How Do You Calculate pH After Titration of HOBr with NaOH?

  • Thread starter gerry73191
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In summary, to calculate the pH of the resulting solution when 20.0 mL of a 0.1 M solution of sodium hydroxide is titrated into 50.0 ml of the 0.14 M HOBr solution, you need to use the Henderson-Hasselbalch equation and calculate the concentrations of HOBr and OBr- from the neutralization reaction stoichiometry. The pH of HOBr is given as 4.74 and Ka is 2.3e-9. By plugging in the calculated concentrations, the final pH is determined to be 8.24.
  • #1
gerry73191
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Homework Statement


20.0 mL of a 0.1 M solution of sodium hydroxide is titrated into 50.0 ml of the 0.14 M HOBr solution. Calculate the pH of the resulting solution.

ph of HOBr is 4.74

Ka is 2.3e-9


Homework Equations



[H+] = ka(a/b)

-log[h+] = ph

ka*kb=kw


The Attempt at a Solution



honestly i have no clue how to do this problem. I already know the answer its 8.24, but I haven't a clue how to get it. can someone point me in the right direction?

thanks
 
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  • #2
Calculate concentrations of HOBr and OBr- (just from the neutralization reaction stoichiometry), plug these numbers in the Henderson-Hasselbalch equation.
 
  • #3


Based on the given information, it seems like this is a buffer problem involving a weak acid (HOBr) and its conjugate base (OBr-). The first step would be to calculate the initial concentrations of both the acid and base in the solution. Since we know the volume and molarity of each solution, we can use the formula M1V1 = M2V2 to find the number of moles of each substance present.

For HOBr:

M1 = 0.14 M

V1 = 50.0 mL = 0.050 L

M2 = unknown (let's call it x M)

V2 = 50.0 mL = 0.050 L

Using the formula, we get:

(0.14 M)(0.050 L) = (x M)(0.050 L)

x = 0.14 M

So the initial concentration of HOBr in the solution is 0.14 M.

For OBr-:

M1 = 0.1 M

V1 = 20.0 mL = 0.020 L

M2 = unknown (let's call it y M)

V2 = 50.0 mL = 0.050 L

Using the formula, we get:

(0.1 M)(0.020 L) = (y M)(0.050 L)

y = 0.04 M

So the initial concentration of OBr- in the solution is 0.04 M.

Next, we can use the Henderson-Hasselbalch equation to calculate the pH of the resulting solution:

pH = pKa + log([base]/[acid])

pH = 4.74 + log(0.04/0.14)

pH = 4.74 + log(0.2857)

pH = 4.74 + (-0.544)

pH = 4.196

However, this is not the final answer because the titration of a weak acid with a strong base creates a buffer solution, and the pH of a buffer solution is affected by the ratio of the concentrations of the acid and base. In this case, the initial concentrations of HOBr and OBr- are not equal, so we need to use the Henderson-Hasselbalch equation again, but this time with the final concentrations of the acid and base.

To calculate the final concentrations, we need to use the neutral
 

What is a buffer solution?

A buffer solution is a solution that resists changes in pH when small amounts of acid or base are added. It is usually composed of a weak acid and its conjugate base or a weak base and its conjugate acid.

What is the purpose of a buffer solution?

The purpose of a buffer solution is to maintain a stable pH in a solution, even when small amounts of acid or base are added. This is important in many biological and chemical processes where a constant pH is necessary for proper functioning.

How do you prepare a buffer solution?

A buffer solution can be prepared by mixing a weak acid or base with its conjugate salt in specific ratios. The ratio will depend on the desired pH of the buffer solution. The components can also be adjusted to create a buffer with a specific pH range.

What is a titration?

A titration is a laboratory technique used to determine the concentration of a substance in a solution by reacting it with a known concentration of another substance. This is often used in acid-base titrations to determine the concentration of an unknown acid or base.

How do you calculate the pH of a buffer solution?

The pH of a buffer solution can be calculated using the Henderson-Hasselbalch equation: pH = pKa + log([conjugate base]/[weak acid]). The pKa value represents the acid dissociation constant and can be found in reference tables. The bracketed terms represent the concentrations of the conjugate base and weak acid in the buffer solution.

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