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I just came to know that pH of rain water will be lesser when it is accompanied by a thunderstorm than when it is not accompanied by thunderstorm.
Why would it be so?
Will the temperature increase due to excessive friction (of air) and lighting and thus more ions of water will dissociate reducing the pH?
Or will it be because of production of acids in air during thunderstorm? Help.

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  • $\begingroup$ Where'd you see this? Do you have a reference or two you can share? $\endgroup$
    – Todd Minehardt
    Feb 15 '16 at 20:26
  • $\begingroup$ @ToddMinehardt It was a question asked in a national level examination. I don't have much of a reference though. $\endgroup$
    – Quark
    Feb 15 '16 at 20:32
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    $\begingroup$ books.google.co.in/… Here is the question $\endgroup$
    – N A
    Feb 16 '16 at 2:23
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Though I've never heard of that statement before, it's possible that nitrogen oxides formed by electrostatic discharge (not only lightning, but even glow discharge) dissolve in the rainwater, producing nitrous and nitric acids.

There is a 1997 citation describing comparison of pH in rain accompanied (or not) by lightning.

That said, presence of fly-ash from coal burning, sulfur oxides from volcanic activity and carbon dioxide from natural as well as anthropogenic sources probably have greater influence on pH.

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When lightning and thundering occurs, it oxidizes into $\ce{SO2}$ and $\ce{NO_x}$ gases from pollution. They mix with rain water and reduces pH of water.

Ref.: L.Bruce Railsback, Lower pH of acid rain associated with lightning: evidence from sampling within 14 showers and storms in the Georgia Piedmont in summer 1996, Science of The Total Environment, Volume 198, Issue 3, 1997, DOI: 10.1016/S0048-9697(97)05459-4.

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Traces of nitric acid are the natural component of the rain, originated by the reaction $$\ce{N2(g) + O2(g) -> 2 NO(g)}.$$

The reaction occurs at electric discharges, naturally like during lightnings and silent discharges, or industrially in electric arc ( in obsolete, energy demanding Birkeland-Eyde process, replaced later by catalytic oxidation of ammonia. )

It goes farther as:

\begin{align} \ce{2 NO(g) + O2(g) &-> 2 NO2(g)}\\ \ce{4 NO2(g) + O2(g) + 2 H2O(l) &-> 4 HNO3(aq)}\\ \ce{4 NO2(g) + 2 H2O(l) &-> 2 HNO3(aq) + 2 HNO2(aq)}\\ \ce{2 HNO2(aq) &-> NO2(g) + NO(g) + H2O(l)} \end{align}

Therefore, in the end, pollution aside, the rain during active thunderstorms is slightly more acidic because of $\ce{HNO3}$, compared to the natural baseline acidity caused by dissolved $\ce{CO2}$.

Unfortunately, the rain $\mathrm{pH}$ depends these days much more on $\ce{SOx}$ and $\ce{NOx}$ pollution than on natural processes, leading to industrially induced very weak solutions of $\ce{H2SO4}$ and $\ce{HNO3}$.

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  • $\begingroup$ I'm old........... $\endgroup$ Jul 20 at 22:32
  • $\begingroup$ It can be advantage, you have a lot of classical knowledge. Natural HNO3 based rain acidity is just a little addition to natural CO2 based rain acidity, which both are just little addition to pollution based acidity. $\endgroup$
    – Poutnik
    Jul 21 at 3:47

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