Chemical Forums
Chemistry Forums for Students => High School Chemistry Forum => Topic started by: Kabuze on April 06, 2011, 08:51:32 PM
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A nitric acid solution is found to have a pH of 2.70. Determine each of the following:
a. [H3O] <-2.0 x 10^-3 M (antilog(-pH))
b.[OH] <-5.0 x 10^-12 M (1.0 x 10^-14/2.0 x 10^-3)
c. The number of moles of HNO3 required to prepare 5.50L of this solution.
moles=concentration x volume right? So (2.0 x 10^-3 M) x 5.50 L = 1.1 x 10^-2 moles
d. The mass of the moles of HNO3 in the solution in part c.
1 mole of HNO3=63.01g/mole
1.1 x 10^-2 moles x 63.01 g/mole = 6.93 x 10^-1 grams right?
e.The milliliters of concentrated acid needed to prepare the solution in part c. (Concentrated nitric acid is 69.5% HNO3 by mass and has a density of 1.42 g/mL.)
I have literally no idea of how to do this one, any tips would be helpful and appreciated. (:
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e.The milliliters of concentrated acid needed to prepare the solution in part c. (Concentrated nitric acid is 69.5% HNO3 by mass and has a density of 1.42 g/mL.)
I have literally no idea of how to do this one, any tips would be helpful and appreciated. (:
You already know the mass of acid - calculate mass of the solution (you will need to use definition of %w/w), then convert to volume (using definition of density).