how do H1 and OH ions are given off in aqueous solution acid and base explain each with example
Answers
A common means of expressing quantities, the values of which may span many orders of magnitude, is to use a logarithmic scale. One such scale that is very popular for chemical concentrations and equilibrium constants is based on the p-function, defined as shown where “X” is the quantity of interest and “log” is the base-10 logarithm:
pX
=
−
log\;X
The pH of a solution is therefore defined as shown here, where [H3O+] is the molar concentration of hydronium ion in the solution:
pH
=
−
log[H
3
O
+
]
Rearranging this equation to isolate the hydronium ion molarity yields the equivalent expression:
[
H
3
O
+
]
=
10
−
pH
Likewise, the hydroxide ion molarity may be expressed as a p-function, or pOH:
pOH
=
−
log[OH
−
]
or
[
OH
−
]
=
10
−
pOH
Finally, the relation between these two ion concentration expressed as p-functions is easily derived from the Kw expression:
K
w
=
[
H
3
O
+
]
[
OH
−
]
−
log
K
w
=
−
log([H
3
O
+
]
[
OH
−
]
)
=
−
log[H
3
O
+
]
+
−
log[OH
−
]
p
K
w
=
pH
+
pOH
At 25 °C, the value of Kw is 1.0 × 10−14, and so:
14.00
=
pH
+
pOH
As was shown in Example 1 in Chapter 14.1 Brønsted-Lowry Acids and Bases, the hydronium ion molarity in pure water (or any neutral solution) is 1.0 × 10−7M at 25 °C. The pH and pOH of a neutral solution at this temperature are therefore:
pH
=
−
log[H
3
O
+
]
=
−
log
(
1.0
×
10
−
7
)
=
7.00
pOH
=
−
log[OH
−
]
=
−
log
(
1.0
×
10
−
7
)
=
7.00
And so, at this temperature, acidic solutions are those with hydronium ion molarities greater than 1.0 × 10−7M and hydroxide ion molarities less than 1.0 × 10−7M (corresponding to pH values less than 7.00 and pOH values greater than 7.00). Basic solutions are those with hydronium ion molarities less than 1.0 × 10−7M and hydroxide ion molarities greater than 1.0 × 10−7M (corresponding to pH values greater than 7.00 and pOH values less than 7.00).