Important
words and concepts from Chapter 2, Campbell & Reece, 2002 (1/14/2005):
by Stephen T. Abedon (abedon.1@osu.edu)
for Biology 113 at the Ohio State University
|
|
Course-external links are
in brackets Click [index] to access site index Click here to access
text’s website Vocabulary
words
are found below |
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(1) Chapter title: The Chemical Context of
Life
(a)
Found at this site (i.e., at www.phage.org) are
additional pages of possibly related interest including: [fundamentals of chemistry]
[history
of earth] [origin of the universe]
(b)
[chemical context of life
(Google Search)] [index]
(a)
In this chapter we will consider the chemical context of life
(b)
To those students interested in going on in biology it makes sense to
get as good a background in chemistry as you can; this lecture will be based as
much on why a given concept is important to your understanding of biology as on
teaching the concept (obviously this is not a chemistry class/I am not a
chemistry instructor so there is only so much that can be taught)
(c)
Note how chemistry (and physics) pervades biology, representing a
dominant subject over chapters 3, 4, 5,
6,
9,
10,
11
and 26, plus this chapter of your text
(d)
This lecture will consider
(i)
The stuff you are made of (elements)
(ii)
Energy
(iii)
Chemical bonding
(iv)
That, in biology, form consistently follows
function
(vi)
(for those of you with some interest, I will additionally include links
to material considering the history of chemistry – and no, not the
history of chemists that you learn
about in chemistry class – but, briefly, the history of the chemicals and,
especially, the elements themselves)
(e)
Note: though we won’t be covering chapter 2 in its entirety in this
lecture, there nevertheless are a number of concepts covered in the text chapter
that are worth either learning or reminding yourself of, because these concepts
are considered in subsequent chapters; it’s a short chapter (12 pages with lots
of pictures), so do yourself a favor and read the whole thing
(f)
Note: a number of you have not completed a college chemistry sequence
so will not have taken enough college chemistry to fully grasp all of the
chemistry that we will be covering in this course (i.e., that is covered by
your text) – translation: don’t assume that you know enough chemistry to blow
off the earlier chapters of this text (and to those who have taken their
majors’ chemistry sequence as well as some or all of organic chemistry: Enjoy!)
(g)
[chemistry for biology
(Google Search)] [index]
CHEMICALS
(a)
The essential elements are what organisms are made of
(b)
The most prevalent elements in your body (96%) are Carbon, Oxygen, Hydrogen, Nitrogen (essentially the elements
that make up water plus the organic compounds that together make up the bulk
of organisms)
(c)
An additional 7 elements make of the bulk of the remaining 4% (Ca, P,
K, S, Na, Cl, Mg—no need to memorize); the rest of the elements found in your
body are considered trace elements
(d)
Additional elements, called trace elements, are
found in smaller amounts but are nevertheless essential to continued and
healthful existence (understand concept, don’t memorize list found in text);
these are many of the minerals found in your daily multivitamin and mineral supplements
(e)
See Table 2.1, Naturally
occurring elements in the human body
(f)
["essential elements"
and chemistry (Google Search)] [index]
CHEMICAL BONDS
(4)
Energy (waterfall
analogy)
(a)
Energy = capacity to do work
(b)
Potential energy = stored energy
(c)
Waterfall analogy:
(i)
Water at the top of a waterfall has potential energy which is realized
(as kinetic energy) when the water flows over the falls
(ii)
It is possible to capture some of that energy, transducing (changing)
it into a different form (e.g., mechanical energy if you place a turbine or
water wheel in the path of the flowing water, or electrical energy if you
attach a generator to the turbine)
(d)
The transduction of energy from one form to another, and the use of
that energy to move or to create complex structures (particularly babies) is
what life is all about
(e)
Bioenergetics is the study of the movement of
energy within living organisms
(f)
[energy and "capacity to
do work", waterfall and energy
(Google Search)] [index]
(a)
Within organisms, stored energy is associated with the electrons
found within biomolecules
(b)
More energy associated with an electron = greater distance that
electron is found from the nucleus the electron is associated with
(c)
Note that this concept is analogous to the waterfall analogy:
(i)
Water that is further from the center of the Earth possesses more
potential energy than water that is closer to the center of the Earth;
electrons that are further from the center of the nucleus they are associated
with possess more energy than electrons that are closer to the center of the
nucleus they are associated with
(ii)
During the movement of water toward the center of the Earth the water
possesses/releases energy that may be captured; during the movement of
electrons toward the nucleus of an atom, energy is also released (and this
energy may also be captured)
(iii)
Movement of water away from the center of the Earth requires an input
of energy; movement of electrons away from the center of a nucleus also
requires an input of energy
(d)
See Figure 2.9, Energy
levels of an atom’s electrons
(e)
Less simplistically, electrons actually exist within specific energy
levels or electron shells which exist as probabilistic clouds surrounding
nuclei and it is the quantum changes in the shape of these clouds that correspond
to changes in the energy associated with an electron (bigger cloud = more
energy, smaller cloud = less energy associated with that electron; thus, the
collapse of a larger cloud to a smaller cloud is associated with a release of
energy, etc.)
