Solved: Explain why small, highly charged metal ions are

In order for a species to act as a Brønsted base, an atom in the species must possess a lone pair of electrons. Explain why this is so.

Calculate the concentration of \(O H^{-}\) ions in a HCl solution whose hydrogen ion concentration is 1.3 M.

(a) List in order of decreasing concentration of all the ionic and molecular species in the following acid solutions: (i) \(\mathrm{HNO}_{3}\) and (ii) HF. (b) List in order of decreasing concentration of all the ionic and molecular species in the following base solutions: (i) \(\mathrm{NH}_{3}\) and (ii) KOH.

Write the formulas of the conjugate bases of the following acids: (a) \(\mathrm{HNO}_{2}\), (b) \(\mathrm{H}_{2} \mathrm{SO}_{4}\), (c) \(\mathrm{H}_{2} \mathrm{S}\), (d) HCN, (e) HCOOH (formic acid).

The \(O H^{-}\) ion concentration of a blood sample is \(2.5 \times 10^{-7}\ M\). What is the pH of the blood?

Predict whether the equilibrium constant for the following reaction is greater than or smaller than 1: \(\mathrm{CH}_{3} \mathrm{COOH}(a q)+\mathrm{HCOO}^{-}(a q)\ \leftrightharpoons\ \mathrm{CH}_{3} \mathrm{COO}^{-}(a q)+\mathrm{HCOOH}(a q)\)

What is the ion-product constant for water?

The pH of a 0.060 M weak monoprotic acid is 3.44. Calculate the \(K_{a}\) of the acid.

What an equation relating \(\left[H^{+}\right]\) and \(\left[\mathrm{OH}^{-}\right]\) in solution at \(25^{\circ} \mathrm{C}\).

Calculate the pH of a 0.26 M methylamine solution (see table 15.4)

Calculate the concentrations of \(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\), \(\mathrm{HC}_{2} \mathrm{O}_{4}^{-}\), \(\mathrm{C}_{2} \mathrm{O}_{4}^{2-}\), and \(H^{+}\) ions in a 0.20 M oxalic acid solution.

Define pH. Why do chemists normally choose to discuss the acidity of a solution in terms of pH rather than hydrogen ion concentration, \(\left[H^{+}\right]\)?

Which of the following acids is weaker: \(\mathrm{HClO}_{2}\) or \(\mathrm{HClO}_{3}\)?

Calculate the pH of a 0.24 M sodium formate solution (HCOONa).

Define pOH. Write the equation relating pH and pOH.

Predict whether the following solutions will be acidic, basic or nearly neutral: (a) \(\mathrm{LiClO}_{4}\), (b) \(N a_{3} P O_{4}\), (c) \(B i\left(N O_{3}\right)_{3}\), (d) \(\mathrm{NH}_{4} \mathrm{CN}\).

Calculate the concentration of \(O H^{-}\) ions in a \(1.4 \times 10^{-3}\) M HCl solution.

Calculate the concentration of \(H^{+}\) ions in a 0.62 M NaOH solution.

Calculate the pH of each of the following solutions: (a) \(2.8 \times 10^{-4}\ \mathrm{M\ Ba}(\mathrm{OH})_{2}\), (b) \(5.2 \times 10^{-4}\ \mathrm{M\ HNO}_{3}\).

Use \(\mathrm{NH}_{3}\) to illustrate what we mean by the strength of a base.

Which of the following has a higher pH? (a) 0.20 M \(\mathrm{NH}_{3}\), (b) 0.20 M NaOH

The diagrams here represent three different weak base solutions of equal concentration. List the bases in order of increasing \(K_{b}\) value. (Water molecules are omitted for clarity.)

Calculate the pH for each of the following solutions: (a) 0.10 M \(\mathrm{NH}_{3}\), (b) 0.050 M \(\mathrm{C}_{5} \mathrm{H}_{5} \mathrm{N}\) (pyridine).

The pH of a 0.30 M solution of a weak base is 10.66. What is the \(K_{b}\) of the base?

What is the original molarity of a solution of ammonia whose pH is 11.22?

In a 0.080 M \(\mathrm{NH}_{3}\) solution, what percent of the \(\mathrm{NH}_{3}\) is present as \(\mathrm{NH}_{4}^{+}\)?

Write the equation relating \(K_{a}\) for a weak acid and \(K_{b}\) for its conjugate base. Use \(\mathrm{NH}_{3}\) and its conjugate acid \(\mathrm{NH}_{4}^{+}\) to derive the relationship between \(K_{a}\) and \(K_{b}\).

From the relationship \(K_{a} K_{b}=K_{w}\), what can you deduce about the relative strengths of a weak acid and its conjugate base?

Carbonic acid is a diprotic acid. Explain what that means.

Write all the species (except water) that are present in a phosphoric acid solution. Indicate which species can act as a Brønsted acid, which as a Brønsted base, and which as both a Brønsted acid and a Brønsted base.

The first and second ionization constants of a diprotic acid \(\mathrm{H}_{2} \mathrm{A}\) are \(K_{a_{1}}\) and \(K_{a_{2}}\) at a certain temperature. Under what conditions will \(\left[A^{2-}\right]=K_{a_{2}}\)?

Compare the pH of a 0.040 M HCl solution with that of a 0.040 M \(\mathrm{H}_{2} \mathrm{SO}_{4}\) solution. (Hint: \(\mathrm{H}_{2} \mathrm{SO}_{4}\) is a strong acid; \(K_{a}\) for \(\mathrm{HSO}_{4}^{-}=1.3 \times 10^{-2}\).)

What are the concentrations of \(\mathrm{HSO}_{4}^{-}\), \(\mathrm{SO}_{4}^{2-}\), and \(H^{+}\) in a 0.20 M \(\mathrm{KHSO}_{4}\) solution?

Calculate the concentrations of \(H^{+}\), \(\mathrm{HCO}_{3}^{-}\), and \(\mathrm{CO}_{3}^{2-}\) in a 0.025 M \(\mathrm{H}_{2} \mathrm{CO}_{3}\) solution.

List four factors that affect the strength of an acid.

