Comprehensive Overview of Buffer Solutions and Their Role in Body Fluids

BUFFER, BUFFER CAPACITY, BUFFERS OF BODY
FLUIDS
Dr. Hafiz M Abdul Qayyum Lecture - 06
CHEM2114
Introduction to Biochemistry
2

Key Concepts
• Buffer Solution
• Types of Buffer Solution
• Applications of Buffer
• How Buffer Works
• Buffer Capacity
• Buffer System in Body Fluids
• Proteins Buffer System
• Phosphate Buffer System
• Carbonic Acid bicarbonate Buffer System
3

Buffer solutions
• A buffer is an aqueous solution that resists changes in pH when acids or bases are added
to it.
• In other words, a buffer solution is a solution that resists changes in pH either when
diluted or when limited amounts of acid or base are added to it.
• Such a solution can be prepared by combining a weak acid and its salt with a strong base
(conjugate base) or, a weak base and its salt with a strong acid (conjugate acid ).
• Weak acid + salt of weak acid with strong base
• CH3COOH + CH3COOH + NaOH
• CH3COOH + CH3COONa -------Acidic buffer
• Weak base+ salt of weak base + Strong acid
• NH4OH+NH4OH+ HCl
• NH4OH + NH4Cl---------- Basic buffer
• Acidic buffer: This buffer is made by mixing weak acid and the salt (preferably Na+ or
K+) of its conjugate base e.g.,
• Acetate buffer: CH3COOH (the weak acid) + CH3COONa (the salt, conjugated base)
Basic Buffer: This is made by mixing of weak base and the salt of its conjugate acid.
Ammonium buffer: NH4OH (the weak base) + NH4Cl (the salt, conjugate acid)----- Basic
buffer
4

Buffer solutions
Together the two species (conjugate acid and conjugate base) resist large
changes in pH by absorbing the H+ ions or OH– ions added to the system.
When H+ ions are added to the system they will react with the conjugate
base in the buffer. When OH– ions are added they will react with the
conjugate acid in the buffer.
• Some chemical reactions proceed only at a certain pH. Many household
and cosmetic products need to control their pH values such as, shampoo
to counteract the alkalinity of the soap and prevent irritation, baby lotion
to maintain a pH of about 6 to prevent bacteria multiplying , washing
powder , eye drops ,etc.,
• The Henderson-Hasselbalch equation allows us to calculate the pH of a
buffered system.
𝑝𝐻 = 𝑝𝐾𝑎 + 𝑙𝑜𝑔 ( [Conjugate base] /[Weak acid] )
• Systems with the greatest buffer capacity will have [conjugate
base]=[weak acid] causing the pH=pKa. Therefore, when choosing a
buffer system you should choose a system that has a pKa close to what
the pH of the overall solution should be.
5

Applications of Buffer
Buffers have many applications. Some of these are
• Maintenance of life
Most biochemical processes work within a relatively small pH range. The
body uses buffers solution to maintain a constant pH. For example, blood
contains a carbonate/bicarbonate buffer that keeps the pH close to 7.4.
• Biochemical assays
Enzyme activity depends on pH, so the pH during an enzyme assay must
stay constant.
• In shampoos
Many shampoos use a citric acid/sodium citrate shampoo to maintain a
slightly acidic pH balance. This counteracts the basicity of the detergents
present in the shampoo.
• In baby lotions
Baby lotions are buffered to a pH of about 6. This hinders the growth of
bacteria within the diaper and helps prevent diaper rash.
• In the brewing Industry
Buffer solutions are added before fermentation begins. This prevents the
solutions becoming too acidic and spoiling the product.
6

• In Textile Industry.
Many dyeing processes use buffers to maintain the correct pH for various
dyes.
• In Laundry Detergents.
Many laundry detergents use buffers to prevent their natural ingredients
from breaking down.
• In Food Industry
Buffers are also used in foods to maintain the activity of the food in order
to preserve the flavor and appearance of food. Buffers maintain the
physical, chemical and microbiological stability of foods. For examples,
potassium citrate is an antioxidant and buffering additives that is found
in a number of food products including cakes, biscuits, cheese and jam.
• In Dyeing Industry
Dyes in textile industries play an important role in giving color to different
fabrics. Color strength of dyes related to narrow pH range which is
maintained by using different buffer systems.
7

• Printing Industries
• Buffers are used during printing to ensure the link maintains its normal
properties. Incorrect pH may affect how the ink penetrates and dries on
the paper and altering the final result.
• Electroplating Industries
• Some alloys can be plated, if very strict pH control is maintained.
• Leather Industries
• Narrow range of pH control of tanning and dyeing baths determine the
texture and color of the finished product.
• In Pharmaceutical Industry
• Most of the medicines are prepared in aqueous solution of different
chemicals so these aqueous solutions require a constant pH in order to
assure the stability and clinical effectiveness of a medicines and this is
done through buffers. Buffer are also added in pharmaceutical to improve
patient comfort and to make longer transportation of medicines possible.
• In Fermentation
• Fermentation reactions such as yogurt is highly affected by varying pH.
This means it is essential to use buffer solutions to avoid harsh changes
and allow fermentation to progress to obtain maximum yield.
8