(f)
Quantum mechanics is (at least in part) the study of the discrete
storage of energy by electrons
(g)
An important part of understanding life is understanding how energy is
stored and moved from molecule to molecule (in fact, bioenergetics along with cell biology, genetics, evolution, and ecology arguably
are the five most important general concepts you will learn in introductory biology)
(h)
["energy storage" and
"chemical bonds" (Google Search)] [index]
(a)
When a chemical reaction occurs, what is happening is the making or
breaking (or both) of chemical bonds
(b)
Chemical bonds consist of electrons that are shared, more-or-less,
between the nuclei of the bonded atoms
(c)
Chemical bonds come in a variety of types that may be characterized in
terms of the shapes of the probabilistic clouds and the related concept of the
degree to which the electrons are shared evenly between the atoms (or otherwise
“hogged” by one atom relative to another)
(d)
The degree of sharing impacts on the energy associated with a chemical
bond (and living things store most of their energy
within chemical bonds); greater “hogging” by one atom relative to another
results in a decrease in the electron’s distance from an atomic nuclei
(relatively so, at least) and therefore a decrease in the amount of energy
stored by the electron; the degree of sharing also impacts on the strength of
the chemical bond
(e)
Chemical bonds vary in strength ranging from very strong to very weak
(i)
Covalent bonds (strong)
(ii)
Polar covalent bonds (strong)
(iii)
Ionic bonds (weaker, at least within the aqueous environments
found in organisms, where ions are surrounded by hydration shells)
(iv)
Hydrogen bonds (weak)
(f)
FAQ: What do you mean by "Energy in bonds"? When
electrons are locked into chemical bonds, there is a certain amount of energy
associated with those electrons. This is the (chemically available) energy that
exists within, for example, the food you eat. Recall that the farther an
electron is from the atomic nucleus, the more energy it contains. This distance
from an atomic nucleus can be locked into an electron when that electron is
locked into a chemical bond. Indeed, one can think of the energy required to
drive forward the endergonic dehydration synthesis reaction as energy that
becomes trapped in chemical bonds and associated with electrons that are now farther
from atomic nuclei than they otherwise might be (in fact, were). Finally, note
that all else held constant, an electron that is shared between two atoms
possessing relatively equal electronegativity will be trapped at a further
distance from the two atomic nuclei than an atom locked between two atoms
having dissimilar electronegativities. For example, an electron found between H
and O will be much closer to an atomic nuclei (i.e., that of O) than an
electron found between C and C, or even O and O.
