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Chapter 17: Problem 35

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

Short Answer

Expert verified
(a) The equivalence point for the titration of \(\mathrm{NaHCO}_{3}\) (a weak acid) with \(\mathrm{NaOH}\) (a strong base) will be above pH 7. (b) The equivalence point for the titration of \(\mathrm{NH}_{3}\) (a weak base) with \(\mathrm{HCl}\) (a strong acid) will be below pH 7. (c) The equivalence point for the titration of \(\mathrm{KOH}\) (a strong base) with \(\mathrm{HBr}\) (a strong acid) will be at pH 7.

Step by step solution

01

(a) NaHCO3 titrated with NaOH

\(\mathrm{NaHCO}_{3}\) is a weak acid, and \(\mathrm{NaOH}\) is a strong base. When a weak acid is titrated with a strong base, the equivalence point pH will be greater than 7. So, the equivalence point for this titration will be above pH 7.
02

(b) NH3 titrated with HCl

In this case, \(\mathrm{NH}_{3}\) is a weak base, and \(\mathrm{HCl}\) is a strong acid. When a weak base is titrated with a strong acid, the equivalence point pH will be less than 7. Thus, the equivalence point for this titration will be below pH 7.
03

(c) KOH titrated with HBr

Here, \(\mathrm{KOH}\) is a strong base, and \(\mathrm{HBr}\) is a strong acid. When a strong base is titrated with a strong acid, the equivalence point pH will be 7. Therefore, the equivalence point for this titration 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
Acid-base titration is a technique used to determine the concentration of an acid or base in a solution. It involves gradually adding a solution of known concentration, called the titrant, to a solution of unknown concentration, until the chemical reaction reaches the equivalence point.
This is where the amount of titrant added is exactly enough to completely react with the unknown solution. During a titration, an indicator is often used to signal when the equivalence point has been reached. Indicators change color at a certain range of pH levels.
However, the selection of an appropriate indicator depends on the expected pH at the equivalence point.
  • For a titration involving strong acids and strong bases, the equivalence point occurs at pH 7.
  • In titrations of strong acids with weak bases, the equivalence point is typically below pH 7.
  • Conversely, when a weak acid is titrated with a strong base, the equivalence point is above pH 7.
Weak Acids and Bases
Weak acids and bases do not completely dissociate in solution. This means only a small fraction of their molecules ionize to produce hydrogen ions (H+) or hydroxide ions (OH-), impacting their behavior during titration.
For instance, a weak acid reacting with a strong base will result in an equivalence point that is basic due to the formation of its conjugate base, which can further react with water to produce OH- ions.The behavior of weak acids and bases during titration requires understanding their dissociation constants (Ka for acids and Kb for bases).
This allows us to predict the pH changes as the titrant is added.
  • When titrating a weak base like ammonia (\(\text{NH}_3\), a relatively small amount of its molecules react with the acid, resulting in an equivalence point with a low pH.
  • For a weak acid such as \(\text{NaHCO}_3\), the reaction with a strong base results in an equivalence point higher than pH 7.
Strong Acids and Bases
Strong acids and bases completely dissociate in water, meaning they release a high concentration of ions. This characteristic makes their behavior predictable when they undergo titration.
For example, when a strong acid titrates a strong base, or vice versa, the resulting solution at the equivalence point is neutral, having a pH of 7.
  • Strong acids, such as hydrochloric acid \( ext{(HCl)}\), can completely dissociate, reacting efficiently with bases.
  • Strong bases, like potassium hydroxide \( ext{(KOH)}\), also fully break down into ions, allowing full interaction with acids.
Strong acids and bases are perfect for precise titrations, usually resulting in a sharp and distinct equivalence point.While using these substances, safety precautions are necessary due to their highly reactive nature.

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Most popular questions from this chapter

Seawater contains \(0.13 \%\) magnesium by mass, and has a density of \(1.025 \mathrm{~g} / \mathrm{mL}\). What fraction of the magnesium can be removed by adding a stoichiometric quantity of \(\mathrm{CaO}\) (that is, one mole of \(\mathrm{CaO}\) for each mole of \(\mathrm{Mg}^{2+}\) )?

An unknown solid is entirely soluble in water. On addition of dilute \(\mathrm{HCl}\), a precipitate forms. After the precipitate is filtered off, the \(\mathrm{pH}\) is adjusted to about 1 and \(\mathrm{H}_{2} \mathrm{~S}\) is bubbled in; a precipitate again forms. After filtering off this precipitate, the \(\mathrm{pH}\) is adjusted to 8 and \(\mathrm{H}_{2} \mathrm{~S}\) is again added; no precipitate forms. No precipitate forms upon addition of \(\left(\mathrm{NH}_{4}\right)_{2} \mathrm{HPO}_{4} .\) The remaining solution shows a yellow color in a flame test. Based on these observations, which of the following compounds might be present, which are definitely present, and which are definitely absent: \(\mathrm{CdS}, \mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2}, \mathrm{HgO}, \mathrm{ZnSO}_{4}, \mathrm{Cd}\left(\mathrm{NO}_{3}\right)_{2}\) and \(\mathrm{Na}_{2} \mathrm{SO}_{4} ?\)

(a) Explain the difference between solubility and solubility-product constant. (b) Write the expression for the solubility-product constant for each of the following ionic compounds: \(\mathrm{MnCO}_{3}, \mathrm{Hg}(\mathrm{OH})_{2}\), and \(\mathrm{Cu}_{3}\left(\mathrm{PO}_{4}\right)_{2}\).

Explain why a mixture of \(\mathrm{CH}_{3} \mathrm{COOH}\) and \(\mathrm{CH}_{3} \mathrm{COONa}\) can act as a buffer while a mixture of \(\mathrm{HCl}\) and \(\mathrm{NaCl}\) cannot.

Two buffers are prepared by adding an equal number of moles of formic acid (HCOOH) and sodium formate (HCOONa) to enough water to make \(1.00\) L of solution. Buffer A is prepared using \(1.00\) mol each of formic acid and sodium formate. Buffer \(B\) is prepared by using \(0.010 \mathrm{~mol}\) of each. (a) Calculate the \(\mathrm{pH}\) of each buffer, and explain why they are equal (b) Which buffer will have the greater buffer capacity? Explain. (c) Calculate the change in \(\mathrm{pH}\) for each buffer upon the addition of \(1.0 \mathrm{~mL}\) of \(1.00 \mathrm{M} \mathrm{HCl}\). (d) Calculate the change in \(\mathrm{pH}\) for each buffer upon the addition of \(10 \mathrm{~mL}\) of \(1.00 \mathrm{M} \mathrm{HCl}\). (e) Discuss your answers for parts (c) and (d) in light of your response to part (b).
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