8.1 ACID-BASE EQUILIBRIUM: The Nature of Acid-Base Equilibria
Arrhenius acid/base Theory:
Acid - H+ ions in water,
Bases – OH- ions in water
acids – sour, turn litmus red, conduct electricity
bases – bitter, soapy taste, slippery, conduct electricity, turns litmus blue
hydronium ions: H2O(l) + H+(aq) -> H3O+(aq) hydronium
It is important to realize that whenever you talk about hydrogen ions in
solution, H+(aq), what you are actually talking about are hydronium ions.
Bronsted-Lowry Theory:
•
An acid is a proton (hydrogen ion) donor.
•
A base is a proton (hydrogen ion) acceptor.
Advantage of B-L definition:
1. reactions can be identified as acid-base neutralization reactions without water
2. salts that form acidic or basic solutions when dissolved can be explained
Example:
•
Either way, the ammonia acts as a base by accepting a hydrogen ion from an acid.
Reversible Acid-Base Reactions:
Weak Acid-Base reactions can proceed in both forward & reverse directions, having both an acid and a
base on both sides of the reaction arrow
Example:
1
Members of a conjugate pair differ from each other by the presence or absence of the transferable
hydrogen ion.
Example:
HC2H3O2(aq) + H2O(l) ->
C2H3O2-(aq) + H3O+(aq)
only about 1.3% of a 0.1 mol/LHC2H3O(aq) have reacted at SATP…this explains why the equilibrium
lies far to the left and why we call acetic acid, a weak acid
A Competition for Protons
•
•
•
stronger acid succeeds in donating the most protons, has a weaker conjugate base
stronger base succeeds in accepting the most protons, has a weaker conjugate acid
the strong acid and strong base are always on the same side of the reaction arrow
Amphoteric substances
A substance which can act as either an acid or a base is described as being amphoteric.
Example:
The Autoionization of Water
Water does ionize into H+ and OH- ions to a very small extent, however, a process we can represent most
simply by the following equilibrium:
…but at room temperature only about one out of every 109 molecules is ionized at any given instant.
Memorize this expression:
2
Example 1:
Calculate the values of [H+(aq)] and [OH-(aq)] in a neutral solution at 25°C.
In an acid solution [H+(aq)] is greater than 1.0 × 10-7 M; in a basic solution [H+(aq)] is less than 1.0 × 10-7 M.
Example 2:
Indicate whether each of the following solutions is neutral, acidic, or basic:
(a) [H+(aq)] = 2 × 10-5 M;
(b) [OH-(aq)] = 3 × 10-9 M;
(c) [OH-(aq)] = 1 × 10-7 M.
Strong Acids/ Strong Bases
A strong acid/base is one which is virtually 100% ionized in solution.
•
ionic hydroxides are all strong bases
•
only a few strong acids
Examples:
hydrochloric acid, HCl(aq),
hydrobromic acid, HBr(aq),
hydroiodic acid, HI(aq),
sulphuric acid, H2SO4(aq),
perchloric acid, H3ClO4(aq),
nitric acid, HNO3(aq),
3
pH (Power of Hydrogen ) and pOH Scales
With this scale, calculating the pOH can be done in the same manner as the pH scale.
Furthermore, because
pKW = -logKW
(See text p. 542-3)
Then…
And you can quickly convert from pH to pOH and vice versa.
Example:
To find the pH of 0.00500 mol L-1 calcium hydroxide solution:
Because the calcium hydroxide is fully ionic, each mole of it gives that double the number of moles of
hydroxide ions in solution.
[OH-] = 0.0100 mol L-1
Note: Ca(OH)2.produce twice as many hydroxide ions in solution.
Method 1:
[H+] [OH-] = 1.00 x 10-14
Method 2:
pOH = -log[OH-(aq)]
4
Example:
What is the pH, pOH and [OH-(aq)] of 0.0850 M HNO3(aq) solution?
Solution
Nitric acid is a monoprotic strong acid.