(g)
[chemical bonds (Google Search)] [index]
(7) Representing chemical bonds
(a)
We will be looking at a number of structural formulas of molecules in
which chemical bonds are shown (we will be showing a lot of these
especially as we introduce carbohydrates, lipids, proteins, and nucleic acids)
(b)
(as an example I will draw glucose on board/show as overhead [glucose
model])
(c)
We will have a tendency in this course to not explicitly show C-H bonds
(d)
We will less-frequently employ molecular formulas, i.e.,
representations in which atoms are listed; e.g., C6H12O6
, a.k.a., C6(H2O)6 (a.k.a., a hexose, e.g.,
glucose)
(e)
Single, double, and triple covalent bonds will represent as in the following
examples: C-C, C-H, C=O, CºC, NºN (not all H’s shown)
(a)
Valence electrons are electrons found in the outer shells of elements
(b)
Knowledge of valence electrons is fundamentally helpful for
understanding chemistry in general, and the chemistry of organic molecules in
particular (and understanding how organic molecules work is fundamental to understanding
how life works (as we will see in chapter 4)
(c)
See Figure 2.10, Electron
configurations of the first 18 elements
(d)
Note the valence of the following atoms: Hydrogen = 1, Oxygen = 2, Nitrogen = 3, Carbon = 4, Phosphorus = 5
(note, though, that phosphorous is weird, not always following the octet rule)
(e)
See Figure 2.12, Covalent
bonding in four molecules
(f)
FAQ: What is the difference between valence, valence electrons, and
a valence shell? Valence electrons are indeed the electrons found in the
outer shell of an atom. This outer shell is referred to as the valence shell.
The valence of an atom, however, is it's bonding capacity. For example: Carbon
has 4 valence electrons, 4 unpaired electrons, and a valence of 4. Hydrogen has
1 valence electron, 1 unpaired electron, and a valence of 1. Oxygen has 6
valence electrons, 2 unpaired electrons, and a valence of 2. Nitrogen has 5
valence electrons, 3 unpaired electrons, and a valence of 3. Phosphorus also
has 5 valence electrons and 3 unpaired electrons. However, in the phosphate ion
it is actually exhibiting a valence of 5 since it forms a total of five bonds
with four molecules of oxygen. The important take home message is simply that
hydrogen tends to form 1 bond, oxygen 2, nitrogen 3, carbon 4, and phosphorus,
in the phosphate ion, 5.
(g)
[valence electrons (Google Search)] [index]
(a)
Covalent bonds are the strongest of bonds
(b)
Covalent bonds involve a sharing of electrons between atoms
(c)
Example: C-C (carbon-to-carbon) bonds that form the basis of most
biomolecules [carbon and the molecular diversity of life, the
structure and function of macromolecules (MicroDude)] [index]
(d)
Covalent bond in which electrons are somewhat evenly shared (e.g., C-C,
C-H, O=O) are important for understanding hydrophobicity as
well as the structure of most organic molecules
(e)
See Figure 2.12, Covalent
bonding in four molecules
(f)
[covalent bonds (Google Search)] [index]
(a)
The concept of elctronegativity is important for understanding
properties of water, polarity, hydrogen bonding, etc.
(b)
Electrons show a greater attraction for atoms (or ions) that display
greater electronegativity
(c)
If two atoms are chemically bonded together, then the atom
with the greater electronegativity will pull the electron associated with that
bond closer to it (i.e., it will “hog”
the electron)
(d)
Note that such bonds, consequently, will have less energy associated
with them than an otherwise equivalent bond in which the electrons are shared
evenly between the two atoms
(e)
FAQ: How can I derive the electronegativity values from my
understanding of the periodic table of elements? Going from left to right
in the periodic table, atoms increasingly fill their outer shell while also
gaining an increased nuclear charge. The increase in the volume of their outer
shell is not as great as their increase in charge because they are filling
equivalent outer shells rather than forming new ones. This means that electrons
similarly distant from the nucleus are exposed to a nucleus with a greater
positive charge. In fact, far from gaining in size, atoms actually decrease
significantly in size going from left to right on the periodic table. Electrons
consequently are not only held more tightly, the nucleus possesses an increased
propensity to attract additional electrons. The Column 1 elements display the
least electronegativity because they have nuclei with the least positive charge
in their row. This means that they tend to readily lose their single electron
(which also serves to complete their now outer shell). In contrast, Fluorine
displays the most electronegativity, readily filling its outer shell at the
expense of other atoms. The exception to these rules is Hydrogen, which is a
column 1 element but which also possesses comparatively significant
electronegativity. This exceptional behavior results from hydrogen only
possessing only a single electron and only a single proton. That is, though
hydrogen’s nuclear charge is small, nevertheless its electrons are held
relatively close to its nucleus. Furthermore, hydrogen, like Carbon, Nitrogen,
Oxygen, etc. but unlike other column 1 elements, can complete it's outer shell
by gaining only a single electron. Finally, as you go down in columns in the
periodic table, elements become less electronegative. This is due to the
increasing size of the outer electron shell.
