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Lecture of Chemical Equilibrium for the Autodissociation of Pure Water
One of the great things about college courses is that they tend to shake the foundation for the stuff you learned to be true in
grade school.
Like when you learned that the pH of pure water is 7.
Today you will learn the truth, water rarely has a pH of 7.
Now the tool I will be using to come to this conclusion is our new found friend, chemical equilibrium.
And pretty much all I
am going to be doing in this lecture is taking a look at a single equilibrium in which water forms ions on its own.
Here is the chemical equilibrium:
H
2
O
⇌
H
+
+ OH

which has an equilibrium expression
K
w
= [H
+
][OH

]
(remember to not include the [H
2
O] term)
To learn the truth about water’s pH, w will use this equilibrium expression
with some things you were taught in CH301 and the
recent material on the RICE calculation.
Time out for pH:
As you have seen, K values can often become unwieldy with large positive and negative exponents.
Likewise, the
concentrations derived from calculations can be rather cumbersome—like 5.7 x 10
13
M.
Now these days people can handle
such numbers in calculators pretty easily, but in the old days, we tended to use log functions to simplify presentation and
calculation.
So the following idea was hatched to make things easier:
let
[ ]
x
pX
log
−
≡
as a way to simplify presentation of numbers and calculations.
So a number like 1 x 10
7
when subjected to the p(1 x 10
7
)
function yielded the number 7 by simply removing it from the exponent and changing the sign.
And in general the pX function
takes anything
and turns it into –log, so
pH=log[H
+
]
pOH=log[OH

]
pK=log[K]
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 Spring '07
 Holcombe
 Chemistry, Equilibrium, pH

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