Pharmaceutical Biochemistry- Chapter 2
pick the correct pair: _______ is more soluble in water than _______. 1. Nitrogen, H 2 S 2. Nitrogen, CO 2 3. Oxygen, ammonia 4. CO 2 , NaCl 5. Ammonia, nitrogen
5.
what is normal blood pH
7.36-7.44; intracellular 7.1 (6.9-7.4)
H2CO3⇌ H++HCO3− describes a __________ buffer system.
carbonate
the _________ - __________ equations; describes the relationship between pH and the pKa of a buffer
henderson-hasselbalch
What is a condensation reaction?
reaction in which two reactants combine to form a single product with the elimination water.
water is often referred to as the "universal ______" because of its ability to hydrate molecules and screen charges
solvent
can gases be soluble in water?
some are soluble in water and some are not
what are micelles?
stable structures formed by lipids in water, which are held together by hydrophobic interactions
the _______ radius is approximately twice the distance of a covalent radius for a single bond. (3 words)
van der Waals
describe a hypotonic solution
water moves in, creating outward pressure; cell swells, may eventually burst
why is the ion product of water (Kw) at 25 degrees Celsius always equal to 1 x 10^-14?
pure water is neutral in nature, H+ ion concentration must be equal to OH- ion concentration.
describe osmotic pressure pertaining to water
water tends to move across a membrane, so as to decrease the difference in osmotic pressure **to move from the side with lower solute concentration to the side with higher solute concentration
what are non-covalent bonds/interactions?
weak interactions that are crucial to the structure and function of macromolecules
formula for osmotic pressure
Π= icRT
What are the colligative properties of water?
• Solutes decrease conc. of water • Properties depend only on the number of solute particles, not their size ** vapor pressure, boiling point, melting/freezing points, osmotic pressure
define buffer
a mixture of a weak acid and its conjugate base
describe "weak" (non-covalent) interactions
* continually broken and reformed *though weak individually, the cumulative effect of many such bonds is great. ***Velcro
what is the buffering region?
+1/-1 of the pH
a carboxyl group is .... 1.Polar 2.Non polar 3.Amphipathic
1.
What other functional groups of biomolecules can form hydrogen bonds? Why don't -CH groups participate?
Hydrogen bonds can only occur with oxygen, nitrogen, and fluorine. C-H groups do not hydrogen bond because carbon is only slightly more electronegative than hydrogen and thus the C-H bond is only very weakly polar.
A- is a proton __________
acceptor
what are some examples of biomolecules that can function as buffers?
amino acids, proteins, ATP, phosphate (pKa= 6.86), Bicabonate (pKa= 3.77, but effectively~ 6.1)
why is buffering so important?
enzyme activity is pH sensitive
Describe isotonic solution
no net water movement
Henderson-Hasselbalch equation
pH = pKa + log [A-]/[HA]
hydro _________ molecules can form energetically favorable interactions with water molecules.
-philic
hydro ________ molecules decrease the entropy of an aqueous system by causing water molecules to become more ordered.
-phobic
which of these processes involves the greatest decrease in entropy?* 1. Dissolving O2 in water 2. Dissolving NH3 in water *Hint: one way to think about this: Remember ∆G = ∆H - T∆S. What happens when either molecule is dissolved in water? The entropy (randomness) of any gas is less when it is dissolved than when it is 'free' ( (∆S for the process is negative), but which is less random, or less evenly dispersed, O2 in water or NH3 in water? The less evenly dispersed, the more negative the ∆S, which tends to make the ∆G more positive, making the process less favorable
1.
which of the following would have the greatest effect on the freezing point of a liter of water: the addition of 2 mol of NaCl or 2 mol of glucose?
2 mol of NaCl *The salt solute is able to depress the freezing point more than the sugar solute because the salt is ionically bonded while the sugar solute is covalently bonded. Because salt is ionically bonded, its ions are able to fully dissociate in solution.
a methyl group is..... 1. Polar 2. Non polar 3. Amphipathic
2.
for the dissolving of a gas in water, the entropy of the system.... 1. does not change. 2. decreases. 3. increases.
2.
the strength of a typical H-bond is.... 1.~4 kJ/mol 2.~20 kJ/mol 3.~100 kJ/mol 4.~350 kJ/mol 5.~500 kJ/mol
2.
a typical hydrogen bond in biomolecules is between H and _______. 1. C or O 2. C or P 3. N or O 4. N or P 5. O or P
3.
for the dissolving of a polar solid in water, the entropy of the system.... 1. does not change. 2. decreases. 3. increases.