How does the strength of an oxoacid depend on the electronegativity and oxidation number of the central atom?

Predict the acid strengths of the following compounds: \(\mathrm{H}_{2} \mathrm{O}\), \(\mathrm{H}_{2} \mathrm{S}\), and \(\mathrm{H}_{2} \mathrm{Se}\).

Compare the strengths of the following pairs of acids: (a) \(\mathrm{H}_{2} \mathrm{SO}_{4}\) and \(\mathrm{H}_{2} \mathrm{SeO}_{4}\), (b) \(\mathrm{H}_{3} \mathrm{PO}_{4}\) and \(\mathrm{H}_{3} \mathrm{AsO}_{4}\).

Which of the following is the stronger acid: \(\mathrm{CH}_{2} \mathrm{ClCOOH}\) or \(\mathrm{CHCl}_{2} \mathrm{COOH}\)? Explain your choice.

Consider the following compounds: Experimentally, phenol is found to be a stronger acid than methanol. Explain this difference in terms of the structures of the conjugate bases. (Hint: A more stable conjugate base favors ionization. Only one of the conjugate bases can be stabilized by resonance.)

Define salt hydrolysis. Categorize salts according to how they affect the pH of a solution.

Explain why small, highly charged metal ions are able to undergo hydrolysis.

\(A l^{3+}\) is not a Brønsted acid but \(\mathrm{Al}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}^{3+}\) is. Explain.

Specify which of the following salts will undergo | hydrolysis: \(\mathrm{KF},\ \mathrm{NaNO}_{3},\ \mathrm{NH}_{4} \mathrm{NO}_{2},\ \mathrm{MgSO}_{4},\ \mathrm{KCN},\ \mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COONa},\ \mathrm{RbI},\ \mathrm{Na}_{2} \mathrm{CO}_{3},\ \mathrm{CaCl}_{2},\ \mathrm{HCOOK}\).

Predict the pH \((>7,\ <7,\ \text { or } \approx\ 7)\) of aqueous solutions containing the following salts: (a) KBr, (b) \(A l\left(\mathrm{NO}_{3}\right)_{3}\), (c) \(\mathrm{BaCl}_{2}\), (d) \(B i\left(\mathrm{NO}_{3}\right)_{3}\).

Predict whether the following solutions are acidic, basic, or nearly neutral: (a) NaBr, (b) \(\mathrm{K}_{2} \mathrm{SO}_{3}\), (c) \(\mathrm{NH}_{4} \mathrm{NO}_{2}\), (d) \(\mathrm{Cr}\left(\mathrm{NO}_{3}\right)_{3}\).

A certain salt, MX (containing the \(M^{+}\) and \(X^{-}\) ions), is dissolved in water, and the pH of the resulting solution is 7.0. Can you say anything about the strengths of the acid and the base from which the salt is derived?

In a certain experiment a student finds that the pHs of 0.10 M solutions of three potassium salts KX, KY, and KZ are 7.0, 9.0, and 11.0, respectively. Arrange the acids HX, HY, and HZ in the order of increasing acid strength.

Calculate the pH of a 0.36 M \(\mathrm{CH}_{3} \mathrm{COONa}\) solution.

The pH of a 0.0642 M solution of a monoprotic acid is 3.86. Is this a strong acid?

Like water, liquid ammonia undergoes autoionization: \(\mathrm{NH}_{3}+\mathrm{NH}_{3} \rightleftharpoons \mathrm{NH}_{4}^{+}+\mathrm{NH}_{2}^{-}\) (a) Identify the Brønsted acids and Brønsted bases in this reaction. (b) What species correspond to \(\mathrm {H^+}\) and \(\mathrm {OH^-}\) and what is the condition for a neutral solution?

HA and HB are both weak acids although HB is the stronger of the two. Will it take a larger volume of a 0.10 M NaOH solution to neutralize 50.0 mL of 0.10 M HB than would be needed to neutralize 50.0 mL of 0.10 M HA?

A solution contains a weak monoprotic acid HA and its sodium salt NaA both at 0.1 M concentration. Show that \(\mathrm {[OH^?]}=K_\mathrm w/K_\mathrm a\).

Use the data in Table 15.3 to calculate the equilibrium constant for the following reaction: \(\mathrm{HCOOH}(a q)+\mathrm{OH}^{-}(a q) \rightleftharpoons \mathrm{HCOO}^{-}(a q)+\mathrm{H}_{2} \mathrm{O}(l)\)

Use the data in Table 15.3 to calculate the equilibrium constant for the following reaction: \(\mathrm{CH}_{3} \mathrm{COOH}(a q)+\mathrm{NO}_{2}^{-}(a q) \rightleftharpoons \mathrm{CH}_{3} \mathrm{COO}^{-}(a q)+\mathrm{HNO}_{2}(a q)\)

Most of the hydrides of Group 1A and Group 2A metals are ionic (the exceptions are \(\mathrm{BeH_2}\) and \(\mathrm{MgH_2}\), which are covalent compounds). (a) Describe the reaction between the hydride ion \(\mathrm{(H^?)}\) and water in terms of a Brønsted acid-base reaction. (b) The same reaction can also be classified as a redox reaction. Identify the oxidizing and reducing agents.

Calculate the pH of a 0.20 M ammonium acetate \(\mathrm{(CH_3COONH_4)}\) solution.

Novocaine, used as a local anesthetic by dentists, is a weak base \((K_\mathrm b = 8.91 \times 10^{?6})\). What is the ratio of the concentration of the base to that of its acid in the blood plasma (pH = 7.40) of a patient?

Which of the following is the stronger base: \(\mathrm {NF_3}\) or \(\mathrm {NH_3}\)? (Hint: F is more electronegative than H.)

Which of the following is a stronger base: \(\mathrm{NH}_3\) or \(\mathrm{PH}_3\)? (Hint. The N - H bond is stronger than the P-H bond.)

The ion product of \(\mathrm{D_2O}\) is \(\mathrm{1.35 \times 10^{?15}}\) at \(\mathrm{25^\circ C}\). (a) Calculate pD where \(\mathrm{pD = ?\log [D^+]}\). (b) For what values of pD will a solution be acidic in \(\mathrm{D_2O}\)? (c) Derive a relation between pD and pOD.

