that large doses of vitamin C would prevent disease, including the common cold. Most clinical research failed to show a connection, but Pauling continued to take large doses daily. He died in 1994, having spent a lifetime establishing a scientific legacy that few will ever equal.

Linus Pauling was one of the most influential chemists of the 20th century.

### Example 2

Describe the electronegativity difference between each pair of atoms and the resulting polarity (or bond type).

1. C and H
2. H and H
3. Na and Cl
4. O and H

Solution

[reveal-answer q=”979132″]Show Answer[/reveal-answer]
[hidden-answer a=”979132″]

1. Carbon has an electronegativity of 2.5, while the value for hydrogen is 2.1. The difference is 0.3, which is rather small. The C–H bond is therefore considered nonpolar.
2. Both hydrogen atoms have the same electronegativity value—2.1. The difference is zero, so the bond is nonpolar.
3. Sodium’s electronegativity is 0.9, while chlorine’s is 3.0. The difference is 2.1, which is rather high, and so sodium and chlorine form an ionic compound.
4. With 2.1 for hydrogen and 3.5 for oxygen, the electronegativity difference is 1.4. We would expect a very polar bond, but not so polar that the O–H bond is considered ionic.[/hidden-answer]

1. C and O

2. K and Br

3. N and N

4. Cs and F

### Concept Review Exercises

1. What is the name for the distance between two atoms in a covalent bond?

2. What does the electronegativity of an atom indicate?

3. What type of bond is formed between two atoms if the difference in electronegativities is small? Medium? Large?

### Answers

[reveal-answer q=”557987″]Show Answer[/reveal-answer]
[hidden-answer a=”557987″]

1. bond length
2. Electronegativity is a qualitative measure of how much an atom attracts electrons in a covalent bond.
3. nonpolar; polar; ionic[/hidden-answer]

### Key Takeaways

• Covalent bonds between different atoms have different bond lengths.
• Covalent bonds can be polar or nonpolar, depending on the electronegativity difference between the atoms involved.

### Exercises

1. Which is longer—a C–H bond or a C–O bond? (Refer to Table 1.1. “Approximate Bond Lengths of Some Single Bonds”.)

2. Which is shorter—an N–H bond or a C–H bond? (Refer to Table 1.1. “Approximate Bond Lengths of Some Single Bonds”.)

3. A nanometer is 10−9 m. Using the data in Table 1.1. “Approximate Bond Lengths of Some Single Bonds” and Table 1.2. “Comparison of Bond Lengths for Single and Multiple Bonds”, determine the length of each bond in nanometers.

1. a C–O bond
2. a C=O bond
3. an H–N bond
4. a C≡N bond
4. An angstrom (Å) is defined as 10−10 m. Using Table 1.1. “Approximate Bond Lengths of Some Single Bonds” and Table 1.2. “Comparison of Bond Lengths for Single and Multiple Bonds”, determine the length of each bond in angstroms.

1. a C–C bond
2. a C=C bond
3. an N≡N bond
4. an H–O bond
5. Refer to Exercise 3. Why is the nanometer unit useful as a unit for expressing bond lengths?

6. Refer to Exercise 4. Why is the angstrom unit useful as a unit for expressing bond lengths?

7. Using Figure 1.7. “Electronegativities of Various Elements”, determine which atom in each pair has the higher electronegativity.

1. H or C
2. O or Br
3. Na or Rb
4. I or Cl
8. Using Figure 1.7. “Electronegativities of Various Elements”, determine which atom in each pair has the lower electronegativity.

1. Mg or O
2. S or F
3. Al or Ga
4. O or I
9. Will the electrons be shared equally or unequally across each covalent bond? If unequally, to which atom are the electrons more strongly drawn?

1. a C–O bond
2. an F–F bond
3. an S–N bond
4. an I–Cl bond
10. Will the electrons be shared equally or unequally across each covalent bond? If unequally, to which atom are the electrons more strongly drawn?

1. a C–C bond
2. a S–Cl bond
3. an O–H bond
4. an H–H bond

### Answers

[reveal-answer q=”165622″]Show Answer[/reveal-answer]
[hidden-answer a=”165622″]

1. A C–O bond is longer.

3. a. 0.143 nm b. 0.120 nm c. 0.100 nm d. 0.116 nm

5. Actual bond lengths are very small, so the nanometer unit makes the expression of length easier to understand.

7. a. C b. O c. Na d. Cl

9. a. unequally toward the O b. equally c. unequally toward the N d. unequally toward the Cl[/hidden-answer]

### Further REading

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