B.sc. biochemistry sem 1 introduction to biochemistry unit 1 foundation of biochemistry

Foundation of Biochemistry
Course: B.Sc. Biochemistry
Sub: introduction to biochemistry
Unit 1

• Fifteen to twenty billion years ago, the universe arose as a
cataclysmic eruption of hot, energy-rich subatomic particles.
Within seconds, the simplest elements (hydrogen and helium)
were formed. As the universe expanded and cooled, material
condensed under the influence of gravity to form stars. Some
stars became enormous and then exploded as supernovae,
releasing the energy needed to fuse simpler atomic nuclei
into the more complex elements. Thus were produced, over
billions of years, the Earth itself and the chemical elements
found on the Earth today. About four billion years ago, life
arose—simple microorganisms with the ability to extract
energy from organic compounds or from sunlight, which they
used to make a vast array of more complex biomolecules
from the simple elements and compounds on the Earth’s
surface.

Biochemistry asks how the remarkable properties
of living organisms arise from the thousands of
different lifeless biomolecules. When these
molecules are isolated and examined individually,
they conform to all the physical and chemical laws
that describe the behavior of inanimate matter—as
do all the processes occurring in living organisms.
The study of biochemistry shows how the
collections of inanimate molecules that constitute
living organisms interact to maintain and
perpetuate life animated solely by the physical and
chemical laws that govern the nonliving universe.

• A high degree of chemical
complexity and microscopic
organization. Thousands of
different molecules make up a cell’s
intricate internal structures. Each
has its characteristic sequence of
subunits, its unique three-
dimensional structure, and its highly
specific selection of binding partners
in the cell.
• Systems for extracting,
transforming, and using energy
from the environment, enabling
organisms to build and maintain
their intricate structures and to do
mechanical, chemical, osmotic, and
electrical work. Inanimate matter
tends, rather, to decay toward a
more disordered state, to come to
equilibrium with its surroundings.

• A capacity for precise self-replication and self-assembly. A
single bacterial cell placed in a sterile nutrient medium can
give rise to a billion identical “daughter” cells in 24 hours.
• Each cell contains thousands of different molecules, some
extremely complex; yet each bacterium is a faithful copy of
the original, its construction directed entirely from
information contained within the genetic material of the
original cell.
• Mechanisms for sensing and responding to alterations in
their surroundings, constantly adjusting to these changes
by adapting their internal chemistry.
• Defined functions for each of their components and
regulated interactions among them.

A history of evolutionary change
• Organisms change their inherited life strategies to survive
in new circumstances. The result of eons of evolution is an
enormous diversity of life forms, superficially very different
but fundamentally related through their shared ancestry.
• Despite these common properties, and the fundamental
unity of life they reveal, very few generalizations about
living organisms are absolutely correct for every organism
under every condition; there is enormous diversity.
• The range of habitats in which organisms live, from hot
springs to Arctic tundra, from animal intestines to college
dormitories, is matched by a correspondingly wide range of
specific biochemical adaptations, achieved within a
common chemical framework.

• Biochemistry describes in molecular terms the
structures, mechanisms, and chemical processes
shared by all organisms and provides organizing
principles that underlie life in all its diverse forms,
principles we refer to collectively as the
molecular logic of life. Although biochemistry
provides important insights and practical
applications in medicine, agriculture, nutrition,
and industry, its ultimate concern is with the
wonder of life itself.

B.sc. biochemistry sem 1 introduction to biochemistry unit 1 foundation of biochemistry

Cellular Foundations
• The unity and diversity of organisms become
apparent even at the cellular level. The smallest
organisms consist of single cells and are
microscopic. Larger, multicellular organisms
contain many different types of cells, which vary
in size, shape, and specialized function. Despite
these obvious differences, all cells of the simplest
and most complex organisms share certain
fundamental properties, which can be seen at the
biochemical level.