Give an example of (a) a weak acid that contains oxygen atoms, (b) a weak acid that does not contain oxygen atoms, (c) a neutral molecule that acts as a Lewis acid, (d) a neutral molecule that acts as a Lewis base, (e) a weak acid that contains two ionizable H atoms, (f) a conjugate acid-base pair, both of which react with HCl to give carbon dioxide gas.

What is the pH of 250.0 mL of an aqueous solution containing 0.616 g of the strong acid trifluorometh ane sulfonic acid \(\mathrm{(CF_3SO_3H)}\)?

(a) Use VSEPR to predict the geometry of the hydronium ion, \(\mathrm{H_3O^+}\). (b) The O atom in \(\mathrm{H_2O}\) has two lone pairs and in principle can accept two \(\mathrm{H^+}\) ions. Explain why the species \(\mathrm{H_4O^{2+}}\) does not exist. What would be its geometry if it did exist?

HF is a weak acid, but its strength increases with concentration. Explain. (Hint: \(\mathrm{F^-}\) reacts with HF to form \(\mathrm{HF_2^-}\). The equilibrium constant for this reaction is 5.2 at \(\mathrm{25^\circ C}\).)

When chlorine reacts with water, the resulting solution is weakly acidic and reacts with \(\mathrm{AgNO_3}\) to give a white precipitate. Write balanced equations to represent these reactions. Explain why manufacturers of household bleaches add bases such as NaOH to their products to increase their effectiveness.

When the concentration of a strong acid is not substantially higher than \(1.0 \times 10^{?7}\) M, the ionization of water must be taken into account in the calculation of the solution’s pH. (a) Derive an expression for the pH of a strong acid solution, including the contribution to \([\mathrm H^+]\) from \(\mathrm{H_2O}\). (b) Calculate the pH of a \(1.0 \times 10^{?7}\) M HCl solution.

Calculate the pH of a 2.00 M \(\mathrm {NH_4CN}\) solution.

Calculate the concentrations of all the species in a 0.100 M \(\mathrm{Na_2CO_3}\) solution.

Explain the action of smelling salt, which is ammonium carbonate \(\mathrm{[(NH_4)_2CO_3]}\), (Hint: The thin film of aqueous solution that lines the nasal passage is slightly basic.)

About half of the hydrochloric acid produced annually in the United States (3.0 billion pounds) is used in metal pickling. This process involves the removal of metal oxide layers from metal surfaces to prepare them for coating. (a) Write the overall and net ionic equations for the reaction between iron(III) oxide, which represents the rust layer over iron, and HCl. Identify the Brønsted acid and base. (b) Hydrochloric acid is also used to remove scale (which is mostly \(\mathrm{CaCO_3}\)) from water pipes (see p. 126). Hydrochloric acid reacts with calcium carbonate in two stages; the first stage forms the bicarbonate ion, which then reacts further to form carbon dioxide. Write equations for these two stages and for the overall reaction. (c) Hydrochloric acid is used to recover oil from the ground. It dissolves rocks (often \(\mathrm{CaCO_3}\)) so that the oil can flow more easily. In one process, a 15 percent (by mass) HCI solution is injected into an oil well to dissolve the rocks. If the density of the acid solution is 1.073 g/mL. what is the pH of the solution?

Which of the following does not represent a Lewis acid-base reaction? (a) \(\mathrm{H}_{2} \mathrm{O}+\mathrm{H}^{+} \longrightarrow \mathrm{H}_{3} \mathrm{O}^{+}\) (b) \(\mathrm{NH}_{3}+\mathrm{BF}_{3} \longrightarrow \mathrm{H}_{3} \mathrm{NBF}_{3}\) (c) \(\mathrm{PF}_{3}+\mathrm{F}_{2} \longrightarrow \mathrm{PF}_{5}\) (d) \(\mathrm{Al}(\mathrm{OH})_{3}+\mathrm{OH}^{-} \longrightarrow \mathrm{Al}(\mathrm{OH})_{4}^{-}\)

True or false? If false, explain why the statement is wrong. (a) All Lewis acids are Brønsted acids (b) the conjugate base of an acid always carries a negative charge (c) the percent ionization of a base increases with its concentration in solution (d) a solution of barium fluoride is acidic

How many milliliters of a strong monoprotic acid solution at pH = 4.12 must be added to 528 mL of the same acid solution at pH = 5.76 to change its pH to 5.34? Assume that the volumes are additive.

Calculate the pH and percent ionization of a 0.80 M \(\mathrm{HNO_2}\) solution.

Consider the two weak acids HX (molar mass = 180 g/mol) and HY (molar mass = 78.0 g/mol). If a solution of 16.9 g/L of HX has the same pH as one containing 9.05 g/L of HY. which is the stronger acid at these concentrations?

A 1.294-g sample of a metal carbonate \(\mathrm {(MCO_3)}\) is reacted with 500 mL of a 0.100 M HCl solution. The excess HC1 acid is then neutralized by 32.80 mL of 0.588 M NaOH. Identify M.

Prove the statement that when the concentration of a weak acid HA decreases by a factor of 10, its percent ionization increases by a factor of \(\sqrt {10}\). State any assumptions.

Calculate the pH of a solution that is 1.00 M HCN and 1.00 M HF. Compare the concentration (in molarity) of the \(\mathrm{CN^-}\) ion in this solution with that in a 1.00 M HCN solution. Comment on the difference.

Use the van’t Hoff equation (see Problem 14.119) and the data in Appendix 3 to calculate the pH of water at its normal boiling point.

At \(\mathrm {28^\circ C}\) and 0.982 atm, gaseous compound HA has a density of 1.16 g/L. A quantity of 2.03 g of this compound is dissolved in water and diluted to exactly 1 L. If the pH of the solution is.5.22 (due to the ionization of HA) at \(\mathrm {25^\circ C}\), calculate the \(K_\mathrm {a}\) of the acid.