How buffer works
Suppose HCl is added to a buffer system containing NaHCO3 and H2CO3. The following reaction will
take place.
NaHCO3 + HCl H2CO3 + NaCl
Thus HCl which is a strong acid and is expected to raise the H+ concentration reacts with a base to
yields H2CO3 (a weak acid due to little dissociation) and a neutral salt, NaCl; the result is that if
HCl is not added in large amounts, there will be only a little change in the original pH of the buffer
solution.
• In the same way if NaOH is added to this buffer solution, it will react with H2CO3 as follows:
• NaOH + H2CO3 NaHCO3 + H2O
• Since NaHCO3 is a much weaker base than NaOH, the resulting rise in pH will be quite small. In
above reaction, other things (like water and sodium chloride) which are present are not important to
the argument.
Buffer solution act best at a certain pH, where their buffering power is the greatest. For example, acetic
acid/acetate buffer has maximum buffering power around pH 4.76. The two buffer systems
described above, i.e., H2CO3/NaHCO3 and NaH2PO4/Na2HPO4 and some other buffer systems in
the body are the first line of defense against any strong acid or base which might get entry into the
living body. The above mentioned examples of buffers apply to weak acids and their salts.
However, a solution containing a weak base, e.g., NH4OH and its salt NH4Cl can also act as a
buffer system.
The other body buffers include hemoglobin (a protein in red blood cells) and proteins present in the
blood plasma. All these buffer systems help in the maintenance of the blood pH at a remarkably
constant level, around 7.4 in spite of large variations in individuals intake of foods and drinks. It is
apparent that this pH is alkaline.
9

Some common weak acids and bases and their salts usually used for the
preparation of buffer solutions
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Acid or base Salt
Acetic acid Sodium acetate
Phosphoric acid Potassium phosphate
Oxalic acid Lithium oxalate
Carbonic acid Sodium carbonate
Ammonia Ammonium chloride

Buffer Capacity
• Buffer capacity (β) is defined as the moles of an acid or base necessary to
change the pH of a solution by 1, divided by the pH change and the volume
of buffer in liters; it is a unit less number.
• It tells us that how much acid or base we can add before pH of the buffer does
changing. The capacity of the buffer solution is a measure of acid or base. If
higher concentration of acid/base in the buffer solution, the buffer capacity
will be higher. If lower concentration of acid/base in the buffer solution, the
buffer capacity will be lower.
• It defined as the number of moles of H+/OH– ions that must be added to one
liter of the buffer in order to decrease /increase the pH by one unit
respectively.
• Buffer capacity (β) is defined as the amount of a strong acid or a strong base
that has to be added to 1 litre of a buffer to cause pH change of 1.0 pH unit.
• It is the number of moles or m moles of a strong base (y), which when added
to one liter of a buffer raises his own pH by one unit , or it is the number of
moles or m moles of a strong acid (x) which, when added to one liter of this
buffer reduced its pH by one unit.
11

Buffer Capacity
• The higher the capacity, the higher the amounts of a strong acid or a strong
base which can be added to the buffer without significantly changing his own
pH.
• Buffer Capacity= No of moles of acid or base added per liter of the buffer/
Change in pH
• Buffer capacity is a measure of the efficiency of a buffer in resisting changes
in pH. Conventionally, the buffer capacity (β) is expressed as the amount of
strong acid or base, in gram-equivalents, that must be added to 1 liter of the
solution to change its pH by one unit.
β= ΔB/ ΔpH
• In which β = Buffer capacity,
• Δ (delta) =a finite change
• ΔB = the small increment in gram equivalents (gEq)/Liter of strong base added
to the buffer solution to produce a pH change of ΔpH. According to equation,
the capacity of a solution has a value of 1 when the addition of 1 g Eq of
strong base (or acid) to 1 Liter of the buffer solution results in a change of 1
pH unit. The higher the buffer capacity the less the buffer solution changes its
pH.
12

Buffer system in Body Fluids
• Buffer solutions are used as a means of keeping pH at a nearly constant value in a
wide variety of chemical applications.
• Hydrogen ions (H+) are present in all body compartments. Maintenance of
appropriate concentration of hydrogen ion (H+) is critical to normal cellular
function. The acid base balance or pH of the body fluids is maintained by a
closely regulated mechanism.
• There are variety of chemicals in body fluids that prevents the fluids from
undergoing large changes in pH. These chemicals are regulate fluctuations in H+
concentration. These chemicals buffer: Bind to H+ ions when there are too many in
a solution so pH remains normal. Release H+ ions when there are too few in a
solution so pH remains normal.
• The pH level of the blood is important in ensuring the proper functionality of
biological system. The addition of an acid or base to a substance changes its pH
level. A buffer is a solution that has the ability to maintain pH and bring it back to
its optimal value. It does this by the additional or removal of hydrogen ions.
13

Buffer system in Body Fluids
• In the human body, the main buffers in blood are bicarbonate,
hemoglobin, plasma proteins and phosphates. For example, the
bicarbonating buffer system is used to regulate the pH of blood. The
normal pH of human blood is 7.4.
• Some people suffer from alkalosis when experiencing severe anxiety.
• Alkalosis is a condition in which the pH of the blood is too high.
• The opposite condition - a blood pH lower than 7.4 is called acidosis
• The problem for the body is that normal metabolism is associated with
continuous production of hydrogen ions (H+) and carbon dioxide (CO2),
both of which tend to reduce pH.
• The mechanism which overcomes this problem and serves to maintain
normal blood pH (i.e. preserve acid-base homeostasis) is a complex
synergy of action involving chemical buffers in blood, the red cells
(erythrocytes), which circulate in blood, and the function of three
organs: lungs, kidneys and brain.
• Three buffer system in body fluids
• 1) Protein buffer system
• 2) Phosphate buffer system
• 3) Carbonic acid bicarbonate buffer system 14

