how are atoms held together what are the three characteristics of atoms that contribute to bonding

Affiliate 7. Chemic Bonding and Molecular Geometry

vii.ii Covalent Bonding

Learning Objectives

By the cease of this section, y'all will be able to:

• Describe the formation of covalent bonds
• Ascertain electronegativity and assess the polarity of covalent bonds

In ionic compounds, electrons are transferred betwixt atoms of different elements to grade ions. Merely this is not the only fashion that compounds can be formed. Atoms can also make chemical bonds by sharing electrons equally betwixt each other. Such bonds are called covalent bonds. Covalent bonds are formed between two atoms when both accept similar tendencies to attract electrons to themselves (i.e., when both atoms have identical or fairly similar ionization energies and electron affinities). For example, 2 hydrogen atoms bail covalently to grade an H₂ molecule; each hydrogen atom in the H₂ molecule has two electrons stabilizing it, giving each atom the same number of valence electrons as the noble gas He.

Compounds that contain covalent bonds exhibit dissimilar concrete properties than ionic compounds. Considering the attraction between molecules, which are electrically neutral, is weaker than that between electrically charged ions, covalent compounds generally have much lower melting and boiling points than ionic compounds. In fact, many covalent compounds are liquids or gases at room temperature, and, in their solid states, they are typically much softer than ionic solids. Furthermore, whereas ionic compounds are expert conductors of electricity when dissolved in water, most covalent compounds are insoluble in water; since they are electrically neutral, they are poor conductors of electricity in any state.

Formation of Covalent Bonds

Nonmetal atoms often class covalent bonds with other nonmetal atoms. For example, the hydrogen molecule, H_ii, contains a covalent bond betwixt its 2 hydrogen atoms. Figure i illustrates why this bond is formed. Starting on the far right, we have two separate hydrogen atoms with a particular potential energy, indicated by the ruby line. Along the x-axis is the distance between the two atoms. As the two atoms arroyo each other (moving left along the x-centrality), their valence orbitals (1southward) begin to overlap. The single electrons on each hydrogen atom and so interact with both atomic nuclei, occupying the infinite around both atoms. The strong allure of each shared electron to both nuclei stabilizes the system, and the potential energy decreases as the bond distance decreases. If the atoms continue to arroyo each other, the positive charges in the ii nuclei brainstorm to repel each other, and the potential energy increases. The bail length is determined by the altitude at which the lowest potential energy is accomplished.

Effigy 1. The potential energy of two split hydrogen atoms (correct) decreases as they approach each other, and the single electrons on each atom are shared to class a covalent bail. The bail length is the internuclear distance at which the everyman potential energy is achieved.

It is essential to think that free energy must exist added to break chemical bonds (an endothermic process), whereas forming chemical bonds releases energy (an exothermic process). In the instance of H₂, the covalent bond is very potent; a large corporeality of energy, 436 kJ, must exist added to break the bonds in one mole of hydrogen molecules and cause the atoms to separate:

[latex]\text{H}_2(thousand) \longrightarrow 2\text{H}(chiliad) \;\;\;\;\; \Delta H = 436\;\text{kJ}[/latex]

Conversely, the same corporeality of free energy is released when one mole of H₂ molecules forms from 2 moles of H atoms:

[latex]ii\text{H}(g) \longrightarrow \text{H}_2(g) \;\;\;\;\; \Delta H = -436 \;\text{kJ}[/latex]

Pure vs. Polar Covalent Bonds

If the atoms that grade a covalent bond are identical, equally in H₂, Cl_two, and other diatomic molecules, then the electrons in the bond must be shared every bit. We refer to this equally a pure covalent bail. Electrons shared in pure covalent bonds have an equal probability of beingness near each nucleus.

In the case of Cl_ii, each cantlet starts off with seven valence electrons, and each Cl shares one electron with the other, forming one covalent bond:

[latex]\text{Cl} + \text{Cl} \longrightarrow \text{Cl}_2[/latex]

The full number of electrons around each private atom consists of six nonbonding electrons and two shared (i.e., bonding) electrons for eight total electrons, matching the number of valence electrons in the noble gas argon. Since the bonding atoms are identical, Cl₂ also features a pure covalent bond.

