B
bobby34
Guest
WHAT CHEMICAL CONSTITUENTS CONTROL pH?
a) In Scheme Waters and Bore Waters
Because of its solubility in water, the presence of carbon dioxide (CO2) in the atmosphere has a major influence on the chemistry and pH of water.
The pH of natural, i.e. uncontaminated, waters with pH values between 4.5 - 8.2 is controlled by the concentrations of bicarbonate anion, HCO3- (sometimes referred to as combined CO2) and free carbon dioxide - where the presence of free carbon dioxide in water lowers pH and bicarbonate elevates pH. It is the amount of bicarbonate (i.e. its alkalinity) in a natural water that determines its buffering capacity. Buffering capacity is the amount of resistance the pH of a water shows to additions of acid or alkali.
The presence of free, i.e. uncombined, carbon dioxide tends to lower the pH because it reacts with water to form carbonic acid thus:
CO2 + H2O = H2CO3
Contrarily, the presence of bicarbonate anion elevates pH because it mops up hydrogen ion thus:
H+ + HCO3 - = H2CO3
The overall reaction is represented by:
CO2 + H2O = H2CO3 = H+ + HCO3-
Thus at high CO2 concentrations the reaction is pushed to the right with the production of more H+ (i.e. pH is lowered). High bicarbonate levels (compared to CO2) mop up H+ with the result the reaction shifts to the left and a higher pH value is produced.
However, a complicating factor is that free carbon dioxide concentrations above about 0.5 mg/L in water are unstable when such waters are exposed to the atmosphere at sea level pressures. Under that condition carbon dioxide in excess of 0.5 mg/L will slowly escape from the water into the atmosphere. This is particularly the case with groundwater's which typically have carbon dioxide contents around 50 - 200 mg/L - as a result of biological activity within the aquifer. When these waters are pumped to the surface, the observed pH rises because the excess (acidic) carbon dioxide escapes. The pH will then rise to a stable value solely dependent on the water's bicarbonate content. For example, a bore water with 100 mg/L bicarbonate and 100 mg/L of free carbon dioxide will have an initial pH of 6.3 gradually rising to 8.2 after it has been exposed to the atmosphere and after which the carbon dioxide content has dropped to around 0.5 mg/L.
The same phenomenon although to a much lesser extent (because of their much lower CO2 contents), occurs with scheme (tap) water. Thus the conclusion - because the pH of natural waters are only stable after aeration, it is only the "after aeration" pH value which is stable and has any interpretative significance. To determine that value, aerate the water by tumbling a sample of it from one container to another, 30-40 times prior to measuring its pH.
In conclusion: interpret pH values with caution because a natural water with a lower pH than another may produce the higher pH after both are aerated!!
Basically i got from this, the PH down in your nutes react with the CO2 causing the ph to rise significantly.
a) In Scheme Waters and Bore Waters
Because of its solubility in water, the presence of carbon dioxide (CO2) in the atmosphere has a major influence on the chemistry and pH of water.
The pH of natural, i.e. uncontaminated, waters with pH values between 4.5 - 8.2 is controlled by the concentrations of bicarbonate anion, HCO3- (sometimes referred to as combined CO2) and free carbon dioxide - where the presence of free carbon dioxide in water lowers pH and bicarbonate elevates pH. It is the amount of bicarbonate (i.e. its alkalinity) in a natural water that determines its buffering capacity. Buffering capacity is the amount of resistance the pH of a water shows to additions of acid or alkali.
The presence of free, i.e. uncombined, carbon dioxide tends to lower the pH because it reacts with water to form carbonic acid thus:
CO2 + H2O = H2CO3
Contrarily, the presence of bicarbonate anion elevates pH because it mops up hydrogen ion thus:
H+ + HCO3 - = H2CO3
The overall reaction is represented by:
CO2 + H2O = H2CO3 = H+ + HCO3-
Thus at high CO2 concentrations the reaction is pushed to the right with the production of more H+ (i.e. pH is lowered). High bicarbonate levels (compared to CO2) mop up H+ with the result the reaction shifts to the left and a higher pH value is produced.
However, a complicating factor is that free carbon dioxide concentrations above about 0.5 mg/L in water are unstable when such waters are exposed to the atmosphere at sea level pressures. Under that condition carbon dioxide in excess of 0.5 mg/L will slowly escape from the water into the atmosphere. This is particularly the case with groundwater's which typically have carbon dioxide contents around 50 - 200 mg/L - as a result of biological activity within the aquifer. When these waters are pumped to the surface, the observed pH rises because the excess (acidic) carbon dioxide escapes. The pH will then rise to a stable value solely dependent on the water's bicarbonate content. For example, a bore water with 100 mg/L bicarbonate and 100 mg/L of free carbon dioxide will have an initial pH of 6.3 gradually rising to 8.2 after it has been exposed to the atmosphere and after which the carbon dioxide content has dropped to around 0.5 mg/L.
The same phenomenon although to a much lesser extent (because of their much lower CO2 contents), occurs with scheme (tap) water. Thus the conclusion - because the pH of natural waters are only stable after aeration, it is only the "after aeration" pH value which is stable and has any interpretative significance. To determine that value, aerate the water by tumbling a sample of it from one container to another, 30-40 times prior to measuring its pH.
In conclusion: interpret pH values with caution because a natural water with a lower pH than another may produce the higher pH after both are aerated!!
Basically i got from this, the PH down in your nutes react with the CO2 causing the ph to rise significantly.
