08_Lecture.ppt

Chapter 8 Lecture
General, Organic, and Biological
Chemistry: An Integrated Approach
Laura Frost, Todd Deal and Karen Timberlake
by Richard Triplett
© 2011 Pearson Education, Inc.
Acids, Bases, and Buffers
in the Body
Chapter 8

Chapter 8 2
Chapter Outline
8.1 Acids and Bases—Definitions
8.2 Strong Acids and Bases
8.3 Chemical Equilibrium
8.4 Weak Acids and Bases
8.5 pH and the pH Scale
8.6 pKa
8.7 Amino Acids: Common Biological Weak Acids
8.8 Buffers: An Important Property of Weak Acids
and Bases

Chapter 8 3
Introduction
• Citrus fruits taste sour because they contain
acid.
• Our stomach produces acid to aid in digestion,
and our muscles produce lactic acid when we
exercise.
• An acid can be neutralized by a base. Soaps are
mild bases, and, like other bases, feel slippery to
the touch.

Chapter 8 4
Introduction, Continued
• pH refers to the acidity of a solution.
• Amino acids, which are the building blocks of
proteins, will change form if the acidity of a
solution changes. Proteins change their shape
and functionality if the pH of a solution is
changed.
• Our bodily fluids contain compounds that
maintain pH. These compounds are called
buffers.

Chapter 8 5
Acids
• The Swedish chemist, Svante Arrhenius, first
described acids as substances that dissociate,
producing hydrogen ions (H+) when dissolved in
water.
• The hydrogen ions give acids their sour taste
and allow acids to corrode some metals.

Chapter 8 6
8.1 Acids and Bases—Definitions, Continued
• In the early twentieth century, Brønsted and
Lowry redefined acids as a compound that
donates a proton.
• H+ is the same as a proton since hydrogen has
lost its electron.
• A free proton rarely exists in an aqueous
solution. The proton is attracted to the partial
negative charge on the oxygen atom of water.

Chapter 8 7
The attraction is strong and the oxygen atom in water
forms a third covalent bond, giving water a positive,
ionic charge creating the hydronium ion, H3O+.

Chapter 8 8
Bases
• Arrhenius described bases as compounds that
dissociate to form a metal ion and a hydroxide
ion (OH-) when dissolved in water. Arrhenius
bases are formed from Group 1A and 2A metals.
Hydroxide bases are characterized by a bitter
taste and a slippery feel.
• The Brønsted–Lowry definition of a base is a
compound that accepts a proton.

Chapter 8 9
Sodium hydroxide, NaOH, is an example of a
metal hydroxide that dissociates into a metal ion
and a hydroxide ion.

Chapter 8 10
Acids and Bases Are Both Present in Aqueous
Solution
• The Brønsted–Lowry definitions of acids and
bases imply that a proton is transferred in an
acidic or basic solution.
• Water can act as an acid or base by donating or
accepting a proton. For example, when a
hydrochloric acid solution is prepared, water
accepts a proton, and is acting as a base.

Chapter 8 11
In another reaction, ammonia (NH3) reacts with
water, and water is acting as an acid by donating
a proton to NH3.

Chapter 8 12
• Acids and bases are classified by their ability to
donate or accept protons, respectively.
• Strong acids completely dissociate into ions when
placed in water, forming hydronium ions and
anions. Examples of strong acids are as follows:

Chapter 8 13
8.2 Strong Acids and Bases, Continued
Weak acids only partially dissociate when placed
in water. Acetic acid, the main component of
vinegar, is an example of a weak acid.

Chapter 8 14
• Sodium hydroxide, NaOH, is a strong base. It is
used in household products such as oven
cleaners and drain openers.
• NaOH and other bases like LiOH, KOH, and
Ca(OH)2 are strong bases that completely
dissociate in water. They dissociate to give a
metal ion and a hydroxide ion.
• Weak bases, like weak acids, only partially
dissociate in water.

Chapter 8 15
Neutralization
• When a strong acid and strong base are mixed,
both completely dissociate to form ions in water.
• For example, when HCl and NaOH are mixed,
sodium ions and chloride ions are formed, as
well as hydroxide ions and hydronium ions.
• The protons in the hydronium ion are attracted
to the hydroxide ion to form water molecules.

Chapter 8 16
• As a result of this chemical reaction, a lot of heat
is produced and is considered to be an
exothermic (exo means “to give off”; thermo
means “heat”) reaction.
• The sodium and chloride ions remain in solution.
• The reaction of a strong acid with a strong base
produces a salt and water.

Chapter 8 17
The formation of water and a salt is called a
neutralization reaction.

Chapter 8 18
Antacids
• Antacids are substances that neutralize excess
stomach acid. Some antacids are bases that are
not very soluble in water.
• Aluminum hydroxide and magnesium hydroxide
are examples of antacids that are not very
soluble in water. They are used in combination
to prevent unpleasant side effects.
• Calcium carbonate is an antacid that will cause
an increase in serum calcium. It is not
recommended for people with peptic ulcers or
for those that have a tendency to form kidney
stones.

Chapter 8 19
• Sodium bicarbonate is an antacid that will affect
the acidity level of blood and will elevate sodium
levels in bodily fluids.
• The following table shows some antacid
preparations.

Chapter 8 20
• After forming products, some chemical reactions
will reverse and reform reactants. These type of
reactions are reversible reactions.
• The generation of ammonia is a reversible
reaction. Once ammonia is formed, the reaction
will reverse and reform nitrogen and hydrogen.

Chapter 8 21
8.3 Chemical Equilibrium, Continued
• The rate of ammonia formation and reforming
nitrogen and hydrogen will eventually become
equal. This is called chemical equilibrium.
• An equilibrium arrow is used in this type of
reaction to indicate that both the forward and
reverse reactions take place simultaneously.
• Because the rates of forward and reverse
reactions are equal, there is no net change in
the amounts of product or reactants.

Chapter 8 22
The Equilibrium Constant, K
In this reaction, if we measure the concentrations
of ammonia, nitrogen, and hydrogen, the fraction
of products to reactants would be a constant
value. This constant value is called the
equilibrium constant, K.

Chapter 8 23
• K is defined as:
• The brackets, [ ], mean “concentration of.”
• This equation states that the equilibrium
constant is equal to the molar concentration of
the products divided by the molar concentrations
of the reactants.

Chapter 8 24
• The expression for K for the generation of
ammonia is:
• The superscripts in the expression come from
the coefficients found in the chemical equation.

Chapter 8 25
• For an equilibrium reaction of the form:
• The equilibrium expression is given as:

Chapter 8 26
• Only substances whose concentrations change
appear in the equilibrium expression.
• Substances like solid and pure liquids have
constant concentrations, so they do not appear
in the expression.

Chapter 8 27
• Values for K vary greatly depending on
temperature and reaction.
• If K has a value of 1, the ratio of
products:reactants is 1:1 or the [products] =
[reactants].
• A K value greater than (>) 1 means the amount
of products is larger than the amount of
reactants or [products] > [reactants].

Chapter 8 28
• A K value less than (<) 1 means that the amount
of reactants is larger than the amount of
products or [products] < [reactants].
• The following table summarizes the
interpretations of K values.

Chapter 8 29
Effect of Concentration on Equilibrium—
Le Châtelier’s Principle
• Reconsider the production of ammonia.
• What would happen if more nitrogen was
injected in the reaction vessel?

Chapter 8 30
• According to Le Châtelier’s principle, if stress
is applied to the equilibrium, the rates of the
forward and reverse reaction will change to
relieve the stress, and equilibrium will be
regained.
• If more N2 is added, the rate of the forward
reaction will increase, shifting the equilibrium to
produce more products.

Chapter 8 31
If one of the reactants, like H2, is removed, the
reverse reaction will increase faster than the
forward reaction, allowing H2 to be replenished.
The equilibrium will shift to the left, forming more
of the reactants.

Chapter 8 32
Effect of Temperature on Equilibrium
• What affect would the change in temperature
have on the ammonia reaction?
• The reaction is known as an exothermic
reaction (produces heat) as opposed to an
endothermic reaction (absorbs heat from the
surrounding).