When the atoms linked by a covalent bond are different, the bonding electrons are shared, just no longer equally. Instead, the bonding electrons are more attracted to 1 atom than the other, giving rising to a shift of electron density toward that atom. This unequal distribution of electrons is known as a polar covalent bond, characterized by a partial positive charge on i atom and a partial negative charge on the other. The atom that attracts the electrons more strongly acquires the partial negative accuse and vice versa. For example, the electrons in the H–Cl bond of a hydrogen chloride molecule spend more time near the chlorine atom than near the hydrogen cantlet. Thus, in an HCl molecule, the chlorine atom carries a partial negative charge and the hydrogen atom has a partial positive accuse. Figure 2 shows the distribution of electrons in the H–Cl bail. Note that the shaded area around Cl is much larger than it is around H. Compare this to Figure one, which shows the fifty-fifty distribution of electrons in the H_two nonpolar bond.

We sometimes designate the positive and negative atoms in a polar covalent bail using a lowercase Greek letter "delta," δ, with a plus sign or minus sign to indicate whether the cantlet has a fractional positive charge (δ+) or a partial negative charge (δ–). This symbolism is shown for the H–Cl molecule in Effigy two.

Effigy 2. (a) The distribution of electron density in the HCl molecule is uneven. The electron density is greater around the chlorine nucleus. The small, black dots point the location of the hydrogen and chlorine nuclei in the molecule. (b) Symbols δ+ and δ– indicate the polarity of the H–Cl bond.

Electronegativity

Whether a bond is nonpolar or polar covalent is determined past a property of the bonding atoms called electronegativity. Electronegativity is a measure of the trend of an atom to concenter electrons (or electron density) towards itself. It determines how the shared electrons are distributed between the 2 atoms in a bond. The more strongly an atom attracts the electrons in its bonds, the larger its electronegativity. Electrons in a polar covalent bail are shifted toward the more electronegative atom; thus, the more electronegative atom is the one with the fractional negative charge. The greater the divergence in electronegativity, the more polarized the electron distribution and the larger the partial charges of the atoms.

Figure 3 shows the electronegativity values of the elements equally proposed past ane of the almost famous chemists of the twentieth century: Linus Pauling (Effigy 4). In general, electronegativity increases from left to right across a menstruation in the periodic table and decreases down a group. Thus, the nonmetals, which lie in the upper right, tend to have the highest electronegativities, with fluorine the about electronegative chemical element of all (EN = 4.0). Metals tend to be less electronegative elements, and the group 1 metals have the lowest electronegativities. Note that noble gases are excluded from this figure because these atoms usually practice not share electrons with others atoms since they take a full valence shell. (While element of group 0 compounds such as XeO₂ do exist, they can only exist formed nether extreme conditions, and thus they do not fit neatly into the general model of electronegativity.)

Effigy 3. The electronegativity values derived by Pauling follow predictable periodic trends with the higher electronegativities toward the upper right of the periodic table.

Electronegativity versus Electron Analogousness

Nosotros must exist careful not to confuse electronegativity and electron affinity. The electron affinity of an element is a measurable physical quantity, namely, the energy released or absorbed when an isolated gas-phase cantlet acquires an electron, measured in kJ/mol. Electronegativity, on the other hand, describes how tightly an atom attracts electrons in a bond. Information technology is a dimensionless quantity that is calculated, not measured. Pauling derived the get-go electronegativity values by comparing the amounts of energy required to pause different types of bonds. He chose an arbitrary relative scale ranging from 0 to iv.

Linus Pauling

Linus Pauling, shown in Figure 4, is the only person to have received two unshared (individual) Nobel Prizes: one for chemistry in 1954 for his work on the nature of chemical bonds and one for peace in 1962 for his opposition to weapons of mass devastation. He developed many of the theories and concepts that are foundational to our electric current understanding of chemistry, including electronegativity and resonance structures.

Effigy 4. Linus Pauling (1901–1994) made many of import contributions to the field of chemical science. He was as well a prominent activist, publicizing bug related to health and nuclear weapons.

Pauling also contributed to many other fields besides chemistry. His inquiry on sickle jail cell anemia revealed the crusade of the disease—the presence of a genetically inherited abnormal protein in the blood—and paved the mode for the field of molecular genetics. His piece of work was likewise pivotal in curbing the testing of nuclear weapons; he proved that radioactive fallout from nuclear testing posed a public health risk.

