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Lecture

BIOL 112 Lecture Notes Part 1.doc


Department
Biology
Course Code
BIOL 112
Professor
Joseph Dent

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BIOL 112/Winter 2011/Lecture Notes Part I
Small molecules
The below is mostly review of basic chemistry
Atomic structure: protons, neutrons, electrons
Mass of proton = standard unit of measure called the dalton OR amu (1 amu is roughly 1.7 x 10-24 g)
Matter: anything that occupies space, and has mass (which is a quantity of matter, circular definitions
ftw)
Inertia: resistance to change of state of motion
Classification of matter:
Separable by physical means = mixture
Homogeneous (uniform throughout, e.g. salt solution) or heterogeneous (not, e.g.
milk - fat in a water-based medium)
Not separable by physical means = substance
Compound (decomposable by chemical process) or element (not)
Atomic mass = mass of neutrons, protons and electrons in amu (atomic mass units)
Neutrons and protons have mass 1, electrons have a mass of something like 0.0005
Note that atomic mass is different from atomic weight - atomic weight is the weighted average
of the atomic mass of the isotopes of an element
Relative atomic mass is a synonym for atomic weight, and this is rarely a whole number
Also note that neutrons and protons don't have exactly the same mass; we approximate
Furthermore, due to binding energy ...
But in any case the lesson here is that atomic mass and weight are rarely whole numbers for
several reasons
The explanation in the lecture itself is a bit misleading, and doesn't really account for the
difference between atomic mass and weight
The periodic table, classifies elements in groups and families, primarily
Isotopes: different number of neutrons, but identical chemical properties (determined by protons and
neutrons)
However, due to the number of neutrons, can be unstable (i.e. radioactive - emit alpha, beta
or gamma radiation from the nucleus)
Chemical bonding: Aufbau principle, electrons fill up s orbitals before p orbitals

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Atoms can form bonds until their outermost shell is filled
Covalent bonds - sharing of electrons (to some degree), strong, predictable lengths and
angles, can be double or TRIPLE
Polar covalent bonds - when electrons are more attracted to one nucleus than the
other, e.g. with water
This results in a dipole
On the other hand, in a C-C bond, for instance, the electrons are shared quite equally
Ionic bonds - attraction of opposite charges (basically covalent but electronegativity difference
is greater, usually >1.8 - not in scope of course)
You'll learn, if you haven't already, that a lot of explanations in basic chem/bio courses are
simplifications (i.e. lies)
van der Waal's forces: interactions of electrons of nonpolar substances
Hydrophobic interactions: nonpolar substances in the presence of polar substances
Bond energy: energy needed to separate two "bonded" atoms under physiological conditions
(?)
Electronegativity: tendency of an atom to attract electrons (when it occurs as part of a compound)
Increases as atomic radius decreases and nuclear charge increases; ionization energy is
similar, but slightly different
Large molecules
Acids and bases
Acids donate hydronium ions (H+, simplified as hydronium ions)
Strong acids vs. weak acids - strong = release many H+ ions (simplified as, they release all their ions),
weak = only some
Depends on pKa value; like covalent/ionic bonding, there is a continuum, but just pretend that what is
said in this class is true
Bases accept H+, release OH-
Ex, NaOH is a strong base
Amino group (NH2): important part of many biological compounds (principal component of protein),
weak base
Water: weak base AND acid, pH of pure water is 7, but almost all water is acidic due to dissolved ions,
gases (e.g. CO2) etc
pH is concentration of hydrogen, measured in logarithms - pH of 7 ==> 10-7 mol / L

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Buffers: make the solution resistant to pH change (reacts with added bases and acids)
Can be composed of a weak acid and its conjugate base. The actual reactions are probably
out of the scope of this course
Buffers can have different ranges over which they act
But once that range has been overshot, they no longer act as buffers
The relation of functional groups to acid/base theory:
Proteins are all amino acids - they have an amine group so they can accept, carboxyl group
so they can donate
Others: ketones (O in the middle), aldehydes, etc
Side chains - ex, phosphates can add charges, sulfhydryl groups can combine, etc
Phosphate groups can store energy, release it; that's why they're part of the energy currency
(ATP)
Isomerism
Isomers: same chemical formula, different geometric structure
Structural: different structural formula; a functional group is attached to a different carbon
atom
Stereoisomers: same chemical formula, different arrangement of atoms in space
Optical isomers (enantiomers), chiral compounds, mirror images of each other
Geometric isomers, ex cis & trans isomers (different sides of a double bond, which
cannot rotate)
Components of a cell
IN A CELL, WE HAVE:
Lipids: good at forming barriers and compartments, as they resist the flow of aqueous
solutions
Nucleic acids and proteins
Carbohydrates: energy storage, surface properties, rigid structure
We are 70% water - most of the rest is biological polymers etc
WE ARE ALL MADE OF THE SAME FUNCTIONAL UNITS
Condensation reactions - one way to make a polymer
Monomers, H + OH --> H2O, then we get polymers
Hydrolysis: breaking a polymer, consumes a molecule of water for every bond broken
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