Protein Buffer System
• Proteins are the most important and widely operating buffers in the body
fluids. They are the most powerful buffers in plasma and cells. The
proteins buffer system is important in regulating blood plasma and intra
cellular fluid pH .
• Proteins are excellent buffer because they contain both acid and base
groups that can give up and take up H+ ions. Protein buffer system helps
to maintain acidity in and around the cells. Protein buffers are either
intracellular and extracellular. Proteins are mainly intracellular found
and include hemoglobin (Hb). Hemoglobin is by far most important
protein buffer in red blood cells. Whereas albumin is main plasma
protein buffer. Plasma proteins are extracellular but their amount is
small. Other proteins containing amino acid are histidine and also good
at buffering.
• Hb is the protein that functions to transport oxygen within the body.
Protein buffers include basic group and acidic group that act as hydrogen
ion or donors to maintain acceptor. Amino acid has the ability to respond
to pH changes by accepting H+(acting as a base) or releasing hydrogen
(acting as an acid).
• Hemoglobin makes an excellent buffer by binding to small amounts of
acids in the blood, before they can alter the pH of the blood.
15

If plasma or extracellular fluid pH drops due to excessive H+ ions, then H+
ions bind to amino groups and pH rises.
If plasma or extracellular fluid rises due to a shortage of H+ ions then acid
groups dissociate or H+ ions releases and pH rises.
Amino acids contains at least one carboxyl group (-COOH), which is a
functional component of the buffer and at least one amino group (-
NH2) .
The free carboxyl group at one end of proteins acts like an acid by
releasing H+ ions when pH rises. When it dissociates, the H+ can react
with excess OH– in the solution to form water.
Amino group acts like a base and combines with H+ ions
Hemoglobin acts as a buffer in blood by picking up CO2 or H+
In free amino acids, both the main structural chain and the side chain can
act as buffers. In a protein, most of the carboxylic and amino groups in
the main chain are tied up in peptide bonds. Only the exposed amino
group and carboxyl group at either end of a protein are available as
buffers. So most of the buffering capacity of proteins is provided by the
R groups of amino acids. Plasma proteins contribute to the buffering
capability of blood
16

• In acidic medium, proteins acts as a base amino group takes up H+ ions from the medium forming
NH3
+ and proteins become positively charged.
• In other words, if pH decreases in a solution: The carboxylate ion and the amino group (-NH2) can
act as weak bases and accepts additional hydrogen ions, forming a carboxyl group (-COOH) and an
amino ion (-NH3), respectively.
• In a neutral solution, and amino acid: Carries both positive and negative charges. This state is
called a zwitterion (hybrid ion).
• In alkaline medium: protein acts as an acid (COOH-) group dissociates and gives H+ forming
COO–so proteins become negatively charged.
• In other words, if pH increases in a solution: The carboxyl group of the amino acid can dissociate,
acting as a weak acid and releasing a hydrogen ion. The carboxyl group then becomes a carboxylate
ion.
• Some amino acid's, notably histidine and cysteine, have R groups (side chains) that will donate
hydrogen ions if the pH increases outside the normal range.
17

Phosphate Buffer System
• Sodium dihydrogen phosphate and disodium hydrogen phosphate constitute the
phosphate buffer (NaH2PO4/Na2 HPO4 ).
• NaH2PO4 acts as the acid and Na2HPO4 acts as its salt because the former one has
one more H+ than the latter.
• Its components, dihydrogen phosphate (H2PO4
–) a weak acid and monohydrogen
phosphate (HPO4
2 –), a weak base.
• The phosphate buffer system consists of two ions hydrogen phosphate ion (weak
base) and dihydrogen phosphate ion (weak acid).
• This system is an effective buffer in urine and intracellular fluid. Although
the phosphate buffer system is not important as an extracellular fluid buffer, yet it
plays a major role in buffering renal tubular fluid and intracellular fluids.
• The main elements of the phosphate buffer system are H2PO4
– and HPO4
2–.
• When a strong acid such as HCl is added to a mixture of these two substances, the
hydrogen is accepted by the base HPO4
2– and converted to H2PO4
–
• The result of this reaction is that the strong acid, HCl, is replaced by an additional
amount of a weak acid, NaH2PO4, and the decrease in pH is minimized.
• When a strong base, such as NaOH, is added to the buffer system, the OH– is
buffered by the H2PO4
– to form additional amounts of HPO4
2– + H2O. In this case, a
strong base, NaOH, is traded for a weak base, NaH2PO4, causing only a slight
increase in pH.
18

• The phosphate buffer system has a pK of 6.8, which is not
far from the normal pH of 7.4 in the body fluids; this
allows the system to operate near its maximum buffering
power. However, its concentration in the extracellular
fluid is low, only about 8 per cent of the concentration of
the bicarbonate buffer. Therefore, the total buffering
power of the phosphate system in the extracellular fluid is
much less than that of the bicarbonate buffering system.
• Hydrogen phosphate ions accepts all addition of H+ ions to
maintain pH.
• Dihydrogen phosphate ions release additional hydrogen
ion to maintain pH.
• Phosphate buffer system operates in the internal fluids of
all cells. It consists of dihydrogen phosphate ions as the
hydrogen ion donor (acid) and hydrogen phosphate ion as
the ion acceptor (base) . 19