A 10.0-g sample of white phosphorus was burned in an excess of oxygen. The product was dissolved in enough water to make 500 mL of solution. Calculate the pH of the solution at \(\mathrm {25^\circ C}\).

Calculate the pH of a 0.20 M \(\mathrm {NaHCO_3}\) solution. (Hint: As an approximation, calculate hydrolysis and ionization separately first, followed by partial neutralization.)

(a) Shown here is a solution containing hydroxide ions and hydronium ions. What is the pH of the solution? (b) How many \(\mathrm {H_3O^+}\) ions would you need to draw for every \(\mathrm {OH^-}\) ion if the pH of the solution is 5.0? The color codes are \(\mathrm {H_3O^+}\) (red) and \(\mathrm {OH^-}\) (green). Water molecules and counter ions are omitted for clarity.

In this chapter. HCl, HBr, and HI are all listed as strong acids because they are assumed to be ionized completely in water. If, however, we choose a solvent such as acetic acid that is a weaker Brønsted base than water, it is possible to rank the acids in increasing strength as HCl < HBr < HI. (a) Write equations showing proton transfer between the acids and \(\mathrm {CH_3COOH}\). Describe how you would compare the strength of the acids in this solvent experimentally. (b) Draw a Lewis structure of the conjugate acid \(\mathrm {CH_3COOH_2^+}\).

Use the data in Appendix 3 to calculate the \(\Delta H\mathrm {^\circ_{rxn}}\) for the following reactions: (a) \(\mathrm{NaOH}(a q)+\mathrm{HCl}(a q) \rightarrow \mathrm{NaCl}(a q)+\mathrm{H}_{2} \mathrm{O}(l)\) and (b) \(\mathrm{KOH}(a q)+\mathrm{HNO}_{3}(a q) \rightarrow \mathrm{KNO}_{3}(a q)+\mathrm{H}_{2} \mathrm{O}(l)\). Comment on your results.

Malonic acid \(\mathrm{[CH_2(COOH)_2]}\) is a diprotic acid. Compare its two \(K_\mathrm a\) values with that of acetic acid \(\mathrm{(CH_3COOH)}\) \((K_\mathrm{a})\). and account for the differences in the three \(K_\mathrm a\) values.

Look up the contents of a Turns tablet. How many tablets are needed to increase the pH of the gastric juice in a person’s stomach from 1.2 to 1.5?

Phosphorous acid, \(\mathrm{H_3PO_3}(aq)\), is a diprotic acid with \(K_\mathrm {a_1}=3\times10^{-2}\) (a) After examining the \(K_\mathrm a\) values in Table 15.5, estimate \(K_\mathrm {a_2}\) for \(\mathrm{H_3PO_3}(aq)\) and calculate the pH of a 0.10 M solution of \(\mathrm{Na_2HPO_3}(aq)\). (b) The structure of \(\mathrm{H_3PO_3}\) is given in Figure 15.5. Explain why \(\mathrm{H_3PO_4}(aq)\) is a triprotic acid, but \(\mathrm{H_3PO_3}(aq)\) is only a diprotic acid.

Chicken egg shells are composed primarily of calcium carbonate, \(\mathrm{CaCO_3}\). In a classic demonstration carried out in chemistry and biology classes, vinegar is used to remove the shell from an uncooked egg, revealing the semipermeable membrane that surrounds the egg keeping it intact. Refer to the Chemical Mystery on p. 776 to see a schematic diagram of a chicken egg. Estimate the minimum amount of vinegar required to remove the entire shell from the egg.

Identify the conjugate acid-base pairs of the reaction \(\mathrm{CN}+\mathrm{H}_{2} \mathrm{O}\ \leftrightharpoons\ \mathrm{HCN}+\mathrm{OH}^{-}\)

Which of the following does not constitute a conjugate acid-base pair? (a) \(\mathrm{HNO}_2 \mathrm{NO}_2^{-}\), (b) \(\mathrm{H}_2 \mathrm{CO}_{3-} \mathrm{CO}_3^{2-}\), (c) \(\mathrm{CH}_3 \mathrm{NH}_{3-}^{+} \mathrm{CH}_3 \mathrm{NH}_2\).

Nitric acid \(\left(\mathrm{HNO}_{3}\right)\) is used in the production of fertilizer, dyes, drugs, and explosives. Calculate the pH of a \(\mathrm{HNO}_{3}\) solution having a hydrogen ion concentration of 0.76 M.

The pH of a certain orange juice is 3.33. Calculate the \(H^{+}\) ion concentration.

Calculate the pH of a \(1.8 \times 10^{-2}\ \mathrm{M\ Ba}(\mathrm{OH})_{2}\) solution.

What is the pH of a 0.122 M monoprotic acid whose \(K_{a}\) is \(5.7 \times 10^{-4}\)?

The “concentration” of water is 55.5 M. Calculate the percent ionization of water.

Consider the following two acids and their ionization constants: \(H C O O H\ \ \ \ \ \ \ \ \ K_{a}=1.7 \times 10^{-4}\) \(H C N\ \ \ \ \ \ \ \ \ K_{a}=4.9 \times 10^{-10}\) Which conjugate base \(\left(H C O O^{-} \text {or }\ \mathrm{CN}^{-}\right)\) is stronger?

Which of the diagrams shown here represents a solution of sulfuric acid? Water molecules have been omitted for clarity.

The diagrams shown here represent solutions of three salts \(\mathrm{NaX}(\mathrm{X}=\mathrm{A}, \mathrm{B}, or \mathrm{C})\). (a) Which \(\mathrm{X}^{-}\) has the weakest conjugate acid? (b) Arrange the three \(X^{-}\) anions in order of increasing base strength. The \(\mathrm{Na}^{+}\) ion water molecules have been omitted for clarity.

Identify the Lewis acid and Lewis base in the reaction \(\mathrm{Co}^{3+}(a q)+6 \mathrm{NH}_{3}(a q)\ \leftrightharpoons\ \mathrm{Co}\left(\mathrm{NH}_{3}\right)_{6}^{3+}(a q)\)

Define Brønsted acids and bases. Give an example of a conjugate pair in an acid-base reaction.

