Unit 2 Compound Structure and Properties

AP Chemistry Covalent Bonds

What a covalent bond is, how polarity is decided, and how to read the potential energy curve behind bond length and bond energy, in four worked examples.

A covalent bond is the electrostatic attraction between two nuclei and the pair of electrons they share. Exam questions on it come in three kinds: classify a bond as polar or nonpolar and name the partial negative atom, read a potential energy graph for the bond length and bond energy, and compare bond strengths. Which atoms form covalent bonds in the first place is covered in ionic vs covalent bonds, and how bond polarity adds up to the polarity of a whole molecule is in molecular polarity. This page covers the bond itself.

What a Covalent Bond Is and How the Energy Curve Shows It

Nonpolar covalent: the two atoms attract the shared pair about equally, as in Cl–Cl, and approximately in C–H.
Polar covalent: one atom is more electronegative and pulls the pair closer, so it carries a partial negative charge (δ−) and the other a partial positive charge (δ+).
Single, double, triple: one, two, or three shared pairs, which is the bond order.

When two atoms approach, attraction between each nucleus and the shared electrons lowers the potential energy, while repulsion between the nuclei and between the electron clouds raises it sharply at short distance. The balance gives a minimum on a graph of potential energy against internuclear distance. The distance at the minimum is the bond length, and the depth of the minimum below the energy of the separated atoms is the bond energy. The graph shows the curves for three nitrogen–nitrogen bonds.

Fig. 1 · Potential energy against internuclear distance

Worked Example 1: Ranking Bonds by Polarity and Naming the Partial Negative Atom

Rank Cl–Cl, C–H, N–H, O–H, and H–F from least to most polar, and name the partial negative atom in each polar bond. Electronegativities: H 2.20, C 2.55, N 3.04, O 3.44, F 3.98, Cl 3.16.

Differences: Cl–Cl 0; C–H 0.35; N–H 0.84; O–H 1.24; H–F 1.78.
Order: Cl–Cl < C–H < N–H < O–H < H–F.
Partial negative atom: N in N–H, O in O–H, F in H–F (the more electronegative atom). Cl–Cl has none, and C–H is so close to nonpolar that it is usually treated as nonpolar.

A bigger difference means a more polar bond; there is no sharp dividing line, so a ranking is what a question normally asks for. The reason these values trend as they do, and more bond comparisons, are in electronegativity.

Worked Example 2: Reading Bond Length and Bond Energy From a Potential Energy Curve

For the nitrogen–nitrogen triple bond curve above, find the bond length and bond energy, and describe what happens to the energy at distances shorter and longer than the bond length.

Bond length: the minimum is at about 110 pm.
Bond energy: the minimum is about 945 kJ/mol below the separated atoms (energy zero), so about 945 kJ/mol must be supplied to break one mole of these bonds.
Shorter than 110 pm: the energy rises steeply because the nuclei and the electron clouds repel.
Longer than 110 pm: the energy rises gradually toward zero as the atoms separate and stop interacting.

The bond is most stable at the minimum, and both stretching and squeezing it cost energy. The curve gives the same information as the bond enthalpies used in endothermic and exothermic calculations, where breaking bonds absorbs energy and forming bonds releases it.

Worked Example 3: Comparing Single, Double, and Triple Bonds Between the Same Atoms

The N–N, N=N, and N≡N bonds have lengths of about 145, 125, and 110 pm and bond energies of about 160, 418, and 945 kJ/mol. Rank them by length and by strength, and explain the pattern.

Length, longest to shortest: N–N (145 pm) > N=N (125 pm) > N≡N (110 pm).
Strength, strongest to weakest: N≡N (945) > N=N (418) > N–N (160).
Explanation: more shared electron pairs between the same two nuclei pull them closer together and hold them more tightly, so a higher bond order gives a shorter, stronger bond.

The strength rises by more than a factor of 2 from a single to a double bond, and by more than a factor of 2 again to a triple bond, which is why nitrogen gas is so unreactive. Counting bond order in molecules with resonance is the subject of bond order.

