Acids and Bases

The acid ionization constant is also known as the: (A) Acid association constant (B) Acid dissociation constant (C) Acid solubility constant (D) Acid precipitation constant

(B) Acid dissociation constant The acid ionization constant is also known as the acid dissociation constant.

What are Bronsted-Lowry acids and bases?

Bronsted-Lowry acids are molecules or ions that are proton donors. Bronsted-Lowry bases have molecules or ions that are proton acceptors.

What is the autoionization constant expression? What are the experimental expression of each component?

Kw = [H3O+][OH-] 1.0 * 10^-7

Which is more acidic, HF or HI, H2S or H2O? Why?

LARGER ATOM INCREASES ACIDITY HI, H2S

Which is more basic, NH3 or HF, CH3NH2 or CH3OH? Why?

LESS ELECTRONEGATIVE ATOM INCREASES BASICITY NH3, CH3NH2

In the reaction of water and BF3, what is the Lewis acid and Lewis base? What is the product?

Lewis acid: electron pair acceptor Lewis Base: electron pair Donor Lewis base: water Lewis acid: BF3 H2-O+-BF3

Which is more basic, OH- or H2O, HCO3- or CO32-? Why?

MORE NEGATIVE CHARGE INCREASES BASICITY OH-, CO32-

Which compound is more acidic, NH3 or NH4+? Why?

MORE POSITIVE CHARGE INCREASES ACIDITY NH4+

An basic salt is

a salt that contains an ion that is a weak base

An acidic salt is

a salt that has an ion that is a weak acid

Carbonic anhydrase is found on both the cell membrane and in the cytoplasm of cells lining the renal tubules. If a new drug were developed that selectively inhibited only the membrane-bound form of carbonic anhydrase, how effective would it be in preventing AMS? A. More effective; more HCO3- would remain in the tubular lumen and decrease blood pH. B. More effective; HCO3- would remain trapped in renal cells and decrease blood pH. C. Less effective; Cl - ion would accumulate in the blood and decrease pH. D. Less effective; H2O and CO2 diffusing into the cell would be converted to HCO3- and could increase blood pH.

In Figure 1 we see carbonic anhydrase (CA) is found both on the cell membrane and in the cytoplasm. CA in the membrane converts bicarbonate into water and carbon dioxide that diffuse into the cell. CA in the cell converts water and carbon dioxide back into bicarbonate where it is exchanged for chloride ion in the blood. If only the membrane CA were blocked, residual carbon dioxide and water could still diffuse into the cell where they would be converted to bicarbonate and exchanged into the blood. Since bicarbonate is a base, this could initially increase pH instead of lowering it (choices A and B are eliminated). Even if the bicarbonate/chloride ion exchange slowed over time, chloride ion (the conjugate of a strong acid) does not have acid-base properties and would not affect pH (eliminates choice C).

Describe a Lewis acid.

It is either neutral or a cation that has an empty orbital

What is the half equivalence point? How is it determined?

It is when exactly one half of the acid has been neutralized by the base; in other words, [HA] = [A-] ->pKa -half of the volume at the equivalence point

An indicator is

A substance which differentiates between acid and base solutions is weak acid and has pKa

Write the acid dissociation constant expression for the following reaction: HA <-> H+ + A-

Ka = [H+][A-]/[HA]

What is the Ka expression for the reaction: H2O + HA -><- H3O+ + A-?

Ka = [H3O+][A-]/[HA]

What is the base ionization constant expression?

Kb = [BH+][OH-]/[B]

What is the pH when you have a [OH-] concentration equal to 9.84 * 10^-8M? a) 8.67 b) 7.92 c) 6.99 d) 5.43

c) 6.99

What is the purpose of an indicator?

- Changes color when adequate temperature and time reached; allows monitoring without probes or computers

Compare the dissociation of strong acids and weak acids through equilibrium.

-strong acids dissociate more completely and favor the products -weak acids dissociate less completely and favor the reactants (Ka < 1)

What is the value of Kw at 25 degrees Celsius?

1 x 10^-14

Define a Bronsted-Lowry Base vs. Bronsted-Lowry Acid? Define a Lewis Base vs. Lewis Acid?

A Bronsted-Lowry Base is a proton acceptor, and a Bronsted-Lowry Acid is a proton donator. A Lewis Base is an electron donor, and a Lewis Acid is an electron acceptor.

If a solution with a pH of 11 is diluted with pure water until the volume is increased by a factor of 10, the pH will then equal: A. 10 B. 9 C. 12 D. 1

A. 10 Each pH unit corresponds to a 10-fold change in the concentration of H+ (or OH-). Therefore, diluting the basic solution by a factor of ten would result in a solution one pH unit closer to neutrality.

