# Why is it considered acid rain with pH <5.6?

I recently read in a book that rain is considered acid rain if the pH falls below 5.6. However a substance is acidic when the pH is below 7; so why is the boundary for acid rain 5.6?

I was thinking pH between 5.6-7 would be too diluted to have an effect on limestone and other materials, and that's why the pH has to be under 5.6. But I wasn't sure if this is true so I wanted to find out.

• Because "normal rain" still is a bit acidic. – Mithoron Jul 27 at 22:11
• Actually, just the carbon dioxide the rain picks up on the way down is enough to have an effect on limestone - that's how limestone is dissolved with no human activity, and that's how all those wonderful caves with stalagmites and stalactites are formed. – Luaan Jul 29 at 10:13
• Would you consider something at a pH of 6,9 acidic? How about 6,99? What's the point of that if the substance is relatively neutral? Keep in mind pH is logarithmic, a couple of points is a major difference. A line is drawn, it had to be drawn somewhere, and this is where they put it. – Mast Jul 30 at 6:17
• @Luaan This, of course, begs the question of how those big limestone masses had managed to form in the first place. :-) – oakad Jul 31 at 1:42
• @oakad It's quite interesting, actually - rain water is acidic enough to dissolve limestone, yes. But the deposits form out of seawater, which is basic rather than acidic. And then you have life, of course. Those deposits then either get uplifted, or release their carbon dioxide as they get subducted. If the deposit gets enough rain, it will slowly be dissolved away back into the oceans (building up caves and sinkholes in the meantime). – Luaan Jul 31 at 7:22

You are forgetting an important component of the air: carbon dioxide. When it dissolves in pure water (=rain water), it makes it acidic. It is not considered that harmful.

Acid rain has a negative connotation; it is mainly caused by anthropogenic activities. The low pH of acid rain is due to sulfur oxides and nitrogen oxides and it is indeed below 5.7.

IUPACs definition "Rain with pH values < about 5; commonly results from acids formed from pollutants. 'Pure' rain water equilibrated with atmospheric CO2 and naturally occurring acids in relatively clean air usually has a pH>5."

• but why is it considered acid rain only below pH 5.7? Why not below 7? – muhammad haider Jul 27 at 19:38
• Because natural rain, which is free from all types of anthropogenic pollution has a pH < 7. You need to "distinguish" natural acidity vs. excessive acidity due man's activities. Otherwise what is the point of calling each an every rain as acid rain. Also remember there is little bit of nitric acid formed from lightning which is the plant food. That will not be termed as acid rain either. As I said, acid rain has a negative meaning. – M. Farooq Jul 27 at 21:37
• "Why not below 7?" I think this has to do with how you parse the term "acid rain". You seem to consider it to mean "an acid (anything with a pH < 7) which is raining". As the other answers have explained, acid rain means rain which is more acid than normal rain. – Philip Ngai Jul 28 at 20:48
• @M.Farooq So why choose 5.7 as a boundary value? Acidic soils having pH < 5 or even 4.5 are not uncommon, so such a rain would not harm them. Of course, such pH values are harmless for humans. – trolley813 Jul 29 at 8:10
• @trolley813, see Loong's answer. A pH of 5.67 is the border pH of CO2 in water. Soil chemistry is a complete different story. Acid rain was not named on the basis of its effect on soil by Angus Smith in 1872. He saw deleterious effects on plants and chlorophyll from "acid gases". His book is available online. – M. Farooq Jul 29 at 13:04

The $$\mathrm{pH}$$ of pure water (rain as well as distilled water) in equilibrium with the atmosphere ($$p_{\ce{CO2}}= 10^{-3.5}\ \mathrm{atm}$$) can be calculated as follows.

$$[\ce{H2CO3^*}]=K_\mathrm H\cdot p_{\ce{CO2}}$$

where $$[\ce{H2CO3^*}]$$ is the total analytical concentration of dissolved $$\ce{CO2}$$, i.e. $$[\ce{H2CO3^*}]=[\ce{CO2(aq)}]+[\ce{H2CO3}]$$, and
$$K_\mathrm H= 3.39\times10^{-2}\ \mathrm{mol\ l^{-1}\ atm^{-1}}$$ is Henry's law constant for $$\ce{CO2}$$.

\begin{align} \log[\ce{H2CO3^*}]&=\log K_\mathrm H+\log p_{\ce{CO2}}\\ &=-1.5-3.5\\ &=-5.0 \end{align}

The commonly used first acid dissociation constant of carbonic acid $$\mathrm pK_{\mathrm a1}=6.3$$ (at $$25\ \mathrm{^\circ C}$$) actually is a composite constant that includes both the hydration reaction $$\ce{H2O + CO2(aq) <=> H2CO3}$$ and the protolysis of true $$\ce{H2CO3}$$ $$\ce{H2CO3 <=> H+ + HCO3-}$$ For a weak acid \begin{align} \log[\ce{H+}]&\approx\frac12\left(\log K_\mathrm a+\log[\ce{H2CO3^*}]\right)\\ &=\frac12\left(-6.3-5.0\right)\\ &=-5.65\\ \mathrm{pH}&=5.65 \end{align}

Thus, pure rain in equilibrium with the atmosphere has about $$\mathrm{pH}=5.65$$. Any acid rain with lower $$\mathrm{pH}$$ would be caused by additional acids.

Under atmospheric pressure, dissolved carbon dioxide can reach an equilibrium state in water that yields a pH of as low as 5.7