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Chapter 15: Problem 111

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

Short Answer

Expert verified
The ratio of the concentration of the base to that of its acid in the blood plasma of the patient is approximately 1.1223.

Step by step solution

01

Calculate the pOH of the blood plasma

Given that the pH is 7.4, calculate the pOH using the relationship pH + pOH = 14. So, pOH = 14 - 7.4 = 6.6.
02

Convert pOH to hydroxide ion concentration

Now convert pOH to the hydroxide ion concentration using the definition of pOH, which is \(-\log[OH^-]\). This implies that the concentration of hydroxide ions [OH^-] in the solution can be found by using the antilog, i.e., \(10^{-pOH}\). Hence, the [OH^-] = \(10^{-6.6}\) M.
03

Calculate the ratio of the base to the acid

Based on the equilibrium reaction \(B + H2O = HB+ + OH-\) where B stands for base (Novocaine) and HB+ is its conjugate acid, the \(K_b\) relation is \(K_b = [HB+] [OH-]/ [B]\). Isolating [B]/[HB+] to find the ratio gives \([B]/[HB+] = [OH-]/ K_b\). Substituting the given values gives the ratio = \(10^{-6.6} / 8.91 * 10^{-6} = 1.1223\). This implies that the concentration of the base is 1.1223 times that of the acid in the blood plasma of the patient.

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Most popular questions from this chapter

What is the \(\mathrm{pH}\) of \(250.0 \mathrm{~mL}\) of an aqueous solution containing \(0.616 \mathrm{~g}\) of the strong acid trifluoromethane sulfonic acid (CF \(_{3} \mathrm{SO}_{3} \mathrm{H}\) )?

In terms of orbitals and electron arrangements, what must be present for a molecule or an ion to act as a 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)?

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: $$ \begin{aligned} \mathrm{Ca}(\mathrm{OCl})_{2}(s)+2 \mathrm{H}_{2} \mathrm{O}(l) & \rightleftharpoons \\ & 2 \mathrm{HClO}(a q)+\mathrm{Ca}(\mathrm{OH})_{2}(s) \end{aligned} $$ which ionizes as follows: \(\mathrm{HClO}(a q) \rightleftharpoons \mathrm{H}^{+}(a q)+\mathrm{ClO}^{-}(a q)\) $$ K_{\mathrm{a}}=3.0 \times 10^{-8} $$ As strong oxidizing agents, both \(\mathrm{HClO}\) 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.

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

Which of the following is the stronger base: \(\mathrm{NF}_{3}\) or \(\mathrm{NH}_{3} ?\) (Hint: \(\mathrm{F}\) is more electronegative than \(\left.\mathrm{H} .\right)\)
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