Chapter 2
QUICKIE BASIC
CHEMISTRY (Review)
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Structure and function of all living things are governed by the
laws of chemistry
Understanding the basic principles of chemistry will give you a
better understanding of all living things and how they function!
“We are all connected; To each other, biologically. To the earth,
chemically. To the rest of the universe atomically.”
― Neil deGrasse Tyson
QUESTION: What examples can you give of how chemistry is
involved in biology?
• Photosynthesis
– Is an example of a
chemical reaction
CHEMISTRY- The science of the composition, structure, properties,
and reactions of matter
CHEM REVIEW
• Matter• Mass• Is MASS the same thing as WEIGHT?
– - The pull of gravity on an object is what gives an object its
"weight"
– - A bowling ball having the same mass on Earth and the
moon would weigh less on the moon than it does on Earth
due to a gravitational pull present on the Moon.
– How much do you weigh on other planets??? Find out
here: http://www.exploratorium.edu/ronh/weight/
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States of Matter
• SOLID: molecules are tightly linked; little movement and
definite shape
• LIQUID: molecules are less tightly linked; moves more freely
than solids; conforms to container
• GAS: molecules are usually not attracted to one another;
move very fast; fills the entire volume of a container
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Elements
• Pure substances that
cannot be broken
down into simpler
substances
• Periodic Table
• created in 1869 by
Mendeleev
• categorizes elements
and shows trends
• 118 elements, 92
occurring naturally
•Noble or "inert" gases on far right
•Alkali metals on far left (minus H)
•Atomic radii increase from left to right and
decrease from top to bottom
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Biologists love CHONPS most of all!
• 96% of all living matter: Sulfur, Phosphorus, Oxygen,
Nitrogen, Carbon and Hydrogen (CHONPS).
• From these six elements, you can make almost any
combination of organic molecules!
–
–
–
–
Carbs
Lipids
Nucleic acids
Amino acids, proteins
• Living organisms still need 14 other elements, but in
smaller amounts.
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Trace elements: Are required by an organism in only minute quantities
Make up the remaining 4% of living matter
Table 2.1
• The effects of essential element deficiencies
Figure 2.3
(a) Nitrogen deficiency
(b) Iodine deficiency
Atoms
• Simplest part of an element that retains all properties of that element
• Too small to see so we make up models to help us understand the structure
of atoms and predict how they will act
• Nucleus• Electrons (-) orbit around nucleus; very fast
• Farther an E- is from the nucleus- More ENERGY!
• Atomic Mass• Atomic Number• # of electrons = # of protons
• - Electrons balance out + charge of protons and have very little mass
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2.1 From atoms to molecules
Isotopes
• Isotopes are atoms that Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
have the same atomic
number but a different
atomic mass because the
number of neutrons differ
larynx
• Examples: 14C/12C, 127I, 131I
thyroid gland
trachea
• Radioactive isotopes are
useful in dating old objects,
imaging body organs and
tissues through x-rays and
killing cancer cells
a.
• Radiation can be harmful by
b.
damaging cells and DNA
and/or causing cancer a: © Biomed Commun./Custom Medical Stock Photo; b(patient): Courtesy National Institutes of Health
(NIH); b(brain scan): © Mazzlota et al./Photo Researchers, Inc.
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Compounds
• Most elements do not exist by themselves in nature
but rather like to combine with other elements
• A molecule is formed when two or more atoms join
together chemically.
• A compound is a molecule that contains at least two
different elements. All compounds are molecules but
not all molecules are compounds.
– Chemical properties of compounds are different than the
elements alone (2 gases = liquid. Ex: Na + Cl = table salt)
• “Chemical rxn’s”- chemical bonds can be broken, atoms
can be rearranged, and new chemical bonds are
formed!
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Bonding
•
•
Most bonding takes place because atoms are
most chemically stable when their
outermost energy levels are filled
“The Octet Rule”
Valence electrons
• Outer shell electrons are called the valence
electrons and are the ones involved in all the
bonding and chemical reactions of each
atom.
