I would have said that the answer was no, because lithium is less reactive than potassium is, so you would expect the reaction to go the other way: with potassium and lithium chloride reacting to produce lithium metal and potassium chloride.
However, the YouTube video Cody'sLab — Potassium Metal From Bananas! appears to show that the answer is yes.
Can anyone explain this result, which contradicts what I thought I had learned in high school chemistry?
Note that element electronegativity $\chi$ is a combined indicative quantity from covalent bond context, so its comparative values do not always apply in scenario analysis what happens or should be preferred.
Reactivity (kinetics) and tendency of reaction outcome (thermodynamics) are independent. TNT is much more reactive than coal, but released specific energy during coal burning is 10 times bigger than during TNT explosion.
Potassium reacts more violently with water than lithium. But lithium has more negative standard redox potential (which applies in water context only). That is because higher ionization energy is more than compensated by more negative hydration enthalpy of $\ce{Li+(aq)}$, compared to $\ce{K+(aq)}$, being a smaller ion with stronger electrostatic potential gradient. Hydration enthalpy contributes to the reaction Gibbs energy of $$\ce{M(s) + H+(aq) -> M+(aq) + 1/2H2(g)}$$ and therefore to the metal standard redox potential.
$\ce{Li}$ is less electropositive/more electronegative than $\ce{K}$ and has higher ionization energy than $\ce{K}$. But $\ce{LiCl}$ has higher lattice energy ( more negative lattice formation enthalpy) than $\ce{KCl}$.
$\ce{K}$ boiling point is near $\ce{KCl}$ melting point, while $\ce{Li}$ boiling point is much higher, so $\ce{K}$ boils away, what supports reaction.
There is equilibrium $$\ce{Li(\ell) + KCl(\ell) <=> K(g) + LiCl(\ell)}$$
with potassium vapors being eliminated by condensation on colder surfaces of the apparatus, used in the video.
Thermodynamic parameters are shown in the table below. higher ionization of lithium is more than compensated by the higher lattice energy and pushing the equilibrium as potassium vapor partial pressure is being decreased by vapor condensation.
| Compound | Melting Point | boiling point | Ioniz. energy | Lattice energy | Cl EA |
|---|---|---|---|---|---|
| $\ce{LiCl}$ | $\pu{605–614 °C}$ | $\pu{1382 °C}$ | $\ce{829 kJ/mol}$ | $\pu{349 kJ/mol}$ | |
| $\ce{KCl}$ | $\pu{770 °C }$ | $\pu{1420 °C }$ | $\pu{698 kJ/mol}$ | $\pu{349 kJ/mol}$ | |
| $\ce{K}$ | $\pu{63.5 °C}$ | $\pu{759 °C}$ | $\pu{419 kJ/mol}$ | $\pu{349 kJ/mol}$ | |
| $\ce{Li}$ | $\pu{180.50 °C}$ | $\pu{1330 °C}$ | $\pu{520 kJ/mol}$ | $\pu{349 kJ/mol}$ |
In the video, you can hear the narrator talk about “potassium vapor” condensing to form the product. Based on that, I’m going to assume that he’s heating the vessel above the boiling point of potassium, and at a slightly higher temperature (above around 770 °C) , potassium chloride and lithium are both liquid.
At this point, I would think that all that needs to happen is for the molten lithium to reduce some of the $\ce{K+}$ in the liquid salt to potassium metal, some of which would vaporize, escape and condense, as Cody seems to describe.
Since the process presumably takes place above the boiling point of potassium, the equilibrium would not even need to favor the formation of $\ce{LiCl}$ for the reaction to progress, although as a commenter before me pointed out, the reduction potential of $\ce{Li+}$ is lower than that of $\ce{K+},$ indicating that lithium metal could reduce $\ce{K+}$ to potassium.
