Chapter 20+
Acids and Bases
Chapter
20: Acids and Bases
Part 1 – Notes: Introduction to Acids and Bases
Electrolytes
An electrolyte is . . .
Question: How does a strong electrolyte differ from
a weak electrolyte?
Properties of Acids and Bases
Acid Ionization:
Acids are molecular compounds that
WILL form ions when dissolved in water. An acid
will donate a proton (a hydrogen ion) to the water forming the hydronium
ion and a negative ion.
Question: Ions are found in water solution as a result
of two different processes, ionization and dissociation. Describe
the differences between ionization and dissociation.
Arrhenius Definitions of Acids and Bases
Arrhenius Acid:
Ex.
Arrhenius Base:
Ex.
Bronsted-Lowry Definitions of Acids and Bases
Bronsted Acid:
Ex.
Bronsted Base:
Ex.
Lewis Definitions of Acids and Bases
Lewis Acid:
Ex.
Lewis Base:
Ex.
Salt:
Oxyacids: acids that contain hydrogen, oxygen,
and at least one additional element
Ex.
Binary Acid: acids that contain only two elements
Ex.
Monoprotic Acid: any acid that contains ONE ionizable hydrogen/proton
Ex.
Diprotic Acid: any acid that contains TWO ionizable
hydrogens/protons
Ex.
Triprotic Acid: any acid that contains THREE
ionizable hydrogens/protons
Ex.
Amphoteric:
Example:
The pH Scale
The pH scale shows the relationship between pH and the hydrogen-ion concentration.
Notice that acids have low pHs and bases have high pHs. A pH of 7 means
a substance is neither acidic or basic – it is neutral.
pH
|
0
|
1
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2
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3
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4
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5
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6
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7
|
8
|
9
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10
|
11
|
12
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13
|
14
|
|
|
|
|
|
|
|
|
|
|
|
|
|
|
|
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[H+]
|
100
|
10-1
|
10-2
|
10-3
|
10-4
|
10-5
|
10-6
|
10-7
|
10-8
|
10-9
|
10-10
|
10-11
|
10-12
|
10-13
|
10-14
|
Chapter 20: Acids and Bases
Part 1 – Assignment: Introduction to Acids and Bases
Determine whether the following acids are monoprotic, diprotic, or triprotic.
(M, D, or T???)
1. HCl
4. H2SO4
7. H3PO4
2. HClO4
5.
HClO3
8. HNO3
3. H3X
6. H2X
9.
HX
Are the following substances amphiprotic (amphoteric)? Write YES
or NO.
10. H2O
14.
H2PO4-1
18. HCO3-1
11. HCl
15. HPO2-2
19.
H2SO4
12. H I
16. H3PO4
20.
HSO4-1
13. NaOH
17. H2CO3
21.
SO4-2
22-24 Write an equation for the ionization of each of
the following acids in water solution.
22. HCl
23. H2CO3
24. H3PO4
25-27 Write an equation for the dissociation of each
of the following in water solution.
25. NaOH
26. Ca(OH)2
27. Sc(OH)3
28. A Bronsted acid is a substance that donates a proton
to another substance. Is a proton removed from the nucleus of one of
the atoms in an acid? If yes, how does this happen since it is not
a nuclear reaction? If no, how does the substance donate a proton?
29. What is the range for the pH scale?
Chapter
20: Acids and Bases
Part 2 – Notes: Conjugate Acid-Base Pairs and the Water Constant
Proton-Transfer Reactions and Conjugate Acid-Base Pairs
The Bronsted definition of acids and bases provides a basis for the study
of PROTOLYSIS, or proton-transfer reactions.
Recall: An Bronsted acid is a species capable of donating
a proton. A Bronsted base is a species that can accept a proton
Let’s suppose an acid gives up a proton. What is true of the piece
that remains after the proton is transferred?
Conjugate base:
What is the conjugate
base of HCl?
Conjugate acid:
What is the conjugate
acid of NH3?
You need to be able to identify conjugate acid-base pairs.
Example 1: HClO4
+ H2O -->
H3O+ +
ClO4-
Example 2: NH3
+ H2O
--> OH- + NH4+
Key point:
In a Bronsted acid-base reaction, the
stronger acid has the weaker conjugate base. In Example
1, ClO4-
ions are the conjugate base of HClO4.