Chapter 8 33
• The reaction can be written as follows to signify
that heat is produced:
• If the temperature of the reaction is raised, the
rate of the reverse reaction increases to offset
the stress of adding heat. This causes the
equilibrium to shift to the left.
• If the reaction is cooled, the rate of the forward
reaction will increase to replenish the heat
produced, shifting the equilibrium to the right.

Chapter 8 34
The effects of temperature changes on ammonia
production are shown in the following:

Chapter 8 35
Consider an endothermic reaction in which heat is
absorbed from its surroundings. Heat is a reactant,
so the opposite shifts occur.

Chapter 8 36
The effects of changes on equilibrium are
summarized here.

Chapter 8 37
Equilibrium
• The principles of equilibrium apply to weak acids
and bases because they only partially dissociate
into ions.
• The dissociation of acetic acid into acetate ions
and hydronium ions is shown as:

Chapter 8 38
8.4 Weak Acids and Bases, Continued
• The equilibrium constant expression for this
reaction would be:
• Pure liquids like water are present in large
amounts, do not change significantly, and are
not included in the equilibrium expression.

Chapter 8 39
The Equilibrium Constant, Ka
• Weak acids dissociate much less than 100%
and have an equilibrium constant called an acid
dissociation constant, Ka.
• The strength of a weak acid can be determined
from the Ka value. The larger the Ka value, the
stronger the acid (the more protons dissociated).

Chapter 8 40
This table shows Ka values for substances acting
as weak acids.

Chapter 8 41
Two common organic functional groups that act as
weak acids are the carboxylic acid group and
protonated amines.

Chapter 8 42
Conjugate Acids and Bases
• The Brønsted–Lowry theory states that the
reaction between an acid and base involves a
proton transfer. If a weak acid is mixed with
water, water will act as a base and accept a
proton from the weak acid.
• Consider the dissociation of acetic acid.

Chapter 8 43
• Acetic acid, CH3COOH, donates a proton to a
molecule of water to form the hydronium ion,
H3O+.
• After the donation of a proton, CH3COO-, a
carboxylate called acetate anion remains and is
called the conjugate base of CH3COOH.
• The hydronium ion, H3O+, is the conjugate acid
of water, which is acting as a base.

Chapter 8 44
• Molecules or ions related by the loss or gain of a
proton are referred to as conjugate acid–base
pairs.
• The functional groups carboxylic acid and
carboxylate are conjugate acid–base pairs of the
same functional group.
• Weak acids are usually designated HA and their
conjugate bases as A-.

Chapter 8 45
• A weak base, such as an amine, will accept a
proton to form a protonated amine. In this case,
water acts as an acid.
• The protonated amine and amine are conjugate
acid–base pairs of the same functional group.

Chapter 8 46
The Autoionization of Water, Kw
• Water can act as either a weak acid or a base
depending on whether a base or acid is present
in solution.
• In pure water, the water molecules
spontaneously react with each other as shown.

Chapter 8 47
8.5 pH and the pH Scale, Continued
• This reaction is called the autoionization of
water.
• The equilibrium constant expression for water,
Kw, can be written as:
Kw = [OH-][H3O+]
• Pure water will not appear in this expression.
• The Kw for water is 1 x 10-14, which means there
are such small amounts of ions in pure water
that water will not conduct electricity.

Chapter 8 48
[H3O+], [OH-], and pH
• Pure water has an equal number of hydroxide
and hydronium ions, so [H3O+] = [OH-]. At 25 oC
both these values are 1 x 10-7 M.
• When these concentrations are equal, the
solution is said to be neutral.
• If acid is added to water, there is an increase in
[H3O+] and a decrease in [OH-], which makes the
solution acidic.

Chapter 8 49
If a base is added, [OH-] increases and [H3O+]
decreases, which makes the solution basic.

Chapter 8 50
• Most aqueous solutions are not neutral,
meaning they have unequal concentrations of
H3O+ and OH-.
• The range of hydronium ion in an aqueous
solution can range from 18 M to 1 x 10-14 M.
• Because of this large range, it is more useful to
compare [H3O+] by a log scale because it gives
values that fall between 0 and 14.

