ph and buffers

1. pH, Buffer, Maintenance of blood pH and pH of
gastric juice
Presented by
SALMAN KHAN

2. WHAT DOES PH MEAN?
• Measure of hydrogen ion concentration.
• Measure of the acidity or alkalinity of a solution.
• pH scale ranges - 0 to 14.
• Aqueous solutions at 25°C with a pH less than 7
are acidic.
• pH greater than 7 are basic or alkaline.
• pH level of 7.0 is neutral because the concentration of
H3O+ equals the concentration of OH− in pure water.

3. 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."

4. ACIDS AND BASES
• Bronsted-Lowry concept
• Acid- donate proton
• Base- accept proton
• When a proton donor loses a proton it becomes the corresponding proton acceptor.
• This pair makes up conjugate acid-base pair.

5. BUFFER
• A buffer solution is an aqueous solution consisting of a mixture of a weak acid and its
conjugate base, or vice versa. Its pH changes very little when a small amount of
strong acid or base is added to it.
• Buffer solutions are used as a means of keeping pH at a nearly constant value in a
wide variety of chemical applications. In nature, there are many systems that use
buffering for pH regulation like the bicarbonate buffering system is used to regulate
the pH of blood.
• Examples: Biochemical Assays- Enzyme activity depends on pH, so the pH during an
enzyme assay must stay constant.
• BUFFERING CAPACITY: The efficiency of a buffer in maintaining a constant pH on
addition of acid or base is referred to as buffering capacity.

6. PRINCIPLE
• Buffer solutions achieve their resistance to pH change because of the presence of
an equilibrium between the weak acid HA and its conjugate base A−:
HA H+ + A−
• When strong acid is added to a mixture of the weak acid and its conjugate base,
hydrogen ions (H+) are added. The A- combines with the H+ and makes HA
therefore removes H+ from the solution.
(OR)
• When strong base OH- is added to the mixture. The HA protonates the oxygen of
OH- turning that into water and also A- is formed. In this the HA removes the OH-
from the solution. Therefore pH remains constant.
OH- + HA → H2O + A−

7. TYPES OF BUFFER
• ACIDIC BUFFERS
• BASIC BUFFERS
Generally buffers are of two types

8. ACIDIC BUFFERS
• An acidic buffer is a combination of weak acid and its salt with a strong base.
I.e., weak acid and salt with strong base (conjugate base).
• Examples,
CH3COOH/CH3COONa
H2CO3/NaHCO3
H3PO4/NaH2PO4
HCOOH/HCOONa

9. MECHANISM OF ACTION OF ACIDIC
BUFFERS
• Consider a buffer system of CH3COOH (Weak electrolyte) and CH3COONa (Strong electrolyte). There will be a large
concentration of Na+ ions, CH3COO – ions, and un- dissociated CH3COOH molecules.
• When an acid is added
• If a strong acid (HCl) is added in CH3COOH / CH3COONa buffer, the changes that will occur may be represented as:
• The hydrogen ions yielded by the HCl are quickly removed as unionized acetic acid, and the hydrogen ion
concentration is therefore only slightly affected (because acetic acid is produced is very weak as compared to HCl
added).

10. • When a base is added
• If a strong base (NaOH) is added in CH3COOH / CH3COONa buffer, the changes that will occur may be represented
as:
• The hydroxyl ions yielded by the NaOH are therefore removed as water. The supply of hydrogen ions needed for this
purpose being constantly provided by the dissociation of acetic acid.
2

11. BASIC BUFFERS
• A basic buffer is a combination of weak base and its salt with a strong acid. i.e. Weak
base & salt with strong acid (conjugate acid).
• EXAMPLES:
NH4OH / NH4Cl
NH3 / NH4Cl
NH3 / (NH4)2CO3

12. MECHANISM OF ACTION OF BASIC
BUFFERS
• Consider a buffer system of NH4OH (Weak electrolyte) and NH4Cl (Strong electrolyte). There will be a large
concentration of NH4
+ ions, Cl– ions, and un- dissociated NH4OH molecules.
• When an acid is added
• If a strong acid (HCl) is added in NH4OH / NH4Cl buffer, the changes that will occur may be represented as:
• The hydrogen ions yielded by the HCl are therefore removed as water. The supply of OH- ions needed for this is
constantly provided by the ammonium hydroxide.
2

13. PREPARING BUFFER SOLUTIONS
HANDERSON HASSELBALCH EQUATION
• Determines the exact amount of acid and
conjugate base needed to make a buffer of
a certain pH, using the Henderson-Hassel
Bach equation:
pH= pKa + log (
[A−]
[HA]
)
where pH is the concentration of [H+], pKa is
the acid dissociation constant, and [ log{A-}]
and [ log{HA}] are concentrations of the
conjugate base and starting acid.

14. IMPORTANT BUFFERS IN LIVING
SYSTEMS
• Buffering is important in living systems as a means of
maintaining a fairly constant internal environment, also
known as homeostasis.
• Small molecules such as bicarbonate and phosphate
provide buffering capacity as do other substances, such
as hemoglobin and other proteins.

