Download How do we control chemical processes?

UNIT 6
How do we
control
chemical
processes?
Chemical thinking can be applied to devise strategies to control the
extent and rate of different types of chemical reactions. This control can be achieved by various means such as selecting the proper
reactants, changing temperature and pressure, or inducing parallel
chemical process with shared reactants or products.
In this Unit we will apply our understanding of chemical reactions to
explore strategies to control an important type of chemical processes.
In particular, the central goal of Unit 6 is to analyze how to control
acid-base reactions in aqueous solutions. Acid-base reactions play a
critical role in many relevant systems, including our planet’s atmosphere,
hydrosphere, and lithosphere, human blood, and all cells in our bodies.
These types of reactions occur at a rather fast rate, which allows us to focus
our efforts on learning to control reaction extent.
To illustrate the core ideas of the Unit, we will explore the acid-base properties of
common pharmaceutical drugs looking to address questions such as:
350
By Nina Matthews (Own work)
[Generic 2.0]
via Flickr Commons
How do acids and bases interact with water? How do we characterize acid and base strength? How do we control the extent of acid-base reactions?
Chemical Thinking
UNIT 6 MODULES
M1. Characterizing Interactions
Recognizing the types of interactions between
main components in a chemical system.
M2. Analyzing Chemical Stability
Analyzing energetic and entropic factors that affect the stability of reactants and products.
M3. Influencing Chemical Equilibrium
Determining and controlling the effect of different factors on the extent of chemical reactions.
351
352
A great variety of chemical processes in our surroundings and inside our bodies lead to changes in chemical composition that may be beneficial or harmful
for the systems involved. To better explain, predict, and control
these chemical reactions chemists have devised strategies that
are based on the careful analysis of how different reactants
interact with each other. The nature of these interactions
depends on the composition and structure of the chemical
compounds, but there are classes of substances that tend to
react in similar ways with one another. Acids and bases are
two prototypical examples of classes of chemical compounds
that undergo similar types of transformations when interacting
Fruits contain
diverse acids
with each other. Water behaves as an acid or as a base in the presence of different substances. Given that many systems of interest contain water,
understanding and controlling acid-base reactions in aqueous environments is of
central importance in many fields, from earth and environmental sciences to biochemistry. Acid-base reactions tend to be quite fast, which facilitates their analysis
because one can focus on the thermodynamic factors that influence reaction extent without worrying about kinetic issues.
THE CHALLENGE
Drug Problem
Many pharmaceutical drugs are acids or bases that react with water, or with other acids and bases present in our bodies. These reactions
may alter the structure of drug molecules or their charge distribution.
•
Why would it be important to control the acid-base reactions that drugs
undergo inside our body ? How could these reactions be controlled?
Share and discuss your ideas with one of your classmates.
This module will help you develop the type of chemical thinking that is used
to answer questions similar to those posed in the challenge. In particular, the
central goal of Module 1 is to help you understand the nature and effect of the
interactions between acids and bases.
By kanenas.net (Own work)
[Generic 2.0] via Flickr Commons
U6: MODULE 1
Characterizing
Interactions
Chemical Thinking
U6
353
How do we control chemical processes?
Water as a Reactant
Many important chemical reactions occur in aqueous environments. When we
think about these processes, it is common to consider water as an inert solvent that
does not participate in the reactions. There are many chemical processes, however,
in which water is not only the solvent but actually reacts with the substances dissolved in it. In this Unit, we are particularly interested in exploring those reactions
that lead to the transfer of H+ ions, commonly referred to as “protons,” between
water molecules and the molecules of other chemical compounds. This proton
transfer is the result of the interaction between sites with partial negative (d–)
or partial positive (d+) charges in the water molecules, and molecular sites with
opposite charge in neighboring particles. Consider, for example, the interaction
between the water molecule and the acetic acid molecule represented in Figure
6.1. The interaction between the positive site in the OH group in the acetic acid
molecule and the negative site (oxygen atom) in the water molecule may result in
the transfer of one proton from the acid to water, resulting in the formation of a
hydronium ion (H3O+) and an acetate ion (CH3COO–):
[
By convention, the arrow representing the chemical interaction between molecules (red arrow in Figure 6.1) is drawn from the site with a partial negative charge
in one molecule to the site with a partial positive charge in the other molecule.
