Senior High School
General Chemistry 2
Quarter 4 - Module 2
Acid-Base Equilibria and Buffer Solutions
General Chemistry 2
Alternative Delivery Mode
Quarter 4 - Module 2: Acid-Base Equilibria and Buffer Solutions
First Edition, 2020
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Author:
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Senior igh School
General
Chemistry 2
Quarter 4 - Module 2
Acid-Base Equilibria and Buffer Solutions
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Table of Contents
What This Module is About ............................................................................................... i
What I Need to Know ........................................................................................................ i
How to Learn from this Module..........................................................................................ii
Icons of this Module ..........................................................................................................ii
What I Know.....................................................................................................................iii
Lesson 1:
Acid-Base Equilibria ................................................................. 1
What I Need to Know ................................................................................ 1
What’s New: Crossword! ........................................................................... 1
What Is It ................................................................................................... 2
What’s More:Let’s Do This! ....................................................................... 4
What I Have Learned:Identify Me! ............................................................. 4
What I Can Do: Explain Me! ...................................................................... 4
Lesson 2:
Buffer Solutions ........................................................................ 5
What I Need to Know ................................................................................ 5
What’s New: Match Me!................................................................................... 5
What Is It ................................................................................................... 6
What’s More: Let’s Do This!....................................................................... 8
What I Have Learned: True or False!......................................................... 8
What I Can Do: Explain Me! ...................................................................... 8
Summary ..........................................................................................................................9
Assessment: (Post-Test) ................................................................................................... 10
References ....................................................................................................................... 14
What This Module is About
The following are the lessons contained in this module:
1. Acid-Base Equilibria
2. Buffer Solutions
What is to be discussed in this module is the continuation of the study of acid-base
reactions with a discussion of buffer action. We will also look at another type of
aqueous equilibrium – that between slightly soluble compounds and their ions in
solution.
You are expected to answer the activities given in each lesson. You may write your
answers on the answer sheets provided. Remember to strictly follow the instructions.
If you have any questions and clarifications about the lessons, feel free to contact me
via cellphone number 09264702108 or via email; anncoleen.hallazgo@deped.gov.ph.
What I Need to Know
At the end of this module, you should be able to:
1. Define Bronsted acids and bases (STEM_GC11ABIVf-g-153);
2. Discuss the acid-base property of water (STEM_GC11ABIVf-g-154);
3. Calculate pH from the concentration of hydrogen ion or hydroxide ions in aqueous
solutions (STEM_GC11ABIVf-g-156);
4. Describe how a buffer solution maintains its pH (STEM_GC11ABIVf-g-160);
5. Calculate the ph of a buffer solution using the Henderson Hasselbalch equation
(STEM_GC11ABIVf-g-161).
i
How to Learn from this Module
To achieve the objectives cited above, you are to do the following:
•
Take your time reading the lessons carefully.
•
Follow the directions and/or instructions in the activities and exercises diligently.
•
Answer all the given tests and exercises.
Icons of this Module
What I Need to
Know
This part contains learning objectives that
are set for you to learn as you go along the
module.
What I know
This is an assessment as to your level of
knowledge to the subject matter at hand,
meant specifically to gauge prior related
knowledge
This part connects previous lesson with that
of the current one.
What’s In
What’s New
An introduction of the new lesson through
various activities, before it will be presented
to you
What is It
These are discussions of the activities as a
way to deepen your discovery and understanding of the concept.
What’s More
These are follow-up activities that are intended for you to practice further in order to
master the competencies.
What I Have
Learned
Activities designed to process what you
have learned from the lesson
What I can do
These are tasks that are designed to showcase your skills and knowledge gained, and
applied into real-life concerns and situations.
ii
What I Know
Multiple Choice. Encircle the letter of the best answer from among the given choices.
