A double bond is a covalent bond in which two atoms share two pairs of electrons (one sigma bond plus one pi bond), giving it a bond order of 2, which makes it shorter and stronger than a single bond but longer and weaker than a triple bond.
A double bond forms when two atoms share two pairs of electrons instead of one. Structurally, it's one sigma (σ) bond plus one pi (π) bond. You'll see it in molecules like O₂ (O=O), CO₂ (two C=O bonds), and ethylene (C=C).
The AP Chem CED cares about double bonds mostly through bond order. Per essential knowledge 2.2.A.2, bond length depends on the size of the atoms' cores and the bond order. Bonds with higher order are shorter and have larger bond energies. So for the same two atoms, a double bond sits lower on the potential energy curve than a single bond does. The atoms pull in closer (shorter equilibrium bond length) and it takes more energy to rip them apart (larger bond energy). Think of bond order as how many rubber bands are holding two balls together. Two bands pull tighter and resist breaking more than one.
Double bonds live in two units. In Unit 2 (Topic 2.2, Intramolecular Force and Potential Energy), learning objective 2.2.A asks you to read and interpret potential energy vs. internuclear distance graphs. A double bond's curve has a deeper well (more energy to escape) and a minimum at a shorter distance than a single bond between the same atoms. In Unit 6 (Topic 6.7, Bond Enthalpy), learning objective 6.7.A has you calculate ΔH from average bond energies, and you can't do that correctly unless you spot which bonds in your Lewis structures are double bonds. Counting C=O in CO₂ as a single bond, or forgetting that O₂ contains an O=O double bond, wrecks the whole calculation. Double bonds are the connective tissue between structure (Unit 2) and energy (Unit 6).
Keep studying AP Chemistry Unit 2
Sigma Bond (σ) and Pi Bond (π) (Unit 2)
Every double bond is one sigma bond plus one pi bond. The sigma bond is the head-on overlap; the pi bond is the side-by-side overlap stacked on top of it. This is why a double bond is stronger than a single bond but not exactly twice as strong. The pi bond's sideways overlap is weaker than the sigma's direct overlap.
Potential Energy & Internuclear Distance (Unit 2)
On a potential energy diagram, a double bond shows a deeper energy well at a shorter internuclear distance than a single bond between the same atoms. Deeper well means more bond energy; shorter minimum means shorter bond length. One graph captures both trends from 2.2.A.2.
Resonance Structures (Unit 2)
When a double bond can be drawn in more than one position (like in O₃ or NO₃⁻), the real molecule averages them out. The actual bonds have a fractional bond order between 1 and 2, with lengths between a typical single and double bond. Experimental bond lengths are how you prove resonance is happening.
Bond Enthalpy Calculations (Unit 6)
ΔH ≈ (energy of bonds broken) − (energy of bonds formed), and double bonds have their own average bond enthalpies. In the classic methane combustion problem, you break O=O bonds in O₂ and form C=O bonds in CO₂. Misidentify those double bonds and your answer is off by hundreds of kJ.
Multiple-choice questions love bond order comparisons. Expect stems like "why is the N≡N bond energy (941 kJ/mol) higher than the O=O bond energy (495 kJ/mol)?" or "why is C≡C in acetylene shorter than C=C in ethylene?" The credited answer always runs through bond order. More shared electron pairs means stronger attraction, shorter bond length, deeper potential energy well, bigger bond energy. You also need double bonds in Unit 6 calculations. A typical problem gives you average bond enthalpies (like C–H 413, O=O 495, C=O 799, O–H 463 kJ/mol) and asks for ΔH of a combustion reaction, which requires drawing or visualizing Lewis structures to count every double bond correctly. On FRQs, you may be asked to compare two potential energy curves and explain which corresponds to the higher bond order, or to justify a bond length trend using Coulombic attraction.
A single bond is one shared pair (one sigma bond); a double bond is two shared pairs (one sigma plus one pi). The double bond is shorter and stronger, but here's the trap. A double bond is NOT exactly twice the strength of a single bond, because the pi bond's side-by-side overlap is weaker than the sigma's head-on overlap. Compare C–C at about 348 kJ/mol with C=C at about 614 kJ/mol. Stronger, yes, but less than double.
A double bond is two shared electron pairs between atoms, made of one sigma bond and one pi bond, with a bond order of 2.
Higher bond order means shorter bond length and larger bond energy, so double bonds are shorter and stronger than single bonds but longer and weaker than triple bonds (EK 2.2.A.2).
On a potential energy vs. internuclear distance graph, a double bond has a deeper well at a shorter distance than a single bond between the same atoms.
A double bond is stronger than a single bond but not twice as strong, because the pi component is weaker than the sigma component.
In bond enthalpy calculations (Topic 6.7), you must count double bonds correctly in Lewis structures, including the O=O bond in O₂ and the two C=O bonds in CO₂.
Resonance structures show that some real bonds are averages with bond orders between 1 and 2, which experimental bond lengths confirm.
A double bond is a covalent bond where two atoms share two pairs of electrons, one in a sigma bond and one in a pi bond. It has a bond order of 2, making it shorter and stronger than a single bond between the same atoms.
No. A double bond is stronger but less than double the strength. C–C is about 348 kJ/mol while C=C is about 614 kJ/mol, because the pi bond's sideways overlap is weaker than the sigma bond's direct overlap. AP questions test this exact misconception.
A triple bond shares three electron pairs (one sigma, two pi) and has bond order 3, so it's even shorter and stronger than a double bond. That's why N≡N (941 kJ/mol) has a much higher bond energy than O=O (495 kJ/mol), and why C≡C in acetylene is shorter than C=C in ethylene.
More shared electron pairs increase the attractive force pulling the two nuclei together, so the atoms settle at a smaller equilibrium internuclear distance. On a potential energy diagram, the minimum of the curve shifts left and deepens as bond order increases.
O₂ (O=O), CO₂ (two C=O bonds), and ethylene (C=C) come up constantly. O₂ and CO₂ matter most in Unit 6, because combustion bond-enthalpy problems require you to count their double bonds when totaling bonds broken and formed.