Basic Biochemistry
CHE 242
MTWR 11:00 am – 12:35 pm
Julian Hall 225
Dr. Jon A. Friesen
Office: 318 Science Laboratory Building
phone: (43)8-7850
email: jfriese@ilstu.edu
What Do Biochemists Study?
The periodic table of the elements
Figure 1.1
The periodic table of the elements
Figure 1.1
Bulk elements
97%
The periodic table of the elements
Figure 1.1
Essential ions
The periodic table of the elements
Figure 1.1
Trace elements
Organic compounds in biochemistry
Figure 1.2
Functional groups in biochemistry
Figure 1.2
Linkages in biochemical compounds
Figure 1.2
Types of molecules in biochemistry
1. Proteins
Are composed of twenty different kinds of monomeric units,
the amino acids.
2. Polysaccharides (sugar)
Are constructed of monomeric units called monosaccharides.
Also called carbohydrates.
3. Nucleic acids (DNA and RNA)
Are synthesized from monomeric units called nucleotides.
4. Lipids (Fat)
Water insoluble molecule containing fatty acids.
Used for membrane structure and energy storage.
Energy Flow
Page 11
Prokaryotic Cells
Figure 1.16
Eukaryotic Cells
Figure 1.18
Water, Water Everywhere
Water is a Polar Molecule
Figure 2.1
Polarity of Small Molecules
Figure 2.2
Hydrogen Bonding Between Two Water Molecules
Figure 2.3
Water Can Form Up To Four Hydrogen Bonds
Figure 2.4
Water Molecules Form a Hexagonal Lattice in Ice
Figure 2.5
Sodium Chloride (NaCl) crystal
Figure 2.6
Ionic and Polar
Substances Dissolve
in Water
Example:
Dissolution of
Sodium Chloride
in water
Figure 2.6
Glucose, a sugar, contains polar groups, and is soluble in water
Nonpolar substances are relatively insoluble in water
Noncovalent interactions
in biomolecules
1. Charge-Charge Interactions
2. Hydrogen Bonds
3. Van der Waals Forces
4. Hydrophobic Interactions
Hydrogen bonding is
a common noncovalent
interaction between
biomolecules
Figure 2.11
Hydrogen bonding between bases in DNA
Figure 2.12
Van der Waals forces
are weak noncovalent
forces between atoms
Figure 2.13
Amphipathic molecules,
such as detergents,
have both a polar and
a nonpolar end.
Figure 2.9
Detergents can form monolayers at the air-water interface
Figure 2.10
Detergents can form micelles in aqueous solution
Figure 2.10
Ionization of Water
Water has a slight tendency to ionize
Pages 41 and 42
Strong acids completely dissociate in water.
Example: Hydrochloric acid (HCl)
Weak acids dissociate in water with a characteristic
acid dissociation constant (Ka).
Example: Acetic acid, present in vinegar
Relationship between pH and pKa
Henderson – Hasselbalch equation
Titration of acetic acid with aqueous base
Figure 2.17
Titration of
phosphoric acid,
a polyprotic
acid, with
aqueous base
Figure 2.19
1
2
3
4
5
6
7
8
9
10
1. Write the equilibrium reaction for the ionization of the weak acid.
2. What is the chemical structure of the conjugate base?
3. What is the pH of a solution containing equal amounts of the
weak acid and the conjugate base?
4. What is the pH of a solution containing 10 times more weak acid
than conjugate base?
5. What is the ratio of conjugate base to weak acid at pH = 7?
1. Write the equilibrium reaction for the ionization of the weak acid.
2. What is the chemical structure of the conjugate base?
1. Write the equilibrium reaction for the ionization of the weak acid.
2. What is the chemical structure of the conjugate base?
3. What is the pH of a solution containing equal amounts of the
weak acid and the conjugate base?
4.8
CH3COOCH3COOH
4.8
1
4.8
0
1. Write the equilibrium reaction for the ionization of the weak acid.
2. What is the chemical structure of the conjugate base?
3. What is the pH of a solution containing equal amounts of the
weak acid and the conjugate base?
4.8
CH3COOCH3COOH
4.8
1
4.8
0
4. What is the pH of a solution containing 10 times more weak acid
than conjugate base?
CH3COO-
4.8
CH3COOH
0.1
4.8
4.8
3.8
(-1)
5. What is the ratio of conjugate base to weak acid at pH = 7?
7
2.2
102.2
158
4.8
CH3COOCH3COOH
CH3COOCH3COOH
CH3COOCH3COOH
CH3COOCH3COOH
Titration of acetic acid with aqueous base
Figure 2.17
Titration of acetic acid with aqueous base
Figure 2.17
Buffering Region
1 pH unit from pKa
Maintenance of Blood pH
in Humans
CO2 – Bicarbonate Buffer System
Carbon dioxide – carbonic acid – bicarbonate buffer system
maintains blood pH at 7.4
Figure 2.21
Regulation of
blood pH
in mammals
Figure 2.22
Why is the CO2 – bicarbonate buffer
system used in the human body?
1. The raw materials (CO2 and H2O) for the production of
carbonic acid (H2CO3) are readily available.
2. The lungs and kidneys can easily adjust to ratio alterations
between carbonic acid (H2CO3) and the conjugate base
bicarbonate (HCO3-).
Role of the lungs and kidneys in
regulation of physiological pH
Lungs
Control the supply of H2CO3 in the blood by controlling the
amount of CO2 exhaled.
When the blood level of HCO3- decreases, the breathing rate is
increased, increasing amount of CO2 expelled, decreasing H2CO3.
If H2CO3 (CO2) increases it is called respiratory acidosis. If H2CO3
(CO2) decreases it is called respiratory alkalosis.
Kidneys
Control the concentration of HCO3-. If HCO3- is too high it is called
metabolic alkalosis. If HCO3- is too low, it is called metabolic alkalosis.
Blood Concentrations
Ratio of HCO3- : H2CO3 = 10 : 1  This results in pH = 7.4
HCO3- = 24 - 27 mEq/L (mM)
H2CO3 = 1.20 - 1.35 mEq/L (mM)
Clinicians often monitor blood pH, HCO3- and CO2 concentrations.
Non-graded Homework:
Use Henderson-Hasselbalch equation to convince yourself
this makes sense.
Problem #11 at the end of the chapter.