1. A Brønsted-Lowry acid is any species capable of:
A. Accepting a proton
B. Donating a proton
C. Exchanging a proton
D. All of the above
2. A Brønsted-Lowry base is any species capable of:
A. Accepting a proton
B. Donating a proton
C. Exchanging a proton
D. All of the above
3. All the given are properties of an acid, except:
A. Taste sour
B. pH < 7
C. pH > 7
D. Changes the color of litmus from blue to red
4. All the given are properties of a base, except:
A. pH < 7
B. pH > 7
C. Tastes bitter
D. Feels slippery
5. Depending on the circumstances, water can act as either an acid or base. The term
for this characteristic of water is called:
A. Amphiprotic
B. Amphotoric
C. Amphoteric
D. Amphotic
6. What is a solution of weak acid and conjugate base or weak base and conjugate acid
used to resist pH change with added solute?
A. Baffer
B. Biffer
C. Buffer
D. Beffer
7. The species created when a base accepts a proton.
A. Acid
B. Conjugate acid
C. Base
D. Conjugate base
8. The species left over after an acid donates a proton.
A. Acid
B. Conjugate acid
C. Base
D. Conjugate base
9. What happens to the equilibrium when some strong acid (more H +) is added to an
equilibrium mixture of the weak acid and its conjugate base?
A. The equilibrium is shifted to the left.
B. The equilibrium is shifted to the right.
C. The equilibrium is shifted upwards.
D. The equilibrium is shifted downwards.
10. The equation that connects the measurable value of the pH of a solution with the
theoretical value pKa.
A. Arrhenius equation
B. Acid equation
C. Henderson–Hasselbalch equation
D. Brønsted-Lowry equation
iii
Acid-Base Equilibria
What I Need to Know
This module discusses about the acid-base equilibrium and its applications to
the pH of solutions and the use of buffer solutions.
After going through this module, you are expected to
1. Define Bronsted acids and bases (STEM_GC11ABIVf-g-153);
2. Discuss the acid-base property of water (STEM_GC11ABIVf-g-154);
3. Calculate pH from the concentration of hydrogen ion or hydroxide ions in aqueous
solutions (STEM_GC11ABIVf-g-156);
What’s New
Activity 4.1.1. Crossword: Use the definition as a clue to the word that goes
each of the corresponding blank spaces.
5
6
1
2
8
3
10
9
4
ACROSS
2. The hydrated hydrogen ion
6. A measure of the overall concentration of hydrogen ions in solution
4. The species left over after an acid donates a proton.
9. The species created when a base accepts a proton.
DOWN
1. Potential of Hydrogen ions
3. Referring to the process by which a compound breaks into its constituent ions in solution.
5. The state of a reaction in which the rates of the forward and reverse reactions are equal.
6. Having the characteristics of both an acid and a base
8. Any process that leads to the dissociation of a neutral atom or molecule into charged
particles (ions).
10. A measure of the overall concentration of hydroxide ions in solution
4
What Is It
In the previous lessons we learned about the spontaneous change, entropy, and free
energy. Also, we learned about the chemical equilibrium and Le Chantelier’s Principle. In this
lesson, we will learn all about acids and bases.
A Brønsted-Lowry acid is any species capable of donating a proton; a Brønsted-Lowry
base is any species capable of accepting a proton.
Originally, acids and bases were defined by Svante Arrhenius. His original definition
stated that acids were compounds that increased the concentration of hydrogen ions (H +) in
solution, whereas bases were compounds that increased the concentration of hydroxide ions
(OH–) in solutions. Problems arise with this conceptualization because Arrhenius’s definition
is limited to aqueous solutions, referring to the solvation of aqueous ions, and is therefore not
inclusive of acids dissolved in organic solvents. To solve this problem, Johannes Nicolaus
Brønsted and Thomas Martin Lowry, in 1923, both independently proposed an alternative
definition of acids and bases.
In this newer system, Brønsted-Lowry acids were defined as any molecule or ion that
can donate a hydrogen cation (proton, H+), whereas a Brønsted-Lowry base is a species with
the ability to gain, or accept, a hydrogen cation. A wide range of compounds can be classified
in the Brønsted-Lowry framework: mineral acids and derivatives such as sulfonates, carboxylic
acids, amines, carbon acids, and many more.
