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Chapter 17: Problem 24

A buffer is prepared by adding 10.0 \(\mathrm{g}\) of ammonium chloride \(\left(\mathrm{NH}_{4} \mathrm{Cl}\right)\) to 250 \(\mathrm{mL}\) of 1.00 \(\mathrm{M} \mathrm{NH}_{3}\) solution. (a) What is the pH of this buffer? (b) Write the complete ionic equation for the reaction that occurs when a few drops of nitric acid are added to the buffer. (c) Write the complete ionic equation for the reaction that occurs when a few drops of potassium hydroxide solution are added to the buffer.

Short Answer

Expert verified
\(a)\) The pH of the buffer solution is 9.376. \(b)\) The complete ionic equation for the reaction with nitric acid is: NH3(aq) + H\(^{+}\)(aq) + NO3\(^{-}\)(aq) -> NH4\(^{+}\)(aq) + NO3\(^{-}\)(aq). \(c)\) The complete ionic equation for the reaction with potassium hydroxide is: NH4\(^{+}\)(aq) + OH\(^{-}\)(aq) -> NH3(aq) + H2O(l).

Step by step solution

01

Calculate the moles of NH4Cl and NH3

First, we need to find the moles of ammonium chloride and ammonia present in the buffer solution. Moles of NH4Cl = mass / molecular weight = 10.0 g / 53.5 g/mol = 0.18692 mol NH3 solution's volume = 250 mL = 0.250 L Moles of NH3 = concentration × volume = 1.00 M × 0.250 L = 0.250 mol
02

Apply the Henderson-Hasselbalch equation

The Henderson-Hasselbalch equation is given by: pH = pKa + log ([A-] / [HA]) In this case, the conjugate acid is NH4+ (ammonium ion) and the conjugate base is NH3 (ammonia). The Ka of NH4+ is 5.56 × 10^(-10), so the pKa is given by: pKa = -log(5.56 × 10^(-10)) = 9.25 Now we can substitute the values of the pKa, moles of NH3 (A-) and moles of NH4+ (HA) in the equation: pH = 9.25 + log((0.250 mol) / (0.18692 mol))
03

Calculate the pH

Using the given values, we can now calculate the pH of the buffer solution: pH = 9.25 + log(1.337) = 9.25 + 0.126 = 9.376 So the pH of the buffer solution is 9.376. #b) Complete ionic equation when nitric acid is added#
04

Write the balanced molecular equation

When nitric acid (HNO3) is added to the buffer solution containing NH3 (ammonia) and NH4+ (ammonium ion), it produces NH4+ (ammonium ion) and NO3- (nitrate ion). NH3 + HNO3 -> NH4+ + NO3-
05

Separate the ions and write the complete ionic equation

Separating the strong electrolytes into their respective ions, we get the ionic equation: NH3 + H+ + NO3- -> NH4+ + NO3- The complete ionic equation for this reaction is: NH3(aq) + H+(aq) + NO3-(aq) -> NH4+(aq) + NO3-(aq) #c) Complete ionic equation when potassium hydroxide is added#
06

Write the balanced molecular equation

When potassium hydroxide (KOH) is added to the buffer solution containing NH4+ (ammonium ion) and NH3 (ammonia), it produces K+ (potassium ion), NH3 (ammonia), and H2O (water). NH4+ + OH- -> NH3 + H2O
07

Separate the ions and write the complete ionic equation

Separating the strong electrolytes into their respective ions, we get the ionic equation: NH4+ + OH- -> NH3 + H2O The complete ionic equation for this reaction is: NH4+(aq) + OH-(aq) -> NH3(aq) + H2O(l)

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Most popular questions from this chapter

Calculate the \(\mathrm{pH}\) at the equivalence point in titrating 0.100 M solutions of each of the following with 0.080 \(\mathrm{M}\) NaOH: (a) hydrobromic acid (HBr), (b) chlorous acid (HClO_{2} ) , (c) benzoic acid \(\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COOH}\right)\)

(a) Calculate the \(\mathrm{pH}\) of a buffer that is 0.105 \(\mathrm{M}\) in \(\mathrm{NaHCO}_{3}\) and 0.125 \(\mathrm{M}\) in \(\mathrm{Na}_{2} \mathrm{CO}_{3}\) . (b) Calculate the pH of a solution formed by mixing 65 \(\mathrm{mL}\) of 0.20 \(\mathrm{M} \mathrm{NaHCO}_{3}\) with 75 \(\mathrm{mL}\) of 0.15 \(\mathrm{M} \mathrm{Na}_{2} \mathrm{CO}_{3} .\)

Rainwater is acidic because \(\mathrm{CO}_{2}(\mathrm{g})\) dissolves in the water, creating carbonic acid, \(\mathrm{H}_{2} \mathrm{CO}_{3}\) . If the rainwater is too acidic, it will react with limestone and seashells (which are principally made of calcium carbonate, CaCO_ \(_{3} ) .\) Calculate the concentrations of carbonic acid, bicarbonate ion \(\left(\mathrm{HCO}_{3}^{-}\right)\) and carbonate ion \(\left(\mathrm{CO}_{3}^{2-}\right)\) that are in a raindrop that has a pH of 5.60 , assuming that the sum of all three species in the raindrop is \(1.0 \times 10^{-5} M .\)

Predict whether the equivalence point of each of the following titrations is below, above, or at \(\mathrm{pH} 7 :(\mathbf{a}) \mathrm{NaHCO}_{3}\) titrated with \(\mathrm{NaOH},(\mathbf{b}) \mathrm{NH}_{3}\) titrated with \(\mathrm{HCl},(\mathbf{c}) \mathrm{KOH}\) titrated with HBr.

A 35.0-mL sample of 0.150\(M\) acetic acid \(\left(\mathrm{CH}_{3} \mathrm{COOH}\right)\) is titrated with 0.150 \(\mathrm{M} \mathrm{NaOH}\) solution. Calculate the pH after the following volumes of base have been added: (a) 0 \(\mathrm{mL}\) (b) \(17.5 \mathrm{mL},(\mathrm{c}) 34.5 \mathrm{mL},(\mathbf{d}) 35.0 \mathrm{mL},(\mathbf{e}) 35.5 \mathrm{mL},(\mathbf{f}) 50.0 \mathrm{mL}\)
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