• Cells Are the Structural and Functional Units of All Living
Organisms Cells of all kinds share certain structural features. The
plasma membrane defines the periphery of the cell, separating its
contents from the surroundings. It is composed of lipid and protein
molecules that form a thin, tough, pliable, hydrophobic barrier
around the cell. The membrane is a barrier to the free passage of
inorganic ions and most other charged or polar compounds.
• Transport proteins in the plasma membrane allow the passage of
certain ions and molecules; receptor proteins transmit signals into
the cell; and membrane enzymes participate in some reaction
pathways. Because the individual lipids and proteins of the plasma
membrane are not covalently linked, the entire structure is
remarkably flexible, allowing changes in the shape and size of the
cell. As a cell grows, newly made lipid and protein molecules are
inserted into its plasma membrane; cell division produces two cells,
each with its own membrane. This growth and cell division (fission)
occurs without loss of membrane integrity.

• All cells have, for at least some part of their life, either
a nucleus or a nucleoid, in which the genome the
complete set of genes, composed of DNA—is stored
and replicated. The nucleoid, in bacteria, is not
separated from the cytoplasm by a membrane; the
nucleus, in higher organisms, consists of nuclear
material enclosed within a double membrane, the
nuclear envelope.
• Cells with nuclear envelopes are called eukaryotes
(Greek eu, “true,” and karyon, “nucleus”); those
without nuclear envelopes—bacterial cells—are
prokaryotes (Greek pro, “before”).

• There Are Three Distinct Domains of Life
• All living organisms fall into one of three large groups
(kingdoms, or domains) that define three branches of
evolution from a common progenitor. Two large groups of
prokaryotes can be distinguished on biochemical grounds:
archaebacteria (Greek arche-, “origin”) and eubacteria
(again, from Greek eu, “true”).
• Eubacteria inhabit soils, surface waters, and the tissues of
other living or decaying organisms. Most of the well studied
bacteria, including Escherichia coli, are eubacteria. The
archaebacteria, more recently discovered, are less well
characterized biochemically; most inhabit extreme
environments—salt lakes, hot springs, highly acidic bogs,
and the ocean depths.

Three Distinct Domains of Life

• Cells Build Supramolecular Structures
• Macromolecules and their monomeric subunits differ
greatly in size. A molecule of alanine is less than 0.5 nm
long. Hemoglobin, the oxygen-carrying protein of
erythrocytes (red blood cells), consists of nearly 600
amino acid subunits in four long chains, folded into
globular shapes and associated in a structure 5.5 nm in
diameter. In turn, proteins are much smaller than
ribosomes (about 20 nm in diameter), which are in
turn much smaller than organelles such as
mitochondria, typically 1,000 nm in diameter. It is a
long jump from simple biomolecules to cellular
structures that can be seen with the light microscope.

Structural hierarchy in the molecular
organization of cells

Biomolecules Are Compounds of Carbon with a
Variety of Functional Groups
• The chemistry of living organisms
is organized around carbon,
which accounts for more than
half the dry weight of cells.
Carbon can form single bonds
with hydrogen atoms, and both
single and double bonds with
oxygen and nitrogen atoms. Of
greatest significance in biology is
the ability of carbon atoms to
form very stable carbon–carbon
single bonds. Each carbon atom
can form single bonds with up to
four other carbon atoms. Two
carbon atoms also can share two
(or three) electron pairs, thus
forming double (or triple) bonds.

 Because of its bonding versatility, carbon can produce a broad array
of carbon–carbon skeletons with a variety of functional groups;
these groups give biomolecules their biological and chemical
personalities.
 A nearly universal set of several hundred small molecules is found
in living cells; the interconversions of these molecules in the central
metabolic pathways have been conserved in evolution.
 Proteins and nucleic acids are linear polymers of simple monomeric
subunits; their sequences contain the information that gives each
molecule its three-dimensional structure and its biological
functions.
 Molecular configuration can be changed only by breaking covalent
bonds. For a carbon atom with four different substituents (a chiral
carbon), the substituent groups can be arranged in two different
ways, generating stereoisomers with distinct properties. Only one
stereoisomer is biologically active. Molecular conformation is the
position of atoms in space that can be changed by rotation about
single bonds, without breaking covalent bonds.
 Interactions between biological molecules are almost invariably
stereospecific: they require a complementary match between the
interacting molecules.