Brønsted-Lowry Acids/Bases Reaction
Acids and bases will neutralize one another to form liquid water and a salt. The general
properties of acids and bases are:
✓ Acid – Tastes sour, acts corrosive, changes the color of litmus from blue to red, pH <7
e.g vinegar, lemon juice, gastric juice, soft drinks
✓ Base – Tastes bitter, feels slippery, changes the color of litmus from red to blue. pH >
7 e.g milk of magnesia, bleach, ammonia, detergents
We should keep in mind that acids and bases must always react in pairs. This is
because if a compound is to behave as an acid, donating its proton, then there must
necessarily be a base present to accept that proton. The general scheme for a BrønstedLowry acid/base reaction can be visualized in the form:
𝑎𝑐𝑖𝑑 + 𝑏𝑎𝑠𝑒 ⇄ 𝑐𝑜𝑛𝑗𝑢𝑔𝑎𝑡𝑒 𝑏𝑎𝑠𝑒 + 𝑐𝑜𝑛𝑗𝑢𝑔𝑎𝑡𝑒 𝑎𝑐𝑖𝑑
Here, a conjugate base is the species that is left over after the Brønsted acid donates
its proton. The conjugate acid is the species that is formed when the Brønsted base accepts
a proton from the Brønsted acid. Therefore, according to the Brønsted-Lowry definition, an
acid-base reaction is one in which a conjugate base and a conjugate acid are formed (note
how this is different from the Arrhenius definition of an acid-base reaction, which is limited to
the reaction of H+ with OH– to produce water). Lastly, note that the reaction can proceed in
either the forward or the backward direction; in each case, the acid donates a proton to the
base.
Consider the reaction between acetic acid and water:
𝐻3𝐶𝐶𝑂𝑂𝐻(𝑎𝑞) + 𝐻2𝑂(𝑙) ⇄ 𝐻3𝐶𝐶𝑂𝑂− + 𝐻3𝑂+
(𝑎𝑞)
(𝑎𝑞)
Here, acetic acid acts as a Brønsted-Lowry acid, donating a proton to water, which
acts as the Brønsted-Lowry base. The products include the acetate ion, which is the conjugate
base formed in the reaction, as well as hydronium ion, which is the conjugate acid formed.
Note that water is amphoteric; depending on the circumstances, it can act as either an
acid or a base, either donating or accepting a proton. For instance, in the presence of
ammonia, water will donate a proton and act as a Brønsted-Lowry acid:
5
Here, ammonia is the Brønsted-Lowry base. The conjugate acid formed in the reaction
is the ammonium ion, and the conjugate base formed is hydroxide.
The Acid-Base Properties of Water
Water is capable of acting as either an acid or a base and can undergo self-ionization.
It acts as an acid in reactions with bases, whereas it acts as a base for reactions involving
acids.
Under standard conditions, water will self-ionize to a very small extent. The selfionization of water refers to the reaction in which a water molecule donates one of its protons
to a neighboring water molecule, either in pure water or in aqueous solution. The result is the
formation of a hydroxide ion (OH–) and a hydronium ion (H3O+). The reaction can be written
as follows:
𝐻2𝑂 + 𝐻2𝑂 ⇌ 𝐻3𝑂+ + 𝑂𝐻−
This is an example of autoprotolysis (meaning “self-protonating”) and it exemplifies the
amphoteric nature of water (ability to act as both an acid and a base). The self-ionization of
water produces hydronium and hydroxide ions in solution.
The Water Ionization Constant, Kw
Note that the self-ionization of water is an equilibrium reaction:
𝐻2𝑂 + 𝐻2𝑂 ⇌ 𝐻3𝑂+ + 𝑂𝐻−𝐾𝑤 = 1.0 𝑥 10−14
Like all equilibrium reactions, this reaction has an equilibrium constant. Because this
is a special equilibrium constant, specific to the self-ionization of water, it is denoted KW; it has
a value of 1.0 x 10−14. If we write out the actual equilibrium expression for K W, we get the
following:
𝐾𝑤 = [𝐻+][𝑂𝐻−] = 1.0 𝑥 10−14
However, because H+ and OH– are formed in a 1:1 molar ratio, we have:
[𝐻+] = [𝑂𝐻−] = √1.0 𝑥 10 −14 = 1.0 𝑥 10−7𝑀
Now, note the definition of pH and pOH:
𝑝𝐻 = −𝑙𝑜𝑔[𝐻+]
𝑝𝑂𝐻 = −𝑙𝑜𝑔[𝑂𝐻−]
[𝑂𝐻−] = 10−𝑝𝑂𝐻
[𝐻3𝑂+] = 10−𝑝𝐻
The term pH refers to the "potential of hydrogen ion." It was proposed by Danish
biochemist Soren Sorensen in 1909 so that there could be a more convenient way to describe
hydronium and hydroxide ion concentrations in aqueous solutions since both concentrations
tend to be extremely small.
