Last updated: September 7th, 2022 |
Alkene Stability (And Instability)
What factors affect alkene stability? If you’ve studied elimination reactions, no doubt you’ve learned about Zaitsev’s Rule – about how elimination reactions generally favor the “more substituted” alkene.
In this post we explore how increasing substitution at carbon increases the stability of alkenes, as well as the effects of conjugation and strain.
Table of Contents
- Heat Of Hydrogenation As A Measure Of Alkene Stability
- Stability of Alkenes Increases With Increasing Substitution
- Heats Of Hydrogenation For Some Monosubstituted Alkenes
- The Relative Stability of cis- and trans- Alkenes
- Alkenes Stabilized By Conjugation: Resonance Energy
- Alkene Stability: Summary
- Bonus Topic #1: Why Is Alkyl Substitution Stabilizing?
- Bonus Topic #2: trans-Cycloalkenes
- Quiz Yourself!
- (Advanced) References and Further Reading
We might not spend as much discussing thermodynamics in here organic chemistry as you did in general chemistry, but that doesn’t mean the concepts have just gone away!
One area where we’ve previously seen the usefulness of thermodynamic data is the use of heat of combustion data to quantify ring strain. [See: Cycloalkanes – How To Calculate Ring Strain]. The heat of combustion for cyclopropane works out to about 166 kcal/mol per CH2 compared to the heat of combustion for unstrained cyclohexane [157 kcal/mol per CH2]. That “extra” heat of combustion seen in cyclopropane is attributed to the instability arising from the strain of bent C-C bonds far away from their ideal angle of 109.5°. That’s angle strain.
Another area of organic chemistry where thermodynamic studies are useful in the stability of alkenes.
Back in 1935, Prof. Kiasatakowsky and co-workers at Harvard published a method for measuring the heat of hydrogenation of ethylene (aka “ethene”) as it was passed over a finely divided metal catalyst containing adsorbed hydrogen. [Note 1] Because hydrogenating a molecule is considerably more gentle than, say, BURNING it, the method tends to be more sensitive for determining subtle differences in enthalpies.
In a hydrogenation reaction, a C-C bond is broken, and two new C-H bonds are formed.
It was found that hydrogenation of ethylene released 32.5 kcal/mol (136 kJ/mol) of heat. [Note 2]
Once the heat of hydrogenation of ethene was obtained, the next logical step was to measure the heat of formation for a huge variety of other alkenes, and to see what patterns emerged from the data.
Well, as you might imagine from someone who had invented a new technique, Kiastakowsky went to town on this, investigating the heat of hydrogenation of a huge variety of alkenes. [Note 3] In the following decades, even more data has been accumulated, which is easily obtainable (with references) from the NIST Chemistry Web Book.
For our purposes, there are six substitution patterns on an alkene (seven if you count ethene).
The most notable trend that was found is that the heat of hydrogenation decreases as C-H bonds are replaced with C-C bonds.
So what does that mean?
Since the same bonds are formed and broken in every hydrogenation reaction, the heat of hydrogenation is measuring the stability of each type of alkene.
This means that the lower the heat of hydrogenation, the greater the stability of the alkene.
The way to visualize “stability” here is to compare it to potential energy, much like a ball becomes more “unstable” with increasing height.
So what we’re really saying here is that alkene stability increases with increasing substitution of hydrogen for carbon.
[The image above uses heat of hydrogenation data for the series hex-1-ene, trans-hex-2-ene, cis hex-2-ene, 2-methylpent-1-ene, 2-methyl-pent-2-ene, and 2,3-dimethylbutene, which all share the molecular formula C6H12. ]
OK, you might ask. So, why does this happen?
The short answer is that substitution of alkyl groups on the alkene allows for donation of electron density between (full) C-C sigma orbitals and the (empty) C-C pi star orbital. It’s often not addressed in introductory courses, so we’ll push the explanation down to this footnote. [Bonus topic one]
Just for fun, let’s look at a series of mono-substituted alkenes. Nothing weird here, we’ll just go from propene up to hex-1-ene.