Electronegativity and Bail Type

The absolute value of the difference in electronegativity (ΔEN) of two bonded atoms provides a crude measure of the polarity to be expected in the bond and, thus, the bond type. When the difference is very small-scale or zero, the bond is covalent and nonpolar. When it is large, the bond is polar covalent or ionic. The absolute values of the electronegativity differences betwixt the atoms in the bonds H–H, H–Cl, and Na–Cl are 0 (nonpolar), 0.9 (polar covalent), and 2.ane (ionic), respectively. The degree to which electrons are shared between atoms varies from completely equal (pure covalent bonding) to not at all (ionic bonding). Figure 5 shows the relationship between electronegativity departure and bail type.

Figure 5. As the electronegativity difference increases betwixt two atoms, the bond becomes more ionic.

A crude approximation of the electronegativity differences associated with covalent, polar covalent, and ionic bonds is shown in Figure 5. This tabular array is just a general guide, however, with many exceptions. For case, the H and F atoms in HF have an electronegativity difference of one.nine, and the N and H atoms in NH₃ a departure of 0.9, yet both of these compounds form bonds that are considered polar covalent. Likewise, the Na and Cl atoms in NaCl accept an electronegativity difference of 2.1, and the Mn and I atoms in MnI_two take a difference of ane.0, yet both of these substances form ionic compounds.

The best guide to the covalent or ionic grapheme of a bail is to consider the types of atoms involved and their relative positions in the periodic tabular array. Bonds between ii nonmetals are generally covalent; bonding between a metallic and a nonmetal is oft ionic.

Some compounds contain both covalent and ionic bonds. The atoms in polyatomic ions, such every bit OH^–, NO_three ⁻, and NH₄ ⁺, are held together by polar covalent bonds. Yet, these polyatomic ions course ionic compounds by combining with ions of contrary charge. For example, potassium nitrate, KNO_three, contains the Grand⁺ cation and the polyatomic NO₃ ⁻ anion. Thus, bonding in potassium nitrate is ionic, resulting from the electrostatic attraction between the ions K⁺ and NO_iii ⁻, as well as covalent between the nitrogen and oxygen atoms in NO₃ ⁻.

Example 1

Electronegativity and Bond Polarity
Bond polarities play an important role in determining the structure of proteins. Using the electronegativity values in Figure 3, arrange the post-obit covalent bonds—all commonly found in amino acids—in order of increasing polarity. Then designate the positive and negative atoms using the symbols δ+ and δ–:

C–H, C–N, C–O, N–H, O–H, S–H

Solution
The polarity of these bonds increases as the absolute value of the electronegativity difference increases. The cantlet with the δ– designation is the more electronegative of the two. Table 1 shows these bonds in lodge of increasing polarity.

Bond	ΔEN	Polarity
C–H	0.iv	[latex]\overset{\delta -}{\text{C}} - \overset{\delta +}{\text{H}}[/latex]
Southward–H	0.4	[latex]\overset{\delta -}{\text{S}} - \overset{\delta +}{\text{H}}[/latex]
C–Northward	0.5	[latex]\overset{\delta +}{\text{C}} - \overset{\delta -}{\text{N}}[/latex]
N–H	0.9	[latex]\overset{\delta -}{\text{N}} - \overset{\delta +}{\text{H}}[/latex]
C–O	one.0	[latex]\overset{\delta +}{\text{C}} - \overset{\delta -}{\text{O}}[/latex]
O–H	1.four	[latex]\overset{\delta -}{\text{O}} - \overset{\delta +}{\text{H}}[/latex]
Table one. Bond Polarity and Electronegativity Difference

Check Your Learning
Silicones are polymeric compounds containing, amidst others, the following types of covalent bonds: Si–O, Si–C, C–H, and C–C. Using the electronegativity values in Figure 3, suit the bonds in club of increasing polarity and designate the positive and negative atoms using the symbols δ+ and δ–.

Answer:

Bond	Electronegativity Divergence	Polarity
C–C	0.0	nonpolar
C–H	0.4	[latex]\overset{\delta -}{\text{C}} - \overset{\delta +}{\text{H}}[/latex]
Si–C	0.7	[latex]\overset{\delta +}{\text{Si}} - \overset{\delta -}{\text{C}}[/latex]
Si–O	ane.seven	[latex]\overset{\delta +}{\text{Si}} - \overset{\delta -}{\text{O}}[/latex]
Tabular array 2.

