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Chapter 13: Problem 24

Milk of Magnesia has a pH of \(10.5\). (a) Calculate \(\left[\mathrm{H}^{+}\right]\). (b) Calculate the ratio of the \(\mathrm{H}^{+}\) concentration of gastric juice, \(\mathrm{pH} 1.5\), to that of Milk of Magnesia.

Short Answer

Expert verified
Answer: The concentration of H+ ions in Milk of Magnesia is 3.16 × 10^{-11} M, and the concentration in gastric juice is 3.16 × 10^{-2} M. The ratio of the concentration of H+ ions of gastric juice to that of Milk of Magnesia is 1,000,000,000:1 or 10^9:1.

Step by step solution

01

Calculate the concentration of H+ ions in Milk of Magnesia

Since pH = -log[H+], we can rewrite the formula to find the concentration of H+ ions ( [H+] ): [H+] = 10^(-pH) The question gives the pH of Milk of Magnesia as 10.5, so we can plug this value into the formula: [H+] = 10^(-10.5) Now, calculating the concentration of H+ ions: [H+] = 3.16 × 10^{-11} M
02

Calculate the concentration of H+ ions in gastric juice

Repeat the same process for gastric juice, with a given pH value of 1.5: [H+] = 10^(-1.5) Now, calculating the concentration of H+ ions in gastric juice: [H+] = 3.16 × 10^{-2} M
03

Calculate the ratio of H+ concentration of gastric juice to H+ concentration of Milk of Magnesia

To calculate the ratio, divide the concentration of H+ ions in gastric juice by the concentration of H+ ions in Milk of Magnesia: Ratio = (3.16 × 10^{-2} M) / (3.16 × 10^{-11} M) The units of Molarity (M) will cancel out: Ratio = 10^9 So, the ratio of the concentration of H+ ions of gastric juice to that of Milk of Magnesia is 1,000,000,000:1 or 10^9:1.

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Most popular questions from this chapter

Arrange the following \(0.1 M\) aqueous solutions in order of decreasing \(\mathrm{pH}\) (highest to lowest). \(\begin{array}{lllll}\mathrm{Ba}\left(\mathrm{NO}_{3}\right)_{2}, & \mathrm{HNO}_{3}, & \mathrm{NH}_{4} \mathrm{NO}_{3}, & \mathrm{Al}\left(\mathrm{NO}_{3}\right)_{3}, & \mathrm{NaF}\end{array}\)

Rank the following solutions in order of increasing \(\left[\mathrm{H}^{+}\right]\). \(\begin{array}{llll}0.1 & M \mathrm{HBr}, & 0.1 \mathrm{M} \mathrm{HF}, & 0.1 \mathrm{MHCHO}_{2} & 0.1 \mathrm{MHCN}\end{array}\)

What is the \(\mathrm{pH}\) of a \(0.200 \mathrm{M}\) solution of \(\mathrm{H}_{2} \mathrm{SO}_{4}\) ? You may assume that the first ionization is complete. The second ionization constant is \(0.010\).

Find the value of \(K_{\mathrm{a}}\) for the conjugate acid of the following bases. (a) pyridine, a pesticide; \(K_{\mathrm{b}}=1.5 \times 10^{-9}\) (b) aniline, an important dye intermediate; \(K_{\mathrm{b}}=3.8 \times 10^{-10}\)

Calculate \(\left[\mathrm{H}^{+}\right]\) and \(\left[\mathrm{OH}^{-}\right]\) in solutions with the following \(\mathrm{pH}\). (a) \(4.0\) (b) \(8.52\) (c) \(0.00\) (d) \(12.60\)
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