# Do Na or Li need heat in order to react with Cl₂?

I've read how putting a burning piece of sodium in chlorine would produce white NaCl fumes. However, I can't find out whether Na or Li can react with Cl₂ without heating or burning.

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From personal experience, Na(s) and Cl2(g) react at room temperature without added energy. Li(s) and Cl2(g) should do so also.

Consider the following heat of formation data harvested from the NIST Chemistry Webbook:

For NaCl: $$\ce{Na(s) + } \frac 1 2\ce{Cl2(g)->NaCl(s)}\quad \Delta H_f^o=-411.12 \; kJ/mol$$

For LiCl: $$\ce{Li(s) + } \frac 1 2\ce{Cl2(g)->LiCl(s)}\quad \Delta H_f^o=-408.27 \; kJ/mol$$

These data suggest that both reactions should occur without added heat. However, the rate of reaction is not determinable from these data. The reaction between Li(s) and Cl2(g) will likely be slower.

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Just what I needed - thanks! (: –  egasimus Jun 11 '12 at 12:16
In general, the NIST Chemistry Webbook is a great resource for physical constants, phase change and thermochemistry data, and spectroscopy (IR, MS, but not NMR). –  Ben Norris Jun 11 '12 at 12:18
Presumably the thickness of the surface hydroxide/carbonate coating that forms in air will affect whether any particular metal sample will react at room temperature. –  Aesin Jun 11 '12 at 12:47