The p Ka of carbonic acid (H2CO3) at standard state is 3.6. Given this value, what is the pKb of carbonate ion (CO32-) under the same conditions? A. 3.7 B. 10.4 C. 11.3 D. 12.7

A. 3.7 It is tempting to simply substitute into the formula pKa+ pKb = 14, but this can only be done for conjugate acid-base pairs. H2CO3 and CO32- are not a conjugate pair because an intermediate step involving dissociation of HCO3- is in between them: The pKb of HCO3- (not CO32-) will be 10.4 (eliminate choice B). Choices C and D are both higher than this, and choice A is lower. Successive H+ dissociations of polyprotic acids get weaker, with subsequently higher pKa values. The pKa of HCO3- will be significantly higher than that of H2CO3. Therefore, the pKb of CO32- must be lower than the pKb of HCO3-. Another way to approach the problem is to note that choices C and D are both higher than the pKb of HCO3- with fairly close numerical values. The MCAT is unlikely to have a problem requiring a calculation to differentiate between C and D, so A is likely the best choice.

Because it dissolves carbon dioxide from the atmosphere to make mild carbonic acid, natural rain has a pH of around 5.5. However, due to pollutants such as sulfates from coal-fired power plants and nitrates from car exhausts, the pH of rain can drop to as low as 2. This decrease in pH represents a change in H+ concentration by approximately what factor? A. 3000 B. 300 C. 3.5 D. 35

A. 3000

What is the Brønsted-Lowry base in the following reaction? HCl(g) + H2O(l) H3O+(aq) + Cl-(aq) A. H2O B. Cl- C. HCl D. H3O+

A. H2O

Red wines are generally acidic, with proper pH values near 3.6. At this pH, which of the following organic acids commonly found in wine have the highest proportion of molecules in the -1 charge-state? (note: pKa1 denotes the pKa of the most acidic proton on the compound; the pKa values of all subsequent protons on the compounds of interest are all near pH = 6) A. Maleic acid (pKa1 = 1.90) B. Citric acid (pKa1 = 3.09) C. Succinic acid (pKa1 = 4.05) D. Tartaric acid (pKa1 = 2.95)

A. Maleic acid (pKa1 = 1.90) As this is a question asking for extremes, it is safe to eliminate tartaric and citric acids for having the non-extreme pKa values. The Henderson-Hasselbalch equation states that pH = pKa + log ([A-]/[HA]), where A- is the molecule in the -1 charge state, and HA is the protonated, neutral molecule. Rearranging this equation and substituting the given pH yields: log ([A-]/[HA]) = 3.6 - pKa. The acid that gives the largest positive number for the right side of this equation is the one with the smallest pKa, maleic acid.

Addition of sodium acetate to a solution of acetic acid will cause the pH to: A. increase due to the common ion effect. B. remain constant because sodium acetate is a buffer. C. remain constant because sodium acetate is neither acidic nor basic. D. decrease due to the common ion effect.

A. increase due to the common ion effect. Sodium acetate is a basic compound, because acetate is the conjugate base of acetic acid, a weak acid ("the conjugate base of a weak acid acts as a base in water"). The addition of a base to any solution, whether it is buffered or not, will increase the pH. The answer is "increase due to the common ion effect."

What is the relationship between the proton dissociations of polyprotic acids and pKa and pKb?

As the dissociations of a proton of polyprotic acids increases and gets weaker, the pKa increases in value and pKb decreases in value.

Ethylenediaminetetraacetic acid (EDTA) is a commonly used scavenger of heavy metals. The molecule has 4 carboxylic acid moieties, with successive pKa values of 1.99, 2.67, 6.16, and 10.26. At neutral pH, what is the predominant charge state of EDTA? A. -1 B. -3 C. -2 D. -4

B. -3

Chlorine is used to preserve the quality of the water in swimming pools. Since this water is not used for consumption purposes, higher concentrations of chlorine are maintained. If the pH of swimming pool water is 5, what is the concentration of HOCl? A. 10-10 M B. 10-5 M C. 10-2.5 M D. 105 M

B. 10-5 M First of all, choice D should be eliminated right away. The highest concentration of any ordinary aqueous solution is about 20 M; to have a solution which is 105 = 100,000 M is impossible. The correct answer is determined from the reaction given in the passage. For every one molecule of HOCl produced, we get one molecule of HCl. So the concentrations of the HOCl and HCl should be the same. The concentration of HCl is indirectly given—the pH is 5, so the concentration of HCl is 10-5 M—thus the concentration of HOCl must also be 10-5 M. Note that while HOCl is an acid and thus will contribute to the total concentration of H+, it is a weak acid, so its real contribution is tiny compared to HCl.