Chapter 8 51
• Values between 1 and 14
describes the pH scale.
• Mathematically, pH can be
determined from the [H3O+]
using the following expression:
pH = - log [H3O+]

Chapter 8 52
The relationship between pH and [H3O+] are
shown below.

Chapter 8 53
Measuring pH
The pH of a solution can be measured
electronically using a pH meter. It can also be
measured visually using pH paper, which is
embedded with indicators that change color based
on the pH of a solution.

Chapter 8 54
Math Matters: Logarithms
• The exponent of 10 is the logarithm, or log, of
these numbers. For example, the log of 102 = 2;
log of 103 = 3, etc.
• The purpose of the log function is to show the
number of tens places included in a really large
or a really small number.

Chapter 8 55
Math Matters: Logarithms, Continued
Negative numbers do not have a log value. The log
function is described as the log of base 10 because
the logs of integers come from numbers that are
whole number multiples of 10.

Chapter 8 56
Math Matters: Inverse Logarithms
• Suppose we want to solve for x in the following:
4 = log x
• To solve for x, the equation must be rearranged.
• To do this, we must take the inverse log of both
sides of the equation, and applying the inverse
log function, we can solve for x:

Chapter 8 57
Calculating pH
• Calculate the pH of a 0.050 M HCl solution.
Because strong acids completely dissociate in
solution, [HCl] = [H3O+],
• The number of significant figures in the [H3O+]
will be the number of decimal places in the pH
value.

Chapter 8 58
Calculating [H3O+]
• If we measure the pH of a solution to be 3.00,
how do we find the corresponding [H3O+] if we
know that pH = - log [H3O+]?
• Multiply both sides of the equation by negative 1
and the inverse log function, INV log. The
inverse log function is 10x.

Chapter 8 59
• Solving the equation for [H3O+], we have:
INV log (-) pH = [H3O+]
or alternatively,
10-pH = [H3O+]
• The solution for the [H3O+] of a pH 3.00 solution is:
INV log (-) 3.00 or 10-3.00 = [H3O+]
1.0 x 10-3 M = [H3O+]
• The number of significant figures in the [H3O+] is:

Chapter 8 60
8.6 pKa
• To determine which acids are strongest, the Ka
values can be compared.
• As seen with the pH scale, it is easier to
compare whole numbers than those in scientific
notation. The pKa values are used to measure
acid strength.
• The lower the pKa value, the stronger the acid.

Chapter 8 61
8.6 pKa, Continued
In this table, a comparison is shown between Ka
and pKa. Those Ka values closer to 1 are the
stronger acids. The pKa values for the stronger
acids are low.

Chapter 8 62
8.6 pKa, Continued
The Relationship between pH and pKa
• pH and pKa are two numbers associated with
weak acid solutions.
• pKa gives us the fraction of acid molecules that
will dissociate, that is, pKa gives the ratio of
conjugate base and hydronium ion to weak acid.
• pH tells how much hydronium ion is present in
solution.

Chapter 8 63
8.6 pKa, Continued
• At constant temperature, pKa does not change
if an acid or a base is added to a weak acid
solution, however the pH does change if an
acid or a base is added to a weak acid
solution.
• Consider an acetic acid solution at three
different pH’s:
1. A pH below the pKa
2. A pH equal to pKa
3. A pH above the pKa

Chapter 8 64
8.6 pKa, Continued
A pH Below the pKa
• Extra H3O+ has been added to the equilibrium
solution.
• The equilibrium has shifted to the left,
meaning there is more acetic acid present
than its conjugate base, acetate.

Chapter 8 65
8.6 pKa, Continued
pH Equal to pKa
• The Ka = [H3O+].
• The pH is equal to the pKa.

Chapter 8 66
8.6 pKa, Continued
pH Is Above the pKa
• Less equilibrium concentration of H3O+ is
present.
• The equilibrium has shifted to the right,
producing more acetate than acetic acid.

Chapter 8 67
8.6 pKa, Continued
The relationship between acid (HA), conjugate
base (A-), pH, and pKa is shown in this table.

