Geology Reference
In-Depth Information
decreased (H
1
in solution increased), the equilibrium shifts to left -
some acidity goes on to non-exchangeable acidity. This represents the
buffer capacity of a soil, and acts to avoid large changes in pH. Acidity in
soil is normally corrected by addition of a liming material such as
CaCO
3
. The amount required can be calculated by measuring the
buffering of a soil by titration with a dilute solution of an alkali (often
Ca(OH)
2
is used). The amount of alkali needed to reach the required pH
can be found from the buffer curve.
Worked example 5.3 - soil buffering and liming
To raise the pH of 5 g of the soil in Worked example 5.1 to 6.5,
required addition of 4 cm
3
of 0.005 mol L
1
Ca(OH)
2
solution. How
much calcium carbonate would be required to be added to 1 ha? (Mol
wt CaCO
3
¼
100).
4cm
3
of 0
:
005mol L
1
¼
0
:
02mmol Ca
ð
OH
Þ
2
solution for 5 g soil
¼
0
:
004mmol g
1
¼
0
:
4mg CaCO
3
g
1
ð
1mmol CaCO
3
¼
100mg
Þ
¼
0
:
4 kg CaCO
3
t
1
From Worked example 5.1, there are 2500 t ha
1
in the top 20 cm,
which approximates to rooting depth.
0.4 kg CaCO
3
t
1
2500 t
¼
1000 kg CaCO
3
required ha
1
.
5.4.2 Soil Acidity
The atmosphere is a major source of soil acidity. Even in unpolluted
environments rainwater is slightly acidic, having a pH of about 5.7 due
to the dissolution of atmospheric CO
2
to form the weak carbonic acid
(see Worked example 5.4). The CO
2
concentration in the partially
enclosed soil pore system can be significantly higher (typically up to
about 10 times) than in the free atmosphere due to respiration of soil
microorganisms and plant roots. This results in a lower pH. In areas
affected by industrial pollution, sulfur dioxide and nitrogen oxides
dissolve in rainwater to produce sulfuric and nitric acids (acid rain),
which are both strong acids and cause even more acidity.
