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Chapter 17: Problem 37

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

Short Answer

Expert verified
(a) NaHCO3 titrated with NaOH: At the equivalence point, the pH will be above 7 due to the basic nature of the resulting sodium carbonate salt. (b) NH3 titrated with HCl: At the equivalence point, the pH will be below 7 due to the acidic nature of the resulting ammonium chloride salt. (c) KOH titrated with HBr: At the equivalence point, the pH will be neutral (pH 7) due to the neutral nature of the resulting potassium bromide salt.

Step by step solution

01

(a) NaHCO3 titrated with NaOH

For this titration, sodium bicarbonate (NaHCO3, a weak acid) reacts with sodium hydroxide (NaOH, a strong base). The balanced chemical equation for this reaction is given by: \[ NaHCO_{3}(aq) + NaOH(aq) \rightarrow Na_{2}CO_{3}(aq) + H_{2}O(l) \] At the equivalence point, we have sodium carbonate (Na2CO3) as the final salt. Sodium carbonate is a salt of a strong base and weak acid, and it will hydrolyze in water and lead to a basic solution. Therefore, at the equivalence point, the pH will be above 7.
02

(b) NH3 titrated with HCl

For this titration, ammonia (NH3, a weak base) reacts with hydrochloric acid (HCl, a strong acid). The balanced chemical equation for this reaction is given by: \[ NH_{3}(aq) + HCl(aq) \rightarrow NH_{4}Cl(aq) \] At the equivalence point, we have ammonium chloride (NH4Cl) as the final salt. Ammonium chloride is a salt of a weak base and strong acid, and it will hydrolyze in water and lead to an acidic solution. Therefore, at the equivalence point, the pH will be below 7.
03

(c) KOH titrated with HBr

For this titration, potassium hydroxide (KOH, a strong base) reacts with hydrobromic acid (HBr, a strong acid). The balanced chemical equation for this reaction is given by: \[ KOH(aq) + HBr(aq) \rightarrow KBr(aq) + H_{2}O(l) \] At the equivalence point, we have potassium bromide (KBr) as the final salt. Potassium bromide is a salt of a strong base and strong acid, and it will not hydrolyze in water. Therefore, at the equivalence point, the pH will be neutral, which means it will be at pH 7.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Acid-Base Titration
In an acid-base titration, a known concentration of an acid or base is gradually added to a solution until the other solution is completely neutralized. This process helps us determine the concentration of the unknown solution.
The point at which the amounts of acid and base are stoichiometrically equal is known as the equivalence point.
  • If a strong acid is titrated with a strong base, the equivalence point will be at pH 7 because both components completely neutralize each other, forming water and a neutral salt.
  • Conversely, if a strong acid is titrated with a weak base (or vice versa), the resulting salt can hydrolyze and affect the pH at equivalence.
This principle helps in predicting whether the equivalence point will be acidic, neutral, or basic depending on the strengths of the reactants involved.
pH Calculation
The pH, a measure of hydrogen ion concentration, indicates the acidity or basicity of a solution. Calculating pH during titration involves understanding the reaction dynamics between the acid and base.
During titration, for each milliliter of titrant added, the pH will change depending on the acid-base strength. For example:
  • In a titration of a weak acid and strong base, the conjugate base form of the weak acid is present at equivalence, typically resulting in a basic pH.
  • Conversely, titrating a weak base with a strong acid will usually produce an acidic solution at the equivalence point due to the formation of its conjugate acid.
  • Strong acid and strong base titrations culminate in a neutral solution, as neither component remains to affect the pH.
By understanding these principles, students can predict the approximate pH at different stages of the titration.
Hydrolysis of Salts
Hydrolysis of salts can significantly affect the pH of a solution at equivalence. Hydrolysis refers to the reaction of salt anions or cations with water, altering the solution's pH.
Depending on the strengths of the acids and bases:
  • Salts derived from a strong base and weak acid will undergo anion hydrolysis, producing hydroxide ions and making the solution basic.
  • Conversely, salts from a strong acid and weak base result in cation hydrolysis, generating hydrogen ions that render the solution acidic.
  • Salts from strong acids and bases don't typically hydrolyze, yielding a neutral pH.
Understanding these reactions allows for accurate predictions of pH changes. It clarifies why, at the equivalence point, the pH can vary so broadly across different titrations.

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Most popular questions from this chapter

You are asked to prepare a pH \(=2.50\) buffer solution starting from \(1.50 \mathrm{~L}\) of a \(0.75 \mathrm{M}\) solution of hydrofluoric acid (HF) and any amount you need of sodium fluoride (NaF). (a) What is the \(\mathrm{pH}\) of the hydrofluoric acid solution prior to adding sodium fluoride? (b) How many grams of sodium fluoride should be added to prepare the buffer solution? Neglect the small volume change that occurs when the sodium fluoride is added.

For each statement, indicate whether it is true or false. (a) The solubility of a slightly soluble salt can be expressed in units of moles per liter. (b) The solubility product of a slightly soluble salt is simply the square of the solubility. (c) The solubility of a slightly soluble salt is independent of the presence of a common ion. (d) The solubility product of a slightly soluble salt is independent of the presence of a common ion.

A sample of \(0.2140 \mathrm{~g}\) of an unknown monoprotic acid was dissolved in \(25.0 \mathrm{~mL}\) of water and titrated with \(0.0950 \mathrm{M}\) \(\mathrm{NaOH}\). The acid required \(30.0 \mathrm{~mL}\) of base to reach the equivalence point. (a) What is the molar mass of the acid? (b) After \(15.0 \mathrm{~mL}\) of base had been added in the titration, the \(\mathrm{pH}\) was found to be \(6.50 .\) What is the \(K_{a}\) for the unknown acid?

A buffer is prepared by adding \(3.5 \mathrm{~g}\) of ammonium chloride \(\left(\mathrm{NH}_{4} \mathrm{Cl}\right)\) to \(100 \mathrm{~mL}\) of \(1.00 \mathrm{M} \mathrm{NH}_{3}\) solution. (a) What is the \(\mathrm{pH}\) of this buffer? (b) Write the complete ionic equation for the reaction that occurs when a few drops of hydrochloric acid are added to the buffer. (c) Write the complete ionic equation for the reaction that occurs when a few drops of sodium hydroxide solution are added to the buffer.

Suppose you want to do a physiological experiment that calls for a pH 6.50 buffer. You find that the organism with which you are working is not sensitive to the weak acid \(\mathrm{H}_{2} \mathrm{~A}\left(K_{a 1}=2 \times 10^{-2} ; K_{a 2}=5.0 \times 10^{-7}\right)\) or its sodium salts. You have available a \(1.0 \mathrm{M}\) solution of this acid and a 1.0 \(M\) solution of \(\mathrm{NaOH}\). How much of the \(\mathrm{NaOH}\) solution should be added to \(1.0 \mathrm{~L}\) of the acid to give a buffer at \(\mathrm{pH}\) 6.50? (Ignore any volume change.)
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