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8.9: Covalent Bond Properties: Order, Length, and Energy

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  • Introduction

    The covalent bond involves the electrostatic interactions of valence electrons and multiple nuclei of the atoms that form the bond. There are both attractive forces between opposite charges, and repulsive forces between like charges, and these can be described by Coulomb's law:

    \[ E= k\dfrac{q_1q_2}{(4\pi\epsilon_0) r} \; \; \; \; \; \; \; \; F=k \dfrac{q_1q_2}{(4\pi\epsilon_0) r^2}\]

    • E is the bond energy
    • F is the force
    • r is the distance of separation between charged particles (note that the potential goes to zero when they are separated by infinity)
    • q1, q2 are the charges of the respective species interacting

    The potential well diagram of a homonuclear diatomic is a good way to understand a covalent bond as illustrated for hydrogen in Figure 8.9.1. It needs to be understood that a bond is dynamic, with the atoms oscillating around an equilibrium distance, which we call the bond length. This oscillation can be understood by looking at Figure \(\PageIndex{1}\). Region (1) shows the atoms completely separated. As they come closer together the electrons of each atom start being attracted to the nuclei of the other atom as in region (2), which lowers the energy and causes a bond. As the atoms keep moving closer the nuculear-nuclear repulsion start to become significant, and in region (4) the energy increases and the atoms move apart. Region (3) is the equilibrium bond length which represents the bottom of the potential well and is the length at which the molecules oscillate back and forth around.

    Figure \(\PageIndex{1}\): Potential Well Diagram for H2, where the energy of the H-H bond is plotted as a function of the internuclear distance (radius)

    From eq. 8.9.1 we note that when the atoms are separated by infinity (r=infinity) the energy is being divided by infinity, which means it is infinitesimally small and thus equals zero. This means the bond itself has a negative energy and the bond dissociation energy is the endothermic energy you must add to break a bond. Note, this also indicates why opposite charges attract, as that lowers the energy of the bonded particles when compared to the separated particles. It is important that you remember that any bond can be written as a "formation" (exothermic) or dissociation (endothermic) process.

    Generic Bond Dissociation

    If we look at a generic covalent bond as two electrons being represented as two shared electrons, we see there are two basic ways to break the bond, homolytically and heterolytically. In the former case, each of the separated species has one electron, making them free radicals (species with a lone electron), while the second forms ions, with one species having both electrons and the other having none.

    Homolytic bond cleavage
    \[A-B \rightarrow A\cdot + B\cdot \]

    Heterolytic bond cleavage
    \[A-B \rightarrow A^+ + B:^- \\ or \\ A-B \rightarrow A:^- + B^+ ​\]

    The potential well diagram in Figure 8.9.1 represents homolytic bond cleavage, which is the type we will discuss in this section. Remember, bond formation is the opposite of dissociation, and so can be visualized as the overlap of two orbitals on two different atoms that contribute one electron each (as indicated in fig. 8.9.1). There is another type of bond called the dative coordinate covalent bond, where one atom provides both electrons that are shared through the overlap with a vacant orbital on another atom.

    Bond Order

    The bond order is the number of covalent bonds in a chemical bond and the concept was introduced when we covered resonance structures.

    Table 4.6.1 Bond Lengths and Bond Dissociation Energies for Bonds with Different Bond Orders in Selected Gas-Phase Molecules at 298 K

    Compound Bond Order Bond Length (pm) Bond Dissociation Energy (kJ/mol) Compound Bond Order Bond Length (pm) Bond Dissociation Energy (kJ/mol)
    H3C–CH3 1 153.5 376 H3C–NH2 1 147.1 331
    H2C=CH2 2 133.9 728 H2C=NH 2 127.3 644
    HC≡CH 3 120.3 965 HC≡N 3 115.3 937
    H2N–NH2 1 144.9 275.3 H3C–OH 1 142.5 377
    HN=NH 2 125.2 456 H2C=O 2 120.8 732
    N≡N 3 109.8 945.3 O=C=O 2 116.0 799
    HO–OH 1 147.5 213 C≡O 3 112.8 1076.5
    O=O 2 120.7 498.4

    Bond Length

    A bond is not static but dynamic with the atoms undergoing the attractive and repulsive forces as described in the potential well (fig. 8.9.1), where the bottom of the well represents the equilibrium position of the oscillating atoms, which we call the bond length. Table 8.9.1 shows bond lengths for a variety of bonds. Please note that this particular table is in picometers (10-12), and that a common unit for bond length is the Angstrom Å (10-10), so for example, the O=O bond is 1.207 Å.

    Coulomb's law states that the bond energy is inversely related to the bond length (r), and so factors which influence a bond's strength influence its length. This can allow us to determine some trends in bond lengths.