Since the base, H2O, was able to
win the competition for the protons, it is stronger than ClO4-
ions. Thus, H3O+,
the conjugate acid of H2O,
is weaker than HClO4, the
stronger a Bronsted acid, the weaker its conjugate base.
Also, Example 2 shows that the conjugate
base of a very weak acid is always a strong base. Water is a very
weak acid; hence OH- ions, the conjugate
base of H2O, is a very strong
base and it will remove protons from any acid stronger than water.
Auto-Dissociation of Water
Pure water dissociates (or ionizes) very slightly into ions. Perhaps
collisions between water molecules may result in a transfer of proton from
one molecule to another. The equation for the dissociation (or ionization)
of water may be written as:
It has been found by experimentation that 1 L of pure water contains only
one ten-millionth of a mole of hydronium ions (1 x 10-7
mol) and one ten-millionth of a mole of hydroxide ions (1 x 10-7
mol). If a substance contains more hydronium ions that hydroxide, it
is acidic. If a substance contains more hydroxide ions than hydronium,
it is basic. Water is considered a neutral substance because it
contains equal numbers of hydronium and hydroxide ions.
The ions are in equilibrium in pure water; the concentration
of the hydronium ions is [H3O+]
= 1.00 x 10-7 M and the concentration
of the hydroxide ion is [OH-] =
1.00 x 10-7 M.
Since we know the equation, we can write the equilibrium expression for
the dissociation of water:
The concentration, in molarity of water is:
Since the concentration of the ions (1.00 x 10-7
M) is so small in comparison with the total concentration of water (55.5
M), the concentration of water can be assumed to remain constant.
If the concentration of water is constant, we can multiply both sides of
the equation by this constant without destroying the relationship.
Since both [H2O] and Keq
are constants, their product will also be a constant. We call this constant
the dissociation constant of water, and it is given the symbol
Kw.
Chapter 20: Acids and Bases
Part 2 – Assignment: Conjugate Acid-Base Pairs and the Water Constant
Show the conjugate acid base pairs for the following reactions. Write
A and B on the reactant side above the appropriate substance. Then write,
CA and CB on the product side above the appropriate substance.
1. CO3-2
+ H2O -->
HCO3-1
+ OH-1
2. HCN + H2O
--> H3O+
+ CN-1
3. H2S
+ (CH3)2NH
--> (CH3)2NH2+
+ HS-1
4. HF + H2O
--> H3O+
+ F -1
5. HC2H3O2
+ H2O --> H3O+
+ C2H3O2-1
6. HSO4-1
+ HPO4--2
--> SO4-2
+ H2PO4-1
7. N2H4
+ H2O -->
N2H5+
+ OH-1
8. NH2OH
+ HCl --> NH3OH+
+ Cl-1
9. NH2-1
+ N2H4
--> NH3
+ N2H2-1
10. HNO3
+ H2SO4
--> H3SO4+
+ NO3-1
Chapter
20: Acids and Bases
Part 3 – Notes: pH, pOH, and Kw – Acid Concentrations
Recall: The equation for the ionization of water is:
2 H2O
--> H3O+
+ OH-
Kw = [H3O+]
[OH-] = [1.00 x 10-7]
[1.00 x 10--7] = 1.00 x 10-14
(at 25oC)
The Concentration of Hydronium and Hydroxide
The simple relationship above has been used to construct the pH scale,
used to indicate how acidic or basic a substance is.
What happens when an acid is dissolved in water? An acid is a substance
whose water solution contains a large concentration of hydronium ions and
at 25oC, the above equation must
always be true.
So, what if the concentration of H3O+
goes up?
Similarly, a water solution of a base has a large concentration of OH-
and a small concentration of H3O+.
The above formula may be rearranged to determine the concentrations of
both:
[H3O+]
=
[OH-]
=
The pH Scale
The pH of a solution is the negative logarithm, to base 10, of the
hydronium ion concentration. It is expressed by the equation:
In general, the pH scale is a numerical scale that normally extends from
0 through 14. The numbers on the scale represent the acidity of solutions
and can be converted to hydronium concentrations. The midpoint of the
scale is taken as 7.
