UTC: Thursday, January 17th, 2019

No biologists, breaking bonds does NOT ‘release energy’

An age old problem that I’ve become a little weary of addressing, so let’s clear this up really quickly and really simply.

No biologists, breaking bonds does NOT ‘release energy’, and here’s why.

No biologists, breaking bonds does NOT 'release energy'Bond breaking is an endothermic process (+ve), and bond making is an exothermic process (-ve). All chemical reactions involve both processes, and the enthalpy change for the reaction is always a SUM of those processes. Biologists choose to ignore the individual breaking and making of bonds during a chemical reaction, and they choose to only report the SUM of these processes, i.e., the enthalpy change, which is fine. However, then they go on to incorrectly label the overall enthalpy change as simply ‘bond breaking’ or ‘bond making’ – that’s plain wrong since they are ignoring the fact that BOTH processes contribute to the overall enthalpy change.

I have no problem with the biologists ignoring the individual processes if that suits them and their own needs, but the problem is with their incorrect labeling of a chemical reaction as only one or the other; it isn’t, it’s both.

As an aside, I am in the habit of teaching enthalpy changes that involve bond breaking and bond making in a manner that is slightly different to most chemistry teachers in the USA. By teaching that a positive number is associated with bond breaking, and that a negative number is associated with bond making, and that ∆H is the SUM of the two, I try to emphasize a knowledge of the endothermic/exothermic nature of bond breaking/making. In my experience, most chemistry teachers don’t do that, rather they tend to say that ∆H = bonds broken MINUS bonds made. Of course this ultimately yields the same answer, but I find that my method tends to help cement an idea that would help this biology based misconception.

Comments

  1. Pamela Gardner says:

    Very useful observation.
    I’m interested in physiological chemistry – originating in dietetics – and all aspects of the transference of energy. And, of course, a main energy source comes from the reduction of ATP to ADP: but to describe the energy simply as coming from the breaking of a bond in ATP would be wrong and not further an understanding in any of the complex metabolic, neural or other pathways.

  2. I teach enthalpy changes the same way. I find it prevents students from having to try and memorize a formula, instead making sense of what is actually going on. Many of my AP Chemistry students have already taken AP Biology, and therefore it is an uphill battle convincing them that breaking a bond does not release energy. Using complete biological examples with them seems to help. Also, telling them that making such a statement is incorrect and will have a negative impact on their grade, also works.

  3. This is a fundamental misunderstanding between a .couple of terms. Biochemistry at the basic AP level in high school or undergraduate level would not deal with enthalpy but with free energy. Thus you’ll see a difference in the term endothermic vs. endergonic, or exothermic vs. exergonic. Free energy is released when phosphate groups are broken away from ATP molecules, and it is gained when those bonds reform during cellular respiration.

    This is a good post but I think it could be improved by adding in some details that tie in the actual biology curriculum.

    Another misunderstanding is that many biology teachers think that covalent bonding is the “strongest” type of bonding, but they are used to water environments where ionic bonds fall apart rather easily.

    Another recommendation I could give to Biology teachers is to incorporate heating and cooling curves into your discussions of water as the properties of water are difficult to discuss without.

    • Thanks for the comment. I don’t know much about the biology curriculum other than the title of this post!

      As for comparing the strength of ionic and covalent bonding, I’ve always found it to be a totally pointless exercise since, among other things, one is never comparing like with like. For example, there is only one type of bonding (or at least only one point on the sliding scale of bond type) present in any given compound. As a result, we can’t compare the relative strength of say the bond between sodium and chlorine in both an ionic AND a covalent situation (if that makes sense).

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