The Ka of formic acid (HCOOH) is 1.8 × 10-4. What is the pKb of the formate ion? A. -14 - log (1.8 × 10-4) B. 14 + log (1.8 × 10-4) C. -14 + log (1.8 × 10-4) D. 14 - log (1.8 × 10-4)

B. 14 + log (1.8 × 10-4)

What would be the approximate pH of a 2.0 M acetazolamide (a monoprotic acid) solution? A. 1.2 B. 3.6 C. 7.1 D. 10.4

B. 3.6 ICE table

Which one of the following will decrease the solubility of CO2 in water? A. Increasing the external pressure of CO2 B. Increasing the temperature of the water C. Increasing the pH of the water D. None of the above

B. Increasing the temperature of the water The solubility of a gas in a liquid decreases with increasing temperature (as illustrated by the fact that CO2 readily erupts from a bottle of warm soda). Choices A and C will increase the solubility of CO2 gas in water. As the kinetic energy of the gaseous solute increases, its molecules have a greater tendency to escape the attraction of the solvent molecules and return to the gas phase. Therefore, the solubility of a gas decreases as the temperature increases.

The acidic strength of HF is greatly enhanced when it is used in conjunction with SbF5. The resultant system, known as "magic acid", is one of the strongest acids known, and follows the dissociation equation below. HF + SbF5 → H+ + SbF6- Which of the following best describes the role of SbF5 in magic acid? A. Lewis base B. Lewis acid C. Brønsted base D. Brønsted acid

B. Lewis acid

Which indicator would be best for finding the equivalence point when titrating NH3 with 0.1 M HCl? A. Thymol Blue (pH range = 1.2 - 2.8) B. Methyl Orange (pH range = 3.1 - 4.4) C. Phenolphthalein (pH range = 8.3 - 10.0) D. Alizarin Yellow (pH range = 10.0 - 12.0)

B. Methyl Orange (pH range = 3.1 - 4.4)

Enough HF (Ka = 7.4 × 10-4) is added to water to create a pH = 2.1 solution. The addition of which of the following would have the least impact on the pH of this solution? A. NH4F B. PbF2 C. Na2CO3 D. NaF

B. PbF2 Both NaF and NH4F are soluble salts, and will decrease the solubility (acidity) of HF due to the common ion effect, so both of these answer choices can be eliminated. Na2CO3 is a soluble salt and CO32- is the conjugate base of a weak acid, so it will act to increase the pH. PbF2 is essentially insoluble in water, and therefore will have the least effect on the pH, making it the best answer.

The conjugate bases of HSO4-, CH3OH, and H3O+ are, respectively: A. SO4-, CH2OH-, H2O B. SO42-, CH3O-, H2O C. SO4-, CH3O-, OH- D. SO42-, CH2OH-, OH-

B. SO42-, CH3O-, H2O

Which one of the following best explains why fluoroamine (NH2F) is a weaker Lewis base than chloroamine (NH2Cl) and ammonia (NH3)? A. The chlorine atom decreases the steric hindrance at the basic site. B. The fluorine atom removes more electron density from the nitrogen. C. Chlorine is a stronger base than fluorine. D. Chlorine changes the geometry of the molecule, thereby increasing the acidity.

B. The fluorine atom removes more electron density from the nitrogen. Choices A, C, and D are all incorrect statements. Choice A leads to the conclusion that chloroamine should be a weaker Lewis base (ligand) than fluoramine—but this contradicts what the question tells us! Choice C is incorrect because HCl is more acidic than HF (recall that the acidity of an element increases as one moves to the right or down in the periodic table). Choice D is wrong because the geometry of a molecule is determined by the number of electron groups around the central atom, not the identity of the atoms at the ends of those electron groups. Choice B is true because F is more electronegative than Cl.

Which one of the following statements concerning 50 mL of basic solution composed of 0.2 M NaH2PO4 and 0.2 M Na2HPO4 is correct? A. The concentration of Na+ is 0.4 M. B. The pH of the solution will not change with the addition of 20 mL of water. C. The addition of 1 mL of 0.1 M HCl to the reaction will result in a pH of 2. D. With the addition of ammonia, a precipitate should be formed.

B. The pH of the solution will not change with the addition of 20 mL of water.

Which of the following statements is true about a solution of acetic acid in equilibrium? CH3CH2COOH (aq) -> H+ (aq) + CH3CH2COO- (aq) A. When HI is added, CH3CH2COOH acts as a Brønsted-Lowry acid. B. When HCl is added, CH3CH2COO- acts as a Brønsted-Lowry base. C. When NaOH is added, CH3CH2COOH acts as a Lewis acid. D. When HF is added, CH3CH2COOH acts as an Arrhenius base.

