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Chapter 17: Problem 110

Indicate whether you would expect the equivalence point of each of the following titrations to be below, above, or at \(\mathrm{pH}\) 7. Explain your reasoning. (a) \(\mathrm{NaHCO}_{3}(\mathrm{aq})\) is titrated with \(\mathrm{NaOH}(\mathrm{aq})\) (a) (b) \(\mathrm{HCl}(\mathrm{aq})\) is titrated with \(\mathrm{NH}_{3}(\mathrm{aq}) ;\) (c) \(\mathrm{KOH}(\mathrm{aq})\) is titrated with HI(aq).

Short Answer

Expert verified
(a) The pH at the equivalence point would be above 7. (b) The pH at the equivalence point would be below 7. (c) The pH at the equivalence point would be at 7.

Step by step solution

01

Determining pH for Titration of NaHCO3 and NaOH

Sodium bicarbonate (\(NaHCO_3\)) is a weak base and sodium hydroxide (\(NaOH\)) is a strong base. When a weak base is titrated with a strong base, the pH at the equivalence point is greater than 7. Hence, we would expect a pH above 7.
02

Determining pH for Titration of HCl and NH3

Hydrochloric acid (\(HCl\)) is a strong acid and ammonia (\(NH_3\)) is a weak base. When a strong acid is titrated with a weak base, the pH at the equivalence point is less than 7. Therefore, we would expect a pH below 7.
03

Determining pH for Titration of KOH and HI

Potassium hydroxide (\(KOH\)) is a strong base and hydroiodic acid (\(HI\)) is a strong acid. When a strong base is titrated with a strong acid, the pH at the equivalence point is exactly 7. As such, we would expect a pH of 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 process used to determine the concentration of an acid or a base in a solution. This analytical technique involves the gradual addition of a known concentration solution, called the titrant, to a solution of the substance being analyzed, known as the analyte. During the titration, the pH of the solution is carefully monitored, often with the aid of indicators that change color at specific pH values.

Titration helps to identify the equivalence point, which is the stage at which the number of moles of the titrant is stoichiometrically equal to the number of moles of the analyte. In simpler terms, it's the point where the acid and base have completely reacted with one another. Understanding the properties of the reactants, namely whether the acids and bases are strong or weak, allows us to predict the pH at the equivalence point.

For instance, titrating a weak base with a strong acid will yield an acidic solution at the equivalence point, resulting in a pH below 7. Conversely, titrating a weak acid with a strong base will yield a basic solution at the equivalence point, with a pH above 7. This predictive capability is crucial in a variety of chemical applications, including pharmaceuticals, environmental monitoring, and the food industry.
pH Scale
The pH scale is a measure of the acidity or basicity of an aqueous solution. The scale ranges from 0 to 14, with 7 being neutral. Solutions with a pH less than 7 are considered acidic, while those with a pH greater than 7 are basic or alkaline. The pH scale is logarithmic, meaning each whole number change represents a tenfold increase or decrease in hydrogen ion (\(H^+\)) concentration. For example, a solution with a pH of 4 is ten times more acidic than a solution with a pH of 5.

The pH scale is not only a numerical scale but also a conceptual tool to visualize the relative strength of acids and bases. In the context of titrations, monitoring the pH allows the chemist to determine the equivalence point accurately and to infer information about the relative strengths of the reactants. Indicators, pH meters, or conductometric measurements are commonly used tools to measure pH changes during a titration process.
Strong and Weak Acids and Bases
Acids and bases can be classified as either strong or weak. This classification is based on their ability to disassociate in water. Strong acids and bases completely disassociate into their ions in solution, while weak acids and bases only partially disassociate.

In a titration, strong acids will react fully with bases, and strong bases will react fully with acids, neither leaving behind any unreacted molecules. This complete reaction leads to a neutral solution at the equivalence point with a pH of 7. On the other hand, the presence of weak acids or bases results in an incomplete reaction and an equivalence point pH that is either greater or less than 7, depending on the nature of the substances involved.