The balanced equation for the dissociation of HNO3(g) in water is:
100%
HNO3(aq) -> H+(aq) + NO3-(aq)
H2O(l)
Solve:
8.1 Practice (p.532, 537, 540, 546, 549): 1ab, 2, 3, 4ab, 5, 7(Q only), 10, 12a, 13, 5abc (d optional), 19
8.1 Questions (p.549-550): 1abc, 2, 5ab, 6abcd
8.2 ACID-BASE EQUILIBRIUM: Weak Acids and Bases
Weak Acids and Weak Bases do not ionize fully when dissolved in water.
Percent Ionization of Weak Acids:
Weak acids ionize less than 50% and have pH’s closer to 7
example: acetic acid, HC2H3O2(aq) , ionized only 1.3% in solution at 25C and 0.10 mol/L
What is the pH, 0.10 M acetic acid solution?
(p. 554) Practice # 1, 2
5
Percent Ionization and Concentration:
The more dilute the concentration, the greater the % ionization
This is an application of Le Chatelier’s Principle
Figure 4 on page 556:
Ionization Constants for Weak Acids:
HA (aq) ↔ H+ (aq) + A –1 (aq)
Ka =
Ionization Constants for Weak Bases:
- for B (aq) + H2O (l) ↔ HB+ (aq) + OH –1(aq)
Kb =
(p. 556) Practice # 3, 4, 5.
The Relationship Between Ka and Kb:
Ka x Kb = Kw
Proof: p. 560
6
The pH of Weak Acid Solutions:
Use the equation for Ka
the HUNDRED rule: shortcut for VERY weak acids: from the ICE table – since the % ionization is
very low, [HA] at equilibrium is assumed to equal the original [HA]
The HUNDRED rule if…
The 5 % rule: validate hundred rule …
Polyprotic acids:
They have more than 1 ionizable hydrogen…
Most often: [H ] is due to the first ionizable H only – the only type of calculation we will do with these
acids
(p 568, 570) 8.2 Practice # 7, 8, 9, 10.
In two days: (p 579) Section 8.2: 1-7, 9, 10, 12- 18, 21a, 22abc, 23a.
7
8.3 ACID-BASE EQUILIBRIUM: Acid-Base Properties of Salt Solutions
POINT OF EMPHASIS:
weaker the acid (Ka), stronger the conjugate base (Kb) ….visa versa
Salts that form NEUTRAL Solutions:
salts with cations from strong bases: sodium and potassium ions ( from NaOH and KOH ):
salts with anions from strong acids: chloride and bromide ions (from HCl, and HBr)
Review: Strong acids … Review Section
8.1
Example 1:
Salts derived from a strong acid and a strong base.
NaCl
•
•
or Ca(NO3)2,
no hydrolysis of cation or anion
pH = 7.00
Backtrack… derived from NaOH and HCl and from Ca(OH)2 and HNO3, respectively.
8
Salts that form ACIDIC Solutions:
contain the cations of weak bases ex: NH4+ (aq) salts
contain highly charged metal ions ex: Fe3+ (aq)
cause water to hydrolyze
Example 2: Salts derived from a weak base and a strong acid.
NH4Cl or Al(NO3)3
•
•
hydrolysis of cation from weak base (to produce H3O+(aq) ions )
pH < 7.00 (acidic)
Salts that form BASIC Solutions:
- contain anions of weak acids
Example 3:
Salts derived from a strong base and a weak acid.
Na2CO3, NaClO, or Ba(C2H3O2)2.
•
•
•
anion is a relatively strong conjugate base.
hydrolysis of anion from the weak acid (to produce OH-(aq) ions)
pH > 7.00 (basic)
Problems:
1. RbF
3. KNO3
2. FeCl3
4. NH4C2H3O2
9
Salts that Act as Acids and Bases:
contain cation of a weak base and anion of a weak acid – the ion with the largest K will have the
most effect
Example 4: NH4CN
Therefore basic.
Hydrolysis of Amphoteric Ions
check the Ka and Kb values
the larger value determines the dominant nature of the ion and so you can predict if the solution
will be acidic or basic
Example: Na2HPO4
Na+(aq) , a spectator ion, does not affect the acidity or basicity of a solution,
but HPO4-2(aq) could. The two possible reactions that it could undergo are:
Reaction A:
HPO42-(aq) + H2O(l)
H3O+(aq) + PO43-(aq)
Ka for HPO4-2(aq) is 4.2 x 10-13.
