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I am trying to understand how to account for pH-temperature relations during industrial pH measurements.

Do I have it right?:

  1. The impact temperature has on the pKa of strong acids are in practical terms non existent as the equilibrium at any temperature lays very far to the right.

  2. The impact of temperature on pKa are more significant for weaker acids.

  3. The impact of temperature on pKw is only significant around the point of neutrality. At lower pH's the few hydronium ions produced by H2O + H2O are so few that it bares no meaning on the logarithmic pH scale. In other words, the pH-temperature relationship is important when dealing with neutralization processes, but not so much when having pH setpoints set to anything in the milder acidic region and below.

  4. Compensating for the pH-temperature relationship caused by the temperature dependent pKa is only important when doing measurements on processes where the temperature of the process differs from the temperature of the final product and where the targeted pH level during process control has to match the targeted pH level of the end product (at another temperature). For example: say I want to produce an end product with a pH at 5 at temperature X. In order to make this product the temperature of my fabrication process has to be something other than temperature X. In this case, temperature compensation to account for changes is pKa becomes important - otherwise the end product won't have the desired pH.

  5. in a different scenario the pH setpoint of a process becomes a means to an end. It's the pH of the process that determines the chemical reactions that takes place in that particular process (perhaps the hydronium ions merely acts as catalysts). In this case, temperature compensation in no way should be applied. This because the pH changes casued by changes in pkA are real and if the process demands say pH 3 for an optimal chemical reaction then it doesn't matter what the pH would have been at another temperature. The important thing is what the temperature is in the process right here, right now.

  6. while pH changes that comes from changes to pKa actually do mean that a solution becomes more or less acidic, changes to pKw doesn't result in a solution becoming more or less acidic even though pH changes. This is because an equal amount of hydroxide ions a produced. So only the point of neutrality is changed.

Do I understand things correctly so far?

If so, then my question now is: What happens when we move into the alkaline region? I mean, pKb is, according to Le Châteliers principle, temperature-dependent as well, right? When looking at this graph https://ibb.co/LRs7G5y (algorithm based on NEN6411 (now DS/EN ISO 10523:2012)) it seems as if alkaline solutions ALSO becomes more acidic with increases in temperature. Why is this? Does it mean that the equilibrium constant pKb runs in reverse of pKa when temperature increases (less OH-)? Or does it mean that pKw plays are much larger importance once we wander into the alkaline region? I don't, however, see how pKw play any importance on the temperature-pH relationship of alkaline solutions. I mean, if we have an alkaline solution of say pH 12, then any extra/or less hydronium ions produced by changes in pKw would be dwarfed by the proportional much larger amount of hydroxide ions.

Or do I get this wrong?

Last question: Why does the graph https://ibb.co/LRs7G5y indicates that no significant pH-temperature relationship exists below pH 6? Is it because the acid gets so strong at this point that the changes to pKa becomes miniscule relatively speaking?

Help would be much appreciated.

/Lars

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    $\begingroup$ The followup of temperature-effects-on-pka-and-pkw. $\endgroup$
    – Poutnik
    Jan 27 at 13:48
  • $\begingroup$ Why do you need to know the pH? Do you need to know the hydrogen ion activity, the hydroxide ion activity, or some other parameter? For strongly basic solutions, the pH is often a proxy for the hydroxide ion concentration or activity (a major species) rather than the hydrogen ion concentration (a species with very low concentration in that case). $\endgroup$ Jan 27 at 16:37
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    $\begingroup$ This site is designed to answer a single focused question rather than to check "my understanding of the topic". People often sneak in multiple questions on the side, but it is important to know the focal question (the problem you want to solve) for receiving meaningful answers. $\endgroup$ Jan 27 at 16:39
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The impact temperature has on the pKa of strong acids are in practical terms non existent as the equilibrium at any temperature lays very far to the right.

Practically, yes.

The impact of temperature on pKa are more significant for weaker acids.

Yes.

The impact of temperature on pKw is only significant around the point of neutrality. At lower pH's the few hydronium ions produced by H2O + H2O are so few that it bares no meaning on the logarithmic pH scale. In other words, the pH-temperature relationship is important when dealing with neutralization processes, but not so much when having pH setpoints set to anything in the milder acidic region and below.

$\mathrm{p}K_\mathrm{w}$ changes apply to all $\ce{H3O+}$ and $\ce{OH-}$, not just to those from water autodisscociation. If $\ce{[H3O+]}$ ( [] is the convention for molar conceentration ) is determined by $\ce{[OH-]}$, or if some process depends directly on $\ce{[OH-]}$ and not on $\ce{[H3O+]}$, it is significant.

Compensating for the pH-temperature relationship caused by the temperature dependent pKa is only important when doing measurements on processes where the temperature of the process differs from the temperature of the final product and where the targeted pH level during process control has to match the targeted pH level of the end product (at another temperature).

For example: say I want to produce an end product with a pH at 5 at temperature X. In order to make this product the temperature of my fabrication process has to be something other than temperature X. In this case, temperature compensation to account for changes is pKa becomes important - otherwise the end product won't have the desired pH.