6
If we plug in the above value into our equation for pH, we find that:
𝑝𝐻 = − log(1.0 𝑥 107) = 7.0
𝑝𝑂𝐻 = − log(1.0 𝑥 107) = 7.0
Here we have the reason why neutral water has a pH of 7.0; it represents the condition
at which the concentrations of H+ and OH– are exactly equal in solution.
Since how the equilibrium constant can be expressed was already established, if we
take the negative logarithm of both sides of this equation, we get the following:
−log (𝐾𝑤) = −log ([𝐻+][𝑂𝐻−])
−log (𝐾𝑤) = − log[𝐻+] + −𝑙𝑜𝑔[𝑂𝐻−]
𝑝𝐾𝑤 = 𝑝𝐻 + 𝑝𝑂𝐻
However, because we know that pKW= 14, we can establish the following relationship:
𝑝𝐻 + 𝑝𝑂𝐻 = 14
This relationship always holds true for any aqueous solution, regardless of its level of
acidity or alkalinity. Utilizing this equation is a convenient way to quickly determine pOH from
pH and vice versa, as well as to determine hydroxide concentration given hydrogen
concentration, or vice versa.
What’s More
Activity 4.1.2. Let’s do this! Answer the following problems. Write your full
solution on your answer sheet and box the final answer.
1. What is the pH of stomach acid, a solution of HCl with a hydronium ion concentration of
1.2 × 10−3 M?
2. Air-saturated water has a hydronium ion concentration caused by the dissolved CO2 of 2.0
× 10−6M, about 20-times larger than that of pure water. Calculate the pH of the solution at
25 °C.
3. Calculate the hydronium ion concentration of a solution with a pH of −1.07.
4. What are the pOH and the pH of a 0.0125M solution of potassium hydroxide, KOH?
5. The hydronium ion concentration of vinegar is approximately 4 × 10−3M. What are the
corresponding values of pOH and pH?
What I Have Learned
Activity 4.1.3. Identify me! Complete the table by identifying whether the
sample solution mentioned is either acidic or basic and rank it all from the most acidic to
most basic (1-5), with 5 as the most basic.
Sample Solution
Acidic/Basic
1. 1 M HCl
2. Ocean Water
3. 1 M NaOH
4. Rainwater
5. Baking Soda
7
Rank
What I Can Do
Activity 4.1.4 Explain Me! Explain the chemistry behind the difference of pH
between normal rainwater and acid rain.
8
Buffer Solutions
What I Need to Know
This module discusses about buffer solutions and Henderson Hasselbalch
Equation.
After going through this module, you are expected to
1. Describe how a buffer solution maintains its pH (STEM_GC11ABIVf-g-160);
2. Calculate the ph of a buffer solution using the Henderson Hasselbalch equation
(STEM_GC11ABIVf-g-161).
What’s New
Activity 4.2.1. Match Me! Choose the answer that best matches the definitions below. Write
the letter of your answer on the column entitled “Match”.
Definition
1
2
3
4
5
Match
Consisting mostly of water.
Key Terms
A. Equilibrium
A quantitative measure of the
strength of an acid in solution; a weak
acid has a pKa value in the
approximate range -2 to 12 in water
and a strong acid has a pKa value of
less than about -2.
A weak acid or base used to maintain
the acidity (pH) of a solution near a
chosen value and which prevents a
rapid change in pH when acids or
bases are added to the solution.
Quantitative measure of the strength
of an acid in solution; typically written
as a ratio of the equilibrium
concentrations of products to
reactants.
The state of a reaction in which the
rates of the forward (reactant to
product) and reverse (product to
reactant) reactions are the same.