Reaction B:
HPO42-(aq) + H2O(l)
H2PO4-(aq) + OH-(aq)
Ka for H2PO4-(aq) is 6.2 x 10-8.
The Kb of reaction B is related to the Ka of H2PO4-(aq) by Kb = Kw / Ka:
Kb = { 1.0 x 10-14 / 6.2 x 10-8 } = 1.6 x 10-7 for HPO42-(aq)
Kb > Ka, so the solution will be basic.
p 588 591 practice # 1,2, 5, (8) (optional), p 594 8.3 Question # 1, 2, 3, 4, 5
10
8.4 ACID-BASE EQUILIBRIUM: Acid-Base Titration
pH (TITRATION) CURVES
The equivalence point of a titration
When you carry out a simple acid-base titration, you use an indicator to tell you when you have the acid
and alkali mixed in exactly the right proportions to "neutralize" each other. When the indicator changes
colour, this is often described as the end point of the titration.
In an ideal world, the colour change would happen when you mix the two solutions together in exactly
equation proportions. That particular mixture is known as the equivalence point.
For example, if you were titrating sodium hydroxide solution with hydrochloric acid, both with a
concentration of 1 mol/L, 25 mL of sodium hydroxide solution would need exactly the same volume of
the acid - because they react 1 : 1 according to the equation.
In this particular instance, this would also be the neutral point of the titration, because sodium chloride
solution has a pH of 7.
But that isn't necessarily true of all the salts you might get formed.
For example, if you titrate ammonia solution with hydrochloric acid, you would get ammonium chloride
formed. The ammonium ion is slightly acidic, and so pure ammonium chloride has a slightly acidic pH.
That means that at the equivalence point (where you had mixed the solutions in the correct proportions
according to the equation), the solution wouldn't actually be neutral. To use the term "neutral point" in
To summarize:
•
•
•
11
Titration curves for strong acid vs strong base
We'll take hydrochloric acid and sodium hydroxide as typical of a strong acid and a strong base.
Running strong acid into strong base
Running strong base into strong acid
This is very similar to the previous curve except, of course, that the pH starts off low and increases as you
add more sodium hydroxide solution.
Again, the pH doesn't change very much until you get close to the equivalence point. Then it surges
upwards very steeply.
Titration curves for weak acid vs strong base
12
We'll take ethanoic acid and sodium hydroxide as typical of a weak acid and a strong base.
Titration curves for strong acid vs weak base
This time we are going to use hydrochloric acid as the strong acid and ammonia solution as the weak base.
A summary of the important curves
13
The way you normally carry out a titration involves adding the acid to the alkali. Here are reduced
versions of the graphs described above so that you can see them all together.
More complicated titration curves
Adding hydrochloric acid to sodium carbonate solution
The overall equation for the reaction between sodium carbonate solution and dilute hydrochloric acid is:
The actual graph looks like this:
14
Part A:
Part B:
8.4 Practice (599) 1, 2, 3, 7, 8, 9
8.4 Questions (p 613) 1, 2, 4, 5, 10, 11.
15
8.5 ACID-BASE EQUILIBRIUM: Buffers
buffer: A buffer solution is one which resists changes in pH when small quantities of an acid or an alkali
are added to it.
•
usually weak acid with added anion (conjugate) of the acid,
ex: acetic acid + sodium acetate
•
•
or… weak base with added cation (conjugateof the weak base
ex: aqueous ammonia + ammonium chloride
Explaining Buffers:
A buffer solution has to contain things, which will remove any hydrogen ions or hydroxide ions that might
be add to it.
Acid buffer solutions:
Example: ethanoic acid and sodium ethanoate buffer
Review of conditions:
Ethanoic acid… weak acid… equilibrium to the left
Le Chatelier's Principle, adding sodium ethanoate will tip equilibrium further to the left.