Classify each of the following species as a Brønsted acid or base, or both: (a) \(\mathrm{H}_{2} \mathrm{O}\), (b) \(\mathrm{OH}^{-}\), (c) \(\mathrm{H}_{3} \mathrm{O}^{-}\), (d) \(\mathrm{NH}_{3}\), (e) \(\mathrm{NH}_{4}^{+}\), (f) \(\mathrm{NH}_{2}^{-}\), (g) \(\mathrm{NO}_{3}^{-}\), (h) \(\mathrm{CO}_{3}^{2-}\), (i) \(H B r\), j) \(H C N\).

Identify the acid-base conjugate pairs in each of the following reactions: (a) \(\mathrm{CH}_{3} \mathrm{COO}^{-}+\mathrm{HCN}\ \leftrightharpoons\ \mathrm{CH}_{3} \mathrm{COOH}+\mathrm{CN}^{-}\) (b) \(\mathrm{HCO}_{3}^{-}+\mathrm{HCO}_{3}^{-}\ \leftrightharpoons\ \mathrm{H}_{2} \mathrm{CO}_{3}+\mathrm{CO}_{3}^{2-}\) (c) \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}+\mathrm{NH}_{3}\ \leftrightharpoons\ \mathrm{HPO}_{4}^{2-}+\mathrm{NH}_{4}^{+}\) (d) \(\mathrm{HClO}+\mathrm{CH}_{3} \mathrm{NH}_{2}\ \leftrightharpoons\ \mathrm{CH}_{3} \mathrm{NH}_{3}^{+}+\mathrm{ClO}^{-}\) (e) \(\mathrm{CO}_{3}^{2-}+\mathrm{H}_{2} \mathrm{O}\ \leftrightharpoons\ \mathrm{HCO}_{3}^{-}+\mathrm{OH}^{-}\)

Write the formula for the conjugate acid of each of the following bases: (a) \(H S^{-}\) (b) \(\mathrm{HCO}_3^{-}\), (c) \(\mathrm{CO}_3^{2-}\), (d) \(\mathrm{H}_2 \mathrm{POO}_4^{-}\), (e) \(\mathrm{HPO}_4^{2-}\), (f) \(\mathrm{PO}_4^{3-}\), (g) \(\mathrm{HSO}_4^{-}\), (h) \(\mathrm{SO}_4^{2-}\), (i) \(\mathrm{SO}_3^{2-}\).

Oxalic acid \(\left(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\right)\) has the following structure: An oxalic acid solution contains the following species in varying concentrations: \(\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\), \(\mathrm{HC}_{2} \mathrm{O}_{4}^{-}\), \(\mathrm{C}_{2} \mathrm{O}_{4}^{2-}\), and \(H^{+}\). (a) Draw Lewis structures of \(\mathrm{HC}_{2} \mathrm{O}_{4}^{-}\), and \(\mathrm{C}_{2} \mathrm{O}_{4}^{2-}\). (b) Which of the above four species can act only as acids, which can act only as bases, and which can act as both acids and bases?

Write the formula for the conjugate base of each of the following acids: (a) \(\mathrm{CH}_2 \mathrm{ClCOOH}_3\), (b) \(\mathrm{HIO}_4\), (c) \(\mathrm{H}_3 \mathrm{PO}_4\), (d) \(\mathrm{H}_2 \mathrm{PO}_4^{-}\), (e) \(\mathrm{HPO}_4^{2-}\), (f) \(\mathrm{H}_2 \mathrm{SO}_4\), (g) \(\mathrm{HSO}_4^{-}\), (h) \(\mathrm{HIO}_3\), (i) \(\mathrm{HSO}_3^{-}\), (j) \(\mathrm{NH}_4^{+}\), (k) \(\mathrm{H}_2 \mathrm{~S}\), (l) \(\mathrm{HS}^{-}\), (m) \(\mathrm{HCIO}\).

The ion-product constant for water is \(1.0 \times 10^{-14}\) at \(25^{\circ} \mathrm{C}\) and \(3.8 \times 10^{-14}\) at \(40^{\circ} \mathrm{C}\). Is the forward process \(\mathrm{H}_2 \mathrm{O}(l) \rightleftharpoons \mathrm{H}^{+}(a q)+\mathrm{OH}^{-}(a q)\) endothermic or exothermic?

The pH of a solution is 6.7. From this statement alone, can you conclude that the solution is acidic? If not, what additional information would you need? Can the pH of a solution be zero or negative? If so, give examples to illustrate these values.

Calculate the pH of each of the following solutions: (a) 0.0010 M HCI, (b) 0.76 M KOH.

Calculate the hydrogen ion concentration in mol/L for solutions with the following pH values: (a) 2.42, (b) 11.21, (c) 6.96, (d) 15.00.

Calculate the hydrogen ion concentration in mol/L for each of the following solutions: (a) a solution whose pH is 5.20, (b) a solution whose pH is 16.00, (c) a solution whose hydroxide concentration is \(3.7 \times 10^{-9} \mathrm{~N}\)

Complete the following table for a solution:

Fill in the word acidic, basic, or neutral for the following solutions: (a) pOH > 7; solution is (b) pOH = 7; solution is (c) pOH < 7; solution is

The pOH of a strong base solution is 1.88 at \(25^{\circ} \mathrm{C}\). Calculate the concentration of the base (a) if the base is KOH and (b) if the base is \(B a(O H)_2\).

Calculate the number of moles of KOH in 5.50 mL of a 0.360 M KOH solution. What is the pOH of the solution?

How much NaOH in grams) is needed to prepare 546 mL of solution with a pH of 10.00?

A solution is made by dissolving 18.4 g of HCl in 662 mL of water. Calculate the pH of the solution. (Assume that the volume remains constant.)

Explain what is meant by the strength of an acid.

Without referring to the text, write the formulas of four strong acids and four weak acids.

What are the strongest acid and strongest base that can exist in water?

\(\mathrm{H}_{2} \mathrm{SO}_{4}\) is a strong acid, but \(\mathrm{HSO}_{4}^{-}\) is a weak acid. Account for the difference in strength of these two related species.

