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Chapter Questions
Why is water necessary for life?
Contemplate biochemistry if atoms did not differ in electronegativity.
What is a van der Waals force?
What is an induced dipole?
What is a salt bridge?
Under what circumstance is a molecule that has a dipole not a polar molecule?
Which would you think would be a stronger interaction and why: an interaction between a sodium ion and the partial negative charge on the oxygen in ethanol $\left(\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{OH}\right),$ or the interaction between two ethanol molecules?
List the three types of van der Waals forces in decreasing order of strength.
A hydrogen bond is a special case of what type of intermolecular force?
Why do you think that most textbooks do not consider a hydrogen bond to be an example of a van der Waals force?
What are some macromolecules that have hydrogen bonds as a part of their structures?
How are hydrogen bonds involved in the transfer of genetic information?
Rationalize the fact that hydrogen bonding has not been observed between $\mathrm{CH}_{4}$ molecules.
Draw three examples of types of molecules that can form hydrogen bonds.
What are the requirements for molecules to form hydrogen bonds? (What atoms must be present and involved in such bonds?
Many properties of acetic acid can be rationalized in terms of a hydrogen-bonded dimer. Propose a structure for such a dimer.
How many water molecules could hydrogenbond directly to the molecules of glucose, sorbitol, and ribitol, shown here?
Both RNA and DNA have negatively charged phosphate groups as part of their structure. Would you expect ions that bind to nucleic acids to be positively or negatively charged? Why?
Identify the conjugate acids and bases in the following pairs of substances:(a) $\quad\left(\mathrm{CH}_{3}\right)_{3} \mathrm{NH}^{+} /\left(\mathrm{CH}_{3}\right)_{3} \mathrm{N}$(b) $^{+} \mathrm{H}_{3} \mathrm{N}-\mathrm{CH}_{2} \mathrm{COOH} /^{+} \mathrm{H}_{3} \mathrm{N}-\mathrm{CH}_{2}-\mathrm{COO}^{-}$$(\mathrm{c})^{+} \mathrm{H}_{3} \mathrm{N}-\mathrm{CH}_{2}-\mathrm{COO}^{-} / \mathrm{H}_{2} \mathrm{N}-\mathrm{CH}_{2}-\mathrm{COO}^{-}$(d) $^{-} \mathrm{OOC}-\mathrm{CH}_{2}-\mathrm{COOH} /^{-} \mathrm{OOC}-\mathrm{CH}_{2}-\mathrm{COO}^{-}$(e) $^{-} \mathrm{OOC}-\mathrm{CH}_{2}-\mathrm{COOH} / \mathrm{HOOC}-\mathrm{CH}_{2}-\mathrm{COOH}$
Identify conjugate acids and bases in the following pairs of substances:
Aspirin is an acid with a $\mathrm{p} K_{\mathrm{a}}$ of $3.5 ;$ its structure includes a carboxyl group. To be absorbed into the bloodstream, it must pass through the membrane lining the stomach and the small intestine. Electrically neutral molecules can pass through a membrane more easily than can charged molecules. Would you expect more aspirin to be absorbed in the stomach, where the pH of gastric juice is about $1,$ or in the small intestine, where the pH is about $6 ?$ Explain your answer.
Why does the pH change by one unit if the hydrogen ion concentration changes by a factor of 10 ?
Calculate the hydrogen ion concentration, $\left[\mathrm{H}^{+}\right]$ for each of the following materials:(a) Blood plasma, pH 7.4(b) Orange juice, pH 3.5(c) Human urine, pH 6.2(d) Household ammonia, pH 11.5(e) Gastric juice, pH 1.8
MATHEMATICAL Calculate the hydrogen ion concentration, $\left[\mathrm{H}^{+}\right]$ for each of the following materials:(a) Saliva, pH 6.5(b) Intracellular fluid of liver, pH 6.9(c) Tomato juice, pH 4.3(d) Grapefruit juice, pH 3.2
Calculate the hydroxide ion concentration, $\left[\mathrm{OH}^{-}\right]$ for each of the materials used in Question 24.
Define the following:(a) Acid dissociation constant(b) Acid strength(c) Amphipathic(d) Buffering capacity(e) Equivalence point(f) Hydrophilic(g) Hydrophobic(h) Nonpolar(i) Polar(j) Titration
Look at Figure 2.17 and Table $2.8 .$ Which compound in the table would give a titration curve the most similar to the one shown in the figure? Why?
Look at Figure $2.17 .$ If you did this titration using TRIS instead of phosphate, how would the titration curve look compared to the figure? Explain.
List the criteria used to select a buffer for a biochemical reaction.
What is the relationship between $\mathrm{p} K_{\mathrm{a}}$ and the useful range of a buffer?
What is the $\left[\mathrm{CH}_{3} \mathrm{COO}^{-}\right] /\left[\mathrm{CH}_{3} \mathrm{COOH}\right]$ ratio in an acetate buffer at pH $5.00 ?$
What is the $\left[\mathrm{CH}_{3} \mathrm{COO}^{-}\right] /\left[\mathrm{CH}_{3} \mathrm{COOH}\right]$ ratio in an acetate buffer at pH $4.00 ?$
What is the ratio of TRIS/TRIS-H' in a TRIS buffer at pH $8.7 ?$
What is the ratio of HEPES/HEPES-H $^{+}$ in a HEPES buffer at pH $7.9 ?$
How would you prepare 1 L of a 0.050 M phosphate buffer at $\mathrm{pH} 7.5$ using crystalline $\mathrm{K}_{2} \mathrm{HPO}_{4}$ and a solution of $1.0 \mathrm{MHCl}$ ?