•
•
•
if a small amount of base is added (reacts with hydrogen ions), the acid molecules ionize and [H+ ]
remains virtually unchanged so the pH remains virtually unchanged
if a small amount of acid is added, the anion accepts H+ and becomes molecular acid, the [H ]
remains virtually unchanged
if all the molecular acid or anion is used up, pH will change dramatically with addition of one
drop more of strong acid or base
Therefore buffer will contain:
16
Alkaline buffer solutions
Example: ammonia and ammonium chloride.
Review of conditions:
Ammonia … a weak base, equilibrium to the left
Le Chatelier's Principle, adding ammonium chloride (common ion effect) adds lots of extra ammonium
ions…removes some hydroxide…will tip equilibrium further to the left.
•
if a small amount of acid is added two processes can remove the hydrogen ions:
Removal by reacting with ammonia
Removal by reacting with hydroxide ions
•
if a small amount of base is added, the hydroxide ions from the base are removed by a simple
reaction with ammonium ions.
Therefore contain these important things:
17
Uses of Buffers
•
Living organisms … control pH of blood and other body fluids
•
Maintain aquarium and swimming pool water
•
Prevent internal harm from use of medicine e.g. buffered aspirin
•
Prevent negative effects of acid in digestive system e.g. antacids
•
Major buffer systems of the human body
bicarbonate
hemoglobin
phosphate
protein
CO2 + H2O H2CO3 H+ + HCO3Hb-H
Hb- + H+
H2PO4
H+ + HPO42Pr-H
Pr- + H+
in blood plasma
in red blood cells
in urine
intracellular fluid
Buffer Solution In Blood
Titration Curve for the Bicarbonate-Buffer System in blood
Any change in pH of more than 0.2 is life
threatening. If blood were not buffered, a glass of
orange juice would be deadly.
8.5 Practice (p.620): 1a
8.5 Questions (p.620): 1, 2, 3, 5, 7abcd,
The slope of the curve is flattest where the pH is
equal to the pK value (6.1) for the buffer. Here,
the buffering capacity is greatest because a shift
in the relative concentrations of bicarbonate and
carbon dioxide produces only a small change in
the pH of the solution. However, at pH values
higher than 7.1, the slope of the curve is much higher. Here, a shift in the relative concentrations of
bicarbonate and carbon dioxide produces a large change in the pH of the solution. Hence, at the
physiological blood pH of 7.4, other organs must help to control the amounts of HCO3- and CO2 in the
blood to keep the pH relatively constant, as described above.
18
Course Code: SCH 4U1- _____
Name: _____________________________________
Date: ________________
EQUILIBRIUM PROBLEM SET #1
Calculate the pH, pOH, and [OH-] of a solution with 1.23 x 10-4 M HCl
Calculate the pOH, pH, and [H3O+] of a solution containing 1.23 x 10-4 M Ca(OH)2.
1.00 L of solution was made by dissolving 0.80 g of NaOH in water. Calculate [OH-], [H+] and pH.
Which of the following is usually referred to as a
a) strong acid in water solution? What are the rest?
HF, HNO2, H2CO3, H2S, HSO4-, Cl-, HNO3, HCN
b) strong base in solution? What are the rest?
Zn(OH)2 , Ca(OH)2 , Fe(OH)3 , Al(OH)3 , NH3 , NH4OH
5. Write the equilibrium system (equation) for the reaction between nitrous acid and water, and write the
expression for its acid-dissociation constant.
6. Using Appendix C p. 803 of the text (Monoprotic Acids), calculate the base-dissociation constants, Kb
, for the nitrite ion, NO2-(aq) , and the benzoate ion, C6H5CO2-(aq) .
7. Using Appendix C p. 803 of the text (Weak Bases), calculate the acid-dissociation constants, Ka , for
the conjugate acid ions, CH3CH3NH2+(aq) , and N2H5+(aq) .
8. Using Appendix C p. 803 of the text (Monoprotic Acids), determine which is the stronger acid, formic
acid (an irritant found in ant bites; also called methanoic acid) or phenol (formerly used as an
antiseptic in hospitals)?