Note that the heat of hydrogenation is quite consistent for a series of linear, non-branched, monosubstituted alkenes.
We all know by now that cis and trans alkenes should differ a little bit in stability because in a cis alkene the groups are held closer together (more strain!) and in a trans-alkene they are further apart. [For a good time, amaze your instructor and call it by its proper name: 1,2-strain]
While a difference of 1 kcal/mol might not seem like a lot, it isn’t *that* small – for an equilibrium at 25 °C, a difference of 1 kcal/mol will give you about an 80:20 ratio of products. [Note 4]
For a really good time you can pick something crazy like the cis– and trans- di t-butyl ethylene. [not the correct IUPAC name, but definitely more vivid than cis- and trans- 2,2,5,5-tetramethylhex-3-ene].
Here the trans is more stable than the cis by about 10 kcal/mol.
That’s a lot of strain.
The stability of alkenes is also affected by conjugation. This is a really a topic for another chapter [specifically, see Conjugation and Resonance] where we talk about pi systems, but the bottom line is that the p-orbitals in adjacent pi-bonds can clump together forming larger “pi-systems”, which provides more “room” for electrons to roam, lowering their energy. [Note 5]
Heat of hydrogenation numbers allow us to quantify the effect of resonance stabilization. How so?
Take but-1-ene. As we saw above the heat of hydrogenation is about 30.1 kcal/mol.
Add a double bond, and you might expect the heat of hydrogenation to double as well. But it doesn’t! It’s actually a little bit less. [56.6 kcal/mol] . The difference (that extra 3.6 kcal/mol of additional stabilization) is called “resonance energy“.
The most dramatic example of resonance energy is found in the example of “cyclohexatriene” , which has an extra stabilization energy of 36 kcal/mol. That’s a sure sign that something highly unusual is going on with this molecule, which is better known as “benzene”. That “highly unusual” property is called aromaticity and it warrants its own chapter. [See: Introduction to Aromaticity]
Three key factors affect the stability of alkenes, and the influence of these factors can be measured through the enthalpy of hydrogenation.
- One important factor is the substitution pattern. As C-H bonds are replaced by C-C bonds, the stability of the alkene gradually increases in the order mono (least stable) < di < tri < tetrasubstituted (most stable).
- When hydrogenation liberates more energy than expected given the substitution pattern, that’s likely a sign of strain. This is exemplified in the difference in enthalpy of hydrogenation between cis- and trans- alkenes, where the trans- alkene is more stable by about 1 kcal/mol.
- When hydrogenation liberates less energy than expected given the substitution pattern, that’s a sign that some extra factor is stabilizing the molecule. Among commonly encountered factors, conjugation ranks high. The difference in energy between the “expected” heat of hydrogenation and the measured heat of hydrogenation is called the resonance energy. The conjugation of one pi bond with an additional pi bond is “worth” about 2-3 kcal/mol.
- Elimination Reactions (2): The Zaitsev Rule
- E and Z Notation For Alkenes (+ Cis/Trans)
- Addition Reactions: Elimination’s Opposite
- Reactions of Enols – Acid-Catalyzed Aldol, Halogenation, and Mannich Reactions
Note 1. It was a copper catalyst, after a lot of trial and error. The advantage of measuring the heat of hydrogenation over the heat of combustion is that it is a more sensitive technique for measuring small energies.
Note 3. Standard heats of hydrogenation have been pulled from the NIST Chembook.
Note 5. If you think of electrons as waves, a larger pi-system allows for longer wavelengths, and since energy is inversely proportional to wavelength, this means a lower overall energy of the electron.
And a big thank you to The Kraken for his steady hands in the stability GIF.
So why does increasing substitution at the alkene increase its stability? This is not an easy question to answer to an introductory audience in a few sentences, and given the time constraints of a typical course the answer you will generally get from an instructor will range from “it’s complicated” to “hyperconjugation” to “orbital mixing”. Very rarely you might get an MO diagram.
The unifying principle here is that full orbitals – even those from single bonds – can donate into empty (even antibonding) orbitals, and that this interaction is stabilizing.
