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I’m really stuck on this problem. This involves dissolving Borax: $$\ce{Na2B4O7•10H2O(s)<=>2Na^+(aq) + B4O7^2-(aq) + 10H2O}$$

$$\begin{align}\Delta H&=\pu{109 kJ/mol} \\ \Delta S&= \pu{340 J/mol\:K} \\ K_\mathrm{c} &=0.047 \;@ \; \pu{25 ^\circ C} \end{align}$$

  1. If the starting concentration of $\ce{B4O7}$ is $\pu{0.17 M}$ and the sodium ions are $\pu{0.34 M}$ at $\pu{25^\circ C}$, which direction will the reaction proceed? I said forward because Q < K

  2. At what temperature would a mixture of these same starting concentrations be a saturation solution (i.e. be at equilibrium) And that’s where I’m stuck

  3. At what temperature would a mixture under standard conditions ($[\ce{B4O7}]=\pu{1M}$, $[\ce{Na}]=\pu{1 M}$) be a saturation solution (i.e. be at equilibrium) And I’m stuck here again

I understand that crossover temperature plays a role ($T_c=\Delta H/\Delta S$) but I don’t understand how to work with given concentrations.

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  • $\begingroup$ Hint: Use vant Hoff equation. $\ln(keq(T2)/keq(T1))= ∆H/R [\frac{1}{T1}-\frac{1}{T2}]$ and the equation you got. $\endgroup$ – user600016 Jun 23 at 9:21

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