Notice that there are many other potential interactions between the acetic acid and
the water molecules that could result in proton transfer. For example, between the
oxygen atom in the carbonyl (C=O) group in the acetic acid molecule, which has
a negative partial charge, and a hydrogen atom in a water molecule , which has a
partial positive charge. In this case, the proton would move from the water to the
acetic acid molecule. However, this type of transfer is less likely to occur due to
unfavorable changes in the energy and entropy of the system.
In the reaction between acetic acid and water molecules, a proton is transferred from the acid to water. There are chemical compounds, however, in which
the interaction between molecules leads to proton transfer from water to the other
reactant. This type of interaction is illustrated in Figure 6.2 for the chemical reaction between methyl amine (CH3NH2) and water. The outcome of this process is a
positive methyl ammonium ion (CH3NH3+) and a negative hydroxide ion (OH–).
Notice that in this case, proton transfer results from the interaction between a
nitrogen atom (d–) in the amine and a hydrogen atom (d+) in water:
d+
[
Figure 6.2 Proton transfer
between water and a methyl
amine molecule. In this case,
a water molecule loses a H+
ion and transforms into a hydroxide ion OH–.
+
[
d–
+
[
–
[
[
–
[
d–
between an acetic acid and a
water molecule. In this case,
a water molecule gains a H+
ion and transforms into a hydronium ion H3O+.
[
d+
Figure 6.1 Proton transfer
354
MODULE 1
Characterizing Interactions
LET’S THINK
Proton Transfer
Analyze the composition and structure of the following molecules:
Lactic Acid
•
Caffeine
Alanine
Predict how each of these molecules may interact with water. Which atoms in each molecule
are most likely to be involved in proton transfer? What would the outcome of the reaction
with water be?
Share and discuss your ideas with a classmate, and clearly justify your reasoning.
Acids and Bases
Substances comprised of particles that transfer protons to water molecules are
known as Brønsted-Lowry acids (“acids” hereafter). As a result of the interaction
between acids and water, hydronium ions (H3O+) are formed. On the other hand,
substances made of particles that accept protons from water molecules are classified as Brønsted-Lowry bases (“bases” hereafter), and they lead to the formation of
hydroxide ions (OH–) in aqueous solution. In general, acids are known as proton
donors while bases are known as proton acceptors.
It is common to represent acids dissolved in water using the generic symbol
HA(aq), while bases are depicted as B(aq). A general chemical reaction between an
acid and water can then be represented in the following manner:
HA(aq) + H2O(l)
A–(aq) + H3O+(aq)
This process is commonly known as “acid dissociation.” The extent of this reaction
depends on the composition and structure of the molecules of the acid HA, on
the acid’s concentration, and on the temperature of the system. Similarly, a general
chemical reaction between a base and water can be expressed as:
B(aq) + H2O(l)
HB+(aq) + OH–(aq)
The position of this chemical equilibrium will be determined by the composition
and structure of base B, its concentration, and the temperature.
Most known acids and bases do not react with water to completion. This is,
not all the molecules of the acid donate a proton to a water molecule, and not all
molecules of a base accept a proton when dissolved in water. To understand and
control the extent to which proton transfer occurs we need to further explore the
properties of water as a reactant.
Chemical Thinking
U6
How do we control chemical processes?
355
LET’S THINK
Acids or Bases?
Many pharmaceutical drugs contain carboxylic groups (–COOH) that exhibit acidic behavior or
amine groups (–NH2, –NHR, –NR2) that exhibit basic behavior when interacting with water.
Salicylic
Acid
Ephedrine
•
Predict and represent the structure of the products of the reaction between each of these two
drugs and water.
•
Write the chemical equation between each of the drugs and water.
Share and discuss your ideas with a classmate, and clearly justify your reasoning.
Water’s Autoionization
d–
–
[
[
d+
[
This process, called the autoionization of water, is often represented in the following manner in symbolic form:
(6.1)
H2O(l) + H2O(l)
OH–(aq) + H3O+(aq)
with an associated equilibrium constant:
(6.2)
Kw = [OH–][H3O+]
Remember that the equilibrium constant expression does not include pure liquids
involved in a reaction as long as their concentrations is constant during the process. The extent to which water autoionizes is very small and only a minimal fraction of the water molecules undergo autoionization. Thus, for practical purposes
the concentration of water, [H2O], remains constant in this process.
