Acid-Base Chemistry
An Introduction
• For thousands of years people have known that
vinegar, lemon juice and many other foods taste
sour. However, it was not until a few hundred
years ago that it was discovered why these
things taste sour - because they are all acids.
The term acid, in fact, comes from the Latin term
acere, which means "sour". While there are
many slightly different definitions of acids and
bases, in this lesson we will introduce the
fundamentals of acid/base chemistry.
• In the seventeenth century, the English writer
and amateur chemist Robert Boyle first labeled
substances as either acids or bases (he called
bases alkalies) according to the following
characteristics:
• Acids taste sour, are corrosive to metals, change litmus
(a dye extracted from lichens) red, and become less
acidic when mixed with bases.
• Bases feel slippery, change litmus blue, and become
less basic when mixed with acids.
• While Boyle and others tried to explain why acids and bases behave
the way they do, the first reasonable definition of acids and bases
would not be proposed until 200 years later.
• In the late 1800s, the Swedish scientist Svante Arrhenius proposed
that water can dissolve many compounds by separating them into
their individual ions. Arrhenius suggested that acids are compounds
that contain hydrogen and can dissolve in water to release hydrogen
ions into solution. For example, hydrochloric acid (HCl) dissolves in
water as follows:
• HCl
H2O
H+(aq) + Cl-(aq)
• Arrhenius defined bases as substances that dissolve in
water to release hydroxide ions (OH-) into solution. For
example, a typical base according to the Arrhenius
definition is sodium hydroxide (NaOH):
H2O
• NaOH
Na+(aq) + OH-(aq)
• The Arrhenius definition of acids and bases explains a
number of things. Arrhenius's theory explains why all
acids have similar properties to each other (and,
conversely, why all bases are similar): because all acids
release H+ into solution (and all bases release OH-). The
Arrhenius definition also explains Boyle's observation
that acids and bases counteract each other. This idea,
that a base can make an acid weaker, and vice versa, is
called neutralization.
Neutralization: As you can see from the equations, acids
release H+ into solution and bases release OH-. If we
were to mix an acid and base together, the H+ ion would
combine with the OH- ion to make the molecule H2O, or
plain water:
H+ (aq) + OH-(aq)
H2O
• The neutralization reaction of an acid with a base will
always produce water and a salt, as shown below:
Acid
Base
Water
Salt
HCl
+
NaOH
H 2O
+
NaCl
HBr
+
KOH
H 2O
+
KBr
Definitions of acids and bases
• Arrhenius
acid: generates [H+] in solution
base: generates [OH-] in solution
normal Arrhenius equation: acid + base <---> salt + water
example: HCl + NaOH <---> NaCl + H2O
• Bronsted-Lowery:
acid: anything that donates a [H+] (proton donor)
base: anything that accepts a [H+] (proton acceptor)
normal Bronsted-Lowery equation: acid + base <---> acid +
base
example: HNO2 + H2O <---> NO2- + H3O+
Each acid has a conjugate base and each base has a
conjugate acid. These conjugate pairs only differ by a
proton. In this example: HNO2 is the acid, H2O is the base,
NO2- is the conj. base, and H3O+ is the conj. acid.
• Lewis:
acid: accepts an electron pair
base: donates an electron pair
The advantage of this theory is that many more reactions
can be considered acid-base reactions because they do not
have to occur in solution.
Salts
A salt is formed when an acid and a base are mixed and the acid
releases H+ ions while the base releases OH- ions. This process is
called hydrolysis. The pH of the salt depends on the strengths of the
original acids and bases:
Acid
Base
Salt pH
strong strong
pH = 7
weak
pH > 7
strong
strong weak
pH < 7
weak
depends on which is stronger
weak
This is because the conjugate base of a strong acid is very weak and
cannot undergo hydrolysis. Similarly, the conjugate acid of a strong
base is very weak and likewise does not undergo hydrolysis.
Water
• We typically talk about acid-base reactions in aqueousphase environments -- that is, in the presence of water.
The most fundamental acid-base reaction is the
dissociation of water:
• In this reaction, water breaks apart to form a hydrogen
ion (H+) and a hydroxyl ion (OH-). In pure water, we can
define a special equilibrium constant (Kw) as follows:
• Where Kw is the equilibrium constant for water (unitless)
[H+] is the molar concentration of hydrogen
[OH- is the molar concentration of hydroxide
• An equilibrium constant less than one (1) suggests that
the reaction prefers to stay on the side of the reactants -in this case, water likes to stay as water. Because water
hardly ionizes, it is a very poor conductor of electricity.
pH
Under the Brønsted-Lowry definition, both acids and bases are related to the
concentration of hydrogen ions present. Acids increase the concentration
of hydrogen ions, while bases decrease the concentration of hydrogen
ions (by accepting them). The acidity or basicity of something therefore
can be measured by its hydrogen ion concentration.
In 1909, the Danish biochemist Sören Sörensen invented the pH scale for
measuring acidity. The pH scale is described by the formula:
pH = -log [H+]
. When measuring pH, [H+] is in units of moles of H+ per liter of solution.
For example, a solution with [H+] = 1 x 10-7 moles/liter has a pH equal to 7 (a
simpler way to think about pH is that it equals the exponent on the H+
concentration, ignoring the minus sign). The pH scale ranges from 0 to
14. Substances with a pH between 0 and less than 7 are acids (pH and
[H+] are inversely related - lower pH means higher [H+] and a stronger
acid). Substances with a pH greater than 7 and up to 14 are bases (the
higher the pH, the stronger the base). Right in the middle, at pH = 7, are
neutral substances, for example, pure water.