Example of Acids and Bases Strength

Hydrochloric acid (\(HCl\)) is a classic example of a strong acid, while acetic acid (\(CH_3COOH\)) is a weak acid. Similarly, sodium hydroxide (\(NaOH\) is a strong base, whereas ammonia (\(NH_3\) is a weak base. In a mixture, their state of disassociation influences the pH level and, during a titration, infers the overall strength and behavior of the acids and bases involved.

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Most popular questions from this chapter

The neutralization of \(\mathrm{NaOH}\) by \(\mathrm{HCl}\) is represented in equation (1), and the neutralization of \(\mathrm{NH}_{3}\) by HCl in equation (2). 1. \(\mathrm{OH}^{-}+\mathrm{H}_{3} \mathrm{O}^{+} \rightleftharpoons 2 \mathrm{H}_{2} \mathrm{O} \quad K=?\) 2\. \(\mathrm{NH}_{3}+\mathrm{H}_{3} \mathrm{O}^{+} \rightleftharpoons \mathrm{NH}_{4}^{+}+\mathrm{H}_{2} \mathrm{O} \quad K=?\) (a) Determine the equilibrium constant \(K\) for each reaction. (b) Explain why each neutralization reaction can be considered to go to completion.

The \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}-\mathrm{HPO}_{4}^{2-}\) combination plays a role in maintaining the pH of blood. (a) Write equations to show how a solution containing these ions functions as a buffer. (b) Verify that this buffer is most effective at \(\mathrm{pH} 7.2\) (c) Calculate the \(\mathrm{pH}\) of a buffer solution in which \(\left[\mathrm{H}_{2} \mathrm{PO}_{4}\right]=0.050 \mathrm{M}\) and \(\left[\mathrm{HPO}_{4}^{2-}\right]=0.150 \mathrm{M} .[\)Hint: Focus on the second step of the phosphoric acid ionization.]

You are asked to prepare a \(\mathrm{KH}_{2} \mathrm{PO}_{4}-\mathrm{Na}_{2} \mathrm{HPO}_{4}\) solu- tion that has the same \(\mathrm{pH}\) as human blood, 7.40 (a) What should be the ratio of concentrations \(\left[\mathrm{HPO}_{4}^{2-}\right] /\left[\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\right]\) in this solution? (b) Suppose you have to prepare \(1.00 \mathrm{L}\) of the solution described in part (a) and that this solution must be isotonic with blood (have the same osmotic pressure as blood). What masses of \(\mathrm{KH}_{2} \mathrm{PO}_{4}\) and of \(\mathrm{Na}_{2} \mathrm{HPO}_{4} \cdot 12 \mathrm{H}_{2} \mathrm{O}\) would you use? [Hint: Refer to the definition of isotonic on page \(580 .\) Recall that a solution of \(\mathrm{NaCl}\) with \(9.2 \mathrm{g} \mathrm{NaCl} / \mathrm{L}\) solution is isotonic with blood, and assume that \(\mathrm{NaCl}\) is completely ionized in aqueous solution.]

If an indicator is to be used in an acid-base titration having an equivalence point in the pH range 8 to 10 , the indicator must (a) be a weak base; (b) have \(K_{\mathrm{a}}=1 \times 10^{-9} ;(\mathrm{c})\) ionize in two steps; (d) be added to the solution only after the solution has become alkaline.

Because an acid-base indicator is a weak acid, it can be titrated with a strong base. Suppose you titrate \(25.00 \mathrm{mL}\) of a \(0.0100 \mathrm{M}\) solution of the indicator \(p\) -nitrophenol, \(\mathrm{HOC}_{6} \mathrm{H}_{4} \mathrm{NO}_{2},\) with \(0.0200 \mathrm{M} \mathrm{NaOH}\) The \(\mathrm{p} K_{\mathrm{a}}\) of \(p\) -nitrophenol is \(7.15,\) and it changes from colorless to yellow in the pH range from 5.6 to 7.6 (a) Sketch the titration curve for this titration. (b) Show the pH range over which \(p\) -nitrophenol changes color. (c) Explain why \(p\) -nitrophenol cannot serve as its own indicator in this titration.
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