Yes. E.g. for acetic acid, $\mathrm{p}K_\mathrm{a, \pu{25^{\circ}C}}=4.75$. That is the $\mathrm{pH}$ of solution acetic acid / sodium acetate in molar ratio 1:1. Let imagine for illustrative purposes $\mathrm{p}K_\mathrm{a, \pu{50^{\circ}C}}=4.5$. If you used at $t=\pu{50^{\circ}C}$ for the final process $\mathrm{pH}=4.75$, there would not be reached 1:1 ratio.

in a different scenario the pH setpoint of a process becomes a means to an end. It's the pH of the process that determines the chemical reactions that takes place in that particular process (perhaps the hydronium ions merely acts as catalysts). In this case, temperature compensation in no way should be applied. This because the pH changes casued by changes in pkA are real and if the process demands say pH 3 for an optimal chemical reaction then it doesn't matter what the pH would have been at another temperature. The important thing is what the temperature is in the process right here, right now.

Yes, if I understant the paragraph well.

while pH changes that comes from changes to pKa actually do mean that a solution becomes more or less acidic, changes to pKw doesn't result in a solution becoming more or less acidic even though pH changes. This is because an equal amount of hydroxide ions a produced. So only the point of neutrality is changed.

This depends on what you mean by acidity. If you mean by acidity the activity or concentration of $\ce{H3O+}$, it is affected by $\mathrm{p}K_\mathrm{w}$ changes.

If so, then my question now is: What happens when we move into the alkaline region? I mean, pKb is, according to Le Châteliers principle, temperature-dependent as well, right? When looking at this graph https://ibb.co/LRs7G5y (algorithm based on NEN6411 (now DS/EN ISO 10523:2012)) it seems as if alkaline solutions ALSO becomes more acidic with increases in temperature. Why is this? Does it mean that the equilibrium constant pKb runs in reverse of pKa when temperature increases (less OH-)? Or does it mean that pKw plays are much larger importance once we wander into the alkaline region? I don't, however, see how pKw play any importance on the temperature-pH relationship of alkaline solutions. I mean, if we have an alkaline solution of say pH 12, then any extra/or less hydronium ions produced by changes in pKw would be dwarfed by the proportional much larger amount of hydroxide ions.

For weak acids, the primary parameters is $\mathrm{p}K_\mathrm{a}$. Then for their conjugate bases, $\mathrm{p}K_\mathrm{b}=\mathrm{p}K_\mathrm{w} - \mathrm{p}K_\mathrm{a}$.

For weak bases, the primary parameters is $\mathrm{p}K_\mathrm{b}$. Then for their conjugate acids, $\mathrm{p}K_\mathrm{a}=\mathrm{p}K_\mathrm{w} - \mathrm{p}K_\mathrm{b}$.

For solutions of strong acids, $\mathrm{pH} \approx -\log{[\ce{H3O+}]}$, $\mathrm{pOH} = \mathrm{p}K_\mathrm{w} - \mathrm{pH}$.

For solutions of strong bases, $\mathrm{pOH} \approx -\log{[\ce{OH-}]}$, $\mathrm{pH} = \mathrm{p}K_\mathrm{w} - \mathrm{pOH}$.


The graph you mention is for strong electrolytes, so $\mathrm{p}K_\mathrm{a}$ is not applicable.

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  • $\begingroup$ Ok, so as I understand it the auto dissociation constant of H2O governs ALL solutions and not just solutions that hovers around the point of neutrality. Kw determines the total possible content of hydronium and hydroxide combined. More hydronium results in less hydroxide (and vice versa) according to Le Chatelier. This means that if I dissolve 1 HCl molecule in a solution, not only will a hydronium ion come into existence but a hydroxide ion will disappear (in order to satisfy the constant pKw). But if this is true, where does the OH- get its H+ from (so that it can return to being H2O)? $\endgroup$ Jan 31 at 13:55
  • $\begingroup$ I am still no sure, however, how all of this ties into basic/alkaline solutions being more temperature dependent than highly acidic solutions (and neutral solutions as well for that matter). If I increase the temperature of a basic solution, why exactly is it that this solutions has a larger shift in pH than an acidic solution? What exactly happens chemically speaking? I am also no quite sure what you mean by this: "If [H3O+] is determined by [OH−], or if some process depends directly on [OH−] and not on [H3O+], it is significant" $\endgroup$ Jan 31 at 13:57
  • $\begingroup$ OH- does not have "it's H3O+". It does not matter if its origin if from water auto-dissociation or acid dissociation. H3O+ and OH- appear and disappear all the time. At equilibrium both rates are equal. Out of equilibrium, one direction is faster. $\endgroup$
    – Poutnik
    Jan 31 at 15:22
  • $\begingroup$ We are again at pH + pOH=pKw(T) . For strong acid/base solution, pH/pOH are given, respectively. pOH/pH are then respectively the function of pKw(T) and pH/pOH. // Some reactions involve OH-vand not H3O+.They depend directly on pOH=pKw(T)) - pH. $\endgroup$
    – Poutnik
    Jan 31 at 15:30
  • $\begingroup$ 1 L of water contains 10^-7 mol/L OH-. if 1 mol of HCl(g) is added, about 10^-7 mol of OH- react with any H3O+( mostly from HCl), so just 10^-14 mol OH- would remain. $\endgroup$
    – Poutnik
    Jan 31 at 15:35

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