• pOH gives us another way to measure the acidity of a solution. It is just
the opposite of pH. A high pOH means the solution is acidic while a low
pOH means the solution is basic.
pOH = -log[OH-]
pH + pOH = 14.00
Strong Acids: These acids completely ionize in solution so they are always
represented in chemical equations in their ionized form. There are only
seven (7) strong acids:
HCl, HBr, HI, H2SO4, HNO3, HClO3, HClO4
To calculate a pH value, it is easiest to follow the standard "Start, Change,
Equilibrium“
Example Problem: Determine the pH of a 0.25 M solution of HBr.
Weak Acids: These are the most common type of acids. They follow the
equation:
HA(aq) <---> H+(aq) + A-(aq)
The equilibrium constant for the dissociation of an acid is known as Ka. The
larger the value of Ka, the stronger the acid.
Example Problem: Determine the pH of .30 M acetic acid (HC2H3O2) with
the Ka of 1.8x10-5.
Strong Bases: Like strong acids, these bases completely ionize in
solution and are always represented in their ionized form in
chemical equations. There are only seven (7) strong bases:
LiOH, NaOH, KOH, RbOH, Ca(OH)2, Sr(OH)2, Ba(OH)2
• Example Problem: Determine the pH of a 0.010 M solution of
Ba(OH)2.
Weak Bases: They follow the equation:
Weak Base + H2O <---> conjugate acid + OHexample:
NH3 + H2O <---> NH4+ + OH+
Kb is the base-dissociation constant:
Ka x Kb = Kw = 1.00x10-14
Acid-Base Titrations
An acid-base titration is when you add a base to an acid until the
equivalence point is reached which is where the moles of acid
equals the moles of base. For the titration of a strong base and a
strong acid, this equivalence point is reached when the pH of the
solution is seven (7) as seen on the following titration curve:
For the titration of a strong base with a weak acid, the equivalence
point is reached when the pH is greater than seven (7). The half
equivalence point is when half of the total amount of base needed to
neutralize the acid has been added. It is at this point where the pH =
pKa of the weak acid.
Henderson-Hasselbalch Equation
In an acid-base titration, the base will react with the weak acid and form a solution
that contains the weak acid and its conjugate base until the acid is completely
gone. To solve these types of problems, we will use the weak acid's Ka value and
the molarities in a similar way as we have before. Before demonstrating this way,
let us first examine a short cut, called the Henderson-Hasselbalch Equation.
This can only be used when you have some acid and some conjugate base in
your solution. If you only have acid, then you must do a pure Ka problem and if
you only have base (like when the titration is complete) then you must do a Kb
problem.
Where:
pH is the log of the molar concentration of the hydrogen
pKa is the equilibrium dissociation constant for an acid
[base] is the molar concentration of a basic solution
[acid] is the molar concentration of an acidic solution
This equation is used frequently when trying to find the pH of buffer solutions.
• Buffers:
•
A buffer is a compound that limits the
change in hydrogen ion concentration (and
so pH) when hydrogen ions are added or
removed from the solution. It may be useful
to think of the buffer as being like a sponge.
When hydrogen ions are in excess, the
sponge mops up the extra ions. When in
short supply the sponge can be squeezed
out to release more hydrogen ions!
•
All buffers are in fact weak acids or bases.
Figure 3 shows how as hydrogen ions are
added to a buffer solution they combine
with A- (the conjugate base) and the
reaction is pushed to the left. This creates
more HA whilst removing the excess H+
from the solution. Similarly, as hydrogen
ions are removed from solution by addition
of a strong base the reaction moves to the
right
restoring
the
hydrogen
ion
concentration and reducing the quantity of
HA.
• The effects of buffers can also be illustrated graphically. If a strong acid is
added slowly to a buffer solution and the hydrogen ion concentration [H+]
is measured then a plot similar to the one in figure 4 will be generated.
Notice that during the highlighted portion of the curve a large volume of
acid is added with little change in [H+] or pH.
• As we shall see later buffers are crucial in maintaining hydrogen ions
within a narrow range concentrations in the body.
NORMAL pH
• There is a normal pH value in each body
compartment (i.e. extracellular fluid,
plasma, intracellular fluid etc). Intracellular
pH is difficult to measure and may vary in
different types of cells and in different
parts of cells.
• pH of the plasma (i.e. pH of the plasma of
whole blood = conventional "blood" pH) is
controlled at 7.4 (7.35 - 7.45).
• There are three mechanisms which
diminish pH changes in body fluid: buffers;
respiratory; renal.
THE BUFFER SYSTEMS OF THE BODY
• (a) Proteins are the most important buffers in the body.
They are mainly intracellular and include haemoglobin.
The plasma proteins are buffers but the absolute amount
is small compared to intracellular protein. Protein
molecules possess basic and acidic groups which act as
H+ acceptors or donors respectively if H+ is added or
removed.
• (b) Phosphate buffer (H2PO4- : HPO42-) is mainly
intracellular. The pK of this sytem is 6.8 so that it is
moderately efficient at physiological pH's. The
concentration of phosphate is low in the extracellular
fluid but the phosphate buffer system is an important
urinary buffer.
• (c) H2CO2 : HCO3- is not an important true buffer system
because normal blood pH (7.4) is so far from its pK (6.1).
H2CO3 and HCO3- are involved in pH control but they are
not acting as a buffer system
Examination of "Buffering" Properties of HCO3-:H2CO2 System
Most texts state that the HCO3- : H2CO2 system is an efficient
physiological buffer because the components of the pair are
controlled separately. As it is not a chemical buffer of any
reasonable efficiency at the blood pH use of the term "buffer" in
respect to HCO3- : H2CO2 action introduces considerable confusion.
This is illustrated in the following example.
Plasma has a [HCO3-] of approximately 24meq/l and [H2CO2] of
1.2meq/l, hence: