A 0.4000M solution of nitric acid is used to titrate 50.00mL of 0.237M barium hy

Question

A 0.4000M solution of nitric acid is used to titrate 50.00mL of 0.237M barium hydroxide. (Assume that volumes are additive).

a)Write a balanced net ionic equation for the reaction that takes place during titration.

b)What are the species present at the equivalence point?

c)What volume of nitric acid is required to reach the equivalence point?

d)What is the pH of the solution before any HNO3 is added?

e)What is the pH of the solution halfway to the equivalence point?

f) What is the pH of the solution at the equivalence point?

Explanation / Answer

moles Ba(OH)2 = 0.05000 L x 0.237 M= 0.01185
moles OH- = 2 x 0.01185=0.0237

moles HNO3 required to reach the half equivalence point = 0.0237/2 = 0.01185
Volume HNO3 = 0.01185/ 0.4000 M= 0.0296 L

total volume = 0.05000 + 0.0296=0.0796 L

moles OH- in excess = 0.01185
[OH-]= 0.01185/ 0.0796=0.149 M
pOH = 0.827
pH = 14 - pOH = 13.2

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