Submitted by jeez-gyoza t3_yff8um in askscience
My teacher told me that covalent is the strongest bond because of the shared electrons, but google results keep saying that ionic is stronger because of the force of attraction between cation and anion.
An example he gave us is: a diamond is stronger than limestone (Ca2CO3) calcium carbonate.
Bond strengths are relative energies. That means you must consider two states: what you have with the bond intact and what you have without the bond intact. The latter state especially can be highly environment dependent.
So, for example, the ionic bond between sodium and chloride ions is quite strong in a dry environment (melting pure salt takes a decent amount of energy). But put salt in water and the bond breaks spontaneously. That is because you alter the energy (or free energy to be precise) of the latter state. Similarly, the covalent bond between two oxygen atoms is pretty strong when O2 is floating as gas. But, thankfully for us, inside our bodies, near hemoglobin especially, that bond can be broken at very reasonable energies. Again, the difference is in the latter of the two states.
So a categorical statement that one bond type always is stronger than the other is not possible. What often is the case is that these energies are considered in vacuum as a reference environment. But that says by itself little about what applies in everyday applications. Because salts tend to be more dependent on whatever solvent we surround them with, as a matter of experience, really strong materials we interact with tend to be more covalent in their bonding. However, metals are an interesting challenge, since say Tungsten is relatively easy to deform, but extremely hard to melt… so even in the vacuum reference environment, these material properties are not as easy to place on a single weak—strong axis.
So, the example of diamond is a bit of an outlier for "covalent" bonds. More specifically, diamond is made up of covalent network bonds. These networks buck the *very general* trend of ionic bonds are stronger than covalent bonds.
But as another commenter said, one usually makes more direct comparisons, where you can actually compare energies. It's just easier to introduce trends before the more nuanced takes.
This is a great answer. TBH I've always found it odd that we teach the concept of "ionic bonds" in introductory classes. IMO thinking of them as bonds can be very unintuitive and misleading, for the reasons you reference in your answer.
Yeah, "bonds" is a fuzzy concept. We also have "intermolecular bonds", but they are even farther from a well-defined structure and require analysis in terms of a distribution of relative configurations.
But I suppose these terms arose historically and are based on some phenomenology of human experiences. Say stuff that sticks together at room temperatures might be understood as "bonded together". So then only gases are understood as lacking bonds. With our molecular level understanding nowadays, that puts very different things in the same conceptual bucket. But history is hard to rewrite...
I'm confused by how diamond is an outlier. The bond dissociation energies in diamond are basically the same as any other C-C single bond.
Each individual bond, yes. But the vast network of these bonds adds up in ways that other covalent compounds with C-C bonds simply can't (not apples to apples, but ethane, for example).
This is just one way that talking about trends generally doesn't really tell the true story for a specific compound. Because of the sheer number of bonds, covalent network bonding is considered the strongest type of bonding (even though I'm not sure this comparison should really be made this way). And the trend is that ionic bonds are stronger than covalent bonds are stronger than intermolecular interactions, but when you start comparing actual compounds' values for bond strength, it is really easy to find exceptions.
One thing I would point out is that the difference in energies between the start and end states are the same no matter the situation. The the examples you give are where the bonds can more easily be broken because the intermediate states are more easily reached. I.e. before the O2 is separated out most first have the bonds Bricker and both O2 would be so reactive that in the time they are separate they would just bond back to one another. This requires a lot of energy i.e. to make something catch fire you need to first provide enough energy to do this, usually with something very hot. In our bodies the intermediate state where they are not bound or loosely bound is stabilized by specific molecules in our body making it easier to break the bond in O2. Since this intermediate state is stabilized you don't need the large initial energy to initiate the reaction (some energy is still needed). However, the energy produced is still the same as it depends on the final and initial states only. Edit: while this is true often there is more than one reaction happening such as enthalpy of solvation, solvent- solvent integrations, and solvent solute interactions.
That said the bond energy of NaCl is 787 kj/Mol and O2 is 498 kj/mol.
Is point out as well your teacher is not correct. Ionic bonds are stronger. But, in an aqueous state they can much more easily be solvated due to water stabilizing the intermediates. Ionic bonds since they do not share electrons, once separated are very easily stabilized. In covalent bonds this is generally not the case.
I disagree. That energy you quote for NaCl assumes the two ions are pulled apart in vacuum. If we then move the respective ion into water, there is a favourable solvation free energy. So the free energy of the final state of separated NaCl in aqueous solution is different than had it been done in vacuum. Any intermediate state stabilization, that only alters the kinetics, is not relevant.
The O2 example is, I admit, a bit different in the details. But in vacuum, if I do a homolytic cleavage of the covalent bond in O2, the two radical oxygen atoms of the final state are far from stable. So any other reactant that can combine with the two oxygen atoms make that final state stabler. The tricker part is that the path from start to finish include at least one transition state. In biology and chemistry, we can alter the kinetics of this cleavage, while leaving the final state the same, by adjusting any catalytic component. That is a more subtle point. My argument is much simpler. In vacuum, O2 is a strong bond because the final state after dissociation is unstable, while in the example system, O2 separates more readily mostly because the dissociation does not generate two free radical atoms, but one where the atomic oxygen binds to iron.
Hence, the strength of a bond must either be defined with a common reference state (vacuum typically as I mention), but then bond strength is a more abstract quantity and less informative of practical questions of stability, robustness etc, or we consider the problem in full, i.e. the stability of initial and final states are part of the analysis.