At 25oC, a solution with a pH
of 7.0 . . .
A solution with a pH less than 7.0 . . .
A solution with a pH greater than 7.0 . . .
The actual relationship between the pH and the molar concentration of
the H3O+
is:
The pOH Scale
The pOH scale is used to indicate the hydroxide concentration of a solution.
Note that as pH increases, the pOH decreases, and vice versa, because the
product of the ion concentrations must always be 1.00 x 10-14.
The pOH of a solution is the negative logarithm, to base 10, of the
hydroxide ion concentration. It is expressed by the equation:
The actual relationship between the pOH and the molar concentration of
the OH- is:
The Relationship Between pH and pOH
We know that as pH increases the pOH decreases and that as pH decreases
the pOH increases, we are able to relate the two scales. In any aqueous
solution at 25oC:
Question: What is the [H3O+]
if the acid is 0.75M HCl?
What is the [H3O+]
if the acid is 0.75M H2SO4?
Example 1: What is the pH of a solution
that contains 1.00 x 10-4 M hydronium
ion?
Example 2: Calculate the [H3O+]
of a solution that has a pH of 3.70.
Example 3: Find the pH and pOH of a solution
that contains 0.00350 M H3O+
ion.
Example 4: What is the [OH-]
and the [H3O+]
of a solution if the pOH is 4.40?
Example 5: What is the pH of 0.250 M
H3PO4?
This is a strong acid and dissociates 100%.
Chapter 20: Acids and Bases
Part 3 – Notes: pH, pOH, and Kw – Acid Concentrations
Supply answers for the following questions.
Show all work, set-ups, etc. Use additional paper if necessary to show
all work.
1. Calculate the pH of the solutions with these concentrations:
(a) 0.00010 mol
H3O+/L
(b)
0.0018 mol OH-/L
(c) 1.62 x 10-5
mol OH-/L
(d) 4.09 x 10-2
mol H3O+/L
2. Calculate the pOH of the solutions with these concentrations:
(a) 4.57 x 10-6
mol OH-/L
(b)
5.75 x 10-3 mol H3O+/L
3. Calculate the [H3O+]
of the following solutions.
(a) pH = 3.72
(b)
pOH = 12.0
4. Calculate the [OH-]
of the following solutions.
(a) pOH = 11.0
(b)
pH = 4.25
5. Calculate the (i) pH and the (ii) pOH of solutions
having the following concentrations. Assume 100% ionization.
Remember: 1 mole of H2SO4
produces TWO moles of H3O+
ions, etc.
(a) 0.0025 M NaOH
(b) 0.0025 M H2SO4
(c) 0.075 M H2SO4
(d) 0.048 M HCl
(e) 0.032 M KOH
(f) 0.00017 M Ca(OH)2
6. The pH of human blood is 7.4. What is the [H3O+]?
7. What is the [H3O+]
in a red ripe tomato that has a pH of 4.2?
What is the pOH? What is the
[OH-]?
8. Find the [H3O+]
and the pH of the following solutions:
(a) 0.325 M H2SO4
(b) 0.0375 M Ca(OH)2
9. Find the [OH-]
and the pH of the following solutions:
(a) 0.00862 M Mg(OH)2
(b)
1.25 M H3PO4
Chapter
20: Acids and Bases
Part 4 – Notes: Titration and Neutralization
Titration:
In acid-base titration, a base is added to an acid (or vice versa) until
. . .
This point of neutralization is called the end point
or equivalence point. These ions react to form water.
An indicator tells when the right amount of acid and base have been mixed.
When a hydronium ion combines with a hydroxide ion, the result is the
neutral compound water: H3O+
+ OH- -->
HOH
Strong and Weak Acids and Bases
A strong acid or a strong base will completely ionize or dissociate.
Weak acids and weak bases will not completely ionize.
The strong
acids you will commonly encounter are:
The strong
bases you will commonly encounter are:
The moderately
strong acids you will encounter are:
The weak acids
you will commonly encounter are:
Examples: HCl
--> H3O+
+ Cl-
one mole of HCl produces one mole of H3O+
ions
NaOH -->
Na+ + OH-
one mole of NaOH produces one mole of OH-
ions
H2SO4
--> 2 H3O+
+ SO4-2
one mole of sulfuric acid
produces 2 moles of H3O+
ions
Steps for Solving Titration Problems
1. Write a balanced equation for the neutralization
of the acid and base.