• When pH is rising (look at the direction of arrow, right):
• H2PO4
– H+ + HPO4
2–
• When pH is falling (look at the direction of arrow, left):
• H2PO4
– H+ + HPO4
2–
• The phosphate system is most important in regulating intracellular fluid pH.
• If intracellular fluid pH drops due to excessive H+ ions, then dihydrogen phosphate is formed and
intracellular fluid pH rises.
• If intracellular fluid pH rises due to a shortage of H+ ions, then monohydrogen phosphate ions are
formed and intracellular fluid pH drops.
• If additional hydroxide ions enter the cellular fluid, they are neutralized by the dehydrogenate
phosphate ion. If extra hydrogen ions enter the cellular fluid then they are neutralized by the hydrogen
phosphate ion.
In contrast to its rather insignificant role as an extracellular buffer, the phosphate buffer is especially
important in the tubular fluids of the kidneys, for two reasons: (1) phosphate usually becomes greatly
concentrated in the tubules, thereby increasing the buffering power of the phosphate system, and (2)
the tubular fluid usually has a considerably lower pH than the extracellular fluid does, bringing the
operating range of the buffer closer to the pK (6.8) of the system.
The phosphate buffer system is also important in buffering intracellular fluid because the concentration of
phosphate in this fluid is many times that in the extracellular fluid. Also, the pH of intracellular fluid is
lower than that of extracellular fluid and therefore is usually closer to the pK of the phosphate buffer
system compared with the extracellular fluid.
20

Carbonic acid bicarbonate buffer system
• Carbonic acid (H2CO3) is a weak acid and is therefore in equilibrium with
bicarbonate (HCO3
–) in solution. When significant amounts of both carbonic
acid and bicarbonate are present, a buffer is formed. This buffer system can be
written as:
• H2CO3 (carbonic acid) HCO3
– (conjugate base)+ H+ (hydrogen ion)
H2CO3 + H2O H3O+ + HCO3
–
• Under normal circumstances there is much more bicarbonate present than
carbonic acid (the ratio is approximately 20:1). As normal metabolism produces
more acids than bases, this is consistent with the body's needs. The blood, with
its high base concentration, is able to neutralize the metabolic acids produced.
Since relatively small amounts of metabolic bases are produced, the carbonic
acid concentration in the blood can be lower.
• Since carbonic acid is not stable in aqueous solutions some of it decomposes to
form carbon dioxide and water. The respiratory system is responsible for
removing the carbon dioxide.
H2CO3 H2O + CO2
21

• By combining the two reactions of carbonic acid we can write:
2H2O + CO2 H2CO3 + H2O H3O+ + HCO3
–
•
It is the production of carbon dioxide from this reaction that couples the carbonic
acid/bicarbonate buffer to the respiratory system.
• The carbonic acid - bicarbonate buffer system consists of carbonic acid, a
weak acid, and the bicarbonate anion, its conjugate base. Likewise, if a strong
base is introduced, it will react with the carbonic acid to form
the bicarbonate anion, thus reducing the potential increase in pH.
• The bicarbonate buffer system functions to maintain pH level in the blood of
mammals.
• It plays a major role in the formation of acid in the stomach. The bicarbonate
buffer system manages acid and base imbalances. Cellular respiration produces
CO2 as a waste product. This is immediately converted to bicarbonate ion in the
blood. On reaching the lungs it is again converted to and released as CO2. While
in the blood , it neutralizes acids released due to other metabolic processes. In the
stomach and duodenum, it also neutralizes gastric acids and stabilizes the intra
cellular pH of epithelial cells by the secretions of bicarbonate ions into the gastric
mucosa.
• –
22

• The carbonic acid bicarbonate system is most important in
regulating blood plasma pH.
• If plasma pH drops due to excessive H+ ions, then carbonic acid is
formed and plasma pH rises.
• If plasma pH rises due to a shortage of H+ ions then bicarbonate ions
are formed and plasma Ph drops.
• The major buffer system in the body fluid H2CO3/HCO3. Some CO2,
the end product of cellular metabolism and the rest dissolves in body
fluids forming carbonic acid and dissociate to produce bicarbonate
(HCO3) and hydronium (H3O+) ions. More of the HCO3 is supplied
by kidneys.
CO2+H2O H2CO3
H2CO3 + H2O H3O+ + HCO3
–
23

PH,BUFFERS AND HENDERSON HASSELBALCH
EQUATION
Fatima Saleem Lecture - 05
CHEM2114
Introduction to Biochemistry
2

Key Concepts
➢ pH and Sorensen’s scale
➢ Understanding Ka and pKa
➢ What Is a Buffer?
➢ How Do Buffers Work?
➢ Henderson-Hasselbalch Equation
3

pH and Sorensen’s scale
4
pH is a measure of hydrogen ion concentration, a
measure of the acidity or alkalinity of a solution. The
pH scale usually ranges from 0 to 14. Aqueous
solutions at 25°C with a pH less than 7 are acidic, while
those with a pH greater than 7 are basic or alkaline. A
pH level of 7.0 at 25°C is defined as "neutral" because
the concentration of H3O+ equals the concentration of
OH− in pure water. Very strong acids might have a
negative pH, while very strong bases might have a pH
greater than 14.