Figure 6.3 Proton transfer
between two water molecules
results into a negative ion
(OH–) and a positive ion (H3O+)
+
[
Water molecules interact strongly with one another due to their strong bond dipoles and their overall molecular polarity. The attractive interaction between atoms with a positive partial charge (hydrogens) in one molecule and the atom with
a negative partial charge (oxygen) in another water molecule may result in proton
transfer between two molecules of water. As shown in Figure 6.3, as a result of
this transference one molecule becomes a hydroxide ion (OH–) while the other
molecule becomes a hydronium ion (H3O+):
356
MODULE 1
Characterizing Interactions
The value of the autoinization equilibrium constant depends on temperature;
experiments indicate that at 25 oC Kw = 1.0 x 10–14. In pure water, the autoionization process produces equal amounts of OH– and H3O+ ions (see Figure 6.3).
Consequently, according to Equation (6.2), we should have [OH–] = [H3O+] =
1.0 x 10–7 M in pure water at 25 oC. The molar concentration of H2O is close to
55.5 M under the same conditions, which confirms that only a very small fraction
of the water molecules autoionize (about one molecule per every 500 hundred
million molecules of water). As is the case for all chemical reactions that reach
equilibrium, autoionization is a dynamic process and particles in the system are
constantly undergoing change. Some water molecules autoionize, but simultaneously some H3O+ ions react with OH– ions to form water molecules. Overall, the
average concentration of H2O, OH–, and H3O+ remains constant.
LET’S THINK
Thermodynamic Analysis
The autoionization of water represented in Equation (6.1) is an endothermic process with DHorxn =
+55 kJ per every two moles of water involved in the reaction. The change of entropy for the same
process is DSorxn = –80.5 J/K.
•
•
Explain the signs of DHorxn and DoSrxn for the autoionization of water based on the changes
induced by the process at the molecular level.
Predict whether the value of Kw increase or decrease when the temperature of water increases.
By Datamax (Own work)
[Public domain] via Wikimedia Commons
Share and discuss your ideas with a classmate, and justify your reasoning.
Figure 6.4 The concentration
of [H3O+] ions can be measured with precision using
electronic devices.
The concentrations of OH– ions and H3O+ ions change when acids or bases
dissolve in water. This is expected because acid molecules react with water molecules to produce more H3O+ ions and bases generate more OH– ions. The addition of an acid to pure water should make [H3O+] larger than 1.0 x 10–7 M, while
the addition of a base should make [OH–] larger than this same value. But what
happens to the concentration of OH– when an acid is added to water or to the
concentration of H3O+ ions when a base is added? Consider the case of an acid in
water. When the acid is added and more H3O+ ions are formed in solution, the
likelihood of them reacting with OH– ions to produce water increases. The backwards process in the autoionization equilibrium in Equation (6.1) becomes faster
than the forward process until chemical equilibrium is reestablished. At equilibrium, the concentration of [H3O+] ions will still be larger than 1.0 x 10–7 M but
the concentration of OH– ions will be smaller than 1.0 x 10–7 because some of
these hydroxide ions had been consumed. In fact, according to Equation (6.2),
the product of concentrations [H3O+][OH–] should be equal to Kw as the concentration of [H3O+] and [OH–] ions in water is always regulated by the autoionization
process. Similarly, the concentration of [OH–] will increase and the concentration
of [H3O+] will decrease when a base is added to water.
Although the concentrations of [H3O+] and [OH–] in pure water or in aqueous solutions of acids and bases are very small, they can be measured with great
precision using instruments that analyze the electrical properties of the solution
(Figure 6.4). Small changes in the concentration of hydronium and hydroxide
ions often have a large impact on the properties of water and of the substances
dissolved in it.
Chemical Thinking
[H3O+]and [OH–] U6
How do we control chemical processes?
357
LET’S THINK
The concentration of acetic acid (CH3COOH) in common vinegar is close to 1 M. Not all the
molecules of this acid react with water molecules and produce H3O+ ions. In fact, only around
four per every 1000 molecules of acetic acid dissociate.
•
•
Estimate the concentrations of H3O+ and OH– ions in common vinegar at 25 oC.
Estimate how many times greater or smaller these concentrations are than their corresponding values in pure water at 25 oC.
Share and discuss your ideas with a classmate, and clearly justify your reasoning.
pH Scale
When working with aqueous solutions of acids and bases it is common to express
the concentration of H3O+ and OH– ions using a logarithmic scale. In particular
the pH and the pOH of the solution are defined in the following way:
(6.3)pH = – log [H3O+]
(6.4)pOH = – log[OH–]
Figure 6.5 The pH scale often
expands from 0 (highly acidic
solutions) to 14 (highly basic
solutions).