O
H H
1 molecule of water is
made up of 2 hydrogen atoms
bonded with 1 oxygen atom. The molecular
formula is H2O
STRUCTURE OF WATER

O
STRUCTURE OF WATER
The bond that forms water
is a covalent bond

• Water consists of an
oxygen atom bound to
two hydrogen atoms by
two single covalent bonds.
– Oxygen has unpaired &
paired electrons which
gives it a slightly
negative charge while
Hydrogen has no
unpaired electrons and
shares all others with
Oxygen
– Leaves molecule with
positively and negative
charged ends
Water is a Polar Molecule
-has oppositely charged ends

Properties of Water
Polar Molecule
Cohesion And Adhesion
High Specific Heat
Density – Greatest At
4oc
Universal Solvent Of
Life
Capillary Action
Surface Tension
Buoyancy

Polarity of Water
• In a water molecule two hydrogen atoms
form single polar covalent bonds with an
oxygen atom. Gives water more structure
than other liquids
– Because oxygen is more
electronegative, the region around
oxygen has a partial negative
charge.
– The region near the two hydrogen
atoms has a partial positive charge.
• A water molecule is a polar molecule with
opposite ends of the molecule with opposite
charges.

24
slightly positive
charge
slightly negative charge
hydrogen bond between
(+) and (-) areas of
different water molecules
Water molecules form Hydrogen bonds

– Water has a variety of unusual properties because
of attractions between these polar molecules.
– The slightly negative regions of one molecule are
attracted to the slightly positive regions of nearby
molecules, forming a hydrogen bond.
– Each water molecule
can form hydrogen
bonds with up to
four neighbors.

“Universal” Solvent
• A liquid that is a completely homogeneous mixture of
two or more substances is called a solution.
– A sugar cube in a glass of water will eventually
dissolve to form a uniform mixture of sugar and water.
• The dissolving agent is the solvent and the substance
that is dissolved is the solute.
– In our example, water is the solvent and sugar the
solute.
• In an aqueous solution, water is the solvent.
• Water is not really a universal solvent, but it is very
versatile because of the polarity of water molecules.

• Water is an effective
solvent as it can form
hydrogen bonds.
– Water clings to polar
molecules causing
them to be soluble in
water.
• Hydrophilic -
attracted to water
– Water tends to
exclude non polar
molecules.
• Hydrophobic -
repelled by water

Acids
• Acids dissociate in water to increase the
concentration of H+.
– Have many H+ ions
– Sour taste
– HCl is hydrochloric acid or stomach acid

Bases
• Bases combine with H+ ions when dissolved in
water, thus decreasing H+ concentration.
– Have many OH- (hydroxide) ions
– Bitter taste
– NaOH = sodium hydroxide or baking soda

Acids and Bases
• An acid is a substance that
increases the hydrogen ion
concentration in a solution.
• Any substance that reduces the
hydrogen ion concentration in a
solution is a base.
– Some bases reduce H+ directly by
accepting hydrogen ions.
• Strong acids and bases complete
dissociate in water.
• Weak acids and bases dissociate
only partially and reversibly.
1

• The structure of any molecule is a unique and
specific aspect of its identity.
• Molecular structure reaches its pinnacle in the
intricate complexity of biological
macromolecules, particularly the proteins.
Although proteins are linear sequences of
covalently linked amino acids, the course of the
protein chain can turn, fold, and coil in the three
dimensions of space to establish a specific, highly
ordered architecture that is an identifying
characteristic of the given protein molecule

Weak Forces Maintain Biological Structure
and Determine Biomolecular Interactions
• Covalent bonds hold atoms together so that molecules are formed.
In contrast, weak chemical forces or noncovalent bonds,
(hydrogen bonds, van der Waals forces, ionic interactions, and
hydrophobic interactions) are intramolecular or intermolecular
attractions between atoms. None of these forces, which typically
range from 4 to 30 kJ/mol, are strong enough to bind free atoms
together. The average kinetic energy of molecules at 25°C is 2.5
kJ/mol, so the energy of weak forces is only several times greater
than the dissociating tendency due to thermal motion of molecules.
Thus, these weak forces create interactions that are constantly
forming and breaking at physiological temperature, unless by
cumulative number they impart stability to the structures
generated by their collective action.