pH Equation
• The equation for calculating pH was proposed in
1909 by Danish biochemist Søren Peter Lauritz
Sørensen:
• pH = -log[H+]
• where log is the base-10 logarithm and [H+]
stands for the hydrogen ion concentration in
units of moles per liter solution. The term "pH"
comes from the German word "potenz," which
means "power," combined with H, the element
;symbol for hydrogen, so pH is an abbreviation for
"power of hydrogen.”
5

IUPAC Definition of pH
• The International Union of Pure and Applied Chemistry
(IUPAC) has a slightly different pH scale that is based on
electrochemical measurements of a standard buffer
solution. Essentially, the definition uses the equation:
• pH = -log aH+
• where aH+ stands for hydrogen activity, which is the
effective concentration of hydrogen ions in a solution.
This might be slightly different from the true
concentration. The IUPAC pH scale also includes
thermodynamic factors, which may influence pH.
6

How pH Is Measured
• Rough pH measurements can be made using litmus
paper or another type of pH paper known to change colors
around a certain pH value. Most indicators and pH papers
are useful only to tell whether a substance is an acid or a
base or to identify pH within a narrow range. A universal
indicator is a mixture of indicator solutions intended to
provide a color change over a pH range of 2 to 10.
• More accurate measurements are made using primary
standards to calibrate a glass electrode and pH meter. The
electrode works by measuring the potential difference
between a hydrogen electrode and a standard electrode.
An example of a standard electrode is silver chloride.
7

Uses of pH
• pH is used in everyday life as well as science
and industry. It's used in cooking (e.g.,
reacting baking powder and an acid to make
baked goods rise), to design cocktails, in
cleaners, and in food preservation.
• It's important in pool maintenance and water
purification, agriculture, medicine, chemistry,
engineering, oceanography, biology, and other
sciences.
8

Understanding Ka and pKa
• Ka, pKa, Kb, and pKb are most helpful when predicting
whether a species will donate or accept protons at a
specific pH value. They describe the degree of ionization of
an acid or base and are true indicators of acid or base
strength because adding water to a solution will not change
the equilibrium constant. Ka and pKa relate to acids, while
Kb and pKb deal with bases. Like pH and pOH, these values
also account for hydrogen ion or proton concentration (for
Ka and pKa) or hydroxide ion concentration (for Kb and
pKb).
• Ka and Kb are related to each other through the ion
constant for water, Kw:
• Kw = Ka x Kb
9

Ka and pKa
• Ka is the acid dissociation constant. pKa is simply the -log of this constant.
Similarly, Kb is the base dissociation constant, while pKb is the -log of the constant.
The acid and base dissociation constants are usually expressed in terms of moles
per liter (mol/L). Acids and bases dissociate according to general equations:
• HA + H2O ⇆ A- + H3O+
• HB + H2O ⇆ B+ + OH-
In the formulas, A stands for acid and B for base.
• Ka = [H+][A-]/ [HA]
• pKa = - log Ka
at half the equivalence point, pH = pKa = -log Ka
• A large Ka value indicates a strong acid because it means the acid is largely
dissociated into its ions. A large Ka value also means the formation of products in
the reaction is favored. A small Ka value means little of the acid dissociates, so you
have a weak acid. The Ka value for most weak acids ranges from 10-2 to 10-14.
The pKa gives the same information, just in a different way. The smaller the value
of pKa, the stronger the acid. Weak acids have a pKa ranging from 2-14.
10

Understanding Kb and pKb
Kb is the base dissociation constant. The base
dissociation constant is a measure of how completely a
base dissociates into its component ions in water.
• Kb = [B+][OH-]/[BOH]
• pKb = -log Kb
A large Kb value indicates the high level of dissociation
of a strong base. A lower pKb value indicates a stronger
base.
pKa and pKb are related by the simple relation:
• pKa + pKb = 14
11

What Is a Buffer?
• There are two key terms associated with
buffers. A buffer is an aqueous solution that
has a highly stable pH. A buffering agent is a
weak acid or weak base that helps maintain
the pH of an aqueous solution after adding
another acid or base. If you add an acid or
a base to a buffered solution, its pH will not
change significantly. Similarly, adding water to
a buffer or allowing water to evaporate will
not change the pH of a buffer.
12

How Do You Make a Buffer?
• A buffer is made by mixing a large volume of
a weak acid or weak base together with its
conjugate. A weak acid and its conjugate base
can remain in solution without neutralizing
each other. The same is true for a weak
base and its conjugate acid
13

How Do Buffers Work?
• In order to understand how a buffer works, consider the
example of a buffer solution made by dissolving sodium
acetate into acetic acid. Acetic acid is (as you can tell from
the name) an acid: CH3COOH, while the sodium acetate
dissociates in solution to yield the conjugate base, acetate
ions of CH3COO-. The equation for the reaction is:
• CH3COOH(aq) + OH-(aq) ⇆ CH3COO-(aq) + H2O(aq)
• If a strong acid is added to this solution, the acetate ion
neutralizes it:
• CH3COO-(aq) + H+(aq) ⇆ CH3COOH(aq)
• This shifts the equilibrium of the initial buffer reaction,
keeping the pH stable. A strong base, on the other hand,
would react with the acetic acid.
14

