[nils290479]

Hi everyone,

Two weeks ago, I started learning chemistry by myself and I’m now reading The Cartoon Guide of Chemistry. There is something that I don’t understand about endothermic reactions:

During endothermic reactions, it is said that heat is absorbed, and I’m not sure I truly understand why energy is *absorbed*. My taking is that during an endothermic reaction, the bond(s) formed are weaker than the bond(s) broken, so if taking one endothermic reaction apart there is a *consumption* of energy (I think “consumption” is a very wrong word since energy cannot be created nor destroyed). But since the weaker form bond will need a smaller quantity of heat to be broken, the next reaction is likely to release heat and be exothermic, and then heat will be released.

So, if I understand well, it is said that heat is *absorbed*during an endothermic reaction not because heat is *stored* in the formed bond(s), but because the formed bond(s) are weaker and the next reaction is more likely to be exothermic and then release heat.
Is it so ?

Thanks.

[mjc123]

That's pretty much right. Consider the reaction 2H₂(g) + O₂(g) 2H₂O(g)
We can estimate the heat of reaction by imagining we first break all the bonds, to give 4 H atoms and 2 O atoms, then form bonds to make H₂O molecules. To do this we need to break 2 H-H bonds and one O=O double bond and form 4 O-H bonds. Tabulated bond energies are
H-H 436 kJ/mol
O=O 498
O-H 464
So we need to put in 2*436 + 498 = 1370 kJ/mol
and we get out 4*464 = 1856 kJ/mol
so the change in energy is 1370 - 1856 = -486 kJ/mol: it is exothermic; energy is given out. The reverse reaction would be endothermic.
(Actually this is only an approximation, as bond energies vary from compound to compound, but it serves to illustrate the point. Also there may be other energy terms to consider, e.g if the product was H₂O(l) you'd have to take into account the latent heat of vaporisation, and in condensed phases there are intermolecular interactions that do not constitute chemical bonds but are energetically significant.)
You are correct to say that in an endothermic reaction the bonds formed are weaker than the bonds broken (that is, taken overall; O=O is stronger than O-H, but it's only one to four), therefore you go from a state of lower energy to a state of higher energy, which requires input of heat. It would be wrong to say that "the heat absorbed is stored in the formed bonds", because heat is released when bonds are formed (relative to the separated atoms). Strong bonds means low energy. It would be more correct to say that the heat is stored by putting the system into a state of (relatively) high potential energy by forming weaker bonds.

[nils290479]

Many thanks for your explanation, mjc123. I have 2 more questions if you don't mind.

1- In your example, we need 1370 kJ/mol to break the bonds. Is this energy what is called activation energy ?

2- You wrote that "the heat is stored by putting the system into a state of (relatively) high potential energy by forming weaker bonds", and I don't get to understand how the strength of a bond is negatively related to the energy in the system, and how the heat is stored in the system as potential energy. I think that this thing of potential energy is what I struggle to understand from the beginning

Well, that's 3 questions actually

[mjc123]

1. No, because the reaction doesn't actually happen by splitting all the molecules into separate atoms. Activation energy is the difference between the energy of the transition state and that of the reactants. The transition state is the configuration of highest energy between reactants and products (or reactants and intermediates for a multistep mechanism). For example, if you have a reaction A + BC  AB + C the transition state might be a species A....B....C in which the A-B bond is partially formed and the B-C bond partially broken.
2. I don't get what you don't get here; doesn't the diagram explain it? The y axis is effectively "potential energy", though it may be called other things. Think of the physical analogy of a ball rolling up a hill; kinetic energy is converted to potential energy. So in an endothermic reaction: heat (kinetic energy of the molecules) is converted into (chemical) potential energy; that's why it gets colder unless heat is put in from the surroundings. (There is a thing called "chemical potential", related to this; a system tries to reach a state of minimum chemical potential. "Why do endothermic reactions occur then?" There are other factors involved, but you have to have been learning chemistry for more than 2 weeks to grasp them.)
The diagram shows clearly that (relative to the separated atoms) systems with stronger bonds (requiring more energy to break) have lower energy. Energy is not stored in chemical bonds, you have to put in energy to break them.

[Enthalpy]

My own 2 cents addition...

Endothermic doesn't mean improbable. The motor of a reaction is not the loss of heat - since an other body has to absorb this heat - but something more subtle called entropy that tells how evenly energy spreads over the many degrees of freedom that can store it.

For instance, melting ice is endothermic, but melting is probable and even certain at +25°C.

There are situations where an object is expected to lose energy. A weight falls naturally. This is because its gravitation energy is concentrated in a single degree of freedom: height, which may store several J. Heat instead stores 2 to 4*10^-21J per degree of freedom at room temperature, that is, each direction of movement of each atom, and similar things. Then, height converts into heat spontaneously when the object falls, as the J can spread more evenly into many degrees of freedom, while the opposite process doesn't happen.

So easily lost energy holds for mechanical energy, joules of electricity...

As opposed, it doesn't hold necessarily when the considered energy is already spread among many degrees of freedom. That's the case for heat, possibly for chemical energy - every time the unit energy change is comparable with R*T (for moles) or k*T (for molecules). Then, equilibria are possible, and are predicted numerically (or rather, properly extrapolated from a small number of measurements) through entropy, Gibb's energy, chemical potential, these things.

That's really a matter of energy concentration, not of energy nature. If you build sensitive electronics or accelerometers, you observe a residual fluctuating electrical or mechanical energy called a noise, that stores about kT/2 and won't disappear spontaneously.

[Borek]

The motor of a reaction is not the loss of heat - since an other body has to absorb this heat - but something more subtle called entropy

If anything, it is a Gibbs free energy that "runs" the reaction, so a combination of the enthalpy and entropy. There are reactions driven by the enthalpy change, there are reactions driven by the entropy change, you can't say one is more important than the other.