•Van der Waals
• Van der Waals forces are the result of induced
electrical interactions between closely approaching
atoms or molecules as their negatively-charged
electron clouds fluctuate instantaneously in time.
These fluctuations allow attractions to occur between
the positively charged nuclei and the electrons of
nearby atoms.
• Hydrogen Bonds
• Hydrogen bonds form between a hydrogen atom
covalently bonded to an electronegative atom (such
as oxygen or nitrogen) and a second electronegative
atom that serves as the hydrogen bond acceptor.

• Ionic Interactions
• Ionic interactions are the result of attractive
forces between oppositely charged polar
functions, such as negative carboxyl groups and
positive amino groups. These electrostatic forces
average about 20 kJ/mol in aqueous solutions.
Typically, the electrical charge is radially
distributed, and so these interactions may lack
the directionality of hydrogen bonds or the
precise fit of van der Waals interactions.

Atom – the smallest unit of matter “indivisible”
Helium
atom

electron shells
a) Atomic number = number of Electrons
b) Electrons vary in the amount of energy
they possess, and they occur at certain
energy levels or electron shells.
c) Electron shells determine how an atom
behaves when it encounters other atoms

Electrons are placed in shells according
to rules:
1) The 1st shell can hold up to two electrons,
and each shell thereafter can hold up to 8
electrons.

Octet Rule = atoms tend to gain, lose or share electrons so as to have 8 electrons
C would like to
N would like to
O would like to
Gain 4 electrons
Gain 3 electrons
Gain 2 electrons

Why are electrons important?
1) Elements have different electron
configurations
 different electron configurations mean different
levels of bonding

Electron Dot Structures
Symbols of atoms with dots to represent the valence-shell
electrons
1 2 13 14 15 16 17 18
H He:
          
Li Be  B   C   N   O  : F  :Ne :
       
          
Na Mg  Al  Si  P S :Cl  :Ar :
       

Chemical bonds: an attempt to fill electron shells
1. Ionic bonds –
2. Covalent bonds –
3. Metallic bonds

IONIC BOND
bond formed between
two ions by the
transfer of electrons

Formation of Ions from Metals
 Ionic compounds result when metals react with
nonmetals
 Metals lose electrons to match the number of valence
electrons of their nearest noble gas
 Positive ions form when the number of electrons are
less than the number of protons
Group 1 metals  ion 1+
Group 2 metals  ion 2+
• Group 13 metals  ion 3+

Formation of Sodium Ion
Sodium atom Sodium ion
Na  – e  Na +
2-8-1 2-8 ( = Ne)
11 p+ 11 p+
11 e- 10 e-
0 1+

Formation of Magnesium Ion
Magnesium atom Magnesium ion

Mg  – 2e  Mg2+
2-8-2 2-8 (=Ne)
12 p+ 12 p+
12 e- 10 e-
0 2+

Some Typical Ions with Positive
Charges (Cations)
Group 1 Group 2 Group 13
H+ Mg2+ Al3+
Li+ Ca2+
Na+ Sr2+
K+ Ba2+

Ions from Nonmetal Ions
In ionic compounds, nonmetals in 15, 16, and 17
gain electrons from metals
Nonmetal add electrons to achieve the octet
arrangement
Nonmetal ionic charge:
3-, 2-, or 1-

Fluoride Ion
unpaired electron octet
    1 -
: F  + e : F :
   
2-7 2-8 (= Ne)
9 p+ 9 p+
9 e- 10 e-
0 1 -
ionic charge

Ionic Bond
• Between atoms of metals and nonmetals with
very different electronegativity
• Bond formed by transfer of electrons
• Produce charged ions all states. Conductors
and have high melting point.
• Examples; NaCl, CaCl2, K2O

Ionic Bonds: One Big Greedy Thief Dog!