According to these definitions, the pH and pOH of pure water at 25 C are then
equal to:
o
pH = – log [H3O+] = – log (1.0 x 10–7) = 7
pOH = – log [OH–] = – log (1.0 x 10–7) = 7
[H3O+]
10–14
In general, given the value of the autoionization constant for water Kw,
as expressed in Equation (6.2), we should expect that:
in any aqueous solution at 25 oC (when Kw = 1.0 x 10–14), no matter
whether acids or bases are present in the system.
Pure water is considered to be a neutral environment from the acid-base point of view, and thus the values pH = 7 or pOH =7 are used
as references in judging the acidity or the basicity of an aqueous solution. When an acid is added to water, the concentration of H3O+ ions
becomes greater than 1.0 x 10–7 and the pH of acidic solutions is then
lower than 7 (while pOH > 7). When a base is added to water, the
concentration of H3O+ is less than 1.0 x 10–7 and the pH of basic solutions is then greater than 7 (while pOH < 7). As shown in Figure 6.5,
a great variety of products used in daily life are acidic or basic aqueous
solutions. The sour taste of many fruits is due to the presence of acids such as citric acid, while the characteristic smell of many cleaning
products is due to the presence of ammonia (a basic compound).
10–12
10–11
10–10
10–9
10–-8
10–7
10–6
10–5
10–4
10–3
10–2
10–1
100
By OpenStax College (Derivative) [Generic 3.0] via Wikimedia Commons
– log [H3O+] – log [OH–] = – log Kw
pH + pOH = 14
10–13
358
MODULE 1
Characterizing Interactions
LET’S THINK
Antiacid?
The pH of milk of magnesia, a product used as an “antiacid” in the treatment of heartburn, is close
to 10.5. The pH of the fluid in a human stomach is close to 2.0.
•
What is the concentration of H3O+ ions in a human stomach? What is the concentration of
OH– ions in milk of magnesia?
The autoionization of water is not thermodynamically favored (Kw << 1). Consequently, the reverse
process, the reaction between H3O+ and OH– to form water, is strongly favored.
•
What would you expect to happen at the molecular level when equal amounts of milk of magnesi and stomach fluid are mixed?
Estimate the final pH of this mixture.
•
Share and discuss your ideas with a classmate, and clearly justify your reasoning.
Reaction Extent
–
+
–
–
+
+
–
+
–
–
+
Figure 6.6 When a strong
acid, like HCl, is added to water, it reacts to completion
(i.e., it fully dissociates).
Not all acids or bases react to the same extent with water. A few chemical compounds react to completion and generate as many H3O+ ions or
OH– ions as particles of the substance are added to water (Figure 6.6).
These types of chemical compounds are labeled as “strong” acids or bases.
Examples of strong acids include hydrochloric acid (HCl), hydrobromic
acid (HBr), hydroiodic acid (HI), and nitric acid (HNO3). The chemical
reaction between any of these strong acids and water can be represented
in the following general manner:
HA(aq) + H2O(l)
A–(aq) + H3O+(aq)
The equilibrium constant Ka associated with this reaction is very large
(Ka >>1).
+
Chemical compounds that are strong bases tend to be ionic substances comprised of OH– ions (hydroxides) that fully dissolve in water. That is
the case of sodium hydroxide (NaOH) and potassium hydroxide (KOH).
The dissolution in water of these strong bases is often represented as shown below:
XOH(s) + H2O(l)
X+(aq) + OH–(aq)
Most chemical substances with acid-base properties, however, do not react
with water to completion. In fact, in most cases the extent of the reaction with
water is pretty small, with equilibrium constants that are typically smaller than 1
x 10–3. These types of chemical compounds are labeled as “weak” acids or bases.
Carboxylic acids (characterized by the presence of the carboxyl functional group
–COOH), such as acetic acid or citric acid, are weak acids, while amines (–NH2,
–NHR, –NR2) are weak bases, including ammonia NH3. The acid-base properties
of many substances is due to the presence of carboxyl or amine groups.
Chemical Thinking
U6
How do we control chemical processes?
359
The extent of the reaction between a weak acid or a weak base with water is
determined by the value of the equilibrium constant for the process. For weak
acids that react as described by the following chemical equation:
HA(aq) + H2O(l)
A–(aq) + H3O+(aq)
the equilibrium constant Ka, also called the acid dissociation constant, can be
expressed as:
[A–(aq)] [H3O+(aq)]
K
=
(6.5) a
[HA(aq)]
where the concentration of water is assumed to remain constant in the process.
