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Recently, I was telling my friends about the violent reaction that takes place when you throw potassium into water. Soon after, a friend of mine claimed that lithium would react more violently than potassium.

I disagreed with him, because potassium is more electropositive than lithium and thus more reactive.

My friend claimed lithium to be more reactive than potassium due to its position in the reactivity series of metals:

$$\ce{Li > Cs > Rb > K > Ba > Sr > Ca > Na > Mg}$$

Then we found out that potassium reacts indeed more violently in water. But what about his argument? Why isn't he right?

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For the reaction,
$$\ce{M -> M+ + e-}$$

the heat liberated is highest for lithium owing to its high negative $E^\circ$ value so one would think that the reaction must be most vigorous.

The reason behind the more violent reactivity of potassium rather than lithium lies in kinetics and not in thermodynamics.

No doubt, maximum energy is evolved with lithium but the vaporization and ionization will also consume maximum energy (the melting point and ionization energy of lithium are the highest) and so the reaction proceeds gently. On the other hand, potassium has a lower melting point and ionization enthalpy. The heat of reaction is sufficient to melt it. The molten metal spreads over the water and exposes a larger surface to water. Also, the hydrated radius of lithium is the greatest out of all alkali metals. This reduces the ionic mobility which in turn reduces the speed of the molten metal.

That's why potassium gives a more violent reaction with water.


Reference:

  • Kumar, Prabhat Conceptual Inorganic Chemistry; Shri Balaji Publications: Muzaffarnagar, U.P., 2014.
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  • $\begingroup$ Could you please clarify this line, and its relevance with the question? "This reduces the ionic mobility which in turn reduces the speed of the molten metal." $\endgroup$ – William R. Ebenezer Jul 9 at 15:18
  • $\begingroup$ I dunno what's there to elaborate. $\endgroup$ – Avnish Kabaj Jul 9 at 15:29
  • $\begingroup$ Kindly refer to my first comment, first line. $\endgroup$ – William R. Ebenezer Jul 9 at 16:37
  • $\begingroup$ And clarify, not elaborate, please. $\endgroup$ – William R. Ebenezer Jul 9 at 16:55
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More than 30 years ago in high school chemistry, our teacher put a flake of Sodium in a beaker filled halfway with water and the Sodium flitted around the beaker like a little motorboat letting out small wisps of smoke.

Once that was done he put in a small flake of Potassium. It reacted much more violently. The potassium shot around the beaker with smoke and sparks.

The sparks came as a result of the reaction being so energetic that the heat of the reaction ignited the free hydrogen. (2K + 2H2O → 2KOH + H2)

The way he explained it was that the single electron in the outer shell of Potassium was further from the nucleus than Sodium (and of course even further from the nucleus than Lithium) and was thus freed from the nucleus' electrostatic bond much easier.

This might not be the technical answer but it was easy for me to understand and remember all these years later.

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Lithium is very reactive, but reactions with water tend to proceed at a somewhat consistent rate. Sodium and potassium are more prone to concentrate the reaction exclusively on outward projections in such fashion as to make a lump become more and more spherical, until there are no more outward projections, whereupon they will suddenly shoot out spikes in what is called a "coulomb explosion" which causes a massive increase in surface area and consequently reaction rate. Search "sodium coulomb explosion" on YouTube for more information.

I don't know the exact details of why lithium doesn't go "bang", but I suspect that's because--unlike sodium and potassium--the will reaction proceeds in some measure on the entire surface of a lump. The reaction may be concentrated on outward projections, but not exclusively as is the case with sodium or potassium, and thus the conditions don't arise that would force the lump to shoot out spikes.

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The rapid kinetics of the reaction of sodium and potassium with water are well explained by the spikes formed and the "Coulomb explosion". But the unexpectedly slow reaction of lithium in water is still a puzzle.

I have watched a small piece of lithium wire (~2 grams) reacting in about 100 mL of water. The reaction is quite rapid at first, then slows down, and then almost stops. The major changes in the system (besides the evolution of hydrogen) are a moderate temperature rise and a pH increase to about 14. Lithium hydroxide is not very soluble (17%, hot). The temperature could not possibly get hot enough to melt lithium (mp = 180 C), whereas sodium melts at 98 C and potassium at 63 C. I watched an irregular piece of sodium (~20 grams) bubble in water (initially cold) until it melted and became globular - just before it exploded!

Lithium is reminiscent of magnesium in water, where magnesium develops a protective oxide/hydroxide coat which inhibits further reaction with water unless the water is very hot. It is worth noting that metals that corrode in water may be inhibited in alkaline media (like iron in concrete). Sodium and potassium are excepted from this inhibition because of the great solubility of their hydroxides and the possibility of melting.

Lithium may start out moderately fast, but slow down as the pH rises, and overall seem a lot slower than sodium or potassium because it will not melt in water and has a solid surface that becomes inhibited by the high pH after some reaction.

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Potassium is higher up in the reactivity series. Reactivity of metals increases as more shells are added. It becomes easier for electrons to be lost, which actually influences the reactivity of the metal.

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