B. When HCl is added, CH3CH2COO- acts as a Brønsted-Lowry base.

A salt solution contains ammonium cation (Ka = 5.7 × 10-10) and nitrite anion (Kb = 1.4 × 10-11) in equal proportions. The pH of the solution will be: A. acidic because nitrite has a higher Kb value than the Ka of ammonium. B. acidic because ammonium has a higher Ka value than the Kb of nitrite. C. basic because ammonium has a higher Ka value than the Kb of nitrite. D. basic because nitrite has a higher Kb value than the Ka of ammonium.

B. acidic because ammonium has a higher Ka value than the Kb of nitrite. The Ka of ammonium is greater than the Kb of nitrite. Eliminate choices stating "nitrite has a higher Kb value than the Ka of ammonium." The fact that ammonium has a higher Ka means that it, in general, is in a more dissociated state than the nitrite is in a protonated state. Therefore, the solution will be acidic.

In a dilute solution, if molarity of a solution is known, the molality of a solution of known volume can always be calculated provided the: A. density and molecular weight of the solute is known. B. density of the final solution is known. C. solution is aqueous. D. molecular weight of the solvent is known.

B. density of the final solution is known. To calculate molality, the moles of solute and kilograms of solvent must be known. The type of solvent in the solution is irrelevant in this question (eliminate choice C), and the molecular weight of the solvent would only be needed if the number of moles of solvent were important (eliminate choice D). While the molecular weight of the solute is needed for the calculation, the solute density is not (eliminate choice A). If the molarity of a solution is known, then the moles of solute can be determined by simply assuming 1 L of solution. If the density of the solution is also known, then the mass of the solution can be calculated from its volume (density × V = m). By converting the number of moles of solute to the mass of solute, and recognizing that Mass Solvent = Mass Solution - Mass Solute, molality can be determined by finding the mass of the solvent by subtraction.

Which of the following best characterizes the ionization constant of a strong acid? A. 0 < Ka < 0.1 B.Ka > 1 C.Ka < 0 D. 0.1 < Ka < 1

B.Ka > 1

Naproxin's pKa is 4.2. At what value of [H+] in the blood would the conjugate of sodium naproxen be at its lowest concentration? A. 4.1 × 10-8 M B. 7.6 × 10-8 M C. 1.2 × 10-7 M D. 5.9 × 10-6 M

C. 1.2 × 10-7 M Again, questions with words like "lowest," "highest," "most," and "least," will typically have answers that correspond to the greatest or lowest numerical value, eliminating choices B and C. Since sodium naproxen, the deprotonated form of naproxen, is a weak base its conjugate will be a weak acid (naproxen), which will increase in concentration with an increasingly acidic environment. Therefore, the answer will be the choice with the lowest H+ concentration, choice A.

Commercial "concentrated HCl" is a 38% by mass solution in water. Assume the density of concentrated HCl is 1.12 g/mL, what is the molarity of such a solution? A. 2 M B. 8 M C. 12 M D. 20 M

C. 12 M Assume there is 1 mL of solution, which will weigh 1.12 g at the given density. Since the solution is 38% HCl by weight, there will be approximately 0.4 × 1.12 g or about 0.45 g of HCl in the 1.12 g of solution. This 0.45 g of HCl converts to 0.45 g ÷ 37 g/mol or somewhere around 0.012 g HCl in 1 mL. This equates to about 12 M.

The Ka of a buffer is 4.5 × 10-4. If the concentration of undissociated weak acid is equal to the concentration of the conjugate base, the pH of this buffer system is between: A. 5 and 6. B. 2 and 3. C. 3 and 4. D. 4 and 5.

C. 3 and 4. The pH of a buffer can be calculated using the Henderson-Hasselbalch equation, pH = pKa + log ([conjugate base]/[conjugate acid]) When the concentrations of the acid and base are equal, the fraction on the right-hand side is 1; since log 1 = 0, we have pH = pKa. If the Ka of the acid is 4.5 × 10-4, which is between 10-4 and 10-3, the pKa (and, therefore, the pH) is between 3 and 4.

Which one of the following best approximates the pH of a solution when 99.9% of the acid in a pH 1 solution is neutralized? A. 7 B. 1 C. 4 D. 6.7

C. 4 The neutralization of 99.9% of acid tells us that for every one thousand H+ originally in solution, only one remains. A change in [H+] by a factor of 103 corresponds to a change of 3 pH units. If the initial pH was 1.0, then the final pH must be 4.