Typical values of Ka for different weak acids are listed in the table at the bottom
of this page. In all these cases, the small value of the equilibrium constant indicates
that only a very small fraction of the molecules of each substance will dissociate to
generate H3O+ ions when reacting with water (Figure 6.7). Given the small values
of Ka for weak acids, it is common to represent these values using a logarithmic
scale. Similarly to how concentrations of H3O+ and OH– are expressed, the pKa of
a weak acid is defined as:
Figure 6.7 When a weak
acid, like HF, is added to water, only a small fraction of the
molecules dissociate.
(6.6) pKa = -log Ka
Given this definition, the smaller the value of Ka (or the weaker the acid), the
larger the value of the pKa for that substance.
The same type of procedure is applied to define and represent the equilibrium
constant Kb for the reaction of weak bases with water:
B(aq) + H2O(l)
HB+(aq) + OH–(aq)
In this case, the equilibrium constant is expressed as
[HB+(aq)] [OH–(aq)]
[B(aq)]
and the associated pKb is defined as:
(6.7)
Kb =
(6.8) pKb = -log Kb
Weak Acids
Weak Bases
Substance
Ka
pKa
Substance
Kb
pKb
HF – Hydrofluoric Acid
7.2 x 10–4
3.14
CH3NH2 – Methylamine
4.4 x 10–4
3.36
HNO2 – Nitrous Acid
4.0 x 10–4
3.39
NH3 – Ammonia
1.8 x 10–5
4.75
HCOOH – Formic Acid
1.8 x 10–4
3.75
CH3COO– – Acetate ion
5.6 x 10–10
9.25
CH3CH2OCOOH – Lactic Acid
1.4 x 10–4
3.86
C6H5NH2 – Aniline
3.8 x 10–10
9.42
C6H5COOH – Benzoic Acid
6.5 x 10–5
4.19
NO2– – Nitrite ion
2.5 x 10–11
10.61
CH3COOH – Acetic Acid
1.8 x 10–5
4.75
F– – Fluoride ion
1.4 x 10–11
10.86
360
MODULE 1
Characterizing Interactions
LET’S THINK
Comparing Strengths
Among the weak acids listed in the table on page 369, hydrofluoric acid (HF) is the strongest acid.
On the other hand, the fluoride ion (F–) is the weakest base of those included in the same table.
•
Analyze the reactions between HF and water and between F– and water, and explain the relative difference in strength between these acid-base pair.
Share and discuss your ideas with a classmate, and clearly justify your reasoning.
Conjugate Pairs
When an acid or a base react with water, protons are transferred between the reacting species. As a result of this process, new chemical species are formed in solution.
For an acid, the associated product has one less proton and it is said to be “deprotonated.” For a base, the associated product has one more proton and it is said to
be “protonated.” These new substances also have acid-base properties because they
may react with water by gaining or losing a proton. Consider the reaction between
acetic acid (CH3COOH) and water:
CH3COOH(aq) + H2O(l)
CH3COO–(aq) + H3O+(aq)
The acetate ion (CH3COO–) formed through this reaction may react with water
and reform acetic acid, behaving as a base:
CH3COO–(aq) + H2O(l)
CH3COOH(aq) + OH–(aq)
In general, the product of the deprotonation of an acid is a chemical base known
as the “conjugate base” of the original acid. The pair CH3COOH/CH3COO– is
said to be a conjugate acid-base pair. Similarly, when a base like ammonia (NH3)
reacts with water:
NH3(g) + H2O(l)
NH4+(aq) + OH–(aq)
the associate product, ammonium ions (NH4+), may react with water as an acid
to reform NH3:
NH4+(aq) + H2O(l)
NH3(aq) + H3O+(aq)
The product of the protonation of a base is a chemical acid known as the “conjugate acid” of the original base. The pair NH4+/NH3 is another example of a
conjugate acid-base pair.
The stronger an acid, the less likely that its conjugate base will react with water
and gain protons back. Protonation and deprotonation are opposite processes, and
the more one of these processes is thermodynamically favored the more the other
process is disfavored. Strong acids, such as HCl, have extremely weak bases, such
as Cl–.
Chemical Thinking
U6
How do we control chemical processes?
Water molecules also undergo changes when reacting with acids or bases.