Chapter 8 68
8.7 Amino Acids: Common Biological Weak
Acids
• The molecule shown below is alanine, with
functional groups identified as a carboxylate
anion and a protonated amine.
• Alanine belongs to a class of molecules called
amino acids, which are the building blocks of
proteins.

Chapter 8 69
Acids, Continued
• Why is alanine shown with a carboxylate group
and a protonated amine rather than a carboxylic
acid group and amine group, respectively?
• The pKa value for the carboxylic acid is 2.3,
below physiological pH, so the conjugate base
form predominates. The pKa value for the amine
group is 9.7, above physiological pH, so the
conjugate acid predominates.
• This ionic form with no net ionic charge is called
a zwitterion.

Chapter 8 70
Acids, Continued
• Amino acids and other biological molecules
contain more than one weak acid group and
have more than one pKa value.
• With more than one pKa value, these molecules
exist in different acid/conjugate base forms
depending on the pH of the solution.
• These molecules have a unique pH at which the
zwitterion is the predominate form. This point is
called the isoelectric point (pI).

Chapter 8 71
Acids, Continued
• At the pI for alanine, the negative charge on the
carboxylate group is balanced by the positive
charge of the ammonium ion, and the net charge
is zero.
• The pI for alanine is 6.0 and is halfway between
the pKa values for the protonated amine and
carboxylic acid.

Chapter 8 72
8.8 Buffers: An Important Property of Weak
Acids and Bases
• Our bodies operate under strict conditions of
temperature, concentration, and pH.
• How do our bodies maintain pH in our
bloodstream when we consume a variety of
foods at different pH’s?
• Our bodies contain solutions of weak acids,
containing both acids and conjugate bases, to
help neutralize incoming acids and bases.

Chapter 8 73
Acids and Bases, Continued
• A buffer is a solution that contains both a weak
acid and its conjugate base, or a base and its
conjugate acid.
• A buffer will resist a change in pH if small
amounts of an acid or a base are added.
• Buffers are what helps our body maintain the
proper pH in our bloodstream when we consume
a variety foods at different pH’s.

Chapter 8 74
• The bicarbonate buffer system is the main
buffer system in our blood.
• Dissolved CO2, produced during cellular
respiration, is equilibrated through carbonic acid
into bicarbonate ions prior to exhalation at the
lungs.
• The intermediates, carbonic acid and water, are
often omitted since they are short lived in the
reaction.

Chapter 8 75
• The bicarbonate buffer system in our
bloodstream is shown in the following figure.
• The bicarbonate buffer system can be denoted
by showing the acid and its conjugate base like
this: H2CO3/HCO3
-. Sometimes the conjugate
base is shown as an ionic compound like this:
H2CO3/NaHCO3.

Chapter 8 76
Maintaining Physiological pH with Bicarbonate
Buffer: Homeostasis
• The bicarbonate system helps our bodies
maintain its optimal physiological pH.
• The ability of an organism to maintain its internal
environment by adjusting such factors as pH,
temperature, and solute concentration is called
homeostasis.

Chapter 8 77
Changes in Ventilation Rate
• In normal breathing, carbon dioxide is
removed from the bloodstream and blood pH
is maintained.
• A person who hypoventilates may fail to
remove enough carbon dioxide due to shallow
breathing causing carbon dioxide to build up
in the bloodstream.
• A buildup of carbon dioxide in the
bloodstream makes the blood more acidic, a
condition that is known as respiratory
acidosis.

Chapter 8 78
Changes in Ventilation Rate, Continued
• Individuals suffering this condition must be
treated to raise blood pH back to normal.
• A bicarbonate solution can be administrated
intravenously. This will drive the equilibrium to
the left, when the excess bicarbonate present
reacts with the excess acid.

Chapter 8 79
• A person who hyperventilates will exhale too
much carbon dioxide from the lungs.
• This will draw H3O+ from the bloodstream,
making the blood more basic. This condition
is known as respiratory alkalosis.
• This condition necessitates getting more
carbon dioxide back into the bloodstream,
which can be done by having the person
breathe into a paper bag.

Chapter 8 80
• This will enrich carbon dioxide in the
bloodstream, shifting the equilibrium back to
the right, thereby producing more H3O+.