Chlorosulfonic acid (HSO3Cl) is stronger than sulfuric acid when solvated in water, yet it is far less acidic in common organic solvents. When 1.16 g of HSO3Cl is dissolved in 10 mL of hexane, spectroscopic methods determine the [SO3Cl-] to be 2.2 × 10-3 M. Which of the following is closest to the pKa of chlorosulfonic acid in hexane? A. 6.2 B. 2.8 C. 5.4 D. 0.2

C. 5.4 The molecular weight of HSO3Cl is ~116 g/mol, so the initial concentration for HSO3Cl in 10 mL of hexane is 1 M. If [SO3Cl-] is 2.2 × 10-3 M, then the Ka ([SO3Cl-][H+]/[HSO3Cl]) in question can be expressed as (2.2 × 10-3)2/(1 - 2.2 × 10-3). The denominator can be safely simplified to 1, giving Ka = ~4 × 10-6. Since -log (4 × 10-6) is between 5 and 6, 5.4 is the only possible answer.

If the initial pH of a solution is 4.5, what will be the final pH after enough water is added to dilute the original concentration to 1% of its initial value? A. 2.5 B. 5.5 C. 6.5 D. 7.5

C. 6.5 The question indicates that the concentration of H+ has dropped to 1% = 1/100 its original value. Since pH is the negative log of the concentration of H+, the change in pH must be an increase by log 100 = 2; i.e., the final pH should equal the initial pH plus 2, which is 4.5 + 2 = 6.5.

Calculate pH of aqueous solution that has 0.11 g of Ca(OH)2 in total volume of 250 mL.

Ca(OH)2 ->Ca2+ + 2OH- O.11g~ 0.10g 0.10g/(molar mass of Ca(OH)2 = 74)/.250L * 2 = ANS 14-(-log(ANS) = pH

A chemist wishes to titrate a sulfuric acid solution. Which base should she use if she wants to be certain that no precipitate will form during the titration? A. Fe(OH)3 B. Be(OH)2 C. Cs(OH) D. Mg(OH)2

C. Cs(OH) Two substances are produced in a neutralization: water and a salt. In this case a sulfate salt will be produced so we need to find the cation that is soluble when paired with sulfate. Solubility rules tell us Group I metal salts are soluble, and Cs is in Group I. Most inorganic hydroxide compounds are only slightly soluble or insoluble, including all other choices listed.

Which of the following is a contributing factor to the relative acidity of HF compared to HCl? A. HCl exhibits stronger hydrogen bonding. B. HF has a greater boiling point. C. F- is less stable in solution than Cl-. D. HCl has a shorter bond length.

C. F- is less stable in solution than Cl-. All hydrogen halide compounds are strong acids except for HF. There are several reasons for this. HF has significant hydrogen bonding between molecules which hinders the proton from contributing to acidity. None of the other hydrogen halides have hydrogen bonding. HF is the smallest of all the hydrogen halides, which increases the electrostatic interaction between hydrogen and fluorine according to Coulomb's law. HF does have a high boiling point due to its relatively strong intermolecular forces, but this does not directly affect acidity. Since HF is much less likely to dissociate into H+ and F- compared to the dissociation of HCl into H+ and Cl-, F- must be less stable as an ion in solution compared to Cl-.

Which of the following is an amphoteric species? A. H+ B. S2- C. HS- D. H2S

C. HS-

The heat released from the dilution of concentrated HCl comes from the: A. dissociation of H-Cl bonds. B. dissociation of H-OH bonds. C. formation of ion-dipole interactions in solvation shells. D. decrease in entropy.

C. formation of ion-dipole interactions in solvation shells. The dissociation of any covalent bond is an endothermic process (choices A and B are eliminated). When the solution is being diluted, its entropy is increasing because the solute molecules are spreading out (choice D is eliminated). The reason this process is so exothermic is because the addition of more water molecules allows for many more ion-dipole interactions between H+ and water and Cl- and water.

Consider a buffered and a normal solution each with a pH of 10. If they are both diluted with water, then: A. the pH of both solutions will increase. B. the pH of both solutions will decrease. C. only the pH of the normal solution will decrease. D. the pH of neither solution will be affected.

C. only the pH of the normal solution will decrease. The continuous dilution of a normal basic solution will decrease the pH of the solution until a pH near 7 is reached (choices A and D are eliminated). One of the fundamental differences between a normal solution and a buffered solution is that the pH of a buffer is not affected by dilutions

Compare acidity strength and Ka of HCl and CH3COOH.

HCl: strong acid and high Ka CH3COOH: weak acid and low Ka

Which anions will never be basic?

Cl-, Br-, I-

If the enzyme carbonic anhydrase were added to a glass of soda pop which had been allowed to reach equilibrium with the atmosphere, it would: A. produce a large number of carbon dioxide bubbles. B. produce a large amount of oxygen. C. form more carbonic acid. D. have no effect on the equilibrium.