Comparing network solids to small molecule compounds isn't what's confusing. Calling diamond an outlier is confusing, it's not even the most thermodynamically stable allotrope of carbon under ambient conditions. There are also countless other comparable network solids that behave like diamond (like boron nitride). My gripe is that there's absolutely nothing special about diamond from the standpoint of "bond strength," so calling it an outlier is weird.
So we are on the same page for the most part here.
But for the dissolution of salt you still have the bonds or really lattice energy to overcome. The totally enthalpy is negative as you make most of the energy required to break the bond during solvation but it still requires that energy, i.e. it's endothermic. This is what I alluded to in the multiple reactions. Total enthalpy is the sum of the bond/ lattice energy + the sum of solvent-solvent attraction energy - the sum of the solute- solvent energy.
Maybe I'm missing something but you still use the standard bond enthalpy for this equation. I.e. the sum of all the reactions enthalpy. Curious on your thoughts here.
Edit: I will agree that bond strength is not especially representative to the common environments like aqueous one.
Covalent bonds are typically stronger than ionic bonds, but there are some exceptions. For example, some ionic bonds are stronger than covalent bonds because the ions are smaller and have a higher charge density.
>But, thankfully for us, inside our bodies, near hemoglobin especially, that bond can be broken at very reasonable energies.
I'm not quite sure if this is what you're implying, but hemoglobin does not break the double bond in molecular oxygen. In mammals, oxygen (as an intact diatomic molecule) binds hemoglobin in the blood, then binds myoglobin in target tissues, and then is released into solution, where it is reduced to water in the mitochondria. And while I don't want to get bogged down in the definition of "reasonable" energy levels, I will point out that that redox reaction involves free electrons. Nonetheless, you are certainly correct that the bonds in molecular oxygen can be considered weak in many everyday contexts - that fact is synonymous with the fact that oxygen is highly reactive in many common chemical environments on Earth.
In hindsight I should have used a better example to illustrate the point that the apparent durability of a covalent bond also depends on a relative free energy, where that delta can be highly dependent on environment, which in turn means that any ranking of bond strengths implicitly assumes some environment. I could have used, as you note, a host of other examples rather than a biological process with multiple intermediates etc. Since I stand by my bigger point in that comment, and it would be a substantial edit to correct the details, I’ll hope you’re helpful corrective will suffice if a fellow redditor wishes to dig into the details. Thanks.
I don't really think the comment about context-dependency answers your question. They're right that energy is measured relatively, but that doesn't change the fact that there is an objective standard for measuring bond energies, and that definition is the one you're most likely to encounter as you're looking up this topic. They've also missed an important point: what we usually cite as the "bond strength" of an ionic interaction is actually the strength of multiple ionic interactions, while that for a covalent bond is for one discrete bond.
Context is still important though - namely, the context in which ionic "bonds" form compared to that in which covalent bonds do. Covalent bonding is a fundamentally quantum mechanical phenomenon that occurs between pairwise electron interactions. This means that if two electrons participate in one bond, they cannot form another without that first one being broken. Bonds do interact, and electrons do move around, but we can still model a given molecule as having a discrete number of covalent bonds. Consequently, each atom has a characteristic maximum amount of covalent bonds that it can (stably) form.
Ionic "bonds", however, are longer-range electrostatic interactions that do not involve significant overlap of electron energy orbitals. You don't need to invoke quantum mechanics to understand ionic bonding (at least at a basic level), since ionic interactions can be accurately modeled as classic charged-particle interactions of the sort you learn in an introductory physics course on electromagnetism. In contrast to the situation of covalent bonding, the number of ionic "bonds" that an ion can participate in is limited only by the available space around it. This is why we do not refer to ionic solids as "molecules" - in an ionic lattice, there is no clear way to define where one "molecule" ends and another begins. The formula "NaCl" is a molar ratio expressing the fact that a lattice of sodium and chloride will have a 1:1 ratio of those two ions.
How, then, should we describe the strength of a typical sodium-chloride interaction? The most common method is not to describe the potential energy between a single sodium ion and a single chloride ion in a vacuum, but instead to talk about the energy present per mole of lattice. Note that this is different from how we define covalent bond energies, which is per mole of bond. We define the molar lattice energy as the energy required to completely dissociate a mole of an ionic solid. For NaCl, this value is 786 kJ/mol (though this is usually expressed as a negative number, reflecting the energy liberated when the lattice forms from gaseous components), which is higher in magnitude than all but the strongest covalent bonds. It's certainly higher than any single covalent bonds, but it's not really reflecting the strength of one bond - it's more like six bonds added together. Each of the Na-Cl electrostatic interactions is weaker than the typical covalent bond, but the summed strength of all the electrostatic interactions between a sodium ion and the six chloride ions immediately surrounding it in a sodium chloride lattice is far higher.
Note that I'm simplifying a lot here. Ionic and covalent interactions are both emergent phenomena of the same underlying quantum-mechanical processes. These are human-made definitions, and there may be ambiguous edge-cases between them. But if we take concrete examples or ask specific questions like, "If ionic bonds are said to be stronger than covalent bonds, then why are ionic bonds (but not covalent ones) often easily broken in water?" we can find some relevant differences between these bonding behaviors that allow us to answer those questions.
Yes I certainly don't think your fundamental point is wrong. I don't think it actually fully clarifies the commenter's particular misunderstanding, but it is relevant and a good contribution.
To summarize where I think OP's confusion arises from: ionic "bond strength" figures are actually molar lattice energies and therefore reflect the strength of multiple bonds. Molar energies for covalent bonds reflect the strength of individual bonds. I've added a top-level comment explaining that point, which is how it was clarified to me in undergrad chemistry.
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[deleted] t1_iu43k4a wrote
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