2. Use the information given in the problem to find the
moles of either acid or base.
3. Use the mole ratio from the balanced equation to
calculate the moles of the unknown acid or base.
4. Convert the moles of unknown into whatever units
the problem requires.
Example 1:
45.00 mL of a 0.350 M H2SO4 solution
is neutralized by a 0.425 M NaOH solution. Calculate the volume of NaOH
required.
Example 2:
50.0 mL NaOH solution is used to exactly neutralize 2.00 g of a solid
acid with a molar mass of 180. g/mol. This acid has a single ionizable
hydrogen. Find the concentration of the sodium hydroxide solution.
Example 3:
2.50 g of H2X (a strong acid
with two ionizable hydrogen) is exactly neutralized by 57.8 mL of 1.25 M
NaOH solution. What is the molar mass of the acid?
Chapter 20: Acids and Bases
Part 4 – Assignment: Titration and Neutralization
Answer the following questions; show all work, set-ups, units, etc.
Use additional paper if necessary.
1. 30.0 mL of 0.200 M HCl is exactly neutralized by
50.0 mL of Ca(OH)2 solution.
Calculate the concentration of the Ca(OH)2
solution.
2. 30.0 mL of 0.200 M H2SO2
is exactly neutralized by 50.0 mL NaOH solution. Calculate the NaOH
concentration.
3. How many mL of 0.400 M HCl are required to react with
10.0 g of solid calcium hydroxide?
4. 32.50 mL of NaOH solution are required to exactly neutralize
1.50 g of a solid acid having a molar mass of 180. g/mol. The acid
has two ionizable hydrogen. Find the concentration of the NaOH solution.
5. 2.25 g of an unknown acid, H2X,
having two ionizable hydrogens, is exactly neutralized by 42.5 mL of 0.800
M NaOH solution. Calculate the molar mass of the acid, H2X.
6. How many grams of barium hydroxide would be needed
to completely neutralize 30.0 mL of 0.500 M phosphoric acid?
7. 40.0 mL of an acetic acid solution is exactly neutralized
by 30.0 mL of 0.250 M barium hydroxide solution. Calculate the concentration
of the acetic acid solution.
8. What mass, in grams, of calcium hydroxide would be
required to react completely with 25.0 mL of 0.300 M sulfuric acid solution?
Chapter
20: Acids and Bases
Part 5 – Notes: Hydrolysis of Salts
Hydrolysis:
Hydrolysis occurs when certain salts dissolve in water to
form solutions that have acidic or basic properties. A rule for predicting
the properties of solutions of salts is based upon the concept of strong
acids and strong bases in the Arrhenius sense. This rule applies to
salts from different combinations of strong and weak acids and bases except
for those salts formed from a weak acid and a weak base. Below is a
summary of the rules for the formation of acidic or basic solutions when
salts are dissolved in water.
Salt formed
from . . .
|
water solution
exhibits
|
Example
|
strong acid
+ strong base
|
|
|
strong acid
+ weak base
|
|
|
weak acid +
strong base
|
|
|
weak acid +
weak base
|
|
|
Recall: The strong acids commonly used are:
The strong bases commonly used are:
What does an acid generally supply during the formation of a salt?
What does a base generally supply during the formation of a salt?
Example 1:
Does hydrolysis occur when KCl is dissolved in water?
Why? If yes, what type of solution is produced?
Example 2:
Does hydrolysis occur when Al2(SO4)3
is dissolved in water? Why? If yes, what type of solution is produced?
Example 3:
Does hydrolysis occur when K2CO3
is dissolved in water? Why? If yes, what type of solution is produced?
Chapter 20: Acids and Bases
Part 5 – Notes: Hydrolysis of Salts
Supply answers to the following questions.
Predict the hydrolysis effect in a solution of each of the following salts.
For answer, write acidic, basic, or neutral. Also indicate the pH (<7,
>7, or =7) of the resulting solution and list the acid and base that react
to produce the result indicated by the problem.