Universal Buffers
• Most buffers work over a relative narrow pH range. An
exception is citric acid because it has three pKa values.
When a compound has multiple pKa values, a larger pH
range becomes available for a buffer. It's also possible
to combine buffers, providing their pKa values are close
(differing by 2 or less), and adjusting the pH with
strong base or acid to reach the required range. For
example, McIvaine's buffer is prepared by combining
mixtures of Na2PO4 and citric acid. Depending on the
ratio between the compounds, the buffer may be
effective from pH 3.0 to 8.0. A mixture of citric acid,
boric acid, monopotassium phosphate, and diethyl
barbituic acid can cover the pH range from 2.6 to 12!
16

Buffer Key Takeaways
• A buffer is an aqueous solution used to keep the
pH of a solution nearly constant.
• A buffer consists of a weak acid and its conjugate
base or a weak base and its conjugate acid.
• Buffer capacity is the amount of acid or base that
can be added before the pH of a buffer changes.
• An example of a buffer solution is bicarbonate in
blood, which maintains the body's internal pH.
17

Examples of Buffers
• blood - contains a bicarbonate buffer system
• TRIS buffer
• phosphate buffer
18

Henderson-Hasselbalch Equation
• The Henderson Hasselbalch equation is an
approximate equation that shows the relationship
between the pH or pOH of a solution and the pKa or
pKb and the ratio of the concentrations of the
dissociated chemical species. In order to use the
equation, the acid dissociation constant must be
known.
Equation
There are multiple ways to write the equation. Two of
the most common are:
• pH = pKa + log ([conjugate base]/[weak acid])
• pOH = pKb + log ([conjugate acid]/[weak base])
19

Sources
• The Henderson-Hasselbalch equation relates pH, pKa,
and molar concentration (concentration in units of
moles per liter):
pH = pKa + log ([A-]/[HA])
• [A-] = molar concentration of a conjugate base
• [HA] = molar concentration of an undissociated weak
acid (M)
The equation can be rewritten to solve for pOH:
pOH = pKb + log ([HB+]/[ B ])
• [HB+] = molar concentration of the conjugate base (M)
• [ B ] = molar concentration of a weak base (M)
20

History
• An equation to calculate the pH of a buffer
solution was derived by Lawrence Joseph
Henderson in 1908. Karl Albert Hasselbalch
rewrote this formula in logarithmic terms in
1917.
21

Example Problem Applying the
Henderson-Hasselbalch Equation
• Calculate the pH of a buffer solution made from 0.20 M
HC2H3O2 and 0.50 M C2H3O2
- that has an acid dissociation
constant for HC2H3O2 of 1.8 x 10-5.
Solve this problem by plugging the values into the
Henderson-Hasselbalch equation for a weak acid and
its conjugate base.
pH = pKa + log ([A-]/[HA])
pH = pKa + log ([C2H3O2
-] / [HC2H3O2])
pH = -log (1.8 x 10-5) + log (0.50 M / 0.20 M)
pH = -log (1.8 x 10-5) + log (2.5)
pH = 4.7 + 0.40
pH = 5.1
22

Assignments
1. How does bicarbonate buffer system work in
blood?
2. If the H+ concentration is 0.0001 M/l, what is the
pH?
3. If the H+ concentration is 0.00001 M/l,
what is the OH- concentration?
23

Dr. Saima Naz Lecture - 20
CHEM3119
Biometabolism
General ways of Amino Acid Degradation
2

Key Concepts
1. General ways of amino acids metabolism
I. Deamination
II. Transamination
III. Decarboxylation
2. Four types of deamination
4

General ways of amino acids
metabolism
The fates of amino acids:
1)For protein synthesis;
2) For synthesis of other nitrogen containing
compounds (creatine, purines, choline, pyrimidine);
3)As the source of energy;
4)For the gluconeogenesis.
5

 The general ways of amino acids degradation:
1. Deamination
2. Transamination
3. Decarboxilation
 The major site of amino acid degradation is the
 Deamination of amino acids
 Deamination - elimination of amino group from
amino acid with ammonia formation.
 Four types of deamination:
1. Oxidative (the most important for higher animals),
2. Reduction
3. Hydrolytic
4. Intramolecular 6

1. Reduction deamination:
R-CH(NH2)-COOH + 2H+  R-CH2-COOH + NH3
amino acid fatty acid
2. Hydrolytic deamination:
R-CH(NH2)-COOH + H2O  R-CH(OH)-COOH +
NH3
amino acid hydroxyacid
3. Intramolecular deamination:
R-CH(NH2)-COOH  R-CH-CH-COOH + NH3
amino acid unsaturated fatty acid 7

General scheme of oxydative
transamination
HOOC C CH2CH2COOH
O
+
R CH COOH
NH2
aminokyselina 2-oxoglutarát
HOOC CH CH2CH2COOH
NH2
+
R C COOH
O
glutamát
2-oxokyselina
aminotransferasa
pyridoxalfosfát
amino acid
2-oxo acid
2-oxoglutarate
glutamate
aminotransferase
pyridoxal phosphate
8

Glutamate dehydrogenase
(GMD, GD, GDH)
• Requires pyridine cofactor NAD(P)+
• GMD reaction is reversible: dehydrogenation
with NAD+,
• Hydrogenation with NADPH+H+
 Two steps
• Dehydrogenation of CH-NH2 to imino group C=NH
• Hydrolysis of imino group to oxo group and
ammonia
:
9

In short
In transaminations, nitrogen of most AA is
concentrated in glutamate
• Glutamate then undergoes dehydrogenation +
deamination and releases free ammonia NH3
10

4. Oxidative deamination
 L-Glutamate dehydrogenase plays a central role
in aminoacid deamination
 In most organisms glutamate is the only amino
acid that has active dehydrogenase
 Present in both the cytosol and mitochondria of
the liver
11

