Naked Science Forum

Non Life Sciences => Chemistry => Topic started by: taregg on 15/02/2014 19:40:23

Title: how to calculate ph for any type of compound.....??
Post by: taregg on 15/02/2014 19:40:23
example......Co2 and KNo3
Title: Re: how to calculate ph for any type of compound.....??
Post by: Bored chemist on 16/02/2014 09:43:34
It's practically impossible to do that without knowing about the material
There are some general rules, but they don't always give the right answer.
For example,
Oxides of non-metal are generally acidic.
Oxides of metals in oxidation states I and II are generally bases.
Oxides of metals in oxidation states higher than 3 are generally acids.

Salts of strong acids with strong bases are usually neutral.

And so on,
But there's no reliable way of calculating pH without a lot more information.
Title: Re: how to calculate ph for any type of compound.....??
Post by: chiralSPO on 16/02/2014 17:45:33
pH is experimentally determined, and refers only to solutions. It is log(concentration of H+ ions in solution). This depends on how much of the substance of interest is dissolved, what temperature the solution is and the acidity/basicity of the compound. I suspect it is this last bit that is most interesting to you.

pKa is a measurement of the acidity of something. It is log(Ka), where Ka is the equilibrium constant of the equation:

AH + H2O <=> A + H3O+

It is possible to calculate approximate values for pKa given a molecular structure, but a de novo calculation is not trivial.

You can also estimate pKa values by finding reported pKa values for structurally anaologous compounds (there are tables and tables and tables of pKa values, for instance:,d.aWc)
Title: Re: how to calculate ph for any type of compound.....??
Post by: snowyco on 21/02/2014 20:21:13
example......Co2 and KNo3

Determine what happens when you add the reagent to water:
KNO3 simply dissociates into K+ and NO3- in water.   

Next:  Determine whether the resulting species are acids or bases or just ordinary ions.  One way of determining this is to write the ions reactions with OH-  or  H+ to see if either ion will change the hydroxide or hydronium ion concentrations:
Neither K+ nor NO3- are either weak acids or weak bases.     This is shown by trying to combine:
K+  +    OH-   ← KOH     because KOH dissociates completely to K+   and  OH-  in water.

NO3-  +  H+  ←  HNO3   because HNO3 dissociates completely to  NO3-   and  H+  in water.

Since neither K+  nor  OH-   react to form acids or bases in water - the pH stays neutral when KNO3 salt is added to water.

CO2 (gas) is more interesting.
Write the appropriate reaction of the gas CO2 with water:
CO2 (g) + H20 ←→ H2CO3   (H2CO3 ~ Carbonic Acid ~ is an acid)

Next write how the H2CO3 product continues to react:

H2CO3 + H2O ←→  H3O+   +  HCO3-
Ka1 = 2.510−4 mole/liter; pKa1 = 3.6 at 25 C

The net amount of  H+    or   H3O+   in solution   depends on how much gaseous CO2 is over the solution.   If you breathe into a beaker of stirred pure distilled water,   the pH drops from the CO2 in your breath.

30 years ago,  the pH of  normal  pure   distilled   lab water   was 5.65   due to atmospheric CO2.
Back then the atmospheric CO2 levels were lower ( 0.0035 atm ).   Many chemistry texts and chemistry websites incorrectly still cite this value.   A more current value for the partial pressure of CO2 is   0.00398 atm.

Consider the calculation for the partial pressure of CO2 with water: 
CO2(gas) in equilibrium with CO2(dissolved) in water as
CO2(gas) ←→  CO2(dissolved) , described by  (

where the Henry constant  kH=29.76 atm/(mol/L) at 25 C

Since the math and associated equilibrium equations are complex (6 equations and 6 unknowns), we generally ignore/neglect the [CO3 2−], because it is so tiny,   yielding this relationship:

To calculate the current pH of pure distilled water in equilibrium with current atmospheric CO2,  change the pCO2  in the formula  from   the outdated  0.0035 atm  instead to the current    0.00398 atm.

This last calculation shows why the oceans are becoming less alkaline  (losing alkalinity) due to increasing CO2 in the atmosphere.