A/B/N
pH
Acid + Base
1. Na3PO4
2. Na2SO4
3. K2C2O4
4. NH4Cl
5. FeCl3
6. NH4NO3
7. Ca(NO3)2
8. KC2H3O2
9. NaC2H3O2
10. K2CO2
Chapter
20: Acids and Bases
Part 6 – Notes & Assignment: Preparation of Acids, Bases, and Salts
A binary acid can be prepared by a direct combination (synthesis)
reaction. When the product is dissolved in water, the acid solution
is formed.
The two reactions that take place in the preparation of hydrogen bromide
are:
H2(g)
+ Br2(g) -->
2 HBr(g)
HBr(g) + H2O --> H3O+(aq)
+ Br-(aq)
Question 1:
A 40.0 g sample of Br2(g) reacts
with excess hydrogen. The HBr(g) is then dissolved in 1.75 L of solution.
a. What is the number
of moles of HBr produced?
b. What is the molarity
of the resulting solution?
c. What is the number
of H3O+
ions in solution?
d. What is the pH
of the solution?
A ternary acid may be prepared by adding a salt of the desired
acid to concentrated sulfuric acid. The mixture is then heated.
Sulfuric acid is used as a reactant (source of H+) because it has a
high boiling point. Since the desired acid has a relatively low boiling
point, it boils out of the reacting mixture and is collected as a gas.
The gas is then dissolved in water to form the acid solution. When
nitric acid is prepared this way, the two reactions that take place are:
NaNO3(s)
+ H2SO4(l)
--> NaHSO4(s) + HNO3(g)
HNO3(g) + H2O
--> H3O+(aq)
+ NO3-(aq)
Question 2:
A 75.0 g sample of NaNO3 is treated
with excess H2SO4,
forming HNO3(g). The HNO3
gas produced is then dissolved in 0.500 kg water?
a. What mass of
H2SO4
reacts?
b. What is the number
of moles of HNO3 produced?
c. What is the molality
of the solution?
d. What additional
information would be required to determine the molarity of the sample?
An acid anhydride is an oxide of a nonmetal that will dissolved in water
to form an acid as the only product. The reaction that takes place
when P4O10
reacts in this was is:
P4O10(s)
+ 6 H2O(l) --> 4 H3PO4(aq)
Question 3:
An 85.0 g sample of solid P4O10
is dissolved to form 1.50 L of aqueous acid solution.
a. What is the molarity
of H3PO4
in the resulting solution?
b. What mass of
water reacts with the solid P4O10?
c. What is the number
of PO4-3
ions available in the aqueous solution?
d. What is the number
of H3O+
ions available in the aqueous solution?
e. What is the pH
of the acidic solution?
Preparation of Bases
When an active metal is added to water, a solution of a hydroxide compound
is produced. When potassium hydroxide is prepared this way, the reaction
is:
2 K(s) + 2 HOH(l) --> 2 KOH(aq)
+ H2(g)
(Note: Using HOH to represent water in acid/base systems sometimes
makes equations easier to understand and balance.)
Question 4:
A 5.00 g sample of solid potassium is added to enough water to prepare
5.00 L of solution.
(Caution: Addition of an active metal to water is very hazardous
and should be done only under supervision.)
a. What volume of
H2(g), measured at STP, is produced?
b. What is the concentration,
in molarity, of KOH in the resulting solution?
c. What is the pOH
of the solution?
A basic anhydride is a metal oxide that reacts with water
to form a base as the only product. When CaO reacts in this way, the
reaction is:
CaO(s) + HOH(l) --> Ca(OH)2(aq)
Question 5:
The solubility of Ca(OH)2 in
water is 0.509 g/L at room temperature.
a. What is the minimum
mass of solid CaO that must be added to water to produce 2.50 L of a saturated
solution of Ca(OH)2 (aq)?
b. What is the concentration,
in molarity, of OH-(aq) in a saturated
solution of Ca(OH)2 (aq)?
Ionization and Dissociation
In water solutions, acids dissolve and ionize to form H3O+(aq)
ions and corresponding anions. In water solutions, bases (generally,
hydroxide compounds) dissolve and dissociate into OH-(aq)
ions and corresponding cations.
Question 6:
What is the molarity of H3O+
when 45.0 g HCl(g) is dissolved into 400. mL of solution.