1). Ionic bond – electron from Na is transferred to Cl,
this causes a charge imbalance in each atom. The Na
becomes (Na+) and the Cl becomes (Cl-), charged
particles or ions.
4

COVALENT BOND
bond formed by the
sharing of electrons

Covalent Bond
• Between nonmetallic elements of similar
electronegativity.
• Formed by sharing electron pairs
• Stable non-ionizing particles, they are not
conductors at any state
• Examples; O2, CO2, C2H6, H2O, SiC

Bonds in all the
polyatomic ions
and diatomics are
all covalent
bonds

when electrons are
shared equally
NONPOLAR
COVALENT BONDS
H2 or Cl2

when electrons are
shared but shared
unequally
POLAR COVALENT
BONDS
H2O

Polar Covalent Bonds: Unevenly matched,
but willing to share.

- water is a polar molecule because oxygen is more electronegative than hydrogen, and
therefore electrons are pulled closer to oxygen.
6

METALLIC BOND
bond found in
metals; holds metal
atoms together
very strongly

Metallic Bond
• Formed between atoms of metallic elements
• Electron cloud around atoms
• Good conductors at all states, lustrous, very
high melting points
• Examples; Na, Fe, Al, Au, Co

Metallic Bonds: Mellow dogs with plenty
of bones to go around.

Ionic Bond, A Sea of Electrons
7

Metals Form Alloys
Metals do not combine with metals. They form
Alloys which is a solution of a metal in a metal.
Examples are steel, brass, bronze and pewter.

Formula Weights
• Formula weight is the sum of the atomic
masses.
• Example- CO2
• Mass, C + O + O
12.011 + 15.994 + 15.994
43.999

References/Sources
• All images are from Lehninger Principles of biochemistry by Nelson and Cox except
1.https://lh4.ggpht.com/ApC7AoJPJITtO2ClFL_HQnGYBiN7W0K2qwIw8lVJJq-
iaSO4qwZ86u7YrskOad2mmojw=s85
2. https://lh6.ggpht.com/p8Y8LOzbxD1vfv_3y6lxRvz0-42-
Y8ayo7ilfE1gUg4wia77gXThu05L6zCZJohN_Qvb1nQ=s93
3. https://lh4.ggpht.com/6qrTKLqGVSMBk3YS3AtZwReykIFGb7mjw--
wePD6IGe1T7ayJ8keL78slq_JCPeEKlgcUA=s139
4. https://lh6.ggpht.com/eb-7TKOD0yyVDojl2jeF5avaObFxQ5Tl69TzOU4v5Hb_bGctE94GS_1SLoPS-
nKRoWVEcQ=s16
5. https://lh3.ggpht.com/d4km1D6g2jIPIc5teiq7Ef9qix-eh-
zVQyHNfbhJ1tMVtMTDTrP9HG5MKOBKZk_OC49ioWw=s116
6. https://lh3.ggpht.com/qxM-Stmq8qzblrCavlKJ99zDaFqn3txD3iqpHMp1-
mkRU14yNN0viCE7qqp5TJQi4yC2=s88
7. https://lh6.ggpht.com/0Xz3DEdGbquMEoXNxolqT6B__qGJghymupo2z-
caEj81y24_fc6U1F2g_1uXaE8ZhDzc=s125
Books/ Web resources
• Lehninger Principles of biochemistry by Nelson and Cox
• https://www.chem.wisc.edu/deptfiles/genchem/sstutorial/.../tx71.html

B.sc. biochemistry sem 1 introduction to biochemistry unit 1 foundation of biochemistry

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B.sc. biochemistry sem 1 introduction to biochemistry unit 1 foundation of biochemistry