3.
Which of the following is true of hydrogen bonds? A. The attraction between the oxygen atom of a water molecule and the hydrogen atom of another molecule constitutes a hydrogen bond. B. Hydrogen bonds form as covalent bonds between positively and negatively charged ions. C. Hydrogen bonds form between nonpolar portions of biomolecules. D. A and B are true. E. A, B, and C are true.
A
what general types of cellular reactions form water. and which consume water? which of these reactions are endergonic and which exergonic?
ATP/ADP coupling reaction ATP hydrolysis is exergonic ATP synthesis is endergonic
In a solution of pH 4.76, containing both acetic acid and acetate, what can you say about the concentrations of acetic acid (CH3COOH) and acetate (CH3COO−) present?
At pH 4.76, the concentrations of acetic acid and acetate in the solution will be equal.
As climbers approach the summit of a mountain they usually increase their rate of breathing to compensate for the "thinner air" due to the lower oxygen pressures at higher elevations. This increased ventilation rate results in a reduction in the levels of CO2 dissolved in the blood. Which of the following accurately describes the effect of lowering the [CO2]dissolved on blood pH? A. It will result in an increase in the dissociation of H2CO3→H++HCO3− and a drop in pH. B. It will result in an increase in the dissociation of H2CO3 → H2O + CO2 and a drop in pH. C. It will result in an increase in the association of H++HCO3−→H2CO3 and an increase in pH. D. It will result in a decrease in the association of H++HCO3−→H2CO3 and a decrease in pH. E. Lowering [CO2]dissolved has no effect on blood pH.
C
In a typical eukaryotic cell the pH is usually around 7.4. What is the [H+] in a typical eukaryotic cell? A. 0.00000074 M B. 6.6 μΜ C. 4 × 10−8 D. 2.3 nM E. 7.4 × 10−5 M
C
The pKa values for the three ionizable groups on tyrosine are pKa(— COOH)=2.2,pKa(—NH3+)=9.11, and pKa(—R)=10.07. In which pH ranges will this amino acid have the greatest buffering capacity? A. At all pH's between 2.2 and 10.07 B. At pH's near 7.1 C. At pH's between 9 and 10 D. At pH's near 5.7 E. Amino acids cannot act as buffers
C
Consider a weak acid in a solution with a pH of 5.0. Which of the following statements is true? A. The weak acid is a proton acceptor. B. The weak acid has a lower affinity for its proton than does a strong acid. C. At its pKa, the weak acid will be totally dissociated. D. The [H+] is 10−5 M. E. All of the above are true.
D
The pH of a sample of blood is 7.4; the pH of a sample of gastric juice is 1.4. The blood sample has an [H+]: A. 5.29 times lower than that of the gastric juice. B. a million times higher than that of the gastric juice. C. 6,000 times lower than that of the gastric juice. D. a million times lower than that of the gastric juice. E. 0.189 times that of the gastric juice.
D
Water derives all its special properties from its: A. cohesiveness and adhesiveness. B. high boiling point and melting point. C. small degree of ionization. D. polarity and hydrogen-bonding capacity. E. high dielectric constant.
D
Which of the following is true of pH? A. pH is the negative logarithm of [OH−]. B. Lemon juice, which has a pH of 2.0, is 60 times more acidic than ammonia, which has a pH of 12.0. C. Varying the pH of a solution will alter the pKa of an ionizable group in that solution. D. Varying the pH of a solution will alter the degree of ionization of an ionizable group in that solution. E. All of the above are true.
D
You mix 100 ml of solution of pH 1 with 100 ml of a solution of pH 3. The pH of the new 200 ml solution will be: A. 1.0. B. 2.0. C. 3.0. D. between pH 1.0 and pH 2.0. E. between pH 2.0 and pH 3.0.
D
Carbonic acid has a Ka of 1.70 × 10−4 and acetic acid has a Ka of 1.74 × 10−5. Which of the following is true? A. Carbonic acid has the higher Ka of the two and would therefore be the best buffer at pH 6. B. The acid with the larger Ka is a better proton acceptor. C. Carbonic acid is the stronger acid and has a lesser tendency to lose its proton compared to acetic acid. D. Neither carbonic acid nor acetic acid can be effective buffers at any pH. E. Acetic acid is a weaker acid and has a lesser tendency to lose its proton compared to carbonic acid.
E
you are running you first marathon on a very warm day. you start to sweat heavily and realize that you may be in danger of dehydration. why is severe dehydration potentially life-threatening? A. Water is a solvent for many biomolecules. B. Water is a chemical participant in many biological reactions. C. Water is necessary for buffering action in the body. D. Water's attraction to itself drives hydrophobic interactions. E. All of the above are true.
E
Certain insects can "skate" along the top surface of water in ponds and streams. What property of water allows this feat, and what bonds or interactions are involved?
Extensive hydrogen binding among water molecules accounts for the surface tension that allows some insects to walk on water.
(t/f) the oxygen atom in water has a partial positive charge.
False; the oxygen atom has a partial negative charge.
(t/f) H bonds can form only between water molecules.
False; they can form between any electronegative atom (usually oxygen or nitrogen) and a hydrogen atom covalently bonded to another electronegative atom in the same or another molecule.
Formic acid has a pKa of 3.75; acetic acid has a pKa of 4.76. Which is the stronger acid? Does the stronger acid have a greater or lesser tendency to lose its proton than the weaker acid?
Formic acid is the stronger acid. It has a greater tendency to lose its proton than does acetic acid.
what are some of the functional groups of bio-molecules that will interact with water electrostatically?
NaCl, ionized carboxylic acids, protonated amines, phosphate esters, and anhydrides
When you are very warm, because of high environmental temperature or physical exertion, you perspire. What property(ies) of water is your body exploiting when it sweats?
The high heat of vaporization of water, which is a measure of the energy required to overcome attractive forces between molecules, allows your body to dissipate excess heat through the evaporation of the water that is perspired
Why is ice less dense than liquid water?
The hydrogen bonds between the water molecules force the molecules to form a lattice structure when water freezes and fewer water molecules can fit into the same space, making ice less dense
how do hydrogen bonds contribute to the high melting and boiling points of water?
The reason for the high melting and boiling temperatures is the hydrogen bonding between water molecules that causes them to stick together and to resist being pulled apart which is what happens when ice melts and water boils to become a gas.
What is the absolute difference in [H+] between two aqueous solutions, one of pH 2.0 and one of pH 3.0? What is the [OH−] of the solution of pH 2.0?
The solution with a pH of 2.0 has a [H+] of 10−2 M, and an [OH−] of 10−12 M; the solution with a pH of 3.0 has a [H+] of 10−3 M. The difference in [H+] is 10−2M − 10−3M = 0.009 M
(t/f) each hydrogen atom of water bears a partial positive charge.
True; because the oxygen atom is more electronegative than the two hydrogen atoms, the electrons are more often in the vicinity of the oxygen, giving each of the hydrogen atoms a partial positive charge.
(t/f) H bonds are relatively weak compared to covalent bonds.
True; hydrogen bonds have a bond dissociation energy of 23 kJ/mol, whereas covalent single bonds have a stabilization energy of approximately 200-460 kJ/mol.
how does the henderson-hasselbalch equation prove that the pKa of a weak acid is equal to the pH of the solution at the midpoint of its titration?
[HA]=[A-] and pH=pKa + log 1 = pKa + 0 so, pH=pKa
what is the ranking of relative solubility of compounds?
[highest] polar solids, liquids, polar gases < nonpolar solids, liquids < nonpolar gases [lowest]
what is a buffer?
a buffer consists of approximately equal amounts of a weak acid and its conjugate (weak) base keeps pH relatively constant biological processes are very pH-sensitive
what are the two equilibrium reactions that are simultaneously adjusting during an experimental titration of a weak acid?
acetic acid/acetate
What is titration?
add a strong base (usually drop-wise) to a weak acid * pH changes as illustrated
describe a solution with a [H+] of 1 x 10^-8
alkaline
what is the entropy of a gas?
always shows a decrease Entropy always decreases (ΔS<0), because the gas is confined to a small volume ** decrease is smaller if polar (molecules disperse), larger if non-polar (molecules cluster)
what is the driving force behind the formation of micelles?
amphipathic molecules; hydrogen bonding; hydrophobic interactions
how does Kw relate to the pH scale?
basis for the pH scale; designates the concentration of H+ and OH- in any aqueous solution in the range of 1 M H+ and 1 M OH-
how do the four types of weak interactions among biomolecules compare in strength to each other and to covalent bonds?
bonds in order of strength: covalent bonds, ionic, hydrogen, van der Waal's, hydrophobic
what is a amphipathic compound?
