Download PRACTICE TEST for EXAM 10

PRACTICE TEST for EXAM 10 (ACIDS & BASES)
1. List the observable properties by which you could identify a clear, colorless, odorless liquid as an
acid, a base, a neutral electrolyte, or a nonelectrolye.
2. Write balanced equations for each of these reactions:
a. The reaction of iodic acid (HIO3) with water according to the Bronsted model. What is the
characteristic behavior of a Bronsted acid?
b. The reaction of ammonia (NH3) with water, according to the Bronsted model. What is the
characteristic behavior of a Bronsted base?
c. The complete neutralization of H2SO4 with NaOH. Identify acid, base, and salt.
d. The complete neutralization of Mg(OH)2 with HCl. Identify acid, base, and salt.
3. Define these terms:
a. proton, as used in acid/base chemistry
b. monoprotic (include an example)
c. polyprotic (include an example)
d. amphoteric (include an example)
e. hydronium ion
4. Write the formulas for finding pH and for finding [H3O1+], convert these, and identify each solution
as acidic, basic, or neutral:
a. 0.000875 M H3O1+ = ? pH
c. 3.27 x 10–5 M H3O1+ = ? pH
1+
b. pH 6.58 = ? [H3O ]
d. pH 7.40 = ? [H3O1+]
5. What is the pH of pure water? Why does pure water conduct electricity weakly? Include the
chemical equation.
6. Draw a line to represent the pH scale. Label neutral, acid region, and base region. What happens to
the pH of a solution if the concentration of hydronium ion increases? What if it decreases?
7. Compare and contrast a strong acid and a weak acid, in terms of neutralization titration results, pH,
and conductivity, and explain their differences in terms of their reaction with water.
8. A weak acid is described as coming to equilibrium: HA (aq) + H2O  H3O1+ (aq) + A1– (aq)
a. What two processes have come into balance when this acid is at equilibrium?
b. Write the equilibrium constant expression, Ka, for this reaction.
c. How could you use the value of Ka to tell which of two acids is weaker?
9. What are the ingredients for a good buffer? What does a buffer do and how does it do it?
10. 2.00 mL of full-strength vinegar (HC2H3O2) are neutralized by 17.52 mL of 0.0960 M NaOH. What
is the concentration of the acetic acid in the vinegar?
11. How many grams of HC8H7O2 are neutralized by 35.00 mL of 0.100 M KOH?
12. How many mL of 0.0872 M NaOH will neutralize 0.305 g HC7H5O2?
13. What is the pH of 0.25 M hypobromous acid (HOBr)? The Ka of HOBr is 2.5 x 10–9
14. What is the pH of 0.050 M nitrous acid (HNO2)? The Ka of HNO2 is 7.2 x 10–4
15. The pH of 0.50 M chlorous acid, HClO2, is 1.16. What is Ka for chlorous acid?
16. The pH of 0.010 M hydrocyanic acid, HCN, is 5.60. What is Ka for hydrocyanic acid?
ANSWERS to PRACTICE TEST for EXAM 10 (ACIDS & BASES)
[See also Ch 19, sections 1-4]
1. Acid solution: conducts electricity, turns indicator paper red, tastes sour, reacts with most metals to
produce a gas, reacts with carbonates to produce a gas.
Base solution: conducts electricity, turns indicator paper blue, tastes bitter, NR with most metals,
NR with carbonates, dissolves grease.
Neutral electrolyte solution: conducts electricity but has no effect on indicator paper.
Non-electrolyte: does not conduct electricity
2. a.
b.
c.
d.
HIO3 + H2O  H3O1+ + IO31–
NH3 + H2O  NH41+ + OH1–
H2SO4 + 2 NaOH  2 H2O + Na2SO4
Mg(OH)2 + 2 HCl  2 H2O + MgCl2
Bronsted acids donate hydrogen ion
Bronsted bases take (accept) hydrogen ion
acid + base  water + salt
base + acid  water + salt
3. a. A proton in acid/base chemistry is a hydrogen ion, H1+. Acids donate it, bases accept it.
b. Monoprotic means the molecule has only one acidic hydrogen, available to donate to a base. HCl
is monoprotic. So is HC2H3O2: the H at the beginning is acidic, the 3 Hs in the middle are not.
c. Polyprotic means the molecule has more than one acidic hydrogen. H2SO4 and H3PO4 are
polyprotic.
d. Amphoteric means the molecule or ion may act as either an acid (H1+ donor) or a base (H1+
acceptor), depending on what other molecules are present in the solution. HOH is the most
obvious example, but many ions like HCO31– and H2PO41– are amphoteric.
e. Hydronium ion is H3O1+, the ion formed when a water molecule accepts an H1+ from an acid.
4. pH = –log [H3O1+]
[H3O1+] = 10–pH
1+
a. pH 3.06 (A) b. [H3O ] = 2.63 x 10–7 M (A)
c. pH 4.49 (A) d. [H3O1+] = 3.98 x 10–8 M (B)
5. The pH of water is exactly 7. Water reacts with itself to produce a small number of ions:
H2O + H2O  H3O1+ (aq) + OH1– (aq)
6. See pg 598 in text. If [H3O1+] increases, pH goes down. If [H3O1+] decreases, pH goes up.
7. Both acids react with water to give H3O1+ (see 2a). Neutralization titration gives the concentration of
acid, [HX], and pH gives the concentration of hydronium ion, [H3O1+], in each solution. In the
strong acid solution, [HX] = [H3O1+] and the solution is a strong electrolyte (good conductor). In the
weak acid solution, [HX] >> [H3O1+] and the solution is a weak electrolyte (poor conductor). The
explanation for these behaviors is that strong acids react 100% with water, producing lots of
hydronium ions, while weak acids react <<100% with water, producing only a few hydronium ions.
8. a. HA donating H1+ to H2O and A1– accepting H1+ from H3O1+ have come into balance.
 H 3O1   A1 
b. K a 
(note H2O was omitted because it is a liquid)
HA 
c. The smaller the numerical value of Ka, the weaker the acid.
9. A buffer is made of a weak acid and its salt, at equilibrium. A buffer resists pH change because it
can neutralize small additions of acid or base, then return to equilibrium with only a minimal change
in pH of the solution.
10. 0.841 M
13. pH 4.60
11. 0.477 g
14. pH 2.22
12. 28.6 mL
15. Ka = 1.1 x 10–2
16. Ka = 6.3 x 10–10