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From my AP Chem Prep Book:

What should you do if you spill sulfuric acid on the countertop?

A) Neutralize the acid with vinegar.
B) Sprinkle solid NaOH on the spill.
C) Neutralize the acid with NaHCO3.
D) Neutralize the acid with an Epsom salt $(\ce{MgSO4})$ solution.

The answer given was C — stating that it should be $\ce{NaHCO3}$ because it is a weak base solution.

However, I do not understand why the strength of the base is of relevance (given the situation) and how to determine the strength of the base — at least relative to other bases.

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    $\begingroup$ A strong base would later be a problem itself (NaOH is corrosive, you can feel the saponification of your fingers if you touch it). A weak base like bicarbonate (which you can even drink with no relevant problems) is a much wiser choice! $\endgroup$
    – user32223
    Jan 5, 2020 at 7:28
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    $\begingroup$ @The_Vinz If the slippery feel is saponification, why does baking soda also feel slippery? Is it reacting with free fatty acids on the skin, or is there something potentially wrong with the "slippery" explanation? $\endgroup$
    – piojo
    Jan 6, 2020 at 4:26
  • $\begingroup$ chemistry.stackexchange.com/questions/35016/… chemistry.stackexchange.com/questions/71486/… $\endgroup$
    – Mithoron
    Feb 3, 2020 at 22:14

1 Answer 1

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First, why other options are not really the options:

A: vinegar, being a weak acid, doesn't neutralize sulfuric acid and only dilutes it;

B: solid sodium hydroxide, a strong base, does neutralize sulfuric acid, but it does so vigorously releasing substantial amount of heat per unit of time: $$\ce{2 NaOH(s) + H2SO4(aq) -> Na2SO4(aq) + 2 H2O(l)}$$ Using solid NaOH it is also tricky to guarantee it will cover the spill and there won't be any unreacted acid or excessive hydroxide (which is equally unwanted) left behind as there is little to no visual clue whether the neutralization is complete, unless you test various spots with, let's say, pH paper.

D: Magnesium sulfate solution doesn't react with sulfuric acid and only dilutes it.

Second, sodium bicarbonate solution not only neutralizes the acid

$$\ce{2 NaHCO3(aq) + H2SO4(aq) -> Na2SO4(aq) + 2 H2O(l) + 2 CO2(g)},$$

but also can be (and should be!) used in excess to assure complete acid neutralization. Once the acid is neutralized, you are left with solution of sodium bicarbonate and sodium sulfate which is only slightly basic due to cation hydrolysis. Another important advantage of using $\ce{NaHCO3}$ is visual control: once the process is complete, gas evolution stops.

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    $\begingroup$ However, the effervescence might also be a disadvantage if it's distributing fine droplets of the liquid. $\endgroup$
    – user7951
    Jan 5, 2020 at 16:19
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    $\begingroup$ Practically, I thought that. I've used (solid) bicarb on real spills of battery acid, but you really want lots to get rid of the still acidic foamy mess. But still best option compared to NaOH which leaves you either a hazardous acid or a hazardous alkali). $\endgroup$
    – Rich
    Jan 6, 2020 at 3:13
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    $\begingroup$ Avoiding the quick release of heat is also a point of emphasis IMO - you sometimes do not want to burn through the floor $\endgroup$
    – Layman
    Jan 6, 2020 at 14:41
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    $\begingroup$ The other thing you want to be careful of is creating an acid mist. Many reactions with sulfuric acid are very exothermic and result in such an energetic reaction that a rather stable acid mist is formed. As you might imagine, you do not want to inhale sulfuric acid mist. That acid is very nasty. $\endgroup$
    – Flydog57
    Jan 31, 2020 at 21:26

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