[Oh Dear!]

I'm not sure how to phrase this so it might sound a bit rambling. 

It's my very limited understanding that the formation of, say, magnesium chloride involves the establishment of ionic bonds between Mg and two separate Cl atoms.  However, in nature, Cl usually comes in diatomic form and, via covalency, each individual atom already presents as having 8 electrons in its outer shell.  Into this structure busts magnesium, donating its "excess" electrons and so (i'm assuming) rendering the covalent bond redundant. I'm having trouble visualising the mechanism by which this happens.  At what point (if any) during the process Mg + Cl₂ MgCl₂ does the Covalent bond disintegrate and how does this occur?

[Aldebaran]

I always have a bit of a problem with questions like this without knowing what level of knowledge the questioner already has and whether they are in formal study classes or just doing their own thing. Anyway I’m going to give you a simplified answer with some points you can follow-up using the internet or a suitable textbook.
Chemical reactions take place in accordance with thermodynamics. As your question states in a chemical reaction bonds are broken in the reactants and new bonds are made in the products. This all involves energy. Whether the bonds are regarded more as ionic or covalent in character isn’t really the issue. If it is energetically favourable to react then the reaction can occur – in other words it is feasible (chemists often use the word spontaneous rather than feasible). Thus it can happen.
Energy considerations not only involve enthalpy (heat energy) but also entropy. The feasibility of the reaction is summarised by the equation:  ∆G = ∆H - T∆S.
Delta G is the change in Gibbs free energy, delta H is the reaction enthalpy change, T is temperature in Kelvin and delta S is the change in entropy of the reaction.
If ∆G is negative the reaction is spontaneous; thus it can happen. If it is positive it can’t happen.
However just because ∆G is negative does not mean the reaction will happen in any sensible time scale because there may be an energy barrier to be overcome…the activation energy.
If you don’t know about enthalpy, entropy and free energy in this context then you need to access  internet research or a high school text book pitched (in the UK) at 16-18 year olds.
Anyway I hope this answer helps you get started on basic chemical thermodynamics.

[Corribus]

@OP

It seems like you're asking about the reaction mechanism, loosely defined as the various steps that occur to get from point A (the reactants) to point Z (the products). The fact is that most reactions don't occur as a single step, as a simple reaction equation would seem to indicate. Even in solution, where all the reactants are well separated, mechanisms tend to be complicated, with many steps that must occur in sequence. Figuring it out requires complicated experiments and, often, computation modelling. Also, each reaction must be approached separately. Sometimes even the mechanism can be different if you so much as change the amount of the reactants present, the temperature, whether or not there is light, and so on. Things get even more complicated in the solid state (such as the one you have asked about) because reactions really only happen at the surface. Therefore solid-state reactions invariably involve things like diffusion, lattice mismatches, surface energy/geometry, atomic organization on the exposed crystal surface, and even silly mechanical effects like surface layer shedding. You can always search the literature to see what is known about a specific reaction like the one you mentioned, but many seemingly simple reactions have not been studied in this kind of detail to really understand what is going on at the molecular level. It may be this question can only be answered is superficial detail.

If ∆G is negative the reaction is spontaneous; thus it can happen. If it is positive it can’t happen.

I understand you may be trying to simplify things for a high school audience, but take care against painting a misleading picture when simplifying. We can't really talk about the Gibbs energy or spontaneity divorced from the concept of equilibrium, since delta-G means nothing without the standard Gibbs change (delta-G⁰) as a reference point to define the system's equilibrium point. This is because spontaneity only refers to whether the reaction will proceed in a forward direction (as it is written) toward the equilibrium point from a specified starting point, with both the equilibrium point and the starting point defined by chemical potentials under the starting and final conditions. (In dilute solution, these chemical potential are usually understood to be proportional to concentration). If you want to simplify it, better to say that delta-G/spontaneity only tells us whether the system will (given enough time) evolve to a state that favors products compared to what we started with.

When we speak of reaction spontaneity with broad terms like "can" or "cannot" happen, students may get confused by situations like this:

Consider a reaction A --> B.

Asking whether this reaction is favorable or unfavorable, spontaneous or not spontaneous, or will or will not happen, is asking an incomplete question. First, what happens to a system comprised of A and B depends on how much A or B is present initially. Second, it also depends on what equilibrium looks like. If the thermodynamic potentials are such that at equilibrium A and B are of equal concentration, the spontaneity of a system comprised of A and B and moving in the direction A --> B will depend on whether there is initially more A than B (spontaneous) or B than A (not spontaneous). Importantly, if you start with all A, the reaction will ALWAYS be spontaneous, regardless of what the equilibrium point is. Presuming that the equilibrium point always allow for some minuscule amount of B, in that sense there is NO reaction that "cannot happen" if one starts with 100% A because by definition delta-G will always be negative. That does NOT mean that a lot of product will form, or that it will form quickly. Nevertheless, by definition, it is still be formally spontaneous. In other words, asking whether a reaction can or cannot happen is not practically relevant. All reactions formally "happen" under the assumption that the equilibrium point is somewhere in the middle. It is better to ask not whether a reaction happens, but rather how much product is formed when it is done happening. This is an important point for a student to understand as consequence of the concept of equilibrium, which is central to all of chemistry.

Anyway, none of this has much to do with the opening post about mechanisms, and I mean no offense by picking at this. I comment only since spontaneity is a concept that a lot of people try to simplify for people new to chemistry, but it is often done in such a way as to cause more problems understanding the concept than saying nothing at all.

[Aldebaran]

@ Corribus . Fair comment and certainly no offence taken. In fact in my first draft of an answer I'd included a paragraph about equilibrium but then felt it might be getting a bit complicated for a first response. Maybe should have left it in on reflection.

[gavindor]

well, i'm no expert but it seems you're saying that a Chlorine molecule is stable in and of itself.. Why is it going to bond with Magnesium.

I think this picture will help.

Draw the Lewis Structure of MgCl2 (magnesium chloride)
By chemistNATE
2:30



So you see if there was a Chlorine molecule, then that  Chlorine molecule is split up.. into two Chlorine atoms. Each Chlorine atom interested in one electron. And Magnesium  as you see from that diagram,  gives one electron to one Chlorine, and one to the other Chlorine.

I suppose you could ask , supposing there was a Chlorine molecule, why would the Chlorine molecule split in the first place.. I don't know for the specific case of magnesium chloride, and there could be different ways e.g. maybe heat, a molecule can split under sufficient heat. This link https://www.vedantu.com/question-answer/a-name-the-salt-formed-by-the-reaction-between-class-11-chemistry-cbse-5fbac8e3ffe88b10a64aa231   mentions MgCl2 forming as a result of the reaction Mg(OH)2 + 2HCl --> MgCl2 + 2H2O

So that reaction formed MgCl2, but didn't involve a Chlorine molecule splitting up.

I was recently looking into Where does NaCl come from, it's a similar puzzle to yours 'cos Chlorine occurs naturally as a molecule. I read NaCl is in sea water. But how does it get there. Apparently it was only found out recently something about how it gets into the sea water. https://www.youtube.com/watch?v=SPF6cSan6tc   though it doesn't mention the reaction.