When water reacts with an acid, it gains protons and behaves as a proton-acceptor
(or a base). The product of the reaction H3O+ may then act as a proton-donor (or
an acid) when reacting with other substances. Thus, the pair H3O+/H2O can be
thought of as a conjugate acid-base pair. Similarly, when water reacts with a base, it
loses protons and behaves as a proton-donor (or an acid). The product of the reaction OH– may thus act as a proton-aceptor (or a base) when reacting with other
substances. Consequently, the pair H2O/OH– is another example of a conjugate
acid-base pair. In general, when an acid or a base reacts with water, the following
two conjugate acid-base pairs can be identified:
Acid
HA(aq) + H2O(l)
Conjugate Base
A–(aq)
+
Base
Base
B(aq)
+ H2O(l)
H3O+(aq)
Conjugate Acid
Conjugate Acid
HB+(aq)
+
OH–(aq)
Conjugate Base
Acid
Conjugate Acid-Base Pairs
LET’S THINK
The solubility and physiological activity of many pharmaceutical drugs strongly depends
on their acid-base properties. Consider the molecular structure of the following drugs:
Ibuprofen
C13H18O2
•
•
•
Methamphetamine
C10H15N
Lidocaine
C14H22N2O
Predict whether each of these substances will behave as an acid or as a base when
reacting with water.
Determine the structure and chemical formula of the conjugate base or acid for
each of these substances.
Represent the chemical equations for the reactions of each of these substances with
water and identify all the conjugate acid-base pairs involved in each process.
Share and discuss your ideas with a classmate, and clearly justify your reasoning.
361
362
MODULE 1
Characterizing Interactions
Aqueous solutions of acids and bases are comprised of different chemical species that are in constant interaction with each other. In an acid’s solution, for
example, we have the acid HA and its conjugate base A–, we also have water H2O,
and hydronium ions H3O+ and hydroxide ions OH–. At chemical equilibrium,
which is reached very fast in these types of systems, the concentration of each of
these species will be determined by the equilibrium constants of the following
chemical reactions between components:
A–(aq) + H3O+(aq)
Ka =
[A–(aq)] [H3O+(aq)]
[HA(aq)]
HA(aq) + OH (aq)
Kb =
[HA(aq)] [OH–(aq)]
[A–(aq)]
HA(aq) + H2O(l)
A (aq) + H2O(l)
–
2 H2O(l)
–
H3O+(aq) + OH–(aq)
Kw = [H3O+(aq)] [OH–(aq)]
The analysis of these relationships indicates that they are not independent
from each other. For example, if we multiply Ka times Kb we get:
[A–(aq)] [H3O+(aq)]
[HA(aq)] [OH–(aq)]
(6.9) = Kw
Ka x Kb =
x
[HA(aq)]
[A–(aq)]
or, by using the definitions of pKa and pKb in Equations (6.6) and (6.8):
(6.10) pKa + pKb = pKw
where pKw = –log Kw = 14 at 25 oC. The relationship between pKa and pKb as
expressed in Equations (6.9) or (6.10) confirms that the stronger an acid is (larger
Ka, smaller pKa) the weaker its conjugate base would be (smaller Kb, larger pKb),
and vice versa. A weak acid or base, however, always has a weak conjugate pair.
LET’S THINK
An Anesthetic
Procaine (C13H20N2O2) is a local anesthetic used primarily
to reduce the pain from penicillin injection, and it is also
used in dentistry. For basic drugs such as procaine, it is
common to report the value of the pKa of their conjugate
acid instead of the pKb of the actual base. The pKa of the
conjugate acid of procaine has a value close to 8.
•
•
•
Predict the structure of the conjugate acid of procaine.
Calculate the Kb and pKb of procaine.
Discuss which species, procaine or its conjugate acid, is stronger in terms of their acid-base
reactivity with water.
Share and discuss you ideas with a classmate, and clearly justify your reasoning.
Chemical Thinking
FACING THE CHALLENGE
U6
How do we control chemical processes?
cally charged species will be more likely to dissolve
in water and less likely to move across nonpolar
environments. Given that drug molecules will enDrugs in Our Body
counter different chemical environments as they
Most drugs in the market are chemical com- distribute throughout our body, drug designers
pounds that behave as weak acids or bases when need to carefully analyze the acid-base properties
added to water. For example, common over-the- of the compounds that they synthesize to predict
counter- analgesics such as aspirin and ibuprofen how they will behave once ingested.