2. Transamination of amino
acids
• Transamination - transfer of an amino group from
an -amino acid to an -keto acid (usually to -
ketoglutarate)
Enzymes: aminotransferases (transaminases).
 -aminoacid  -ketoacid
 -ketoacid -amino acid
13

There are different transaminases
 The most common:
1. alanine aminotransferase alanine + -
ketoglutarate  pyruvate + glutamate
2. aspartate aminotransferase
aspartate + -ketoglutarate  oxaloacetate +
glutamate
14

3: Aminotransferases
Funnel -amino groups from a variety of
amino acids to -ketoglutarate with glutamate
formation Glutamate can be deaminated with
NH4
+ release
15

Mechanism of
transamination
 All aminotransferases require
the prosthetic group pyridoxal
phosphate (PLP), which is
derived from pyridoxine
(vitamin B6).
16

• Ping-pong kinetic mechanism
• First step: the amino group of
amino acid is transferred to
pyridoxal phosphate, forming
pyridoxamine phosphate and
releasing ketoacid.
• Second step: -ketoglutarate
reacts with pyridoxamine
phosphate forming glutamate
17

Ping-pong kinetic mechanism of aspartate
transaminase
aspartate + -ketoglutarate  oxaloacetate + glutamate
18

amine
• Usually amines have high physiological activity
(hormones, neurotransmitters etc).
• Enzyme: decarboxylases
• Coenzyme – pyrydoxalphosphate
Decarboxylation of amino acids
• Decarboxylation – removal of carbon dioxide from
amino acid with formation of amines.
19

α-decarboxilation
ω-decarboxilation
Decarboxilation with transamination
Decarboxylation Of Amino
Acids
20

• Decarboxylation with conjugation of two
molecules
21

Significance of amino acid
decarboxylation
1. Formation of physiologically active compounds
glutamate gamma-aminobutyric acid
(GABA)
histamine
histidine
22

1) A lot of histamine is formed in inflammatory place; It
has vasodilator action;
• Mediator of inflamation, mediator of pain;
Responsible for the allergy development;
• Stimulate HCI secretion in stomach
.
2 ) Tryptophan Serotonin
• It has Vasoconstrictor action
• Takes part in regulation of arterial pressure, body
temperature, respiration, kidney filtration, mediator of
nervous system 23

3) Tyrosine  Dopamine
• It is precursor of epinephrine and
norepinephrine. mediator of central
nervous system
4) Glutamate   -aminobutyrate (GABA)
• Is is ingibitory mediator of central nervous
system. In medicine we use with
anticonvulsion purpose (action).
24

References
25
1. "Transcription, Translation and Replication". www.atdbio.com.
Retrieved 2019-02-12.
2. Berg JM, Tymoczko JL, Stryer L (2002). Biochemistry (5th ed.). New York: W.H.
Freeman. ISBN 978-0716730514. OCLC48055706
3. "Chemistry for Biologists: Nucleic acids". www.rsc.org. Retrieved 2019-02-20
4. Nelson DL, Cox MM, Lehninger AL (2013). Lehninger principles of
biochemistry(6th ed.). New York: W.H. Freeman and Company. OCLC
824794893.
5. Ahern K, Rajagopal I (2013). Biochemistry Free and Easy (PDF) (2nd ed.).
Oregon State University.
6. N.V. Bhagavan, Chung-Eun Ha, in Essentials of Medical Biochemistry (Second
Edition), 2015

Assignment
Q 1 : Describe four general ways of deamination of
proteins by chemical reactions.
Q2: What is the significance of amino acid
decarboxylation?
26

Dr. Saima Naz Lecture - 03 CHEM3119
Biometabolism
1
Digestion and metabolism of proteins

COURSE: (CHEM3119)
Biometabolism
Topic :Digestion and metabolism of proteins
Presented By
Dr Saima Naz
Assistant Professor
Department of Chemistry
Email: saima.naz@ue.edu.pk
2

Key Concepts
1. Metabolism
2. Protein metabolism
3. Pathway of protein metabolism
4. Nitrogen balance
4

Introduction
 Metabolism is the term used to describe
• The interconversion of chemical compounds in the
body
• The pathways taken by individual molecules,
• Their interrelationships, and the mechanisms that
regulate the flow of metabolites through the
pathways
 It falls mainly in 3 categories: catabolism,
anabolism and amphibolic pathways

Metabolism
 Anabolic pathways
 Involved in the synthesis of larger and more
complex compounds from smaller precursors
 Ex: Synthesis of protein from amino acids and the
synthesis of reserves of tri-acylglycerol and
glycogen.
 Anabolic pathways are endothermic.

7
 Catabolic pathways
 Involved in the breakdown of larger molecules, commonly
involving oxidative reactions;
 They are exothermic, producing reducing equivalents,
and, mainly via the respiratory chain
 Amphibolic pathways
 Occur at the “crossroads” of metabolism, acting as links
between the anabolic and catabolic pathways
 Ex: Citric acid cycle

Metabolism
• A 70-kg adult human being requires about 1920-
2900 kcal from metabolic fuels each day, depending
on physical activity.
• This energy requirement is met from
• Carbohydrates (40%-60%)
• Lipids (mainly triacylglycerol, 30%-40%)
• Protein (10%-15%), as well as alcohol.