D. ENZYME = CATALYST = INCREASES RATE TO REACH EQUILIBRIUM BUT NOT NECESSARY IF ALREADY AT EQUILIBIRUM

Nitric acid can be concentrated to a 70% by volume solution in water. Such a solution, called concentrated nitric acid, is 16 M. Assuming that nitric acid completely dissociates in such a solution, what is its pH? A. 1.2 B. 0.8 C. -0.8 D. -1.2

D. -1.2 First write the concentration of nitric acid in scientific notation: 1.6 × 101. If it is assumed that the acid completely dissociates, this is also equal to the concentration of hydronium ions. pH = -log[H3O+] pH = -log(1.6 x 101) The pH is between -log(101) and -log(102). Therefore, pH is between -1 and -2.

A mixture of acidic wastes is found to be 2 M H2SO4, 0.5 M HI, and 3 M HNO3. What volume of 5 M NaOH solution will be required, per liter of acidic waste, to completely neutralize the acid? A. 3.0 L B. 1.5 mL C. 1.1 L D. 1.5 L

D. 1.5 L The easiest way to answer this question is to determine the molarity of H+ in the acidic wastes. Accounting for all dissociable protons there are 7.5 moles of H+ per liter of solution. Remember, there are two acidic protons on H2SO4. It is then the case that 7.5 moles = 5 MNaOH x ? L solution, resulting in 1.5 L of NaOH solution being necessary for complete neutralization.

In one experiment, a fluffy-white precipitate formed upon bubbling the gases through the water. It is possible that the water was contaminated with a small quantity of which of the following ions? A. K+ B. Li+ C. NH4+ D. Ba2+

D. Ba2+ More solubility rules. Sulfate and carbonate are formed in solution in this experiment. Since all Group I and ammonium salts are soluble, choices A, B, and C would not form precipitates.

Which of the following compounds is a weak electrolyte? A. BaCl2 B. CH3OH C. KOH D. CH3COOH

D. CH3COOH

Which of the following compounds would form the least acidic solution when dissolved in water? A. FeI3 B. NaHSO4 C. H2S D. CaCl2

D. CaCl2 There are three rules that can help identify whether a chemical species will form acidic, basic, or pH neutral solutions. The acidity of an element increases as one moves to the right or down the periodic table. The conjugate base (or acid) of a strong acid (or base) forms pH neutral solutions. The conjugate base (or acid) of a weak acid (or base) is a weak base (or acid). Based upon these rules, we can characterize each species given in the answer choices as either acidic, basic, or pH neutral: Fe3+ weak acid (Fe3+ is the conjugate acid of Fe(OH)3, a weak base - rule 3) I- is pH neutral (I- is the conjugate base of HI, a strong acid (rule 2) HSO4- weak acid (HSO4- is the conjugate acid of SO42-, a very weak base - rule 3) Na+ pH neutral (Na+ is the conjugate acid of NaOH, a strong base - rule 2) H2S weak acid (S is below O - rule 1) Ca2+ essentially pH neutral (Ca2+ is the conjugate acid of Ca(OH)2, a strong base - rule 2) Cl- pH neutral (Cl- is the conjugate base of a strong acid - rule 2) Therefore, the only ionic compound that doesn't make a solution acidic is CaCl2 (choice D).

Carbon dioxide is much more soluble in water than is oxygen. Why? A. Oxygen has a greater dipole moment than carbon dioxide. B. Carbon dioxide has a greater dipole moment than oxygen. C. The kinetic barrier to the dissolution of carbon dioxide is smaller than that of oxygen. D. Carbon dioxide is susceptible to nucleophilic attack

D. Carbon dioxide is susceptible to nucleophilic attack The first three choices are false. Neither O2 nor CO2 has a permanent dipole, so choices A and B can be ruled out. The kinetic factors described in choice C do not affect overall solubility, which is a thermodynamic value. The passage states that the solubility of carbon dioxide is due primarily to its reaction with water where the oxygen atom in a water molecule adds reversibly to the carbon atom in carbon dioxide (Reaction 1). This reaction is a nucleophilic attack on carbon dioxide, making choice D correct.

The addition of 1 g of which of the following compounds to distilled water would cause the least change in pH? A. LiH B. HCl C. LiOH D. CsOH

D. CsOH Since the question posits that 1 g of each strong electrolyte (which undergoes complete dissociation) is added to distilled water, the compound with the highest molecular weight will yield the lowest molar concentration of either H+ or OH-. Cesium hydroxide (MW = ~150) is far heavier than any of the other compounds in question and is therefore the correct answer. LiH is a strong base, and when added to water will undergo the reaction LiH + H2O ? LiOH + H2. Hence the resultant pH from 1 g of LiH will be just slightly higher than 1 g of LiOH.