The intracellular fluid is considerably more
are weak acids, while stronger narcotic analgesic
acidic than the cellular plasma.
such as morphine and oxyThe distribution of drugs incodone are weak bases. The
side and outside cells may thus
acid-base properties of drugs
be different for acid and basic
are determined by the nature
drugs. Basic drug molecules are
of key functional groups (e.g.,
more likely to exist as ions in
carboxyl, amine) present in the
acidic medium than acid drug
molecules of these substances.
molecules are. The intracelluWhen acid or basic drugs
lar concentration of weak bases
react with water, the particles
should thus be higher than the
of the conjugate acid-base
concentration of weak acids.
pairs that are formed have
In general, basic drugs tend to
a different electrical charge
be more active than acid drugs
than the original species. For
because of these differences in
example, the conjugate bases
intracellular concentrations.
of carboxylic acids have negaBy psyberartist (Own work) [Generic 2.0] via Flickr Commons
The charge distribution in drug molecules also
tive charge (anions), while the conjugate bases of
amines have positive charge (cations). The extent affects how these molecules interact with the active
to which these charged species will be produced site of the protein receptors that a drug is designed
depends on the pKa of the drugs and on the con- to target. The presence or absence of charge in key
centration of H3O+ and OH– ions already present areas of drug molecules will affect the strength to
in the medium in which drugs are dissolved. For which these particles can bind to protein recepexample, if the pH of the medium is low (i.e., the tor sites. Low binding strength may lead to drug
concentration of H3O+ is high), like in stomach inactivity while too strong of a binding may result
fluids, it is less probable that acidic drug molecules in high toxicity.
Changes in the pH of urine induced by differwill dissociate to form more H3O+. Thus, the drug
will most likely exist in its neutral form. On the ent types of food may have a profound influence
other hand, basic drug molecules are more likely on the facility with which acid and basic drugs are
to be protonated and acquire a positive charge excreted. Weak acids are more easily excreted in
when the concentration of H3O+ is high. In gen- highly alkaline urine where their molecules will
eral, the pH of the medium has a critical effect on be mostly ionized, and weak bases in acidic urine
the proportion of ionized and unionized acid or where most molecules will be protonated. The
elimination of potentially toxic drugs may thus be
basic drug molecules present in a system.
The electrical charge of drug molecules has a accelerated by adjusting the urinary pH. Changstrong influence on the absorption, distribution, ing the pH of urine often facilitates the detection
metabolism, excretion, and toxicity of the drugs. of recreational drugs. Amphetamines, for example,
The electrical charge affects the solubility of the may be more easily detected in acidic urine where
drug in water and in cell membranes, which are molecules of the drugs are protonated and thus are
mainly comprised of nonpolar substances. Electri- more soluble in water.
363
364
MODULE 1
Characterizing Interactions
Let’s Apply
Over 70% of the drugs in the market have acid-base properties. Of these, close to 60% are weak
bases, around 25% are weak acids, and the rest exhibit both acid and base behavior (amphoteric).
Most of the basic drugs contain amine groups and the pKa of their conjugate acids is close to 9. The
strongest acid drugs tend to be carboxylic acids with pKa values close to 4, and the weakest acid
drugs tend to be barbiturates with pKa values close to 8.
Analgesics
Analgesics are pharmaceutical drugs used to relieve pain (pain killers). They include different
types of drugs, from non-steroidal anti-inflammatory drugs (NSAIDs) such as aspirin and ibuprofen to opioid drugs such as morphine and oxycodone. The molecular structures, chemical
formulas, and pKas of aspirin and morphine are shown below, where main functional groups
with acid-base properties are highlighted:
Aspirin
C7H6O3
pKa = 3.5
•
•
•
•
•
•
Morphine
C17H19NO3
pKa = 8.2
Predict whether each of these drugs will behave as an acid or as a base when reacting
with water. Make a drawing to represent how each drug molecule will interact with a
water molecule.
Write the chemical formula of the conjugate acid or base of
each of these substances, and calculate their associated pKb.
Discuss what information the values of pKa and pKb (or Ka and
Kb) convey for eah of these substances. Identify the strongest
species in each conjugate acid-base pair.
Write the chemical equation that represents the reaction of
each of these drugs with water.
Identify all the conjugate acid-base pairs involved in each of
the reactions.
Express the equilibrium constant for the reaction of aspirin with water.
Share and discuss your ideas with a classmate, and clearly justify your reasoning.
Morphine
is extracted
from
opium
poppies
By Me (Own work) [Public domain]
via Wikimedia Commons
ASSESS WHAT YOU KNOW
Drug Properties
Chemical Thinking
U6
How do we control chemical processes?
365
Drug Absorption
In the stomach, where the pH of the gastric juice is close to 2, drugs such as morphine
and aspirin mainly exist in their acid form. Their structure under these conditions influences how fast the drugs are absorbed through the stomach walls.
•
Which of these two drugs will be absorbed at a faster rate through stomach cell
membranes?
