Answers: 1
, assuming that this acid is monoprotic. 
denote this acid.
.![[{\rm HA}] = 0.730\; \rm mol \cdot L^{-1}](/tpl/images/1377/5071/8d084.png)
.
of that was dissociated, the concentration of both 
and 
(conjugate base of this acid) would become: 
. 
.![[{\rm HA}]](/tpl/images/1377/5071/3042a.png)
, ![[{\rm H}^{+}]](/tpl/images/1377/5071/23c55.png)
, and ![[{\rm A}^{-}]](/tpl/images/1377/5071/9f958.png)
denote the concentration (in 
or 
) of the corresponding species at equilibrium. Calculate the acid dissociation constant 
for ![\begin{aligned}K_{\rm a} &= \frac{[{\rm H}^{+}] \cdot [{\rm A}^{-}]}{[{\rm HA}]} \\ &= \frac{(0.09125\; \rm mol \cdot L^{-1}) \times (0.09125\; \rm mol \cdot L^{-1})}{0.63875\; \rm mol \cdot L^{-1}}\\[0.5em]&\approx 1.30 \times 10^{-2} \end{aligned}](/tpl/images/1377/5071/efa8b.png)
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In a 0.730 M solution, a weak acid is 12.5% dissociated. Calculate Ka of the acid....
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