In ethene (below left) all of the C-H bonds are in the plane of the alkene, and none can overlap with the pi bond.
When a methyl group is added, say, in propene, one of the C-H bonds can now align with the pi-system of the alkene. The pair of electrons from the C-H bond can then donate into the empty pi* orbital.
This can be visualized through “no-bond resonance”, below right, where a “resonance” form is shown with a broken C-H bond and a new C-C pi bond. [The quotation marks are to differentiate it from our traditional view of resonance where only pi-bonds are allowed to form and break].
This mixing results in a stabilization of the molecule. . Although CH3 is in rapid rotation, at any given moment at least one of the C-H bonds will have the proper geometry to allow overlap with the pi system.
Predicted to slightly lengthen C-H and C-C pi and strengthen C-C sigma.
99% of people reading this will never use this so it is going down in the footnotes.
In the vast majority of molecules you will encounter, the double bonds in rings are cis. Why? The most vivid answer is provided by trying to make them with a model kit.
That is not a happy double bond.
However at a ring size of 7, a trans double bond becomes more than transiently stable (albeit very short lived at 0°), and at a ring size of 8 there’s enough floppiness in the ring such that its boiling point can be measured [143°C !] . Larger ring sizes than 8 can easily accommodate a trans double bond.
The heat of hydrogenation can be used to quantify the stability of these rings (note that this is not the whole picture, since it doesn’t take entropy into account, and that can be quite significant).
At ring sizes of 11 and 12 the trans-isomer actually becomes more stable (when allowed to equilibrate with acid) but recall that anything involving equilibrium is ultimately a measure of delta G, and delta G also includes an entropy term (S). It turns out that the main factor in the increased stability of 11- and 12- membered trans-cycloalkenes is their greater entropy. See this reference.
All heat of hydrogenation values cited here were obtained from the NIST Chemistry Web Book. Searching by CAS number never fails. Selected original references below.
- Heats of Organic Reactions. I. The Apparatus and the Heat of Hydrogenation of Ethylene
G. B. Kistiakowsky, H. Romeyn Jr., J. R. Ruhoff, Hilton A. Smith, and W. E. Vaughan
Journal of the American Chemical Society 1935 57 (1), 65-75
Prof. Kistiakowsky’s first (of many) papers on the heat of hydrogenation of organic molecules, where he describes the apparatus required to obtain accurate heat of hydrogenation data in painstaking detail. The results stand up.
- Heats of Organic Reactions. IV. Hydrogenation of Some Dienes and of Benzene
G. B. Kistiakowsky, John R. Ruhoff, Hilton A. Smith, and W. E. Vaughan
Journal of the American Chemical Society 1936 58 (1), 146-153
Contains the heat of hydrogenation for 1,3 butadiene, benzene, and other unsaturated molecules, including allene (71.0 kcal/mol).
- Heats of Hydrogenation. IV. Hydrogenation of Some cis- and trans-Cycloölefins1
Richard B. Turner and W. R. Meador
Journal of the American Chemical Society 1957 79 (15), 4133-4136
- Heats of hydrogenation. IX. Cyclic acetylenes and some miscellaneous olefins
Richard B. Turner, A. D. Jarrett, P. Goebel, and Barbara J. Mallon
Journal of the American Chemical Society 1973 95 (3), 790-792
- RELATIVE STABILITIES OF cis- AND trans-CYCLONONENE, CYCLODECENE, CYCLOUNDECENE AND CYCLODODECENE
Arthur C. Cope, Phylis T. Moore, and William R. Moore
Journal of the American Chemical Society 1959 81 (12), 3153-3153
A.C. Cope reported that when cis– and trans– cycloundecene (11-membered) and cyclododecene (12-membered) are allowed to equilibrate (by heating with catalytic TsOH) the trans-double bond is favored at equilibrium (i.e. has lower Δ G)… even though trans-dodecene has a higher enthalpy (Δ H) than its cis-isomer. This is a helpful reminder that enthalpy (delta H) is just one part of the Gibbs equation (Δ G = Δ H – TΔ S), the trans-cycloalkenes have higher entropy (S) and this explains their greater stability.