B. Buffer
C. Acid dissociation
constant
D. pKa
E. Aqueous
F. pKb
What Is It
9
In the previous lesson, we discover all about Brønsted-Lowry Acids/Bases Reaction.
In this lesson, buffer solutions will be discussed.
A buffer is a solution of weak acid and conjugate base or weak base and conjugate
acid used to resist pH change with added solute. A buffer’s pH changes very little when a
small amount of strong acid or base is added to it. It is used to prevent any change in the pH
of a solution, regardless of solute. Buffer solutions are used as a means of keeping pH at a
nearly constant value in a wide variety of chemical applications. Blood in the human body is a
buffer solution.
Buffer solutions are resistant to pH change because of the presence of an equilibrium
between the acid (HA) and its conjugate base (A–). The balanced equation for this reaction is:
𝐻𝐴 ⇌ 𝐻+ + 𝐴−
When some strong acid (more H+) is added to an equilibrium mixture of the weak acid
and its conjugate base, the equilibrium is shifted to the left, in accordance with Le Chatelier’s
principle. This causes the hydrogen ion (H+) concentration to increase by less than the amount
expected for the quantity of strong acid added. Similarly, if a strong base is added to the
mixture, the hydrogen ion concentration decreases by less than the amount expected for the
quantity of base added. This is because the reaction shifts to the right to accommodate for the
loss of H+ in the reaction with the base. To visualize this, below shows an example diagram.
Buffer solutions are necessary in a wide range of applications. In biology, they are
necessary for keeping the correct pH for proteins to work; if the pH moves outside of a narrow
range, the proteins stop working and can fall apart. A buffer of carbonic acid (H 2CO3) and
bicarbonate (HCO −) is needed in blood plasma to maintain a pH between 7.35 and 7.45.
Industrially, buffer solutions are used in fermentation processes and in setting the correct
conditions for dyes used in coloring fabrics.
3
Preparing a Buffer Solution
There are a couple of ways to prepare a buffer solution of a specific pH. In the first
method, prepare a solution with an acid and its conjugate base by dissolving the acid form of
10
the buffer in about 60% of the volume of water required to obtain the final solution volume.
Then, measure the pH of the solution using a pH probe. The pH can be adjusted up to the
desired value using a strong base like NaOH. If the buffer is made with a base and its
conjugate acid, the pH can be adjusted using a strong acid like HCl. Once the pH is correct,
dilute the solution to the final desired volume.
Alternatively, you can prepare solutions of both the acid form and base form of the
solution. Both solutions must contain the same buffer concentration as the concentration of
the buffer in the final solution. To get the final buffer, add one solution to the other while
monitoring the pH.
In a third method, you can determine the exact amount of acid and conjugate base
needed to make a buffer of a certain pH, using the Henderson-Hasselbach equation:
[𝐴−]
𝑝𝐻 = 𝑝𝐾𝑎 + log (
)
[𝐻𝐴]
Where, pH is the concentration of [H+], pKa is the acid dissociation constant, and [A-]
and [HA] are concentrations of the conjugate base and starting acid, respectively.
The strength of a weak acid is usually represented as an equilibrium constant. The
acid-dissociation equilibrium constant (K a), which measures the propensity of an acid to
dissociate, for the reaction is:
[𝐻+][𝐴−]
𝐾𝑎 =
[𝐻𝐴]
The greater [H+] x [A–] is than [HA], the greater the value of Ka, the more the formation
of H+ is favored, and the lower the pH of the solution.
The Henderson-Hasselbalch Equation
The Henderson–Hasselbalch equation mathematically connects the measurable pH of
a solution with the pKa (which is equal to -log Ka) of the acid. The equation is also useful for
estimating the pH of a buffer solution and finding the equilibrium pH in an acid-base reaction.