9. What is the pH of a 0.100 M acetic acid with Ka = 1.77 x 10-5 ?
10. Determine the pH of 0.150 M ammonia (NH3) with Kb=1.8 x 10-5.
11. Is a solution of sodium acetate acidic, neutral or basic? Explain.
12. Is a solution of ammonium chloride acidic, basic or neutral? Explain.
13. Determine the pH of a solution formed by mixing 0.100 mol of HC2H3O2 with a Ka of 1.8x10-5 and
0.200 mol of NaC2H3O2 in a total volume of 1.00 L.
14. What is the pH of a 0.400 M KBr solution? Explain.
15. Use the following titration curves.
1.
2.
3.
4.
a)
b)
c)
d)
Compare and contrast titration curves A and B.
Provide titles for each titration curve from A to E.
Determine the equivalence points for each titration curve from A to E.
Using Appendix C10 on p. 804 (Acid-Base Indicators), choose the best acid-base indicator for
each titration from A to E.
A
B
19
C
D
E
16. Using Appendix C10 on p. 804 (Acid-Base Indicators), choose the best acid-base indicator for
titrations in which the predicted equivalence point has pH of
(a) 4.5 (b) 5.6 (c) 6.8 (d) 8.1 (e) 10.0
17. What is a buffer? What is a buffer’s composition? Construct the titration curve for the carbonic acid –
bicarbonate buffer system. Explain how it works. Describe at least one each of biological,
pharmacological, commercial, and environmental uses of buffers.
18. Title and label the apparatus.
1) pH =3.910 pOH = 10.090 [OH-(aq)] =8.13 x 10-11 M 2) pOH = 3.609 3) [NaOH (aq)] = [OH-(aq)] = 2.0 x 10-2 M pOH = 1.70
4) a) HNO3 , b) Ca(OH)2 5.) 6) Kb , NO2-(aq) = 1.4 x 10-11
Kb , C6H5CO2-(aq) = 1.6 x 10-10 7.) Ka , CH3CH3NH2+ (aq) = 1.0 x
-11
+
-8
10 Ka , N2H5 (aq) = 1.0 x 10 8.) formic acid 9.) pH = 2.876 10) pH = 14.000 – 2.785 = 11.215 11. basic 12. acidic 13. pH
= - log [H3O+] = - log (9.00 x 10-6) = 5.046 14. neutral 15.) e. A. litmus B. litmus C. bromocresol green D.phenolphthalein E.
litmus 16. (a) bromocresol green (b) methyl red (c) p-nitrophenol & litmus (d) cresol red (e) α-naphtholbenzene 17.
18.
20
Name:______________________ Date:______________
Chemical Systems and Equilibrium Common Ion Effect
Use sigdig rules and scientific notation for calculations. Proper English usage is required.
Part One Solubility Product Constant
Many common salts which have a very limited solubility in water are called slightly soluble salts. A
saturated solution of a slightly soluble salt is a dynamic equilibrium between the solid salt and its ions in
solution. The mass action expression for this equilibrium equals an equilibrium constant, called the
solubility product constant, Ksp, for the salt.
Purpose: to find the Ksp and molar solubility of calcium hydroxide
Question: What are the Ksp and molar solubility of calcium hydroxide?
Materials available: lab apron, safety goggles, buret, glass funnel, 3 beakers (Acid, Base, and Waste),
125-mL Erlenmeyer flask, suction bulb, 10.0 mL transfer pipette, masking tape, bromothymol blue,
distilled water, 0.100 M HCl(aq) , 10.0 mL of saturated Ca(OH)2 (aq) , and paper towels
Safety Precautions: Safety goggles and apron must be worn. If you are unsure of safety concerns,
consult with the teacher first.