Which of the following diagrams best represents a strong acid, such as HCI, dissolved in water? Which represents a weak acid? Which represents a very weak acid? (The hydrated proton is shown as a hydronium ion. Water molecules are omitted for clarity.)

(1) Which of the following diagrams represents a solution of a weak diprotic acid? (2) Which diagrams represent chemically implausible situations? (The hydrated proton is shown as a hydronium ion. Water molecules are omitted for clarity.)

Classify each of the following species as a weak or strong acid: (a) \(\mathrm{HNO}_3\), (b) HF, (c) \(\mathrm{H}_2 \mathrm{SO}_4\), (d) \(\mathrm{HSO}_4^{-}\), (e) \(\mathrm{H}_2 \mathrm{CO}_3\), (f) \(\mathrm{HCO}_3^{-}\), (g) HCl, (h) HCN, (i) \(\mathrm{HNO}_2\).

Classify each of the following species as a weak or strong base: (a) LiOH, (b) \(C N^{-}\), (c) \(\mathrm{H}_{2} \mathrm{O}\), (d) \(\mathrm{CIO}_{4}^{-}\), (e) \(\mathrm{NH}_{2}^{-}\).

Which of the following statements is/are true for a 0.10 M solution of a weak acid HA? (a) The pH is 1.00. (b) \(\left[H^{+}\right]\) \(\left[A^{-}\right]\) (c) \(\left[H^{+}\right]=\left[A^{-}\right]\) (d) The pH is less than 1.

Which of the following statements is/are true regarding a 1.0 M solution of a strong acid HA? (a) \(\left[A^{-}\right]>\left[H^{+}\right]\) (b) The pH is 0.00. (c) \(\left[H^{+}\right]=1.0\ \mathrm{M}\) (d) [HA] = 1.0M

Predict the direction that predominates in this reaction: \(F^{-}(a q)+H_{2} O(l)\ \leftrightharpoons\ H F(a q)+O H^{-}(a q)\)

Predict whether the following reaction will proceed from left to right to any measurable extent: \(\mathrm{CH}_{3} \mathrm{COOH}(a q)+\mathrm{Cl}^{-}(a q)\ \rightarrow\)

What does the ionization constant tell us about the strength of an acid?

List the factors on which the \(K_{a}\) of a weak acid depends.

Why do we normally not quote \(K_{a}\) values for strong acids such as HCl and \(\mathrm{HNO}_{3}\)? Why is it necessary to specify temperature when giving \(K_{a}\) values?

Which of the following solutions has the highest pH ? (a) \(0.40 \mathrm{M} \mathrm{HCOOH}\), (b) \(0.40 \mathrm{M} \mathrm{HCIO}_4\), (c) \(0.40 \mathrm{M} \mathrm{CH}_3 \mathrm{COOH}\).

The \(K_{a}\) for benzoic acid is \(6.5 \times 10^{-5}\). Calculate the pH of a 0.10 M benzoic acid solution.

A 0.0560-g quantity of acetic acid is dissolved in enough water to make 50.0 mL of solution. Calculate the concentrations of \(\mathrm{H}^{+}, \mathrm{CH}_3 \mathrm{COO}^{-}\), and \(\mathrm{CH}_3 \mathrm{COOH}\) at equilibrium. \(\left(K_{\mathrm{a}}\right.\) for acetic acid \(=1.8 \times 10^{-5}\).)

The pH of an acid solution is 6.20. Calculate the \(K_{a}\) for the acid. The initial acid concentration is 0.010 M.

What is the original molarity of a solution of formic acid (HCOOH) whose pH is 3.26 at equilibrium?

Calculate the percent ionization of benzoic acid having the following concentrations: (a) 0.20 M, (b) 0.00020 M.

Calculate the percent ionization of hydrofluoric acid at the following concentrations: (a) 0.60 M, (b) 0.0046 M. (c) 0.00028 M. Comment on the trends.

A 0.040 M solution of a monoprotic acid is 14 percent ionized. Calculate the ionization constant of the acid.

(a) Calculate the percent ionization of a 0.20 M solution of the monoprotic acetylsalicylic acid (aspirin) for which \(K_{a}=3.0 \times 10^{-4}\). (b) The pH of gastric juice in the stomach of a certain individual is 1.00. After a few aspirin tablets have been swallowed, the concentration of acetylsalicylic acid in the stomach is 0.20 M. Calculate the percent ionization of the acid under these conditions. What effect does the nonionized acid have on the membranes lining the stomach? (Hint: See the Chemistry in Action essay on p. 708.)

Calculate the pH of a 0.24 M solution of a weak base with a \(K_{b}\) of \(3.5 \times 10^{-6}\).

Calculate the pH of a 0.42 M \(\mathrm{NH}_{4} \mathrm{Cl}\) solution.

Predict the pH \((>7,\ <7,\ \approx\ 7)\) of a \(\mathrm{NaHCO}_{3}\) solution.

Predict whether a solution containing the salt \(\mathrm{K}_{2} \mathrm{HPO}_{4}\) will be acidic, neutral, or basic.

Classify the following oxides as acidic, basic, amphoteric, or neutral: (a) \(\mathrm{CO}_{2}\), (b) \(\mathrm{K}_{2} \mathrm{O}\), (c) \(\mathrm{CaO}\), (d) \(\mathrm{N}_{2} \mathrm{O}_{5}\), (e) \(\mathrm{CO}\) (f) \(N O\), (g) \(\mathrm{SnO}_{2}\), (h) \(\mathrm{SO}_{3}\), (i) \(\mathrm{Al}_{2} \mathrm{O}_{3}\), (j) \(B a O\).

Write equations for the reactions between (a) \(\mathrm{CO}_{2}\) and \(\mathrm{NaOH}(a q)\), (b) \(\mathrm{Na}_{2} \mathrm{O}\), and \(\mathrm{HNO}_{3}(a q)\).

Explain why metal oxides tend to be basic if the oxidation number of the metal is low and acidic if the oxidation number of the metal is high. (Hint: Metallic compounds in which the oxidation numbers of the metals are low are more ionic than those in which the oxidation numbers of the metals are high.)