Question 7:
A solution is prepared by adding 14.0 g NaOH(s) to 100. mL of 1.50 M KOH(aq).
(Assume no change in volume.) What is the total molarity of the OH-
in solution?
Chapter
20: Acids and Bases
Part 7 – Notes: Acids, Bases and Ionization Equilibria
A strong acid is one that ionizes completely (or almost
completely) in an aqueous system.
What does this mean in terms of the equilibrium? How is it shifted?
What kind of value does the Keq have for this system?
When a weak acid dissolve in water, only partial ionization
occurs.
What does this mean in terms of the equilibrium? How is it shifted?
Consider the ionization of HNO2:
The above equation
represents the state of equilibrium in a solution of the weak acid HNO2.
Although it does not totally ionize, the concentration of the hydronium ions
is sufficient enough to give the solution acidic properties.
Ionization constants:
When strong bases, such as NaOH or KOH, dissolve in water
they undergo the process of dissociation, which is nearly 100% complete.
Like strong acids, strong bases have an equilibrium constant that is so
large that its numerical value is meaningless.
In a solution of a weak base a state of equilibrium exists. For
the weak base NH4OH there is an
equilibrium represented by the equation: NH3
+ H2O ←→ NH4+
+ OH-. Since the hydroxide ion is formed by a reaction in
which dissociation occurs, for NH4OH
the equilibrium constant is known as an ionization constant, like that of
weak acids.
Ionization Constants
For the reaction of a weak acid, such as HNO2,
with water, we can derive the equilibrium constant as follows:
HNO2
+ H2O ←→ H3O+
+ NO2-
Ka or KI =
[H3O+]
[NO2-]
[HNO2]
H2O is a constant and therefore not
included in the Ka.
What does the Ka or the Kb (more generally, the KI)
indicate?
If the value of Ka or Kb is large, what does it indicate?
If the vale of Ka or Kb is small, what does it indicate?
NOTE: All values of K are positive.
Solutions of Weak Acids and Bases
General equation for the ionization of a weak acid is written:
HX + H2O
←→ H3O+(aq)
+ X-(aq)
In this system, the equilibrium is NOT completely displaced to the right;
so the dissociation is NOT complete. All equilibrium systems containing
weak acids or weak bases favor reactants instead of products. You
need to use the dissociation constant of the acid (Ka) or the dissociation
constant of the base (Kb).
Remember the following when working with problems of weak acids or weak
bases:
a. The original
concentration of the weak acid is NOT equal to the concentration of
hydronium. [H3O+]
< [weak acid]
b. The original
concentration of the weak base is NOT equal to the concentration of
hydroxide. [OH-]
< [weak base]
c. Equilibrium systems
having weak acids or weak bases always favor the REACTANTS, and not the products,
as in the case of strong acids or strong bases. So you must use Ka
or Kb in order to find [H3O+]
or [OH-].
d. You can calculate
the percentage dissociation of the acid or the base by using the following
formula:
percentage dissociation
=
[H3O+]
original concentration
of the acid
Example 1 : Find (a) the [H3O+],
(b) pH, and (c) percentage dissociation for 0.100 M acetic acid at 25oC.
Ka = 1.80 x 10-5
Example 2: A 0.500 M solution of a weak
acid, HX, is only partially ionized. The [H3O+]
was found to be 2.63 x 10-3M.
Find the dissociation constant for this acid.
Chapter 20: Acids and Bases
Part 7 – Notes: Acids, Bases and Ionization Equilibria
Answer the following questions on a separate sheet of paper. Show
all work, set-ups, units, etc.
1. Calculate the (a) pH and (b) Ka of a 0.100 M acetic
acid solution if 1.34% of the CH3COOH
molecules in this solution have ionized to form product ions.
2. Calculate (a) the pH and (b) the percent dissociation
of a 0.200 M solution of HCN. Ka = 4.90 x 10-10
3. A 0.100 M solution of acetic acid is only partially
ionized. Using a measure of pH, the [H3O+]
is determined to be 1.34x10-3 M.
What is the acid dissociation constant (Ka) of acetic acid?
4. A 0.375 M solution of a weak acid, H2X,
is only partially ionized. The [H3O+]
was found to be 4.58 x 10-4M.
Find the dissociation constant for this acid.