• Knowing the valence of an electron will give
you a better idea of it’s bonding properties
• To find out how many valence electrons an
atom has, you can do a shell diagram (2, 8,
8, 18) or a dot diagram (shows only valence)
•
http://www.green-planet-solar-energy.com/electron-dotstructure.html
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Bonding
• Most bonding takes place because atoms are
most chemically stable when their outermost
energy levels are filled
Types of Chemical Bonds
• Covalent bonds:
– Strong bonds
– Shared electrons, simulate a full outer orbital
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Polar covalent bond:
• Electrons are shared, but not evenly shared
This creates further potential for hydrogen bonds to form between molecules
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The Significance of Weak Chemical Bonds
• Several types of weak chemical bonds are
important in living systems
• Weak chemical bonds
– Reinforce the shapes of large molecules
– Help molecules adhere to each other
– Hydrogen bonds are such bonds
Hydrogen Bonds
• Relatively weak, singly, but rather strong
collectively
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Van der Waals Interactions
• Van der Waals interactions
– Occur when transiently positive and negative
regions of molecules attract each other
• Molecular shape
– Determines how biological molecules recognize and
respond to one another with specificity
Molecular
Shape often
determines
function
Carbon
Nitrogen
Hydrogen
Sulfur
Oxygen
Natural
endorphin
Morphine
(a) Structures of endorphin and morphine. The boxed portion of the endorphin molecule (left) binds to
receptor molecules on target cells in the brain. The boxed portion of the morphine molecule is a close match.
Natural
endorphin
Brain cell
Figure 2.17
Morphine
Endorphin
receptors
(b) Binding to endorphin receptors. Endorphin receptors on the surface of a brain cell
recognize and can bind to both endorphin and morphine.
Ionic Bonds
• Strong bonds
• Create ions
• Create electrical charges: when + and – charges
attract, ionic bonds are created
• OXIDATION: Na becomes (Na+) = loses e• REDUCTION: Cl becomes (Cl-)= gains an e-
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Energy
• Basic definition?
• Laws of Thermodynamics.
• Energy can neither be created or destroyed but it can be
changed from form to form
• Many energy forms are important in biology:
– chemical energy
– thermal energy
– radiant energy
– electrical energy
– mechanical energy
• Amount of energy in the universe remains the same over time
– Energy = usable energy + dissipates (ex: heat)
• Free energy ΔG: energy in a system that is available for work.
– For example, in a cell, it is the energy available to fuel cell
processes (growth, cell division, metabolism, etc)
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Energy and Chemical Reactions
• Exergonic rxn’s = release energy (products have less
chemical energy than reactants)
ex: AB + CD  AC + DB + energy
• Endergonic rxn’s = absorb energy (products have more
chemical energy than reactants)
ex: AB + CD + energy  AC + DB
• Activation energy: energy added to reactants to
"jumpstart" the rxn
• Catalysts: reduce the amount of activation energy that
is needed to start the rxn. See Figure 2-7 in your book.
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The energetics of chemical
reactions
•
•
Think about rolling a boulder up a hill…it takes
energy input, right? And the boulder at the
top of the hill now has that energy stored in it
(potential). So the product ends up having
MORE energy than the reactant.
Think about a boulder rolling down a hill…it has
lots of energy. And the boulder at the bottom
of the hill now has less (potential) energy. So
the product ends up having LESS energy than
the reactant.
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Solutions
• Water is extremely important to all living things, so the chemistry of
living things often involves the study of solutions!
• Solution: mixture in which substances are uniformly distributed in
another substance
– Solutions can be mixtures of liquids, solids, or gases
• Solute:
• Solvent:
• Concentration [ ]: measurement of the amount of solute dissolved in a
fixed amount of solvent
• Saturated?
• Aqueous solutions:
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Acids and Bases
Dissociation: force between H2O molecules is so
H2O
strong that the O- atom from one H2O molecule
can pull off the H+ atom from another molecule
• Water dissociates into H+ and OH- equally
(hydrogen and hydroxide)
• A Hydronium ion (H3O+)can form when a free H+
ion can react with another H2O molecule
H+ + OH-
• Acidity and Alkalinity is a measure of the
relative amount of OH- and H+ ions in a solution!
• Pure water has equal OH- and H+ ions in
solution; pH of 7.0
• Acidic solutions have H+ > OH- ions
– pH is below 7.0
– sour
• Basic (alkaline) solutions have H+ < OH- ions
– pH is above 7.0
– slippery and bitter
pH= measure of how many H+ ions are in a solution
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Buffers
• …are chemical substances that neutralize
small amounts of either an acid or a base
• Most chemical rxn’s in living organisms are controlled
by pH, therefore...
• Buffers are very important for maintenance of
homeostasis.
– If blood pH drops below 7.0 (acidosis), it could be fatal
– If blood pH goes above 7.7 (alkalosis), it could be fatal
• If our blood did not contain a buffering system
(bicarbonate ion), we would not be able to drink and
eat acidic/basic foods!
• Great chem review:
http://scidiv.bcc.ctc.edu/rkr/Biology101/lectures/pdfs/Chemistry101.pdf
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