Trivia Question of the Week!

speakerguy

Premium Member
I'm bored, so I figured this might be fun.

Pure, neutral water with no dissolved gases:

a) Always has a pH of 7.0

b) Sometimes has a pH above 7.0

c) Sometimes has a pH below 7.0

d) Both B and C

e) Tastes like chicken


First person with the correct answer AND explanation wins an e-cookie.
 
why does ph change with temperature and which way does it change? If memory serves me correctly it is a logarhythmic scale of carbon saturation. does the raised temp release carbon from a solution?

It it possible to have normal water with no disolved gasses in it? seems kinda wierd seeing as how Hydrogen and oxygen in their periodic states are gasses.

sorry to rehash... Im on this whole chemistry thing now and learning.

Brian
 
pH is a measure of the H+ ions in solution. More H+ means lower pH.

The main thing controlling it is the autodissociation of water:

H2O <----> H+ + OH-

As temperature rises, that equilibrium is driven to the right, making more H+ and OH-.

Since H+ rises as the temperature rises, the pH of pure water drops. :)
 
Well I KNEW that you knew the answer :)

Oh well, here's you're e-cookie anyway:
choc%20chip%20cookie.jpg
 
Hey, don't give Rand his cookie yet :D

As temperature rises, that equilibrium is driven to the right, making more H+ and OH-.

Since H+ rises as the temperature rises, the pH of pure water drops.


But if both rised the same the pH would be the same, as the H+ will be canceled out by a OH-.


I think what he meant to say was this;


pH can indeed vary with temperature. The reasons why
depend on the context, but even a simple solution of a weak
acid (HA) will exhibit a (weak) temperature dependence. The pH is
given by the Henderson-Hasselbalch equation:

pH = pKa + log { [A-]/[HA] }

where Ka is the equilibrium constant for the reaction

HA ---> H+ + A-

( ka = [H+][A-] / [HA] )

and pKa = - log Ka .

Ka is itself a function of temperature, since it is related to
the Gibbs free energy of reaction (delta G) by the equation

delta G = - RT ln Ka = -2.303 RT log Ka = 2.303 RT * pKa

so we have

pkA = delta G / (2.303 RT)

delta G is itself given by

delta G = delta H - T * delta S

where delta H is the enthalpy of reaction and delta S is the entropy
of reaction. Combining these, we get

pKa = (delta H / (2.303 RT)) - (delta S / (2.303 R))

If we assume for the sake of simplicity that delta H and delta S
are approximately independent of temperature T (constant), then
the variation with temperature is determined by the sign of delta H.
For example, if delta H is positive (endothermic dissociation),
pKa gets smaller as the temperature gets larger. A
decrease in pKa amounts to an increase in Ka, which means that the
reaction favors dissociation more as temperature increases (in
agreement with LeChatelier's principle). This increases [H+] and
decreases the pH. If the reaction is exothermic the opposite effect will
be observed. Either way, we expect the pH to depend on temperature.

These arguments can be extended to strong acids too. Things get
complicated when there are multiple chemical reactions. Biological
systems can use enzyme-catalyzed reactions to keep the pH constant even
when T varies (within limits, of course).

Boom of course did not wrtie this :lol: :D
 
But if both rised the same the pH would be the same, as the H+ will be canceled out by a OH-.


The pH will be different.

The pKw (autodissociation of water) changes with temperaure (and also very slightly with pressure).

Since we are talking about pure water [H+] = [OH-].

Kw=[H+] [OH-] , combining with above equality gives:

Kw=[H+]^2

so [H+] = Kw^0.5

Which gives: pH = pKw/2

Since pKw changes with temperaure, pH changes too.
 
Too late, Boomer, I already ate the cookie. :D

But if both rised the same the pH would be the same, as the H+ will be canceled out by a OH-.

pH = -log [H+]

If H+ rises (as it does due to changes in Kw as Habib points out), pH drops, regardless of what happens to OH- :)
 
OK, now that is all cleared up ;) as that very tech explanation I posted was a little confusing and I do not know much about Gibbs free energy of reaction. Hab you made it much easier :) That post also says This increases [H+] and
decreases the pH.
, so I figured Randy was right
 
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