Arrange the oxides in each of the following groups in order of increasing basicity: (a) \(\mathrm{K}_2 \mathrm{O}, \mathrm{Al}_2 \mathrm{O}_3, \mathrm{BaO}\), (b) \(\mathrm{CrO}_3, \mathrm{CrO}, \mathrm{Cr}_2 \mathrm{O}_3\).

\(\mathrm{Zn}(\mathrm{OH})_2\) is an amphoteric hydroxide. Write balanced ionic equations to show its reaction with (a) HCl. (b) NaOH [the product is \(\mathrm{Zn}(\mathrm{OH})_4^{2-}\)].

\(A l(O H)_{3}\) is an insoluble compound. It dissolves in excess NaOH in solution. Write a balanced ionic equation for this reaction. What type of reaction is this?

What are the Lewis Definition of an acid and a base? In what way are they more general than the Brønsted definitions?

In terms of orbitals and electron arrangements, what must be present for a molecule or an ion to act as Lewis acid (use \(\mathrm H^+\) and \(\mathrm {BF_3}\) as examples)? What must be present for a molecule or ion to act as a Lewis base (use \(\mathrm {OH^-}\) and \(\mathrm {NH_3}\) as examples)?

Classify each of the following species as a Lewis acid or a Lewis base: (a) \(\mathrm{CO}_{2}\) (b) \(\mathrm{H}_{2} \mathrm{O}\) (c) \(\mathrm{I}^{-}\) (d) \(\mathrm{SO}_{2}\) (e) \(\mathrm{NH}_{3}\) (f) \(\mathrm{OH}^{-}\) (g) \(\mathrm{H}^{+}\) (h) \(\mathrm{BCl}_{3}\)

Describe the following reaction in terms of the Lewis theory of acids and bases: \(\mathrm{AlCl}_{3}(s)+\mathrm{Cl}^{-}(a q) \longrightarrow \mathrm{AlCl}_{4}^{-}(a q)\)

Which would be considered a stronger Lewis acid: (a) \(\mathrm{BF}_{3}\) or \(\mathrm{BCl}_{3}\), (b) \(\mathrm{Fe}^{2+}\) or \(\mathrm{Fe}^{3+}\)? Explain.

All Brønsted acids are Lewis acids, but the reverse is not true. Give two examples of Lewis acids that are not Brønsted acids.

Determine the concentration of a \(\mathrm{NaNO}_{2}\) solution that has a pH of 8.22.

Determine the concentration of a \(\mathrm{NH}_{4} \mathrm{Cl}\) solution that has a pH of 5.64.

The diagrams here show three weak acids HA (A = X, Y, or Z) in solution. (a) Arrange the acids in order of increasing \(K_{a}\) (b) Arrange the conjugate bases in increasing order of \(K_{\mathrm{b}}\) (c) Calculate the percent ionization of each acid. (d) Which of the 0.1 M sodium salt solutions (NaX, NaY, or NaZ) has the lowest pH? (The hydrated proton is shown as a hydronium ion. Water molecules are omitted for clarity.)

A typical reaction between an antacid and the hydrochloric acid in gastric juice is \(\mathrm{NaHCO}_{3}(s)+\mathrm{HCl}(a q) \rightleftharpoons \mathrm{NaCl}(a q)+\mathrm{H}_{2} \mathrm{O}(l)+\mathrm{CO}_{2}(g)\) Calculate the volume (in L) of \(\mathrm{CO}_{2}\) generated from 0.350 g of \(\mathrm{NaHCO}_{3}\) and excess gastric juice at 1.00 atm and \(37.0^{\circ} \mathrm{C}\).

To which of the following would the addition of an equal volume of 0.60 M NaOH lead to a solution having a lower pH? (a) water (b) 0.30 M HCl (c) 0.70 M KOH (d) 0.40 M \(\mathrm{NaNO_3}\)

The three common chromium oxides are \(\mathrm{CrO}\), \(\mathrm{Cr}_{2} \mathrm{O}_{3}\), and \(\mathrm{CrO}_{3}\). If \(\mathrm{Cr}_{2} \mathrm{O}_{3}\) is amphoteric, what can you say about the acid-base properties of \(\mathrm{CrO}\) and \(\mathrm{CrO}_{3}\)?

Calculate the concentrations of all species in a 0.100 M \(\mathrm{H}_{3} \mathrm{PO}_{4}\) solution.

Identify the Lewis acid and Lewis base that lead to the formation of the following species: (a) \(\mathrm{AlCl}_{4}^{-}\) (b) \(\mathrm{Cd}(\mathrm{CN})_{4}^{2-}\) (c) \(\mathrm{HCO}_{3}^{-}\) (d) \(\mathrm{H}_{2} \mathrm{SO}_{4}\)

Very concentrated NaOH solutions should not be stored in Pyrex glassware. Why? (Hint: See Section 11.7.)

In the vapor phase, acetic acid molecules associate to a certain extent to form dimers: \(2 \mathrm{CH}_{3} \mathrm{COOH}(g) \rightleftharpoons\left(\mathrm{CH}_{3} \mathrm{COOH}\right)_{2}(g)\) At \(51^{\circ} \mathrm{C}\) the pressure of a certain acetic acid vapor system is 0.0342 atm in a 360-mL flask. The vapor is condensed and neutralized with 13.8 mL of 0.0568 M NaOH. (a) Calculate the degree of dissociation \((\alpha)\) of the dimer under these conditions: \(\left(\mathrm{CH}_{3} \mathrm{COOH}\right)_{2} \rightleftharpoons 2 \mathrm{CH}_{3} \mathrm{COOH}\) (Hint: See Problem 14.117 for general procedure.) (b) Calculate the equilibrium constant \(K_{P}\) for the reaction in (a).

Henry's law constant for \(\mathrm{CO}_{2}\) at \(38^{\circ} \mathrm{C}\) is \(2.28 \times 10^{-3}~\mathrm{mol/L \cdot atm}\). Calculate the pH of a solution of \(\mathrm{CO}_{2}\) at \(38^{\circ} \mathrm{C}\) in equilibrium with the gas at a partial pressure of 3.20 atm.