Experimental Design:
(a) Design the steps of an experiment to find the Ksp and molar solubility of calicum hydroxide. Use IPTPV. (3 marks)
21
Observations:
Table 1 . Ksp and Molar Solubility of Ca(OH)2
Trial 1
Trial 2
Trial 3
Concentration of standard
HCl solution (mol/L)
Volume of saturated
Ca(OH)2 solution (mL)
Buret reading, initial (mL)
Buret reading, final (mL)
Volume of standard HCl
used (mL)
Concentration of
Ca(OH)2 (M)
[OH-] at equilibrium
(mol/L)
[Ca2+] at equilibrium
(mol/L)
Molar Solubility of
Ca(OH)2
Ksp of Ca(OH)2 at room
temperature
Table 2 . Ksp and Molar Solubility of Ca(OH)2 with added Calcium ions
Trial 1
Trial 2
Trial 3
Concentration of standard
HCl solution (mol/L)
Volume of saturated
Ca(OH)2 solution (mL)
Buret reading, initial (mL)
Buret reading, final (mL)
Volume of standard HCl
used (mL)
Concentration of
Ca(OH)2 (M)
[OH-] at equilibrium
(mol/L)
[Ca2+] at equilibrium
(mol/L)
Molar Solubility of
Ca(OH)2
Ksp of Ca(OH)2 at
room temperature
22
(b) Calculations Show work only for trial 1 in table 1.
(11 marks …)
Evaluation:
(c) What is the average Ksp for each table of observations? (2 marks)
(f) Write the net ionic equation for the dissociation of solid calicum hydroxide.
(g) How does the addition of CaCl2 affect the molar solubility of Ca(OH) 2?
(2 marks)
( 1 mark)
23
(h) How will the transfer of some solid Ca(OH)2 into the Erlenmeyer titrating flask affect the reported Ksp
value of Ca(OH)2? Explain. (2 marks )
(i) As a result of this transfer error, will the molar solubility of Ca(OH)2 be higher or lower than the
accepted value for Ca(OH)2? Explain.
(2 marks.)
(j) Observe and record the meniscus in Figure 1A and 1B to the proper sigdigs.
Figure 1A (initial)
(2 marks)
Figure 1B (final)
1A: ___________________________
1B: ___________________________
24
(k) If the endpoint of the titration is surpassed, will the Ksp value of Ca(OH)2 be higher or lower than the
accepted value? Explain. (2 marks)
(l) How would the use of freshly boiled (uncooled) distilled water affect the calculated Ksp value of of
Ca(OH)2. Explain.
(2 marks
(m) How would the use of distilled water left on the bench overnight affect the calculated Ksp value of of
Ca(OH)2. Explain.
(2 marks)
(n) How would the use of tapwater affect the calculated Ksp value of of Ca(OH)2. Explain.
(2 marks)
25
Name:__________________________________ Date:_____________________
Chapter 8 (8.1 to 8.5) Quiz
Acid-Base Equilibrium
/47
Knowledge
(True/False)
3 marks
Indicate whether the sentence or statement is true or false.
Hb- + H+
____
1. This buffer system operates in human red blood cells. Hb-H
___________________________________
____
2. During weak acid - strong base titrations the equivalence point has a pH less than 7.00.
___________________________________
____
3. An acid base indicator is usually itself a weak acid.
__________________________________
Knowledge
(Multiple Choice)
2 marks
Identify the letter of the choice that best completes the statement or answers the question.
____
4. Identify the conjugate base in this equilibrium system.
CH3COOH(aq) + H2O(l)
CH3COO - (aq) + H3O+(aq)
a. CH3COOH(aq)
d. H3O+(aq)
e. none of the above
b. H2O(l)
c. CH3COO - (aq)
____
5.
a.
b.
c.
Knowledge
Which of the following salts could be combined with hypochlorous acid to form a buffer?
NaClO
d. Zn(CN)2
FeCl2
e. none of the above
MgSO4
(Short Answer)
12 marks
6. Define
(a) a Bronsted-Lowry acid.
(b) equivalence point.
(1 mark)
(1 mark)
.
26
7. Predict whether the following solution is acidic, basic, or neutral without using Ka and Kb
values. Provide derivations of the ions to support your predictions.
(2 marks)
... ammonium sulfate, (NH4)2SO4 (aq) (fertilizer)
8.
9.
.
10.
Inquiry:
17 marks
11. What is the pH of 0.250 M ammonia (NH3) with Kb = 1.8 x 10-5? Use both assumptions and
attempt to validate them.
(8 marks)
NH3 (aq)
+
H2O (l)
NH4+ (aq)
+
OH- (aq)
27
12.
