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I reacted about $0.25 \,\mathrm{g}$ of magnesium, sodium and lithium in $20 \,\mathrm{ml}$ of $2 \,\mathrm{M} \,\ce{HCl}$ at room temperature. This was all done in a $100 \,\mathrm{mL}$ beaker.

I found the biggest temperature change in the reaction mixture with lithium, followed by the reaction mixture with magnesium and finally sodium.

I measured the temperature change over the course of the reaction with a Vernier temperature probe and accounted for the heat loss by interpolating a fitted line based on the heat loss signature after the reaction has completed.

I calculated the heat of the reaction and divided by the number of moles and get this:

  • Li reaction: $120 \,\mathrm{kJ/mol}$
  • Mg reaction: $91 \,\mathrm{kJ/mol}$
  • Na reaction: $33 \,\mathrm{kJ/mol}$

I did notice the evolution of a lot of gas when reacting $\ce{Na}$ with the $\ce{HCl}$. As expected a flame was produced with the sodium reaction.

My questions are:

  • Why is it that the reaction that apparently seemed to be the most vigorous produced the least temperature change?

    • I was initially thinking that the sodium had a lower temperature change as it reacted so vigorously that the hydrogen gas produced had more kinetic energy, and as it leaves the system, the temperature of the system does not change by much. But I figured this explanation to be incomplete as lithium is supposed to be more reactive than sodium, so if this was the case I would have an even lower temperature change for lithium?
  • Why is it that some reactivity series have lithium at the very top while others just omit it entirely or put it below sodium and potassium? Who is right?

I guess what I am really asking is for ideas on where I went wrong, because this data does not seem to be right.

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