(f)
[electronegativity (Google Search)] [index]
(a)
Within an aqueous environment, polar covalent bonds are intermediate in
strength between ionic and covalent bonds
(b)
Polar covalent bonds result when electrons are not shared equally
between atoms
(c)
Example: the O-H (oxygen-to-hydrogen) bonds in water
(d)
Example: the N-H (nitrogen-to-hydrogen) bonds found in nucleic acids and proteins
(e)
Polar covalent bonds are important for understanding hydrogen bonding (as well as the structure of
most organic molecules)
(f)
See Figure 2.13, Polar
covalent bonds in a water molecule
(g)
FAQ: What is the definition of a nonpolar covalent bond?
Covalent (as opposed to ionic) bonds between atoms of similar
electronegativity. The most important nonpolar covalent bonds we have talked
about are C-C bonds and C-H bonds. These, by the way, are also the bonds
associated with reduced (as opposed to oxidized) carbon. Any covalent bond
(i.e., a bond in which electrons are shared between two or more atoms) between
two atoms of similar electronegativity are considered non-polar. With ionic
bonds, electrons are not shared between the two contributing atoms.
Polar-covalent bonds lie somewhere between these two extremes.
(h)
FAQ: How do you know that the bond is polar or nonpolar? Does it
have to do with the elements location on the Periodic Table or is it something
else? I am able to understand that the C-C bond is nonpolar because they are
the same element, but the C-H covalent bond being nonpolar has really confused
me. Can you explain why it is a nonpolar bond? You are correct that it is
easy to understand that carbon has the same electronegativity as itself (as
does H to itself, O to itself, N to itself, ect.). To judge degrees of polarity
you have to know what the electronegativity of the two bonded atoms are. These
values are typically not found on periodic tables though a generalization may
be made: The greater the number of valence electrons, the higher the
electronegativity. In addition, the lower the atomic weight (i.e., going up
columns in the periodic table) the higher the electronegativity. Hydrogen turns
out to be somewhat exceptional, possessing a much higher electrogativity than
the other column 1 elements. This probably has to do with hydrogen possessing
only a single electron, and requiring only a single additional electron to fill
its outer electron shell (recall that hydrogen requires only 2 electrons to
fill its sole electron shell). It turns out that the C-H covalent bond is
indeed slightly polar. However, because of the unusually high electronegativity
of hydrogen, C and H have sufficiently similar electronegativity that the
polarity of the C-H bond falls on the nonpolar end of the continuum. In fact,
there is sufficiently low polarity in this bond that Van der Waal's interactions
between molecules containing numerous C-H bonds are greater than the hydrogen
bonding capability of the H in the C-H bond. This contributes to hydrophobic
exclusion, which we'll consider in more detail when we consider water. I picked
up a random inorganic chemistry text (that is, an intro chemistry text)
and was able to find a table of electronegativity values which are given on
something called a Pauling Scale. The higher the number, the greater the
electronegativity. Here's a sample of values: H = 2.2, C = 2.6, N = 3.1, O =
3.5, F = 4.0, Cl = 3.2, P = 2.2, Na = 0.9, K = 0.8. An immediate observation
would be that the electronegativity difference between C and H (=0.4) is not
exactly trivial. However, the difference between O and C or N and C ranges from
a little more than the C to H difference (N-C; =0.5) to more than twice the
difference (O-C; =0.9). Furthermore, the difference between O and H or N and H
are even greater (=1.7 and =0.9, respectively). Consequently, the C-O, C=0,
O-H, and N-H bonds are considerably more polar than the C-H bond. We call the
former “polar covalent bonds,” and we lump the latter (i.e., C-H) among the
at-best weakly polar bonds, which for our purposes act more non-polar-like than
polar-like. Note that even highly polar but still covalent bonds (e.g., C-O)
only fall about half way on the continuum between truly non-polar covalent
bonds (e.g., C-C) and the extremely polar ionic bonds (e.g., Na-Cl). The
take-home message regardless is that we will be lumping together C-H and C-C
bonds as more or less non-polar with C-N, C-O, C=O, N-H, and especially O-H
considered polar covalent bonds which are capable of participating in hydrogen
bonding. Why this is important will become more obvious as we consider water
and then the various biological molecules.
