A question in my college test was whether $\ce{MgSO4}$ is neutral, acidic or basic. I was told to solve this kind of problem by hydrolyzing any ions that do not "come from" a strong acid or a strong base, and if nothing is hydrolyzed, then the salt is neutral. So by this logic, as $\ce{Mg^2+}$ comes from $\ce{Mg(OH)2}$ and $\ce{SO4^2-}$ comes from $\ce{H2SO4}$, I answered "neutral".

My professor marked it as wrong, and said that the salt $\ce{MgSO4}$ is basic. How can you explain this?

  • 2
    $\begingroup$ Hint: weak bases are usually insoluble. $\endgroup$ – permeakra Apr 20 '16 at 8:13
  • 2
    $\begingroup$ Basic makes no sense at all. The only way for the solution to be basic is for $\ce{SO4^{2-}}$ to hydrolyze to $\ce{HSO4^{-}}$. The pKa2 for sulfuric acid is 1.99, so that isn't going to happen in a solution of $\ce{SO4^{2-}}$. $\endgroup$ – MaxW Mar 1 '17 at 19:48

Looking at the following references, it seems that most agree that a water solution of $\ce{MgSO4}$ is slightly acidic.

"The pH of an aqueous magnesium sulfate solution is related to the molarity of the MgSO4. Typically, the pH is between 5.5 and 6.5 due to magnesium's affinity for hydroxide ion (OH-). As the sulfate goes into solution, hydroxide anions associate with the magnesium, increasing the relative ratio of H+ to OH-. This shift results in more acidic solutions."

"The pH of a pure magnesium sulfate solution is approximately 6.4"

"The pH of [magnesium sulfate] hydrates is average 6.0 (5.5 to 6.5)."

"pH: Neutral to litmus"

Why not perform the experiment in class, using an accurate pH meter and/or indicators (see the pH indicator chart in Wikipedia for suggestions).

  • $\begingroup$ +1 for nice answer and all the references. // I'm not so sure that the pH meter experiment would be illustrative. A very low amount of impurities could shift pH acidic or basic. So how do you obtain salt of "sufficient" purity? $\endgroup$ – MaxW Mar 2 '17 at 1:00

Aqueous solution of magnesium sulfate is going to be slightly acidic, and there is no need to even go to the lab to test it. $\ce{MgSO4}$ is a salt formed by a weak base and a strong acid (both dibasic), so there is a two-step hydrolysis of magnesium cation.

First and primary step is the formation of a basic salt:

$$ \begin{align} \ce{2 MgSO4 + 2 H2O &<=> (MgOH)2SO4 + H2SO4} \tag{R1.1}\\ \ce{2 Mg^2+ + 2 SO4^2- + 2 H2O &<=> 2 MgOH+ + SO4^2- + 2 H+ + SO4^2-}\tag{R1.2}\\ \ce{Mg^2+ + H2O &<=> MgOH+ + H+}\tag{R1.3} \end{align} $$

Second minor (negligible) step:

$$ \begin{align} \ce{(MgOH)2SO4 + 2 H2O &<<=> 2 Mg(OH)2 + H2SO4}\tag{R2.1}\\ \ce{2 MgOH + SO4^2- + 2 H2O &<<=> 2 Mg(OH)2 + 2 H+ + SO4^2-}\tag{R2.2}\\ \ce{MgOH+ + H2O &<<=> Mg(OH)2 + H+}\tag{R2.3} \end{align} $$

Here indices .1, .2, .3 denote molecular, complete ionic, and net ionic equation, correspondingly. As a result of the hydrolysis reaction, some excess amount of hydronium cations appears, therefore, the the medium is expected to be acidic.


$\ce{SO4^2-}$ is actually going to make the pH basic because $\ce{HSO4-}$ is a weak acid (not $\ce{H2SO4}$, only the first hydrogen ionizes completely). By itself aqueous $\ce{Mg^2+}$ will not form magnesium hydroxide in water. But with the $\ce{OH-}$ from the $\ce{HSO4-}$ is will form a suspension with some dissolved and some not so dissolved, but still the solution is basic.


Your Answer

By clicking “Post Your Answer”, you agree to our terms of service, privacy policy and cookie policy

Not the answer you're looking for? Browse other questions tagged or ask your own question.