5. Calculate the acid dissociation constant of a weak
monoprotic acid if a 0.500 M solution of this acid gives a hydronium ion
concentration of 0.000100 M.
6. Determine the (a) [H3O+],
(b) [OH--], (c) pH, and (d) pOH
of a 1.00 x 10-4M HCN solution that
undergoes 3.5% ionization.
7. Determine the (a) [H3O+],
(b) [OH-], (c) pH, and (d) pOH of
a 1.60 x 10-3 oxalic acid solution
that undergoes 0.166% ionization. H2C2O4
+ H2O ⇔ H3O+
+ HC2O4-
8. What would be the ionization constant for acetic
acid at 25oC is a 0.0100 M solution
of this acid is ionized to the extent of 1.38%?
9. A 0.500 M solution of a weak acid, HX, has a pH of
1.75. Calculate the Ka of HX.
10. A weak acid, HX, has Ka = 5.00 x 10-6.
A certain solution of this acid has pH = 2.70. Calculate the concentration
of the acid in this solution.
11. Nitrous acid, HNO2,
has Ka of 5.1 x 10-4.
a. Calculate the pH of 0.25 M HNO2.
b. Calculate the percent ionization
of 0.25 M HNO2.
12. A 0.010 M solution of acid HX is 7.5% ionized.
Calculate the Ka and the pH of this acid.
Chapter 20: Acids and Bases
Part 8 – Notes: Buffers and Acid-Base Wrap Up Information
Buffers:
You can create a buffer by . . .
Example: If acetic acid (HAc) and sodium
acetate (NaAc) were dissolved into the same solvent you would create a buffer.
a. The conjugate base of HAc is Ac-1.
So by dissolving NaAc into water . . .
b. The sodium ions don’t play any
role here so they are left out of the chemical equation.
c. Notice this is the same as the
equilibrium equation for the ionization of acetic acid with a Ka of 1.8x10-5.
d. Given such a small Ka, not many
acetate ions (Ac-1) will form naturally.
However, by adding sodium acetate, a large number of the Ac-1
ions will be present. So, once you set up the buffer, there will be
a large number of undissociated HAc and also a large number of Ac-1
ions present.
e. Now consider adding a strong acid
to this equilibrium system, (addition of a H3O+
donor). According to Le Chatelier, what will happen to this system?
f. Now consider adding a strong base
to this equilibrium system (addition of a H+
acceptor). Anything with OH-1
ions will dissolve to react with the H3O+
ions to form H2O. According to Le Chatelier, what will happen
to this system?
Without buffers, we would not be able to maintain the proper pH of our
blood. Blood pH must remain between 7.35 and 7.45 or we might die.
There are several buffer systems in our bodies that help to maintain our
proper pH levels. One of them involves our breathing.
• Our blood contains significant amounts
of both carbonic acid (H2CO3)
and its conjugate base, bicarbonate ion (HCO3-1).
This creates the equilibrium:
• Carbonic acid also forms the equilibrium:
• By combining these two equations,
you get:
• As our bodies metabolic processes
produce acids, equilibrium is shifted from right to left and we exhale the
CO2, maintaining proper levels of all chemicals.
• You can raise your blood pH by hyperventilating
which removes excessive amounts of carbon dioxide causing the entire system
of equilibrium equations to shift left removing hydronium ions from the system.
If you hyperventilate for too long you could enter alkalosis.
What are some common acids and bases you have encountered in and around
your home?
Baking soda:
Drain cleaner and
oven cleaner:
Vinegar:
Citrus fruits:
Battery acid:
Antacids like
Milk of magnesia:
Household ammonia
(aqueous ammonia):
What are some problems associated with “acid rain”?
Why do some shampoo bottles say “acid-balanced”?
What do antacids do? Are there any problems with taking too much?
What is the difference between acid strength and acid concentration?
Is it true that a neutral solution has a pH of 7?
How does endpoint and neutralization point differ?
MISCONCEPTION: For sulfuric acid, to say that
dissociation is complete means that the acid has lost both protons.
To say that the dissociation of sulfuric acid is complete does not mean
that the species in solution are two hydronium ions and one sulfate ion
for each molecule dissolved. The fact is, sulfuric acid is “strong”
only in regard to the removal of the first hydrogen ion.