Before we go into the details explaining the bong lengths and bond strengths in organic chemistry, let’s put a small summary for these two properties right from the beginning as it stays relevant for all types of bonds we are going to talk about.

So, remember this: the shorter the bond, the stronger it is.

To understand the principles behind bond strength and bond length pertaining to organic molecules, let’s first discuss the data known for the hydrogen halides:

The bond strength increases from HI to HF, so the H-F is the strongest bond while the H-I is the weakest.

Why is this the case? First, looking at the periodic table, we can notice a pattern correlating the bond strength and the atomic size.

Remember that the atomic size increases down the periodic table, and fluorine, for example, uses an sp³ hybrid orbital made of its second shell orbitals to form a bond with hydrogen:

The other halogens are on the 3^rd, 4^th, and 5^th rows of the periodic table and therefore, they use larger orbitals during the hybridization and consequently bond formation. If we put them next to each other, we can use this demonstration of differences in bond length to explain the bond strengths as well:

What we see is that as the atoms become larger, the bonds get longer and weaker as well. Longer bonds are a result of larger orbitals, which presume a smaller electron density and a poor percent overlap with the s orbital of the hydrogen. This is what happens as we move down the periodic table, and therefore, the H-X bonds become weaker as they get longer.

So, keeping this in mind, let’s now see how the length and the strength of C-C and C-H bonds are correlated to the hybridization state of the carbon atom.

Bond Strength and Electronegativity

The correlation between the bond length and bond strength is very accurate and works in most types of bonds you will see in organic chemistry. However, there are, of course, some exceptions. The most notable is probably the effect of electronegativity on the bond strength. For example, the bonds get weaker in the following series, even though we’d expect the C-N bond to be stronger than the C-O bond because nitrogen has a smaller atomic size than oxygen: CH₃–F > CH₃–OH > CH₃–NH₂.

The reason for this is the higher electronegativity of oxygen compared to nitrogen. Remember, as the difference between the electronegativities of bonded atoms increases, we switch from nonpolar to polar covalent bonds, in which the electron density is unequally shared between the atoms and they are characterized by partial positive and negative charges. 

Now, when the atoms have these partial charges, the bonding between them starts to attain some ionic character as well. Ionic bonds are generally stronger than covalent bonds, which we can also see by their significantly higher melting points. 

So, this should explain why the C-O bond (143 pm, 360 kJ/mol) is normally stronger and shorter than the C-N (147 pm, 305 kJ/mol) bond because the atomic sizes are not so different, while the greater electronegativity of the oxygen brings some ionic character to the C-O bond. This is also true when comparing the strengths of O-H (97 pm, 464 kJ/mol )and N-H (100 pm, 389 kJ/mol) bonds. 

Bond Length and Strength in Organic Molecules

Why do you think the bond strength of the C-H bond in alkane, alkene, and alkyne follows the pattern shown below?

We have concluded, in the previous part, that the bond strength is inversely correlated to the bond length, and according to this, the data suggest that the C-C bond in alkanes must be the longest as it is the weakest, and the C-C bond in alkynes is the shortest as it appears to be the strongest.

And this, in fact, is true because remember, the bond length decreases going from sp³ to sp hybridization:

To understand this trend of bond lengths depending on the hybridization, let’s quickly recall how the hybridizations occur. For the sp³ hybridization, there is one s and three p orbitals mixed; sp² requires one s and two p orbitals, while sp is a mix of one s and one p orbital.

Now, there is something called “s character” which refers to the % of the s orbital initially involved in the hybridization process. For example, in the sp³ hybridization, there is a total of four orbitals – one s and three p, and out of these, only one is (was) an s. Therefore, the s character of an sp³ orbital is ¼ = 25%. With the same principle, sp² orbitals are 33%, and sp orbitals have 50% s character:

The next question is – how the s character is related to the bond length and strength. Here, you need to remember that for a given energy level, the s orbital is smaller than the p orbital. A smaller orbital, in turn, means stronger interaction between the electrons and the nucleus, shorter and therefore, a stronger covalent bond. This is why the C-C bond in alkynes is the shortest/strongest, and that of alkanes is the longest/weakest as we have seen in the table above.

The C-C vs C-H Bond Strength

The relative size of the s orbital explains also why the C-C σ bond is weaker than the C-H σ bond. And that is because the hydrogen uses a “pure” s orbital (100% s character), which is closer to the nucleus than is the sp³ orbital of carbon. As a result, the nuclei are held closer in an sp³–s C-H bond than in a sp³–sp³ C-C bond:

Now there are different types of C-H bonds depending on the hybridization of the carbon to which the hydrogen is attached. As in all the examples we talked about so far, the C-H bond strength here depends on the length and thus on the hybridization of the carbon to which the hydrogen is bonded.

The higher the s character in the hybrid orbital connecting the two atoms, the shorter and stronger the C-H bond:

To summarize the information in the table, remember the bond strength order C(sp)-H > C(sp2)-H > C(sp3)-H. The reverse would be true about the bond lengths.

All these values mentioned in the tables are called bond dissociation energies – that is, the energy required to break the given bond. Specifically, we are talking about the homolytic cleavage when each atom gets one electron upon breaking the bond. The bond dissociation energies of the most common bonds in organic chemistry, as well as the mechanism of homolytic cleavage (radical reactions), will be covered in a later article, which you can find here.

The Strength of Sigma and Pi Bonds

There is one important thing we should address when comparing the strength of a single bond with a double or a triple bond. Remember that a multiple bond consists of one σ and one or two π bonds. Now, if we compare the single bond strength with the double bond, we have 88 kcal/mol:152 kcal/mol. This is not a 1:2 ratio, which indicates that σ bonds are stronger than π bonds; otherwise, the double bond would have been 176 kcal/mol strong (2 x 88).

Using the difference of values of C(sp²)- C(sp²) double bond and C(sp²)- C(sp²) σ bond, we can determine the bond energy of a given π bond.