13.
14.
Communication:
6 marks
15. a) Determine the strengths of acid and base in each titration. (2 marks)
b) Determine the equivalence point for each titration.
(2 marks)
c) Choose the best acid-base indicator to use for each titration from Appendix C10 on p. 804.
(2 marks)
A
B
28
strength of acid
strength of base
equivalence point
acid-base
indicator
A
B
Making Connections:
7 marks
16. **8.5**
Discuss the composition and operation of the main buffer system for the pH of human blood.
(4 marks)
Chapter 8 Quiz
Answer Section
Acid-Base Equilibrium
v9
(8.1 to 8.5)
/47
MODIFIED TRUE/FALSE
1. T
2. F, pH more than 7.00
3. T
MULTIPLE CHOICE
4. C
5. A
SHORT ANSWER
6.
ANS:
(a) proton (hydrogen ion) donor
(b) time in titration when amount of H3O+ and OH- are equal
.
7.
Practice#1b and 2, pg 588
derived from weak base and strong acid ... acidic property
8.
Practice #3, pg 599
9.
Practice #7, 8, 9 on pg 611
29
10. ANS:
Practice #1 and 2 pg 620
PROBLEM
11. ANS:
NH4+ (aq) + OH- (aq)
NH3 (aq) + H2O (l)
Kb = (NH4+] [OH-] ) / [NH3] = 1.8 x 10-5
Initial concentrations (mol/L)
Change in concentrations (mol/L)
Equilibrium concentrations (mol/L)
NH3 (aq) +
0.250
-x
0.250 - x
H2O (l)
NH4+ (aq) +
0
+x
x
OH- (aq)
≅0
+x
x
[OH-] = [NH4+] = x
1.8 x 10-5 = {(x) (x)} / (0.250 - x)
First assumption: x is small so that 0.250 - x is 0.250.
1.8 x 10-5 = (x) (x) / 0.250)
x2 = 4.50 x 10-6
x = 2.12 x 10-3 M
pOH = - log [OH-] = - log (2.12 x 10-3) = - (-2.673) = 2.673
pH = 14.000 – 2.673 = 11.327
Checking the 5% rule: ( [x] / [NH3]0 ) x 100% < 5%
2.12 x 10-3 M / 0.250 M x 100% = 0.85%,
This means the assumption that 0.250 - x approximately equals 0.250 is valid
(acceptable).
Second assumption: initial concentration of OH- (aq) due to the autoionization of water
is so small that it is zero. This assumption is valid because Kw (1.0 x 10-14) is much less
than Kb (1.8 x 10-5) or Kb x Cb (1.8 x 10-5)(0.250) is less than 1.0 x 10-13.
12. ANS:
Practice 4ab pg 537
30
13. ANS:
Practice 19, pg 549
14. ANS:
Practice #6, pg 562
15. ANS:
a) Determine the strengths of acid and base in each titration. (2 marks)
b) Determine the equivalence point for each titration. (2 marks)
c) Choose the best acid-base indicator to use for each titration from Appendix C10 on p. 804.
(2 marks)
A
B
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A
B
strength of acid
strength of base
equivalence
point
acid-base indicator
strong
weak
strong
strong
7.0
9.2
litmus
phenolphthalein
OTHER
16. ANS:
composition of buffer system: (either)
1) carbon dioxide - carbonic acid - bicarbonate buffer ... weak acid and its salt OR
2) CO2(aq) + H2O(l) ⇔ H2CO3(aq) ⇔ H+(aq) + HCO3-(aq)
operation: (any two)
1) As acid is added, the pH decreases and the buffer shifts toward greater H2CO3 and
CO2 concentration.
2) Conversely, as base is added, the pH increases and the buffer shifts toward greater
HCO3- concentration.
3) The region of maximum buffering capacity where pH equals pK value for buffer = 6.1
4) The graph of pH vs buffer system shows flattest part of curve in which the composition
is about 50% carbonic acid and 50% bicarbonate ion.
5) It resists changes to pH when small amounts of acid or base are added, especially when
acid is being added.
.
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