In an alternate application, the equation can be used to determine the amount of acid
and conjugate base needed to make a buffer of a certain pH. With a given pH and known pKa,
the solution of the Henderson-Hasselbalch equation gives the logarithm of a ratio which can
be solved by performing the antilogarithm of pH/pKa:
[𝑏𝑎𝑠𝑒]
10𝑝𝐻−𝑝𝐾𝑎 =
[𝑎𝑐𝑖𝑑]
An example of how to use the Henderson-Hasselbalch equation to solve for the pH of
a buffer solution is as follows:
1. What is the pH of a buffer solution consisting of 0.0350 M NH3 and 0.0500 M NH 4+
(Ka for NH4 + is 5.6 x 10-10)? The equation for the reaction is:
𝑁𝐻4+ ⇌ 𝐻+ + 𝑁𝐻3
Assuming that the change in concentrations is negligible, in order for the system to
reach equilibrium, the Henderson-Hasselbalch equation will be:
[𝑁𝐻3]
𝑝𝐻 = 𝑝𝐾𝑎 + log (
+ )
[𝑁𝐻4 ]
𝑝𝐻 = 9.25 + log (
11
0.0350
)
0.0500
𝑝𝐻 = 9.095
What’s More
Activity 4.2.2. Let’s do this! Answer the following problems. Write your full
solution on your answer sheet and box the final answer.
1. What is [H3O+] in a solution of 0.25 M CH3CO2H and 0.030 M NaCH3CO2?
CH3CO2H(aq)+H2O(l)⇌H3O+(aq)+CH3CO2−(aq)
Ka=1.8×10−5
−
2. What is [OH ] in a solution of 0.125 M CH3NH2 and 0.130 M CH3NH3Cl?
CH3NH2(aq)+H2O(l)⇌CH3NH 3+(aq)+OH−(aq)
Kb=4.4×10−4
3. What concentration of NH4NO3 is required to make [OH−] = 1.0 × 10−5 in a 0.200M solution of NH3?
What I Have Learned
Activity 4.2.3. True or False! Write T if the statement is correct. If the
statement is wrong, underline the word/s that make it wrong, then write the
correct answer.
1. If the buffer is made with a base and its conjugate acid, the pH can be
adjusted using a weak acid like HCl.
2. A buffer’s pH changes greatly when a small amount of strong acid or base
is added to it.
3. A good buffer mixture should have about equal concentrations of both of its
components.
4. The strength of a weak acid (buffer) is usually represented as an equilibrium
constant.
5.The Henderson–Hasselbalch equation connects the measurable value of
the pH of a solution with the theoretical value pKa.
12
What I Can Do
Activity 4.2.4 Explain Me! Explain the chemistry behind the buffer system in
blood.
Summary
✓ An acid is a substance that donates protons (in the Brønsted-Lowry definition) or
accepts a pair of valence electrons to form a bond (in the Lewis definition).
✓ The strength of an acid refers to its ability or tendency to lose a proton; a strong acid
is one that completely dissociates in water.
✓ The formation of conjugate acids and bases is central to the Brønsted-Lowry definition
of acids and bases. The conjugate base is the ion or molecule remaining after the acid
has lost its proton, and the conjugate acid is the species created when the base
accepts the proton.
✓ Interestingly, water is amphoteric and can act as both an acid and a base. Therefore,
it can play all four roles: conjugate acid, conjugate base, acid, and base.
✓ A Brønsted-Lowry acid-base reaction can be defined as: acid + base ⇌conjugate base
+ conjugate acid.
✓ The self- ionization of water can be expressed as: 𝐻2𝑂 + 𝐻2𝑂 ⇌ 𝐻3𝑂+ + 𝑂𝐻−
✓ The equilibrium constant for the self-ionization of water is known as Kw; it has a
value of 1.0 x 10-14.
✓ The value of Kw leads to the convenient equation relating pH with pOH: pH + pOH=14.
✓ Buffer solutions are resistant to pH change because of the presence of an equilibrium
between the acid (HA) and its conjugate base (A-).
✓ When some strong acid is added to a buffer, the equilibrium is shifted to the left, and
the hydrogen ion concentration increases by less than expected for the amount of
strong acid added.
✓ Buffer solutions are necessary in biology for keeping the correct pH for proteins to
work.
✓ Buffers can be prepared in multiple ways by creating a solution of an acid and its
conjugate base.
✓ The strength of a weak acid (buffer) is usually represented as an equilibrium constant.
✓ The acid-dissociation equilibrium constant, which measures the propensity of an acid
to dissociate, is described using the equation:𝐾𝑎 =
[𝐻+][𝐴−]
[𝐻𝐴]
✓ Using Ka and the equilibrium equation, you can solve for the concentration of [H+].
