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julsineya [31]
3 years ago
15

A 15.0 mL sample of 0.013 M HNO3 is titrated with 0.017 M CH$NH2 which he Kb=3.9 X 10-10. Determine the pH at these points: At t

he beginning (before base is added) pH = -log (H₂O) = -log(0.013) - 1-800 1.013) = 1.886 After adding 10.0 mL of C2H3NH2
Chemistry
1 answer:
kramer3 years ago
4 0

<u>Answer:</u> The pH of the solution in the beginning is 1.89 and the pH of the solution after the addition of base is

<u>Explanation:</u>

  • <u>For 1:</u> At the beginning

To calculate the pH of the solution, we use the equation:

pH=-\log[H^+]

We are given:

Nitric acid is a monoprotic acid and it dissociates 1 mole of hydrogen ions. So, the concentration of hydrogen ions is 0.013 M

[H^+]=0.013M

Putting values in above equation, we get:

pH=-\log(0.013)\\\\pH=1.89

  • <u>For 2:</u>

To calculate the number of moles, we use the equation:  

\text{Molarity of the solution}=\frac{\text{Moles of solute}\times 1000}{\text{Volume of solution (in mL)}}

  • <u>For nitric acid:</u>

Molarity of nitric acid solution = 0.013 M

Volume of solution = 15 mL

Putting values in above equation, we get:

0.013M=\frac{\text{Moles of }HNO_3\times 1000}{15}\\\\\text{Moles of }HNO_3=1.95\times 10^{-4}mol

  • <u>For methylamine:</u>

Molarity of methylamine solution = 0.017 M

Volume of solution = 10 mL

Putting values in above equation, we get:

0.017M=\frac{\text{Moles of }CH_3NH_2\times 1000}{10}\\\\\text{Moles of }CH_3NH_2=1.7\times 10^{-4}mol

  • The chemical equation for the reaction of nitric acid and methylamine follows:

                       HNO_3+CH_3NH_2\rightarrow CH_3NH_3^++NO_3^-

As, the mole ratio of nitric acid and methyl amine is 1 : 1. So, the limiting reagent will be the reactant whose number of moles are less, which is methyl amine.

By Stoichiometry of the reaction:

1 mole of methyl amine produces 1 mole of CH_3NH_3^+

So, 1.7\times 10^{-4}mol of methyl amine will produce = \frac{1}{1}\times 1.7\times 10^{-4}=1.7\times 10^{-4}\text{ moles of }CH_3NH_3^+

To calculate the pK_b of base, we use the equation:

pK_b=-\log(K_b)

where,

K_b = base dissociation constant = 3.9\times 10^{-10}

Putting values in above equation, we get:

pK_b=-\log(3.9\time 10^{-10})\\\\pK_b=9.41

  • To calculate the pOH of basic buffer, we use the equation given by Henderson Hasselbalch:

pOH=pK_b+\log(\frac{[salt]}{[base]})

pOH=pK_b+\log(\frac{[CH_3NH_3^+]}{[CH_3NH_2]})

We are given:

pK_b=9.41

[CH_3NH_3^+]=\frac{1.7\times 10^{-4}}{10+15}=6.8\times 10^{-6}M

[CH_3NH_2]=\frac{1.7\times 10^{-4}}{10+15}=6.8\times 10^{-6}M

Putting values in above equation, we get:

pOH=9.41+\log(\frac{6.8\times 10^{-6}}{6.8\times 10^{-6}})\\\\pOH=9.41

To calculate pH of the solution, we use the equation:

pH+pOH=14\\pH=14-9.41=4.59

Hence, the pH of the solution is 4.59

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