Chapter 8 81
Changes in Metabolic Acid Production
• Chemical reactions in the body change the pH
of blood by producing too much or too little
H3O+.
• Since diabetics use less glucose for energy
production, they rely on fatty acids as an
energy source. A by-product of fatty acid
breakdown is acid production.
• An imbalance caused by chemical reactions
in the body is termed metabolic acidosis.

Chapter 8 82
Changes in Metabolic Acid Production,
Continued
• A treatment for this condition is to administer
bicarbonate to neutralize the excess acid so
that more carbon dioxide can be exhaled.
• If the blood is basic, the body is losing too
much acid, a condition known as metabolic
alkalosis.

Chapter 8 83
Changes in Metabolic Acid Production,
Continued
• Metabolic alkalosis occurs under conditions of
excess vomiting.
• To lower the pH back to normal, ammonium
chloride (a weak acid) can be administrated.

Chapter 8 84
A summary of acidosis and alkalosis is shown in
this table.

Chapter 8 85
Chapter Summary
• Arrhenius defines acids as producing H+ and
defines bases as producing OH-.
• Brønsted–Lowry defines acids as H+ donors and
bases as H+ acceptors.
• H+, a proton, forms a hydronium ion with a water
molecule.

Chapter 8 86
Chapter Summary, Continued
• Strong acids and bases completely dissociate in
solution.
• Water can act as either an acid or base.
• Neutralization reactions occur when a strong
acid combines with a strong base. Products are
a salt and water.
• Antacids are basic and neutralize stomach acid.

Chapter 8 87
• When forward and reverse reactions occur at the
same rate, a chemical equilibrium is
established.
• K, equilibrium constant, defines the extent of a
chemical reaction.
• If a chemical reaction at equilibrium is stressed,
the reaction regains equilibrium according to Le
Châtelier’s principle.

Chapter 8 88
• Endothermic reactions require heat in
order to occur.
• Exothermic reactions give off heat when
they occur.

Chapter 8 89
• Weak acids partially ionize in solution and have
an equilibrium constant called the acid
dissociation constant, Ka.
• Weak acids produce a conjugate base when
they dissociate, and weak bases produce a
conjugate acid when they dissociate. These are
referred to as conjugate acid–base pairs.
• Strong acids have high Ka values.

Chapter 8 90
• When water ionizes it produces hydronium ions
and hydroxide ions.
• When these ions are of equal concentration, a
neutral solution exists.
• An excess of hydronium ions produces an
acidic solution and an excess of hydroxide ions
produces a basic solution.

Chapter 8 91
• The pH scale is a measure of acidity with values
between 0–14. Neutral solutions have a pH of 7,
acidic solutions have values less than 7, and
basic solutions have values greater than 7.
• pH can be measured with a pH meter or using
pH paper.
• Mathematically written, pH = -log [H3O+].

Chapter 8 92
8.6 pKa
• pKa is a constant for a specific weak acid at a
certain temperature.
• If the pH of a solution is the same as the pKa
value of a weak acid, then the acid and
conjugate base forms are present in equal
amounts.
• If pH is higher than pKa, the conjugate base
predominates in solution.

Chapter 8 93
8.6 pKa, Continued
If pH is lower than pKa, the acid form predominates
in solution.

Chapter 8 94
Acids
• Amino acids contain an acid and base functional group.
• Amino acids are building blocks of proteins.
• Since amino acids contain both carboxylate and
protonated amine functional groups at physiological pH,
they have a net charge of zero and are called a
zwitterion.
• The isoelectric point is the pH where the zwitterion
exists.

Chapter 8 95
Acids and Bases
• Buffers resist changes in pH when an acid or a
base is added to a solution.
• Buffer solutions consist of weak acids and their
conjugate base. Strong acids cannot be buffers.
• If pH of blood drops below the normal range of
7.35–7.45 a condition called acidosis occurs.

Chapter 8 96
• If blood pH becomes elevated, a condition called
alkalosis exists.
• Acidosis can be caused by changes in breathing
or by changes in metabolism, causing acid to
build up in the bloodstream. Alkalosis is caused
by acid being removed from the bloodstream.

08_Lecture.ppt

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