compound containing both polar and nonpolar regions
Why would a zipper or a Velcro strip be an appropriate analogy to weak interactions in biochemical reactions?
continually broken and reformed **though weak individually, the cumulative effect of many such bonds is great
what is hydrolysis?
covalent bond breakage by the addition of water
what is pH?
denotes the concentration of H+ (and therefore of OH-) in an aqueous solution
what is an acidic solution?
describes a solution in which [H+] is greater than [OH-]
what is a basic solution?
describes a solution in which [OH-] is greater than [H+]
the equilibrium constant for the reaction HA ⇌ H+ + A− is also called the ___________ constant, Ka
dissociation
what are solutes?
dissolved molecules
HA is a proton __________
donor
noncovalent bonds have weaker bond __________ thatn covalent bonds.
energy
what is the most sensitive aspect of cell function (mentioned many times in this chapter) in relationship to changes in pH?
enzyme activity?
the electrostatic interactions between the hydrogen and oxygen atoms on adjacent H2O molecules constitute a ________. (2 words)
hydrogen bond
van der Waals interactions are a weak, transient subcategory of which type of noncovalent interaction?
hydrogen bonding; non-covalent bonding
describe what occurs when a crystalline salt dissolves in water in terms of the enthalpy (H) and the entropy (S) of the system
increases the entropy (S) of the system which is largely responsible for the ease of dissolving salts and a decrease of the enthalpy (H)
what is Kw
ion product of water *10^-14
water molecules readily dissolve compounds such as NaCl because they screen _______ interactions between Na+ and Cl-.
ionic
what is the entropy of a polar solid/liquid solute?
it becomes more disordered (ΔS>0)
what is the entropy of a nonpolar solid/liquid solute?
it becomes more ordered polar, becomes more ordered (ΔS<0)
the number 1, 10, 100, and 1000 are placed at equal intervals on a __________ scale.
log scale
What is osmolarity?
measure of total concentration of solute particles
amphipathicity is important to the structure (and therefore the function) of which bio-molecules?
micelles, lipids
enzymes show maximum activity at a characteristic pH ____________
optimum
what is pKa?
pH at which [HAc] = [Ac−]
Define pH and pKa and explain how they are different.
pH is the negative logarithm of [H+] in an aqueous solution. It provides a standard way to measure the H+ concentration in an aqueous solution. pKa is the negative logarithm of an equilibrium constant. It is equal to the pH at which a weak acid is one-half dissociated; i.e., the pH at which there are equal concentrations of a weak acid and its conjugate base. The pKa occurs at the midpoint of the titration curve of a weak acid, the center of the range providing the maximum buffering capacity of the conjugate acid-base pair. The pKa is an integral property of an ionizable group. It is the extent of ionization of an ionizable group that varies with the pH of the solution.
The Henderson-Hasselbalch equation is pH=pKa + log [proton acceptor][proton donor]. Show how it proves that the pKa of a weak acid is equal to the pH of the solution at the midpoint of its titration.
pH=pKa + log [proton acceptor][proton donor] At the midpoint of the titration, [proton acceptor] = [proton donor]. The log of 1 = 0, so pH = pKa + 0; pH = pKa
H2PO4−⇌ H++PO42− describes a _________ buffer system.
phosphate
does a strong acid have a greater or lesser tendency to lose its proton than does a weak acid? does the strong acid have a higher or lower Ka? a higher or lower pKa?
strong acids have a greater tendency to lose its proton as it fully dissociates into a solution and a weak acid only partially dissociates into a solution; stronger acids have a higher Ka and thus a smaller pKa than weak acids
what does the relatively flat zone of a titration curve tell you about the pH changes within that zone?
the addition of the acid/base does not affect the pH of the solution drastically
what is Kw?
the ion product of water; it is 1 x 10^-14 M in aqueous solutions at 25 degrees Celsius
Define equilibrium
the point in a reversible chemical reaction at which the rate of product formation equals the rate of product breakdown to the starting reactants
what is the equilibrium point?
the point in the center of a titration curve where pH=pKa
a plot of pH vs OH- equivalents added is a ________ curve
titration
Describe a hypertonic solution
water moves out and cell shrinks
at a pH equal to the pKa of a weak acid, what can be said about the concentrations of the acid and its conjugate base? what point on a titration curve indicates the pKa of that weak acid?
when pH and pKa equal the concentrations of acid and base are equal to each other; the point on the titration curve when pH and pKa equal each other indicates the pKa of the weak acid
