Question

What volume of an HCl solution with a pH of 1.3 can be neutralized by one dose of milk of magnesia? If the stomach contains 220 mL of pH 1.3 solution, will all the acid be neutralized?

Milk of magnesia is often taken to reduce the discomfort associated with acid stomach or heartburn. The recommended dose is 1 teaspoon, which contains 4.00×102mg of Mg(OH)2.

Answer 1

Consider,

`2HCl + Mg(OH_2) rightleftharpoons MgCl_2 + 2H_2O`

We must first assume that strong acids totally dissociate.

Therefore, we can obtain the concentration of hydrochloric acid in the stomach,

`10^(-"pH") = [H^+] approx 5.01*10^-2"M"`

Moreover, we're given data to support the neutralization of excess hydrochloric acid via milk of magnesia, which contains `Mg(OH)_2`.

Now, one molar equivalent of hydrochloric acid will be neutralized by the moles of hydroxide ions that make it to the stomach acid solution and dissociate from `Mg(OH)_2`.

`4.00*10^2"mg" * "g"/(10^3"mg") * "mol"/(58.3"g") * (2OH^-)/(Mg(OH)_2) approx 1.37*10^-2"mol"`

of hydroxide ions will dissociate into solution.

There are,

`0.220"L" * 5.01*10^-2"M" approx 1.1*10^-2"mol"`

of protons in the solution described.

Hence, since `n_(H^+) < n_(OH^-)`, all of the excess acid will be neutralized.

Note: you can solve this by doing a BCA table for the reaction, but this professor was pulling a trick on you—two hydroxide ions dissociate per mole of `Mg(OH)_2`, which a lot of beginning chemistry students would miss.