What is
Biochemistry?
Chemical Context of Life
Key Elements of Life
 92 naturally existing elements on Earth
 25 play a role in the chemical processes of life
 96 percent of all biological matter contain:




carbon (C)
oxygen (O)
hydrogen (H)
nitrogen (N)
 ALL organic molecules contain carbon
Concept check:
1. Explain why table salt is a compound, while the
oxygen we breath is not.
1. What four chemical elements are most abundant
in the food you ate yesterday?
Atomic structure
Stable
Stable
Stable
Stable
Stable
Unstable
Uses of Isotopes
 Radioisotopes are radioactive and break
down, emitting radiation, over time.
 Radioisotopes can be used to help understand
chemical and biological processes in
organisms.
 They can also be used in radiometric dating
which is useful in determining the age of
fossils
For Example:
A solution of phosphate,
containing radioactive
phosphorus-32, is injected into
the root system of a plant.
 Since phosphorus-32 behaves
identically to that of
phosphorus-31, the more
common form of the element, it
is used by the plant in the same
way.
 A Geiger counter is used to
detect the movement of the
radioactive phosphorus-32
throughout the plant.
 This information helps
scientists understand the
detailed mechanism of how
plants utilized phosphorus to
grow and reproduce.
For Example
 Brachytherapy is a form of
radiation therapy where
radioactive isotopes in the
form of small pellets (called
seeds) are inserted into
cancerous tumours to
destroy cancer cells while
reducing the exposure of
healthy tissue to radiation.
 It is currently approved for
treatment of prostate cancer
and cancers of the head and
neck. There are also studies
underway to see whether it
can be used in the treatment
of lung cancer.
Radioactive Decay:
Using the rate at which radioactive isotopes break
down to determine the age of a specimen
Uses of Stable
Isotopes
How Stable Isotopes Work
Examples of Uses
Chemical Bonding
Valence electrons are those electrons that are available for bonding.
The electrons in the outermost s and p orbitals
Chemical Bonding
Bonds are formed when:
 electrons are shared or exchanged between atoms.
1. Intermolecular Bonds
Bonds formed between molecules
2. Intramolecular Bonds
Bonds formed between atoms WITHIN a
molecule
Intramolecular Bonds:
1. covalent bonds
sharing of electrons between atoms
Atoms are shared between non-metals
Double and triple bonds can result if 2 or 3 pairs are
shared



polar covalent bonds
a)

unequal sharing of electrons between atoms
non-polar covalent bonds
b)

equal sharing of electrons between atoms
2. ionic bonds
electrostatic attraction between ions upon electron
transfer….occurs b/w a metal and a non-metal

 Na+ [ Cl

]

Cl

Na  +





Electronegativity
 nonpolar bond: electrons are shared equally
H 2,
Cl2:
• polar bond: electrons are shared unequally
because of the difference in electronegativity.
HCl:
Bond Polarity
A polar bond can be pictured using partial
charges:
+

H
Cl
2.1
 = 0.9
3.0
Electronegativity
Difference
Bond Type
0 - 0.5
Nonpolar
0.5 - 2.0
Polar
2.0 
Ionic
Intermolecular Bonds
 weak bonds
1. van der Waals interactions

attraction between nearby molecules

only occur when atoms and molecules are very close
together
2. dipole-dipole interactions

attraction of + and - ends of polar molecules
3. hydrogen bonding (H-bonding)

special type of dipole-dipole interaction

generally occur when hydrogen bonded to O, N
•Hydrogen bonds: strongest intermolecular bonds
Homework:
1. Rd. Chapter 1.1 – focus on Water and Acid and
Bases (pg. 16 – 22) (skip molecular shape pg. 15)
2. Answer questions 4, 7-10, 12-15.
Chemical Bonding cont:
Orbital's: volumes of space around the nucleus where
electrons are most likely found
Electrons are so small that they are not able to be exactly
located. Scientists have determined spaces around the nucleus
where electrons are most likely found.
Ionic Compounds
Ionic compounds consist of a lattice of positive and
negative ions.
NaCl:
Double and Triple Bonds
 Atoms can share four electrons to form a double bond or
six electrons to form a triple bond.


N 2:
N N

O
=O



O 2:
• The number of electron pairs is the
bond order.
Electronegativity is a
measure of an atom's ability
to attract a shared electron
pair when it is participating
in a covalent bond with
another atom
By calculating the net
Electronegativity of the
atoms, we can determine the
distribution of electrons and
the nature of the molecule
The Pauling scale of
electronegativity
Electronegativity
 Number assigned to each element (En)
 Difference in En helps determine the nature
of the bond: ionic or covalent
 Covalent bonds can be either nonpolar
covalent or polar covalent –
 This difference has Biological
Consequences
MOLECULAR SHAPE
 When atoms form a covalent bond, the
valence electron pairs arrange themselves to
be as far away from each other as possible.
 This change to the orientation of the
valence electrons is Hybridization.
 Symmetry of shape will also determine
polarity of molecule or functional group.
•Methane: CH4 – tetrahedral shape – equal 109.5 angle
between valence electrons; symmetrical; nonpolar
Ammonia: NH3 – pyramidal shape – equal 107 angle between
valence electrons; asymetrical shape – polar molecule
To determine if a molecule is polar or
nonpolar, must consider both
electronegativity and shape
Which of the following is the essential
characteristic of a polar molecule?
a.Contains double or triple bonds
b.is formed at extremely low temperatures
c.contains ions as part of the structure
d.has an asymmetrical distribution of electrical
charge
e.contains the element oxygen
Van der Waals forces:
•intermolecular forces of attraction
Examples: London forces, dipole-dipole forces, hydrogen bonds
London forces
- weak/temporary/
random charges
-gases at room
temperature
- volatile e.g.; octane
Dipole-dipole
-between polar
molecules
-stronger
- e.g.; HCl