15. BUFFERING SYSTEM OF BLOOD
• Maintaining a constant blood pH is critical for
the proper functioning of our body.
• The buffer that maintains the pH of human
blood involves a carbonic acid (H2CO3) -
bicarbonate ion (HCO3
-) system.
• The bicarbonate buffer system is an effective
physiological buffer near pH 7.4, because the
H2CO3 of blood plasma is in equilibrium with a
large reserve capacity of CO2(g) in the air space
of the lungs.

16. • This buffer system involves three reversible equilibria between gaseous CO2 in the
lungs and bicarbonate (HCO3
- ) in the blood plasma.
• When H+ (from lactic acid produced in muscle tissue during vigorous exercise, for
example) is added to blood as it passes through the tissues, reaction 1 proceeds
toward a new equilibrium, in which the concentration of H2CO3 is increased.

17. • This increases the concentration of CO2(d) in the blood plasma (reaction 2) and thus increases the
pressure of CO2(g) in the air space of the lungs (reaction 3); the extra CO2 is exhaled.
• Conversely, when the pH of blood plasma is raised (by NH3 production during protein catabolism, for
example), the opposite events occur:
• The H+ concentration of blood plasma is lowered, causing more H2CO3 to dissociate into H+ and HCO3
-
. This in turn causes more CO2(g) from the lungs to dissolve in the blood plasma. The rate of
breathing—that is, the rate of inhaling and exhaling CO2—can quickly adjust these equilibria to keep
the blood pH nearly constant.

18. PH OF GASTRIC JUICES
• Many aspects of cell structure and function are
influenced by pH, it is the catalytic activity of enzymes
that is especially sensitive.
• Enzymes typically show maximal catalytic activity at a
characteristic pH, called the pH optimum.
• On either side of the optimum pH their catalytic activity
often declines sharply. Thus, a small change in pH can
make a large difference in the rate of some crucial
enzyme-catalyzed reactions.

19. • ▪ Gastric secretion is a colorless, watery,
acidic, digestive fluid produced in the
stomach
• Physical properties;
• Watery fluid, that has a pale yellow color
• pH is 1-3 ,
• Volume secreted per day is 2-3 L .
• The stomach secretes acid (HCl), which is
important either to the digestive process or
to control of gastric function.
• Parietal cells: Secrete HCl into the stomach
lumen where it establishes an extremely
acidic environment. This acid is important
for activation of pepsinogen and inactivation
of ingested microorganisms such as
bacteria.
• Chief cells: It secrets pepsinogen(zymogen).
Once secreted, pepsinogen is activated by
stomach acid into the active protease
pepsin.

20. REFERENCES
• Lehninger, A. L., Nelson, D. L., & Cox, M. M. (2000). Lehninger
principles of biochemistry. New York: Worth Publishers.
• Text book of medical biochemistry by MN chaterjea , Rana Shinde.
• https://courses.lumenlearning.com/introchem/chapter/ph-buffers-acids-
and-bases/
• https://bio.libretexts.org/Bookshelves/Biotechnology/Lab_Manual%3A_I
ntroduction_to_Biotechnology/01%3A_Techniques/1.07%3A_pH_and_Bu
ffers

21. MCQs
1.If a decinormal solution of NaOH is added in a mixture of weak base and its strong salt then in the following condition which option is
correct?
a) Very high change in OH– ions
b) High change in OH– ions
c) Slight change in OH– ions
d) No change in OH– ions
2. Which of the following is not a simple buffer?
a) CH3COONH4
b) NH4CN
c) H3PO4 + NaH2PO4
d) (NH4)2 CO3
3.Which of the following is not a type of Basic buffer mixture?
a) NH4OH
b) NH4Cl
c) H2CO3+Na2CO3
d) Glycine + Glycine hydrochloride
4. What is the H+ ion concentration in pure water?
a) 1×107 m
b) 1×10-7 m
c) 1×1014 m
d) 1×10-14 m

22. 5.The pH of buffer solution depends upon concentration of?
a) Strong acid
b) Strong base
c) Weak acid
d) Salt
6. . The pH can be kept constant with the help of?
a) Saturated solution
b) Unsaturated solution
c) Buffer solution
d) Non-Saturated solution
7. A solution was prepared by dissolving 0.02 moles of acetic acid (HOAc; pKa = 4.8) in water to give 1 liter of solution. What is the pH?
a) 3.00
b) 3.05
c) 3.15
d) 3.25
8. Buffers usually contain ________________ with its conjugate ____________
a) weak base, base
b) strong base, acid
c) weak acid, base
d) weak acid, acid

23. 9.Carbonic acid and bicarbonate ions buffer which of the following?
a) Cytosol
b) Cytoplasm
c) Blood
d) Lymph
10. In presence of an acid, amino group can be ____________
a) Polarized
b) Washed away
c) Protonated
d) Replaced