    Some trends in bond lengths for homologous series (where one atom is constant and another is changed)

    • Bond length increases going down the periodic table. (smaller atoms make smaller bonds)
      H-F < H-Cl < H-Br < H-I
    • Bond length decreases going across the periodic table. (smaller atoms make smaller bonds)
      C-C > C-N > C-O
    • For bonds between equivalent atoms, the greater bond order the shorter the bond length
      C\(\equiv \)C > C=C < C-C
      Thus the NO bond of Nitrate, (fig. 8.9.2) is greater than N-O and less than N=O

    Bond Dissociation Enthalpy

    The IUPAC definition for bond dissociation energy, D, is "the enthalpy (per mole) required to break a given bond of some specific molecular entity by homolysis, e.g. for \(CH_{4}\rightarrow \cdot CH_{3)} + H\cdot\), (IUPAC Gold Book) and the "mean bond energy" is "the average value of the gas phase bond dissociation energies (usually at a temperature of 298K) for all bonds of the same type within the same chemical species.

    Reaction D (kJ/mol)
    CH4(g) → CH3(g) + H(g) 439
    CH3(g) → CH2(g) + H(g) 462
    CH2(g) → CH(g) + H(g) 424
    CH(g) → C(g) + H(g) 338
    Table \(\PageIndex{2}\) Energies for the successive breakup of C-H bonds in methane Source: Data from CRC Handbook of Chemstry and Physics (2004)K.

    Average bond energies can be tabulated as in table \(\PageIndex{3}\) and it needs to be emphasized that these values are average energies. One uese of these tables is that you can determine the enthalpy of reaction if you know the structure of all species in a balanced chemical equation.

    \[ ΔH_{rxn} \approx \sum{\text{(bond energies of bonds broken)}}−\sum{\text{(bond energies of bonds formed)}}\]

    Single Bonds Multiple Bonds
    H–H 432 C–C 346 N–N ≈167 O–O ≈142 F–F 155 C=C 602
    H–C 411 C–Si 318 N–O 201 O–F 190 F–Cl 249 C≡C 835
    H–Si 318 C–N 305 N–F 283 O–Cl 218 F–Br 249 C=N 615
    H–N 386 C–O 358 N–Cl 313 O–Br 201 F–I 278 C≡N 887
    H–P ≈322 C–S 272 N–Br 243 O–I 201 Cl–Cl 240 C=O 749
    H–O 459 C–F 485 P–P 201 S–S 226 Cl–Br 216 C≡O 1072
    H–S 363 C–Cl 327 S–F 284 Cl–I 208 N=N 418
    H–F 565 C–Br 285 S–Cl 255 Br–Br 190 N≡N 942
    H–Cl 428 C–I 213 S–Br 218 Br–I 175 N=O 607
    H–Br 362 Si–Si 222 I–I 149 O=O 494
    H–I 295 Si–O 452 S=O 532

    Table \(\PageIndex{3}\): Average Bond Energies (kJ/mol) for Commonly Encountered Bonds at 273 K. Source: Data from J. E. Huheey, E. A. Keiter, and R. L. Keiter, Inorganic Chemistry, 4th ed. (1993).

    Trends in Bond Dissociation Energy

    There are 3 trends we can observe with bond energies. You should be able to relate these trends to the bond length trends (above).

    1. Bonds with hydrogen (H-X) decrease in strength as we go down a column of the periodic table.

    H-F > H-Cl > H-Br > H-I

    1. Bonds between like atoms generally become weaker going down a column of the periodic table

    Cl-Cl > Br-Br > F-F > I-I

    Figure \(\PageIndex{1}\) shows trends (1) & (2) as being results of the trends in atomic radii (as you move down a period atoms take up more volume), and since the bond energy from Coulomb's law is inversely proportional to the distance (1/r) between the atoms, the smaller atoms form stronger bonds. By being closer, a greater percent of the orbitals overlap and this results in stronger bonding interactions.

    1. Bond Strength increases as Bond Order Increases

    C≡C < C=C < C-C

    Enthalpies of Reaction from Bond Dissociation Energies

    Since enthalpy is a state function (path independent), we can choose a reaction path where we break all the bonds of the reactants and then form the reactants by creating new bonds. The first step is the summation of their Bond Dissociation Energies (D) for the reactants. The second step is the formation of the products, which is the opposite of their dissociation, so we subtract the sum of their dissociation records.

    \( ΔH_{rxn} \approx \sum{\text{(bond energies of bonds broken)}}−\sum{\text{(bond energies of bonds formed)}}\)

    Why is subtracting the bond dissociation energies the same as their formation energies?

    First Law of Thermodynamics states energy of the universe is conserved, and so if a process is endothermic in one direction (bond dissociation), it must be exothermic by the same quantity in the opposite direction (bond formation).

    There is sort of a three step process, and you must draw Lewis dot structures in order to identify the bonds.

    1. Balance equation
    2. write Lewis dot structures of all species so you can identify bonds
    3. Apply the equation:
      \( ΔH_{rxn} \approx \sum{\text{(bond energies of bonds broken)}}−\sum{\text{(bond energies of bonds formed)}}\)

    Bond Dissociation Problem:

    Using bond dissociation energies calculate the molar enthalpy of combustion for acetic acid.

    Video \(\PageIndex{1}\): 7'49" YouTube going over bond dissociation energies

    Contributors and Attributions

    Robert E. Belford (University of Arkansas Little Rock; Department of Chemistry). The breadth, depth and veracity of this work is the responsibility of Robert E. Belford, You should contact him if you have any concerns. This material has both original contributions, and content built upon prior contributions of the LibreTexts Community and other resources, including but not limited to:

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