The buffer needed for Question 35 can also be prepared using crystalline $\mathrm{NaH}_{2} \mathrm{PO}_{4}$ and a solution of $1.0 \mathrm{MNaOH}$. How would you do this?
Calculate the pH of a buffer solution prepared by mixing $75 \mathrm{mL}$ of $1.0 \mathrm{M}$ lactic acid (see Table 2.6 ) and $25 \mathrm{mL}$ of $1.0 M$ sodium lactate.
Calculate the pH of a buffer solution prepared by mixing $25 \mathrm{mL}$ of $1.0 \mathrm{M}$ lactic acid and $75 \mathrm{mL}$ of $1.0 \mathrm{M}$ sodium lactate.
Calculate the pH of a buffer solution that contains $0.10 M$ acetic acid (Table 2.6 ) and $0.25 M$ sodium acetate.
A catalog in the lab has a recipe for preparing 1 L of a TRIS buffer at $0.0500 \mathrm{M}$ and with pH 8.0 : dissolve $2.02 \mathrm{g}$ of TRIS (free base, $\mathrm{MW}=121.1 \mathrm{g} / \mathrm{mol}$ ) and $5.25 \mathrm{g}$ of TRIS hydrochloride (the acidic form, $\mathrm{MW}=157.6 \mathrm{g} / \mathrm{mol}$ ) in a total volume of 1 L. Verify that this recipe is correct.
If you mix equal volumes of $0.1 \mathrm{M}$ HCl and $0.20 \mathrm{M}$ TRIS (free amine form; see Table 2.8 ), is the resulting solution a buffer? Why or why not?
What would be the pH of the solution described in Question $41 ?$
If you have $100 \mathrm{mL}$ of a $0.10 \mathrm{M}$ TRIS buffer at $\mathrm{pH} 8.3 \text { (Table } 2.8)$ and you add $3.0 \mathrm{mL}$ of $1 \mathrm{MHCl}$, what will be the new pH?
What would be the pH of the solution in Question 43 if you were to add $3.0 \mathrm{mL}$ more of $1 \mathrm{MHCl}$ ?
Show that, for a pure weak acid in water $\mathrm{pH}=\left(\mathrm{p} K_{\mathrm{a}}-\log [\mathrm{HA}]\right) / 2.$
What is the ratio of concentrations of acetate ion and undissociated acetic acid in a solution that has a pH of 5.12 ?
BIOCHEMICAL CONNECTIONS You need to carry out an enzymatic reaction at $\mathrm{pH} 7.5 .$ A friend suggests a weak acid with a $\mathrm{p} K_{\mathrm{a}}$ of 3.9 as the basis of a buffer. Will this substance and its conjugate base make a suitable buffer? Why or why not?
If the buffer suggested in Question 47 were made, what would be the ratio of the conjugate base/conjugate acid?
Suggest a suitable buffer range for each of the following substances:(a) Lactic acid $\left(\mathrm{p} K_{\mathrm{a}}=3.86\right)$ and its sodium salt(b) Acetic acid $\left(\mathrm{p} K_{\mathrm{a}}=4.76\right)$ and its sodium salt(c) $\operatorname{TRIS}\left(\mathrm{p} K_{\mathrm{a}}=8.3 ; \text { see Table } 2.8\right)$ in its protonated form and its free amine form(d) HEPES $\left(\mathrm{p} K_{\mathrm{a}}=7.55 ; \text { see Table } 2.8\right)$ in its zwitterionic form and its anionic form
Which of the buffers shown in Table 2.8 would you choose to make a buffer with a pH of 7.3 ? Explain why.
The solution in Question 35 is called $0.050 M$, even though the concentration of neither the free base nor the conjugate acid is $0.050 M .$ Why is $0.050 M$ the correct concentration to report?
In Section $2-4$ we said that at the equivalence point of a titration of acetic acid, essentially all the acid has been converted to acetate ion. Why do we not say that all the acetic acid has been converted to acetate ion?
Define buffering capacity. How do the following buffers differ in buffering capacity? How do they differ in pH? Buffer a: $0.01 \mathrm{M} \mathrm{Na}_{2} \mathrm{HPO}_{4}$ and $0.01 \mathrm{MNaH}_{2} \mathrm{PO}_{4}$Buffer b: $0.10 \mathrm{M} \mathrm{Na}_{2} \mathrm{HPO}_{4}$ and $0.10 \mathrm{MNaH}_{2} \mathrm{PO}_{4}$Buffer $c: 1.0 M \mathrm{Na}_{2} \mathrm{HPO}_{4}$ and $1.0 \mathrm{MNaH}_{2} \mathrm{PO}_{4}$
If you wanted to make a HEPES buffer at $\mathrm{pH} 8.3,$ and you had both HEPES acid and HEPES base available, which would you start with, and why?
We usually say that a perfect buffer has its pH equal to its $\mathrm{p} K_{\mathrm{a}}$. Give an example of a situation in which it would be advantageous to have a buffer with a pH 0.5 unit higher than its $\mathrm{p} K_{a}.$
What quality of zwitterions makes them desirable buffers?
Many of the buffers used these days, such as HEPES and PIPES, were developed because they have desirable characteristics, such as resisting pH change with dilution. Why would resisting pH change with dilution be advantageous?
Another characteristic of modern buffers such as HEPES is that their pH changes little with changes in temperature. Why is this desirable?
Identify the zwitterions in the list of substances in Question 19
A frequently recommended treatment for hiccups is to hold one's breath. The resulting condition, hypoventilation, causes buildup of carbon dioxide in the lungs. Predict the effect on the pH of blood.