Key Concepts and Summary

Covalent bonds course when electrons are shared between atoms and are attracted past the nuclei of both atoms. In pure covalent bonds, the electrons are shared equally. In polar covalent bonds, the electrons are shared unequally, as one atom exerts a stronger force of attraction on the electrons than the other. The ability of an atom to attract a pair of electrons in a chemical bond is called its electronegativity. The difference in electronegativity betwixt two atoms determines how polar a bond volition be. In a diatomic molecule with two identical atoms, there is no departure in electronegativity, and then the bond is nonpolar or pure covalent. When the electronegativity difference is very large, as is the case between metals and nonmetals, the bonding is characterized equally ionic.

Chemical science End of Chapter Exercises

1. Why is it incorrect to speak of a molecule of solid NaCl?
2. What information can y'all use to predict whether a bail between two atoms is covalent or ionic?
3. Predict which of the following compounds are ionic and which are covalent, based on the location of their constituent atoms in the periodic table:

   (a) Cl₂CO

   (b) MnO

   (c) NCl_iii

   (d) CoBr₂

   (e) G₂S

   (f) CO

   (g) CaF₂

   (h) HI

   (i) CaO

   (j) IBr

   (k) CO₂

4. Explain the deviation between a nonpolar covalent bond, a polar covalent bond, and an ionic bond.
5. From its position in the periodic table, determine which atom in each pair is more than electronegative:

   (a) Br or Cl

   (b) North or O

   (c) Due south or O

   (d) P or Due south

   (eastward) Si or N

   (f) Ba or P

   (thou) Northward or Chiliad

6. From its position in the periodic table, determine which atom in each pair is more electronegative:

   (a) N or P

   (b) Northward or Ge

   (c) S or F

   (d) Cl or S

   (e) H or C

   (f) Se or P

   (thousand) C or Si

7. From their positions in the periodic tabular array, adapt the atoms in each of the following series in society of increasing electronegativity:

   (a) C, F, H, North, O

   (b) Br, Cl, F, H, I

   (c) F, H, O, P, S

   (d) Al, H, Na, O, P

   (due east) Ba, H, N, O, Equally

8. From their positions in the periodic table, suit the atoms in each of the post-obit serial in order of increasing electronegativity:

   (a) As, H, North, P, Sb

   (b) Cl, H, P, S, Si

   (c) Br, Cl, Ge, H, Sr

   (d) Ca, H, 1000, Northward, Si

   (due east) Cl, Cs, Ge, H, Sr

9. Which atoms can bond to sulfur so as to produce a positive partial charge on the sulfur atom?
10. Which is the near polar bail?

    (a) C–C

    (b) C–H

    (c) N–H

    (d) O–H

    (e) Se–H

11. Identify the more polar bond in each of the following pairs of bonds:

    (a) HF or HCl

    (b) NO or CO

    (c) SH or OH

    (d) PCl or SCl

    (e) CH or NH

    (f) SO or PO

    (g) CN or NN

12. Which of the post-obit molecules or ions contain polar bonds?

    (a) O₃

    (b) S_viii

    (c) O22−O22−

    (d) NO3−NO3−

    (e) CO₂

    (f) H₂South

    (g) BH4−BH4−

Glossary

bail length

distance between the nuclei of two bonded atoms at which the everyman potential energy is achieved

covalent bond

bond formed when electrons are shared betwixt atoms

electronegativity

tendency of an atom to attract electrons in a bail to itself

polar covalent bond

covalent bail between atoms of different electronegativities; a covalent bond with a positive end and a negative end

pure covalent bond

(also, nonpolar covalent bond) covalent bond betwixt atoms of identical electronegativities

Solutions

Answers to Chemistry End of Chapter Exercises

1. NaCl consists of detached ions arranged in a crystal lattice, not covalently bonded molecules.

3. ionic: (b), (d), (e), (k), and (i); covalent: (a), (c), (f), (h), (j), and (k)

5. (a) Cl; (b) O; (c) O; (d) S; (eastward) N; (f) P; (g) N

7. (a) H, C, N, O, F; (b) H, I, Br, Cl, F; (c) H, P, S, O, F; (d) Na, Al, H, P, O; (e) Ba, H, As, N, O

ix. Northward, O, F, and Cl

11. (a) HF; (b) CO; (c) OH; (d) PCl; (e) NH; (f) PO; (g) CN

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