Came across this question in a book:

Among aluminum, zinc, tin and magnesium, which metals (if any) will react with an aqueous solution of sodium hydroxide to displace dihydrogen?

The answer to the question is : Aluminum, zinc and tin.

Their reasoning?

Under normal conditions, magnesium metal does not react with an aqueous solution of sodium hydroxide.

Given that the book does not go on to explain the rationale behind this one-liner, naturally my question would be: "Why not?"


Why doesn't magnesium react with an aqueous solution of sodium hydroxide and release dihydrogen?

[@paracetamol thinks]

Initially, I was under the impression that they were talking about the oxide layer that forms over magnesium metal (the very reason we buff the surface of magnesium ribbons with sandpaper before igniting it), which would give the reaction a slow start. Then, the statement could've been: "Magnesium metal does not immediately react with aqueous sodium hydroxide (owing to the presence of an oxide layer)".

But earlier on in the chapter, there is explicit mention of the reaction between aluminum metal ( which also naturally come with an oxide coat) and aqueous sodium hydroxide, to form sodium aluminate and dihydrogen. There was no mention of an oxide layer hampering the reaction.

[@paracetamol thinks harder]

I then (cursorily) looked at the Electrochemistry aspect of the reaction; because whenever we're asked if a metal reacts with a mineral acid (to liberate dihydrogen), we're expected to flip open our table of standard electrode reduction potentials and say:

Metals having a higher electrode reduction potential than that for the reaction $$\ce{2H^{+} -> H_{2}}$$ (which is arbitrarily taken as zero volts) will be able to liberate dihydrogen from the solution. However, metals such as copper, silver, mercury, etc which have a lower standard electrode reduction potential than that for $H^+$ will not liberate dihydrogen.

When I compared the standard electrode reduction potentials of aluminum, zinc, tin and magnesium, I've observed that all of them have electrode reduction potentials that are higher than the electrode reduction potential for $H^+$. Meaning, that all those metals (magnesium included) should be able to displace dihydrogen from the solution, at least, from the high-school electrochemistry perspective. (Yes, I know the textbook mentions an alkali in the question, but what I'm trying to point out is, that since the four metals have higher electrode reduction potentials that $H^+$, then they ought to behave the same way. So if aluminum, zinc and tin can liberate dihydrogen (as the textbook claims) then why not magnesium too?)

[@paracetamol Googles]

I Googled my question, and arrived at this Quora post. Same question as mine, but it's generated such lackluster answers, that they would only be worthy of...well...Quora. O:)

[@paracetamol bamboozled]

I can't for the world figure out why magnesium doesn't react with aqueous sodium hydroxide (if it doesn't). Anyone?

  • 1
    $\begingroup$ NaOH lowers the concentration of H+ in solution, so there is less H+ to begin with and considering that magnesium slowly reacts with water, I'm not astonished that it doesn't react with basic water. $\endgroup$ – AS_1000 Aug 27 '17 at 13:34
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    $\begingroup$ Magnesium does react with hot water (not so much with cold water). I see no reason for it to not react with hot sodium hydroxide (though it might happen slower). Also, magnesium is a very active metal and in mixture of fine powders reacts (actually, burns) with things like sand, many metal oxides, and for example, hydrate of barium chloride. Pretty sure it should burn with powder of solid sodium hydroxide. $\endgroup$ – permeakra Aug 27 '17 at 21:46
  • $\begingroup$ You need to use the Nernst equation to adjust the potentials accordingly to the much higher pH value. The potential of hydrogen gas and protons equalling zero is only true at pH zero. $\endgroup$ – Jan Oct 18 '17 at 12:48
  • $\begingroup$ The reason why aluminium is so reactive to NaOH is because it destroys the oxide layer (the oxide is acidic). The oxide on the Mg might well not have this character. Worth a look. $\endgroup$ – matt_black Dec 17 '17 at 15:29

For a metal to react with sodium hydroxide solution it's oxide must be amphoteric, not just basic. Alkaline earth metals are called that for a reason; except for beryllium the oxides are purely basic (and in fact moderate to strong in basic strength) in water, so sodium hydroxide won't make those metals react.

Magnesium, as a typical alkaline earth metal, can be made to react with acids displacing hydrogen, even when the "acid" is a transition metal salt that hydrolyzes slightly like cupric sulfate.

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    $\begingroup$ But we are looking at the properties of the metal. How are the properties of their oxides relevant to the discussion? $\endgroup$ – Tan Yong Boon Nov 17 '17 at 16:04
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    $\begingroup$ Metals cannot react with aqueous hydroxide ions. Their oxides/hydroxides can, if these are amphoteric. $\endgroup$ – Oscar Lanzi Nov 17 '17 at 18:53

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