Which of the following chemical species will be present at the highest concentration when 1.0 g of Cl2(g) is dissolved in 2.0 L of water? A. HCl B. Cl- C. HOCl D. H3O+

D. H3O+ According to Reaction 1 in the passage, chlorine gas reacts with water to form HCl and HOCl in equimolar amounts. Since the actual mass of dissolved chlorine or volume of water are irrelevant to answering the question, let's assume that enough chlorine was used to make a 1 mM solution. HCl is a strong acid and will completely dissociate in water, (eliminate choice A, since there is no HCl in solution) to yield 1 mM of Cl -and 1 mM of H3O+. The passage states that HOCl is a very weak acd, so it will dissociate to a very limited extent, reducing the HOCl concentration to just below 1 mM (eliminate choice C) and producing a small amount of H3O+, which raises total H3O+ concentration slightly above 1 mM, and making choice D the correct answer.

Ba(NO2)2 is a basic salt due to the interaction of the nitrite ion with water according to the following reaction: NO2- + H2O -> HNO2 + OH- Why does the equilibrium shown favor the left side of the reaction? A. H2O is a stronger acid than HNO2. B. H2O is a weaker base than OH−. C. NO2− is a stronger base than OH−. D. NO2− is a weaker base than OH−.

D. NO2− is a weaker base than OH−. Acid/base equilibria always favor the weaker acid/base pair. Since NO2- has a negative charge and no proton to donate, it must be the base, making water the acid. If the equilibrium favors the left hand side of the reaction, NO2- and H2O must be the weakest base and acid, respectively, eliminating these choices. While the choice, "H2O is a weaker base than OH−" is a true statement, water is not acting as a base in this particular reaction so this choice is not a good justification.

Which is more basic, I- or F-, SCH3- or OCH3? Why?

F-, OCH3-

True or False. Atoms with lone pairs are usually not basic

False. Atoms without lone pairs are usually not basic

True or False. Neutralization reactions are endothermic.

False. They are exothermic. Acid and base combining to become salt is exothermic

Can Group I and II cations be acidic?

Group I and II cations are never acidic

List all the strong acids

H2SO4, HClO4, HClO3, HNO3, HCl, HBr, HI

What is the standard equation for a weak acid? What is the equilibrium constant?

HA + H2O -><- H3O+ + A- Ka = [H3O+][A-]/[HA]

What is an example of an Arrhenius acid reaction?

HCl ->(H2O))<- H+ + Cl-

What is the exception to the Arrhenius definitions of acids and bases?

Molecules with no OH ions can still produce basic solutions and react with acids.

Which is more acidic, H2O or CH4, H2S or HCl? Why?

More electronegative atom increases acidity H2O, HCl

What is an example of weak base?

NH3

Out of the following, which will be acidic? NH4Cl, NaNO3, FeCl3

NH4Cl and FeCl3

Which of the following are basic? NaF, Li2CO3, MgBr2

NaF, Li2CO3

List all the strong bases

O2-, OH-, OR-, NH2-, NR2-, H-, R-

Molality is used instead of molarity in situations where the temperature of a solution is going to change. This is because the molarity of a solution changes when its density changes, but molality does not vary with density. Which of the following measures of concentration also does not vary with density? A. Normality B. Volume-volume C. Mass-volume D. Mole fraction

Since normality is simply an integer multiple of molarity, it also changes with density (eliminate A). Volume-volume and mass-volume solutions are formed identically to molar solutions, by taking a known amount of solute and diluting it to a particular final volume of solution. They will both be dependent on density (eliminate B and C). The mole fraction is a ratio of moles without regard to volume and will not change with temperature or density changes. D

How can you determine the pH from knowing the concentration of hydronium ions?

[H3O+] = 10^-pH

The resulting solution made from the combination of 50 mL of 1.0 M LiOH with 50 mL of 1.0 M HBr will be identical in all respects to 100 mL of: A. 0.5 M LiBr. B. 1 M LiBr. C. a saturated solution of LiBr(s). D. 2 M LiBr.

The best way to approach this question is to ignore OH- and H+, because they will neutralize each other when mixed. Thus, in this question we are mixing 50 mL of 1.0 M Li+ and 50 mL of 1.0 M Br-. Since the solution volume is doubling, the concentration of each ion is expected to decrease by a factor of 2. Therefore, the correct choice is "0.5 M LiBr".

How can Ka be determined to be high or low?

The equilibrium falls to the side where the molecule is less likely to donate ions, increasing the concentration of the molecule. If higher concentration on left side, Ka small. If higher concentration on right side, Ka large

Explain the titration between an AA and strong base.