Aspirin Tablets
When a 500 mg tablet of Aspirin is dissolved in a glass of water, the pH is close to 2.8.
When two 500 mg tablet of Aspirin are dissolved in the same amount of water, the pH
is close to 2.5
•
How would you explain this result? Why does the pH decrease? Why is the change
in pH so small when the amount of aspiring is doubled?
Share and discuss your ideas with a classmate, and clearly justify your reasoning.
Other Analgesics
Other over-the-counter analgesics include ibuprofen and acetominophen (or paracetamol),
substances that behave like acids when reacting with water:
Ibuprofen
C13H18O2
pKa = 4.5
•
Acetominophen
C8H9NO2
pKa = 4.5
Compare the acid-base properties of aspirin, ibuprofen, and acetominophen.
Which is the strongest acid? Which substance has the strongest conjugate base?
Share and discuss your ideas with a classmate, and clearly justify your reasoning.
ASSESS WHAT YOU KNOW
Share and discuss your ideas with a classmate, and clearly justify your reasoning.
366
MODULE 1
Characterizing Interactions
Let’s Apply
Acidic Atmosphere
ASSESS WHAT YOU KNOW
The atmosphere of our planet contains a variety of acidic substances. Understanding the acid-base
properties of these chemical compounds is of central importance to predict and control their interactions.
Atmospheric Acids
Consider the following molecular structures of acids formed in Earth’s atmosphere:
Nitric Acid
HNO3
•
•
•
•
•
•
Carbonic Acid
H2CO3
Sulfuric Acid
H2SO4
Predict how each of these acids may interact with water. Make a drawing to represent
how each of the acid molecules will interact with a water molecule.
Draw the molecular structure of the conjugate bases of each of these acids.
Write the chemical equation that represents the reaction of each of the acids with water.
Identify all conjugate acid-base pairs in each of the processes.
Carbonic acid, H2CO3, and sulfuric acid, H2SO4, are examples of polyprotic acids,
which are substances comprised of molecules that may donate more than one proton to
water molecules. Draw the molecular structure of the bases that would form when each
of these acids donates two protons to water molecules.
Nitric acid is a strong acid (Ka >> 1). Draw a molecular representation of a solution of
nitric acid in water (no need to represent water molecules).
Carbonic acid is a weak acid. The pKa for the dissociation of a first proton is 3.6, while the pKa
for the dissociation of the second proton, represented by this chemical equation, is 6.3:
HCO3–(aq) + H2O(l)
•
CO32–(aq) + H3O+(aq)
Provide a reasonable explanation for the difference in the pKa values for the first and
second dissociations.
Share and discuss your ideas with a classmate, and clearly justify your reasoning.
Chemical Thinking
U6
How do we control chemical processes?
367
Sulfuric Acid
Sulfuric acid, H2SO4, is not only an important atmospheric component but also a widely
used chemical compound with multiple industrial applications. Each molecule of this diprotic acid may donate up to two protons to water molecules. The pKa for the first dissociation has a negative value (strong acid) while the pKa for the second dissociation is 1.9.
•
Share and discuss your ideas with a classmate, and clearly justify your reasoning.
Acid Rain
Effect of acid rain on statues
The rain in our planet tends to be acidic due to the
presence of H2CO3, HNO3, and H2SO4. Rain that
is derived from the Atlantic Ocean can have a pH
as high as 5.6; rain that comes across the continent
from the west can have a pH as low as 3.8.
•
How many times larger is the concentration
of H3O+ ions in the most acidic rain versus in
the less acidic rain?
Share and discuss your ideas with a classmate, and
clearly justify your reasoning.
By Nino Barbieri (Own work) [Generic 2.5]
via Wikimedia Commons
Ammonia Effects
Ammonia, NH3, is a gaseous basic compound (pKb = 4.7) widely used in the production
of fertilizers. This substance can react with nitric acid and sulfuric acid in the atmosphere:
HNO3(aq) + NH3(g)
HSO4(aq) + NH3(g)
•
NO3–(aq) + NH4+(aq)
HSO4–(aq) + NH4+(aq)
Discuss how these reactions may affect the pH of rainwater.
Share and discuss your ideas with a classmate, and clearly justify your reasoning.
ASSESS WHAT YOU KNOW
•
•
Write the chemical equations that represent the first and the second dissociations
of sulfuric acid in water.
Express the equilibrium constants for each of the dissociation processes.
A solution of sulfuric acid with a pH close to –0.6 is used as electrolyte in car batteries. What is the concentration of H3O+ ions in these types of solutions?