Hydrocyanic acid (HCN) is a weak acid and a deadly poisonous compound—in the gaseous form (hydrogen cyanide) it is used in gas chambers. Why is it dangerous to treat sodium cyanide with acids (such as HCl) without proper ventilation?

How many grams of NaCN would you need to dissolve in enough water to make exactly 250 mL of solution with a pH of 10.00?

A solution of formic acid (HCOOH) has a pH of 2.53. How many grams of formic acid are there in 100.0 mL of the solution?

Calculate the pH of a 1-L solution containing 0.150 mole of \(\mathrm{CH_3COOH}\) and 0.100 mole of HCl.

A 1.87-g sample of Mg reacts with 80.0 mL of a HCl solution whose pH is –0.544. What is the pH of the solution after all the Mg has reacted? Assume constant volume.

You are given two beakers, one containing an aqueous solution of strong acid (HA) and the other an aqueous solution of weak acid (HB) of the same concentration. Describe how you would compare the strengths of these two acids by (a) measuring the pH (b) measuring electrical conductance (c) studying the rate of hydrogen gas evolution when these solutions are reacted with an active metal such as Mg or Zn.

Use Le Châtelier's principle to predict the effect of the following changes on the extent of hydrolysis of sodium nitrite \(\left(\mathrm{NaNO}_{2}\right)\) solution: (a) HCl is added (b) NaOH is added (c) NaCl is added (d) the solution is diluted

Describe the hydration of \(\mathrm{SO}_{2}\) as a Lewis acid-base reaction. (Hint: Refer to the discussion of the hydration of \(\mathrm{CO}_{2}\) on p. 707.)

The disagreeable odor of fish is mainly due to organic compounds \(\left(\mathrm{RNH}_{2}\right)\) containing an amino group, \(-\mathrm{NH}_{2}\), where R is the rest of the molecule. Amines are bases just like ammonia. Explain why putting some lemon juice on fish can greatly reduce the odor.

A solution of methylamine \(\left(\mathrm{CH}_{3} \mathrm{NH}_{2}\right)\) has a pH of 10.64. How many grams of methylamine are there in 100.0 mL of the solution?

A 0.400 M formic acid (HCOOH) solution freezes at \(-0.758^{\circ} \mathrm{C}\). Calculate the \(K_{\mathrm{a}}\) of the acid at that temperature. (Hint: Assume that molarity is equal to molality. Carry your calculations to three significant figures and round off to two for \(K_{\mathrm{a}}\).)

Both the amide ion \(\left(\mathrm{NH}_{2}^{-}\right)\) and the nitride ion \(\left(\mathrm{N}^{3-}\right)\) are stronger bases than the hydroxide ion and hence do not exist in aqueous solutions. (a) Write equations showing the reactions of these ions with water, and identify the Brønsted acid and base in each case. (b) Which of the two is the stronger base?

The atmospheric sulfur dioxide \(\left(\mathrm{SO}_{2}\right)\) concentration over a certain region is 0.12 ppm by volume. Calculate the pH of the rainwater due to this pollutant. Assume that the dissolution of \(\mathrm{SO}_{2}\) does not affect its pressure.

Calcium hypochlorite \(\left[\mathrm{Ca}(\mathrm{OCl})_{2}\right]\) is used as a disinfectant for swimming pools. When dissolved in water it produces hypochlorous acid \(\mathrm{Ca}(\mathrm{OCl})_{2}(s)+2 \mathrm{H}_{2} \mathrm{O}(l) \underset{2 \mathrm{HClO}(a q)+\mathrm{Ca}(\mathrm{OH})_{2}(s)}{\rightleftharpoons}\) which ionizes as follows: \(\begin{aligned} \mathrm{HClO}(a q) \rightleftharpoons \mathrm{H}^{+}(a q)+\mathrm{ClO}^{-}(a q) & \\ &K_{\mathrm{a}}=3.0 \times 10^{-8} \end{aligned} \) As strong oxidizing agents, both HClO and \(\mathrm{ClO}^{-}\) can kill bacteria by destroying their cellular components. However, too high a HClO concentration is irritating to the eyes of swimmers and too high a concentration of \(\mathrm{ClO}^{-}\) will cause the ions to decompose in sunlight. The recommended pH for pool water is 7.8. Calculate the percent of these species present at this pH.

Hemoglobin (Hb) is a blood protein that is responsible for transporting oxygen. It can exist in the protonated form as \(\mathrm{HbH}^{+}\). The binding of oxygen can be represented by the simplified equation \(\mathrm{HbH}^{+}+\mathrm{O}_{2} \rightleftharpoons \mathrm{HbO}_{2}+\mathrm{H}^{+}\) (a) What form of hemoglobin is favored in the lungs where oxygen concentration is highest? (b) In body tissues, where the cells release carbon dioxide produced by metabolism, the blood is more acidic due to the formation of carbonic acid. What form of hemoglobin is favored under this condition? (c) When a person hyperventilates, the concentration of \(\mathrm{CO}_{2}\) in his or her blood decreases. How does this action affect the above equilibrium? Frequently a person who is hyperventilating is advised to breathe into a paper bag. Why does this action help the individual?

Teeth enamel is hydroxyapatite \(\left[\mathrm{Ca}_{3}\left(\mathrm{PO}_{4}\right)_{3} \mathrm{OH}\right]\). When it dissolves in water (a process called demineralization), it dissociates as follows: \(\mathrm{Ca}_{5}\left(\mathrm{PO}_{4}\right)_{3} \mathrm{OH} \longrightarrow 5 \mathrm{Ca}^{2+}+3 \mathrm{PO}_{4}^{3-}+\mathrm{OH}^{-}\) The reverse process, called remineralization, is the body's natural defense against tooth decay. Acids produced from food remove the \(\mathrm{OH}^{-}\) ions and thereby weaken the enamel layer. Most toothpastes contain a fluoride compound such as NaF or \(\mathrm{SnF}_{2}\). What is the function of these compounds in preventing tooth decay?