9
• There is a constant requirement for
metabolic fuels throughout the day
• Most people consume their daily intake of
metabolic fuels in two or three meals, so
there is a need to form reserves

Metabolism
• Reserves of
• Carbohydrate: glycogen in liver and muscle
• Lipid: triacylglycerol in adipose tissue
• Labile protein
• If the intake of metabolic fuels is consistently
greater than energy expenditure.
• Surplus is stored, largely as triacylglycerol in
adipose tissue.

11
• Leading to the development of obesity
• If the intake of metabolic fuels is consistently
lower than energy expenditure
• Reserves of fat and carbohydrate, and amino
acids are used for energy-yielding metabolism
• This leads to emaciation, wasting, and,
eventually, death

Metabolism
• All the products of
digestion are
metabolized to
acetyl-CoA -
oxidized by the
citric acid cycle

Carbohydrate
Metabolism
• Glucose - major fuel of most
tissues
• Metabolized to pyruvate
aerobically and lactate
anaerobically by the
pathway of glycolysis which
further gets reduced to
acetyl-CoA
• Enter the Citric acid cycle -
formation of ATP in the
process of oxidative
phosphorylation

Carbohydrate Metabolism
• Glucose may also undergo Glycogenesis - synthesis of the
storage polymer glycogen in skeletal muscle and liver
• It may also enter Pentose Phosphate Pathway (PPP)
– source of reducing equivalents (NADPH) for fatty acid
synthesis
– ribose for nucleotide and nucleic acid synthesis
• Pyruvate and intermediates of the Citric Acid Cycle - carbon
skeletons for the synthesis of nonessential amino acids

15
• Acetyl-CoA is the precursor of fatty acids and
cholesterol (steroid hormones synthesized in
the body)
• Gluconeogenesis is the process of
synthesizing glucose from non carbohydrate
precursors such as, lactate, amino acids,
and glycerol

Lipid Metabolism
• Source of long-chain fatty acids
• either dietary lipid
• Or de novo synthesis from acetyl-
CoA
• Fatty acids may either
• gets oxidized to acetyl- CoA (
β-oxidation)
• esterified with glycerol forming
triacylglycerol - body’s main fuel
reserve

Lipid Metabolism
• Acetyl-CoA formed by β-oxidation of fatty acids may
undergo three fates
• oxidized via the citric acid cycle precursor for
synthesis of cholesterol and other steroids.
• In the liver, it is used to form the ketone
bodies, acetoacetate and 3-hydroxybutyrate
important fuels in prolonged fasting and
starvation.

Amino Acid Metabolism
• AA forms metabolic intermediates by transamination
using the amino group from other amino acids
• Deamination – removal of amino nitrogen as urea via
Urea cycle,
• The carbon skeletons that remain after transamination
• Oxidized via the citric acid cycle,
• used to synthesize glucose (gluconeogenesis)

19
• Form ketone bodies or acetyl CoA,
which may be oxidized or used for
synthesis of fatty acids
• AA may also be precursors of other
compounds
• Purines and pyrimidines,
• Hormones such as epinephrine and thyroxine,
• Neurotransmitters.

• Intake of dietary protein is in the range of
50-100 g/day
• About 30-100 g/day of endogenous protein
is derived form the digestive enzymes and
worn out cells of the digestive tract.
• The digestion and absorption of proteins is
very efficient in healthy humans, hence very
little protein (about 5-10 g/day) is lost
through feces and major quantity get
absorbed in various parts of body to perform
various functions.
Protein metabolism

21
• Total of all chemical changes that occur in body. Includes:
• Anabolism: energy-requiring process where small
molecules joined to form larger molecules • E.g. Glucose +
Glucose maltose
• Catabolism: energy-releasing process where large
molecules broken down to smaller •
• Energy in carbohydrates, lipids, proteins is used to produce
ATP through oxidation-reduction reactions

Oxidation-Reduction Reactions
• Oxidation occurs via the loss of hydrogen or the
gain of oxygen. Removal of electron or removal
of hydrogen.
• Reduction occurs through gain of H, or loss of
oxygen
• Whenever one substance is oxidized, another
substance is reduced
• Oxidized substances lose energy

• Reduced substances gain energy
• Coenzymes act as hydrogen (or electron)
acceptors
• Two important coenzymes are nicotinamide
adenine dinucleotide (NAD+) and flavin adenine
dinucleotide (FAD) take part in biochemical
reactions.

Metabolic Pathways
• The enzymatic reactions of
metabolism form a network of
interconnected chemical
reactions, or pathways.
• The molecules of the pathway
are called intermediates
because the products of one
reaction become the
substrates of the next.
• Enzymes control the flow of
energy through a pathway.
Intermediary Metabolism

Stages of Metabolism
Figure 24.3
 Energy-containing nutrients are
processed in three major
stages:
 Digestion – breakdown of
food; nutrients are transported
to tissues
 Anabolism and formation of
catabolic intermediates where
nutrients are built into lipids,
proteins, and glycogen or
Broken down by catabolic
pathways to pyruvic acid and
acetyl CoA.
 Oxidative breakdown –
nutrients are catabolized to
carbon dioxide, water, and ATP

• A deficiency of even one
amino acid results in a
negative nitrogen balance.
In this state, more protein
is degraded than
synthesized.
 Histidine & arginine are
semi essential. They are
essential only for infants
growth, but not for old
children or adults where
in adults histidine
requirement is obtained
by intestinal flora &
arginine by urea cycle.

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