# How does an anode metal come into solution if it cannot react with an electrolyte?

From what I understand, an anode material is chosen so that it does not react with the anode electrolyte solution.

I was just wondering, then, that when you have a galvanic cell which has a Mg metal anode submerged in sodium chloride solution, how the magnesium metal gives two of its electrons and come into solution (as magnesium ion) even though it does not react with (and cannot react with) sodium chloride?

What I mean is, in a typical single displacement reaction, there is a metal that reacts with the solution and comes into solution. For example, with $\ce{Zn}$ and copper sulfate, the zinc replaces the copper and comes into solution as zinc ion. However, magnesium doesn't react with sodium chloride (because the activity series predicts it can't). How then can we say that we have magnesium ion in solution?

Is it that the metal is more stable in its ion state and thus readily gives up its electrons?

Traditionally, the anion is the one responsible for giving electrons: the chloride ion in the sodium chloride solution should come into contact with magnesium anode and give its electrons, which the magnesium metal should then carry. And thus the reaction should be

$$\ce{ 2Cl- \!(aq) -> Cl2\!(g) + 2e-}$$

and not

$$\ce{Mg\!(s) -> Mg^2+\!(aq) + 2 e-}$$

However, my books don't say this: they say that magnesium gives up two of its electrons and enters the solution.

What am I understanding wrong?

• For understanding the reason, you must have knowledge of electrode potentials and electrochemistry – Amritansh Singhal Sep 26 '16 at 7:02
• RE: From what I understand, anode material is chosen so that it does not react with the anode electrolyte solution. You don't want an electrode which would spontaneously react by itself in the solution. For instance a sodium electrode in a galvanic cell with a water medium would decompose in the water to sodium hydroxide and hydrogen gas. – MaxW Sep 26 '16 at 7:29
• @MaxW Magnesium anodes are used quite frequently in electrochemistry. They do react with the electrolyte, but quite slowly; slowly enough that they are perfectly safe to use, and last sufficiently long to be of practical value. See, e.g., here and here. – hBy2Py Dec 25 '16 at 11:21

I was just wondering then that when you have a galvanic cell which has a $\ce{Mg}$ metal anode submerged in sodium chloride solution, how the magnesium metal gives two of its electrons and come into solution (as magnesium ion) even though it does not react with (and cannot react with) sodium chloride?

Your misunderstanding stems from assuming that the electrode reaction has to be the same sort of single-displacement redox reaction that would occur with no electrodes present at all.

At a dissolving magnesium anode$^\dagger$, the only reaction that needs to occur is the one you wrote involving magnesium:

$$\ce{Mg\!(s) -> Mg^2+\!(aq) + 2e-}$$

The potential applied to the electrode pair is what enables the above reaction to occur without a paired reduction reaction at the anode surface. Instead of being taken up in another chemical reaction, the electrons are drawn into the anode and on into the connected electrical circuit. Those electrons are then "pushed" into the cathodic reaction, whatever it might be.

From what I understand, anode material is chosen so that it does not react with the anode electrolyte solution.

This is true, but does not have to strictly hold. MaxW is right: $\ce{Mg}$ does react directly with water. But, the reaction is slow enough that it's perfectly safe, and that one can conduct useful electrochemical operations even while it's going on (e.g., see here and here).

Is it that the metal is more stable in its ion state and thus readily gives up its electrons?

Despite what one reads from a table of standard reduction potentials or a metal reactivity series, in practical electrochemistry, the above is true for most metals, except for the noblest of the activity series. In my experience, it's usually a lot easier to dissolve metals than to electrodeposit them. You may have to use additional chemistry to make it happen (e.g., including $\ce{HF}$$^\ddagger$ to strip the passivating oxide from $\ce{Ti}$, $\ce{Nb}$, $\ce{Ta}$, etc.), but once the metals are oxidized/dissolved they tend to stay that way.

$^\dagger$ Operating at $100\%$ faradaic efficiency. Other reactions that might occur in this system, depending on the applied potential, include the chloride oxidation reaction you noted and water electrolysis to oxygen gas.
$^\ddagger$ $\ce{HF}$ is very dangerous, do not use at home!

• can you please elaborate on the sentence "The potential applied to the electrode pair is what enables the above reaction to occur without a paired reduction reaction at the anode surface." – user510 Jan 24 '17 at 4:44
• @user510 The potential applied to the two electrodes allows the reduction and oxidation reactions to occur separated in space. Thus, in theory, in this system it is possible for the only reaction at the anode to be magnesium dissolution; the corresponding reduction reaction will take place at the cathode, instead of also at the anode. Without the electrodes and the potential, the reduction reaction would have to occur at the surface of the magnesium. – hBy2Py Jan 24 '17 at 5:04
• if the potential applied to the electrode pair is what enables the above reaction to occur, then why is there resistance in battery even with weak electrolytes (i.e. I could use any metal and it would give up its electron). – user510 Jan 24 '17 at 17:27
• Also, why is it the potential that allows magnesium to give its electron? – user510 Jan 24 '17 at 17:27
• @user510 if the potential applied to the electrode pair is what enables the above reaction to occur, then why is there resistance in battery even with weak electrolytes (i.e. I could use any metal and it would give up its electron). I'm sorry, this question doesn't make sense to me. – hBy2Py Jan 24 '17 at 18:00

Hmmm. Electrochemistry can definitely be confusing. Context is quite important. Anyone who makes a broad claim that anodes don't react is just plain wrong. It depends on the set-up, including the applied current (and/or voltage). Anyone who makes the statement that in measuring electrolytic reactions both electrodes should be inert is going to be right a lot (that is, more often than someone will be able to find exceptions). Magnesium is a VERY reactive metal! The only reason it doesn't explode into flame upon contact with water is that the coating of MgO slows down the reaction. Anyone who thinks putting a piece of Mg metal into a salt solution is a good idea, needs to be very well aware of the hazards. When thinking about electrochemical reactions, you have to be aware of all of the possible oxidations and reductions that can go on. Running current through pure water can generate H2, O2, as well as several other compounds (in trace amounts, usually). As you probably know, the concentration of ions in pure water is very low [H] = [OH] = 1E-7. Which means that the amount of current (work) you can do is quite limited. So, NaCl (or some other electrolyte) is added. This complicates things since potential reactions (is that punny or what?) include Na(+)→Na° (not going to happen), Cl(-)→Cl2 ( one way to make chlorine gas), and things like Cl(-) + H2O→ OCl(-)+... (and a bunch more). Only if you know what the reaction potentials are for all of the possibilities can you make good predictions. Some electrode materials can act as catalysts (Pt electrode, for instance). These lower the EP necessary while remaining "inert". It's been so long since I did any electrochemistry (decades) that I don't recall which electrode is the "anode" and which is the "cathode"...doesn't it depend on the type of cell involved? I've forgotten. I do remember that Leo says Ger (Leo is another name for lion, and Grrr is the English for the sound a threatened (or threatening) animal makes.) LEO = Loss of Electrons is Oxidation. GER = Gain of Electrons is Reduction. You have written one oxidation and one reduction. In a driven electrolytic cell, they will NOT occur at the same electrode!!!!!!!!

• Magnesium will react with water, but it is nearly as reactive as sodium. Even if you cleaned off the typical oxide coating it still won't react violently but very slowly. Any such reaction though means that magnesium would be a poor choice for a battery that used any water in the electrolyte. – MaxW Sep 26 '16 at 16:47
• @MazW do you mean magnesium submerged in aqueous solutions is bad? – user510 Sep 26 '16 at 19:53