✓ The concentration of [H+] can then be used to calculate the pH of a solution, as part
of the equation: pH = -log([H+]).
✓ The Henderson-Hasselbalch equation is useful for estimating the pH of a buffer
solution and finding the equilibrium pH in an acid – base reaction.
✓ The
formula − for
the
Henderson–Hasselbalch
equation
is:
[𝐴 ]
+
+
log
(
)
,where
pH
is
the
concentration
of
[H
],
pKa
is
the
acid
𝑝𝐻 =
[𝐻𝐴]
𝑝𝐾𝑎
dissociation constant, and [A–] and [HA] are concentrations of the conjugate base and
starting acid.
✓ The equation can be used to determine the amount of acid and conjugate base needed
to make a buffer solution of a certain pH.
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Assessment: (Post-Test)
Multiple Choice. Encircle the letter of the best answer from among the given
choices.
1. Who is this person who originally defined the acids and bases and stated that acids
were compounds that increased the concentration of hydrogen ions (H +) in solution,
whereas bases were compounds that increased the concentration of hydroxide ions
(OH–) in solutions?
A. Johannes Nicolaus Brønsted
B. Svante Arrhenius
C. Thomas Martin Lowry
D. Archimedes
2. Aside from water's capability of acting as either an acid or a base, it can also undergo
.
A. Self-immunization
B. Self-protonization
C. Self-conjugation
D. Self-ionization
3. Acids and bases will neutralize one another to form:
A. Ice and a Salt
B. Liquid Water and Sugar
C. Liquid Water and Salt
D. Ice and a Sugar
4. The equilibrium constant of water has a value of:
A. 1.0 x 10−7
B. 1.0 x 10−28
C. 1.0 x 10−21
D. 1.0 x 10−14
5. The term pH refers to the:
A. Potential of hydrogen ion
B. Power of hydrogen ion
C. All of the above
D. None of the above
6. How is a buffer used?
A. It is used to prevent any change in the pH of a solution, regardless of solute
B. It is used as a means of keeping pH at a nearly constant value in a wide
variety of chemical applications.
C. All of the above
D. None of the above
7. In preparing a solution with an acid and its conjugate base, it is done by dissolving the
acid form of the buffer in about how many percent volume of water to obtain the final
solution volume required?
A. 24%
B. 50%
C. 70%
D. 60%
8.
The greater the value of Ka, the the pH of the solution.
A. Higher
B. Lower
C. Same
D. None of the above
9. pKa is equal to:
B. -log (Ka)
C. log (Ka)
D. 10-Ka
A. 10Ka
10. If a strong base is added to the mixture, the hydrogen ion concentration decreases by
less than the amount expected for the quantity of base added. This is because:
A. The reaction shifts to the right to accommodate for the loss of H+ in the
reaction with the base
B. The reaction shifts to the left to accommodate for the loss of H+ in the
reaction with the base
C. The reaction shifts upward to accommodate for the loss of H+ in the
reaction with the base
D. The reaction shifts downward to accommodate for the loss of H+ in the
reaction with the base
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References
Acids and Bases. https://courses.lumenlearning.com/boundless-chemistry/chapter/acidsand-bases/ (retrieved September 15, 2020).
Buffer Solutions. https://courses.lumenlearning.com/boundless-chemistry/chapter/buffersolutions/ (retrieved September 15, 2020).
pH and pOH. https://opentextbc.ca/chemistry/chapter/14-2-ph-and-poh/ (retrieved
September 15, 2020).
Bronsted-Lowry Acids and Bases. https://opentextbc.ca/chemistry/chapter/14-1-bronstedlowry-acids-and-bases/ (retrieved September 15, 2020).
Buffers. https://opentextbc.ca/chemistry/chapter/14-6-buffers/ (retrieved September 15,2020).
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For inquiries and feedback, please write or call:
Department of Education – Division of Cagayan de Oro City
Fr. William F. Masterson Ave., Upper Balulang, Cagayan de Oro
Telefax:
((08822)855-0048
E-mail Address:
cagayandeoro.city@deped.gov.ph
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