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Biochemistry Basics
Section 1.1
Subatomic Particles and the Atom
• Protons (+ charge) and
neutrons (neutral)
– found in the nucleus
• Electrons (- charge)
– Surround the nucleus in
a “cloud” or orbital
• Orbital
– the 3D space where an
electron is found 90% of
the time
– Each orbital can only fit
only 2 electrons
Bonding – Covalent Bonds
Hydrogen atoms (2 H)
• Atoms bond
through
interaction of their
valence (outer
orbital) electrons
• Covalent bond
– electrons are
shared between
atoms and the
valence orbitals
overlap
In each hydrogen
atom, the single electron
is held in its orbital by
its attraction to the
proton in the nucleus.
1
When two hydrogen
atoms approach each
other, the electron of
each atom is also
attracted to the proton
in the other nucleus.
2
3
The two electrons
become shared in a
covalent bond,
forming an H2
molecule.
+
+
+
+
+
+
Hydrogen
molecule (H2)
Name
(molecular
formula)
Water (H2O).
Two hydrogen
atoms and one
oxygen atom are
joined by covalent
bonds to produce a
molecule of water.
Methane (CH4).
Four hydrogen
atoms can satisfy
the valence of
one carbon
atom, forming
methane.
Electronshell
diagram
Structural
formula
O
H
H
H
H
C
H
H
Spacefilling
model
Ionic Bonds
• In some cases, atoms strip electrons away
from their bonding partners
• Ionic bond – electrons are transferred from
one atom to the other, resulting in a negative
ion (anion) and a positive ion (cation), which
are electrostatically attracted to each other
The lone valence electron of a sodium
atom is transferred to join the 7 valence
electrons of a chlorine atom.
Na
Na
Sodium atom
(an uncharged
atom)
Cl
Cl
Chlorine atom
(an uncharged
atom)
Each resulting ion has a completed
valence shell. An ionic bond can form
between the oppositely charged ions.
+
–
Na
Cl
Na+
Sodium on
(a cation)
Cl–
Chloride ion
(an anion)
Sodium chloride (NaCl)
• Covalent bonds are stronger than ionic bonds
• Covalent and Ionic bonds are intramolecular
forces of attraction because they are within
molecules
Polarity
• Electronegativity
– Is the attraction of an atom
for electrons
• The more electronegative
an atom
– The more strongly it pulls
electrons toward itself
• The smaller the atom
– the more electronegative
• to determine the type of bond between two atoms, calculate the
difference between their electronegativity values
=0
covalent
0 < x < 1.7
polar covalent
>= 1.7
ionic
(extreme polarity)
strong
electrons
shared equally
electrons
partially shared
weak
electrons not
shared
• the greater their difference in electronegativity, the greater the
polarity of that substance
• Polar Covalent Bond – electrons are shared unequally between
atoms of different electronegativity; electrons are closer to the
atom with the higher value
Because oxygen (O) is more electronegative than hydrogen (H),
shared electrons are pulled more toward oxygen.
This results in a
partial negative
charge on the
oxygen and a
partial positive
charge on
the hydrogens.
d–
O
d+
H
H
H2O
d+
Intermolecular Forces
• intermolecular forces of attraction exist
between molecules
• London forces
– form when the electrons of one molecule are
attracted to the positive nuclei of neighbouring
molecules; holds large nonpolar molecules
together; very weak
• hydrogen bonds
– form when the slightly negative O or N that is
bonded to a slightly positive H is attracted to the
slightly positive H of a neighbouring molecule;
strongest
Water
(H2O)
Hd +
d –O
H
d+
d–
Ammonia
(NH3)
N
H
d+
Figure 2.15
H
H
d+
A hydrogen
bond results
from the
attraction
between the
partial positive
charge on the
hydrogen atom
of water and
the partial
negative charge
on the nitrogen
atom of
ammonia.
• dipole-dipole forces
– form when the slightly negative end of a polar
molecule is attracted to the slightly positive end of
a neighbouring polar molecule; stronger
– Occurs because electrons are in constant motion
and may accumulate by chance on one part of the
molecule. The result is “hot spots” of positive and
negative charge.
Water
• highly polar because of asymmetrical shape
and polar covalent bond
• The polarity of water molecules results in
hydrogen boding
d–
Hydrogen
bonds
+
H
+
d–
+
Figure 3.2
d–
H
+
d–
“Like Dissolves Like”
• ionic compounds dissolve in water because
the ions separate
• However, molecules do not need to be ionic to
dissolve in water
• polar covalent molecules (eg: sugars, alcohols)
can dissolve in water, but large nonpolar
molecules (eg: oils) do not
• small nonpolar molecules (eg: O2, CO2) are
slightly soluble and need soluble protein
molecules to carry them (eg: hemoglobin
transports oxygen through the blood)
• hydrophilic – “water-loving;” dissolves in
water
– e.g. polar or ionic molecules, carbohydrates, salts
• hydrophobic – “water-fearing;” does not
dissolve in water
– e.g. non-polar molecules, lipids
Acids and Bases
• acid – donates H+ to water; pH 0-7
• base –donates OH- to water (or H3O); pH 7-14
• neutralization reaction – the reaction of an
acid and a base to produce water and a salt
(ionic compound)
Strong and Weak Acids/Bases
• strong acids and bases – ionize completely
when dissolved in water
– HCl(aq) (100% H3O+(aq))
– NaOH(aq) (100% OH-(aq))
• weak acids and bases – ionize only partially
when dissolved in water
– CH3COOH(aq) (1.3%  H3O+(aq))
– NH3(aq) (10%  OH-(aq))
Buffers
• The internal pH of most living cells must
remain close to pH 7
• Buffers
– Are substances that minimize changes in the
concentrations of hydrogen and hydroxide ions in
a solution
– Can donate H+ ions or remove H+ ions when
required
– E.g. carbonic acid creates bicarbonate ions (base)
and hydrogen ions (acid) (reversible reaction)
Functional Groups
• Functional groups
– Are reactive clusters of atoms attached to the carbon
backbone of organic molecules
Group
Chemical
Formula
Structural Formula
Found In
hydroxyl
—OH
alcohols
(eg: ethanol)
carboxyl
—COOH
acids (eg: vinegar)
amino
—NH2
bases
(eg: ammonia)
sulfhydryl
—SH
rubber
phosphate
—PO4
ATP
Carbonyl
(aldehydes) —COH
(keytones)
—CO—
aldehydes (eg:
formaldehyde)
ketones (eg:
acetone)
To Do
• Section 1.1 Questions
– Pg. 23 #1, 2, 4, 6-8, 12, 14, 15