Worked Example 4: Finding the Energy and Light Needed to Break One Bond

The H–H bond energy is 436 kJ/mol. What energy breaks a single H–H bond, and what is the longest wavelength of a photon with enough energy to do it?

Energy per bond: (436,000 J/mol) ÷ (6.022 × 10²³ mol⁻¹) = 7.24 × 10⁻¹⁹ J.
Wavelength: λ = hc/E = (6.626 × 10⁻³⁴ J·s)(2.998 × 10⁸ m/s) ÷ (7.24 × 10⁻¹⁹ J) = 2.74 × 10⁻⁷ m = 274 nm.

Light of 274 nm or shorter is ultraviolet, which is why ultraviolet light, not visible light, breaks covalent bonds. The same conversion between energy per mole and photon energy is the subject of photon energy.

Five Covalent Bond Practice Questions With Answers

Which bond is more polar, C–O or C–N, and which atom is partially negative in it?
Show answer
C–O. The difference is 3.44 − 2.55 = 0.89, against 3.04 − 2.55 = 0.49 for C–N. Oxygen is partially negative.
In H–Cl, which atom has the partial negative charge, and why?
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Chlorine. It is more electronegative than hydrogen, so it attracts the shared pair more strongly.
A potential energy curve has its minimum at 199 pm and 242 kJ/mol below the energy of separated atoms. What do those two numbers represent?
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The bond length is 199 pm, and the bond energy is 242 kJ/mol, the energy needed to break one mole of the bonds.
Why does the potential energy rise very steeply when the atoms are pushed closer than the bond length?
Show answer
The nuclei repel each other, and the electron clouds repel too, and this repulsion grows rapidly as the distance shrinks.
Is the bond in O₂ polar? Explain.
Show answer
No. Both atoms are the same element, so they have equal electronegativity and share the electrons equally.

Common Covalent Bond Mistakes

Saying all covalent bonds share electrons equally. Only bonds between identical or nearly identical atoms do; most are polar.
Confusing bond polarity with molecular polarity. A molecule with polar bonds can still be nonpolar if the dipoles cancel.
Reading the energy curve backward. The depth is the energy needed to break the bond, and the horizontal position of the minimum is the bond length.
Treating bond breaking as releasing energy. Breaking a bond absorbs energy; forming a bond releases it.
Mixing up bond length and bond energy. Length is where the minimum is; energy is how deep it is.
Assuming a longer bond is stronger. For the same pair of atoms, a shorter bond is stronger.

Bond questions combine definitions, reasoning, and a graph, so the same few ideas keep paying off. To see how a stronger Unit 2 would shift your overall estimate, the AP Chem Score Calculator turns section scores into a 1–5 range.

Frequently Asked Questions

What is a covalent bond?

A covalent bond is the attraction between two nuclei and a pair of electrons that they share. It forms when the atoms are close enough that the lowering of potential energy from the shared electrons outweighs the repulsion between the nuclei, and it usually occurs between nonmetal atoms.

What is the difference between a polar and a nonpolar covalent bond?

In a nonpolar covalent bond the shared electrons are attracted about equally, as in Cl–Cl or, nearly, C–H. In a polar covalent bond one atom is more electronegative, so the electron pair is pulled toward it, giving that atom a partial negative charge and the other a partial positive charge.

How do you tell whether a bond is polar?

Compare the electronegativities of the two atoms. A difference near zero means nonpolar, and a larger difference means more polar. There is no sharp cutoff, so the usual task is to rank bonds by polarity and name the partial negative atom.

What does the lowest point on a potential energy curve represent?

The minimum is the most stable arrangement of the two bonded atoms. Its horizontal position is the bond length, and its depth below the energy of the separated atoms is the bond energy, the energy needed to break the bond.

Why does a triple bond have a shorter length than a single bond?

More shared electron pairs hold the nuclei together more strongly, so the minimum on the curve is at a shorter distance and is deeper. A triple bond is both shorter and stronger than a double or a single bond between the same atoms.

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