The first equivalence point occurs at an early stage and then the second occurs at a later stage

What is the relationship between Kb and basicity?

The higher Kb is, the stronger the base is

What is the relationship of the acidity and pKa?

The lower the pKa, the greater the acidity

A 2 M solution of the potent organic pollutant sodium methylthiolate (NaSCH3) is titrated to its endpoint with 2 M HCl. Which of the following best approximates the pH of the resultant solution? (pKb of SCH3- = 3.6) A. -log (1 × 10-3.6) B. -log (1 × 10-5.2) C. -log (2 × 10-5.2) D. -log (1 × 10-1.8)

The result of the titration will be a 1 M solution of methylthiol, HSCH3 since the initial 2 M solution has been doubled in volume during the titration. Since we know the pKb of its conjugate base SCH3- is 3.6, then we know the pKa of HSCH3 is 10.4. Therefore, at equilibrium: 10-10.4 = (x)(x)/(1-x) Where x = [H+] = [SCH3-], at equilibrium. As the resultant value of x is negligibly small compared to 1, we ignore x in the denominator and say that x = [H+] = [SCH3-] = (10-10.4)(0.5) giving 1 × 10-5.2. The answer -log (2 × 10-5.2) would be correct if the resultant solution was 2 M in concentration. The answer -log (1 × 10-1.8) is the result of using the pKb of the methylthiolate anion in the calculation rather than the pKa of HSCH3, and -log (1 × 10-3.6) would be the result of a similar a calculation using this pKb value, while not taking the square root.

What is the relationship between acid strength and ability to donate protons? How does this affect Ka?

The stronger the acid, the more likely the acid is to donate protons, the larger the Ka

What is the relationship of an acid and its conjugate base?

The stronger the acid, the weaker the conjugate base

What is the effect of adding acid to a buffer system?

acid would form with conjugate base and form more of the undissociated acid

What color do litmus paper turn in when acidic, basic or neutral?

acid: pink base: blue neutral: purple

How does water act in acidic conditions and basic conditions in Bronsted-Lowry?

acidic conditions: acts as a base basic conditions: acts as an acid

What are Arrhenius acids and what are Arrhenius bases?

acids: compounds that produce H+ ions in H2O bases: substances that produce OH- ions in H2O

Compounds that acts as acids and bases are called

amphoteric

In the reaction of water and HCl, which is the Bronsted Lowry acid and which is the the Bronsted base? Which are the conjugate acid and the conjugate base?

base: water acid: HCl conjugate acid: H3O+ conjugate base: Cl-

Are Lewis Bases electrophiles or neutrophiles? Lewis acids?

bases: electrophiles acids: nucleophiles

What is an example of a weak acid?

carbonic acid (H2CO3)

The equation pKa + pKb = 14 can be only used for

comparing strengths of conjugate acids and bases

The bond that forms in the reaction of a Lewis base and Lewis acid is called: Why?

coordinate covalent bond The electrons come from a single atom

If pH > pKa, then the acid is

deprotonated

What is pI?

isoelectric point; the point of where the compound is neutral

What is a buffer?

mixture of a weak acid and its conjugate base/solution that resists changes in pH (cannot eliminate)

A titration of a strong acid and a weak acid will have an equivalence point of:

pH < 7

How can the pH for a weak acid be calculated? pOH for weak base?

pH = -1/2log(Ka[WA]) pOH= -1/2log(Kb[WB])

How is the equivalence point determined for a titration with a weak diprotic acid with a strong base?

pH = 1/2 (pK1 + kK2) (pI) pK1 and pK2 are half equivalence points

Calculate the pH of an aqueous ammonia solution with a [OH-] of 2.1 * 10^-3 M.

pH = 14 - (-log [2.1E-3]) = 11.32

What is the Henderson-Hasselbalch equation?

pH = pKa + log [A-]/[HA]

A titration of a weak acid and a strong acid will have an equivalence point of:

pH > 7

If pH < pKa then the acid will be

protonated

What is the effect of adding base to a buffer system?

protons will be taken by base and not much of acid will dissociate as there is an abundance of the conjugate base in solution

Compare the dissociation of strong bases and weak bases through equilibrium.

strong bases dissociate more completely and favor the products -weak bases dissociate less completely and favor the reactants (Kb < 1)

What is buffer capacity?

the amount of acid or base that can be added to a buffer solution before a significant change in pH occurs

What is the equivalence point?

the number of moles of hydrogen ions equals the number of moles of hydroxide ions

How can atoms without hydrogen be identified as acids?

they are electron deficient or with large positive charges

When is the pH scale of 0 to 14 and 7 as neutrality not applicable?

when there is a negative pH or non-7 neutrality