(i)
[polar covalent bonds
(Google Search)] [index]
(a)
Ionic bonds involve less (often much less) sharing of electrons between
atoms
(b)
Ionic bonds result from one atom essentially giving an electron to
another atom
(c)
Ionic bonds are typically not as strong as covalent bonds (at least as
found in an aqueous solution)
(d)
Example: Na-Cl (sodium-to-chlorine) bonds in table salt
(e)
Ionic bonds represent an extreme of polarity and are represented in
biological systems as the salt bridges within proteins, etc.
(many biomolecules are salts at physiological pHs and therefore
capable of forming ionic bonds)
(f)
See Figure 2.14, Electron transfer and ionic bonding
(g)
FAQ: Are ionic bonds polar or
are they nonpolar? If you
think about it, regular ordinary bonds range in their polarity from complete
sharing of electrons (i.e., non-polar covalent bond) to the complete donation
of an electron by one atom to a second atom (i.e., an ionic bond). If it is
only partial donation (due to sufficient differences in electronegativity) then
we might call that bond a polar covalent bond. Therefore, increasing polarity
is observed with increasing donation of electrons, and ionic bonds represent an
extreme example of electron donation. Another way of thinking about this is
that, with a polar covalent bond, one of the atoms takes on a partial negative
charge and the other atom takes on a partial positive charge. From the
existence of these partial charges we infer polarity in the bond (i.e., the
electron is held more closely by one atom than it is by the other) and we would
describe this bond as polar covalent. In an ionic bond the respective atoms
take on not just a partial charge but a full charge. Hence, such bonds
are very polar. So polar, in fact, that we don't even refer to them as
covalently bonded (since covalent bonding implies a sharing of
electrons).
(h)
[ionic bonds (Google Search)] [are ionic bonds polar or are they nonpolar? (MicroDude)] [index]
(a)
Relatively weak chemical bonds are a necessary requirement for
chemical-based dynamic system (such as living things) just as precisely
machined and well lubricated parts are important for mechanical-based dynamic
systems (e.g., a bicycle)
(b)
This is because living things are constantly making and breaking
chemical bonds
(c)
Often the energy required to make or break these bonds may be supplied
solely by the ambient heat of the environment (e.g., your body temperature)
(d)
If chemical bonds were universally too strong, then this making and
breaking of bonds would require too much energy for life to exist (i.e., we
would all be rocks)
(e)
The hydrogen bond is the signature example of
a weak chemical bond that plays numerous and important roles in biological
systems (and which we will consider in much more depth during our water
lecture – pardon the pun)
(f)
(“Covalent bonding alone cannot begin to describe the complexity of
molecular structure in biology. Much weaker interactions are responsible for
most the elegant cellular architecture… These are the noncovalent interactions, also called noncovalent forces or noncovalent
bonds, between ions, molecules, and parts of molecules. ¶ Consider macromolecules…
The linear sequence of the atoms in a strand of DNA is maintained by the
covalent bonds between them. But DNA also has a highly specific
three-dimensional structure, which is stabilized by noncovalent interactions
between different parts of the molecule. Similarly, every kind of protein is
made up of covalently linked amino acids but is also folded into a specific
molecular conformation by noncovalent forces. Proteins interact with other
protein molecules or with DNA to form still more complex structures. All of
this complexity is accounted for by a myriad of noncovalent interactions within
and between macromolecules. Moving up a step in the organization of life, we
note that the cytoplasm of a cell is itself a highly organized structure, also
held together for the most part by noncovalent forces. ¶ What
makes noncovalent interactions so important in biology and biochemistry? The
key is [that] [b]iologically important noncovalent
bonds are 10 to 100 times weaker [than the covalent bonds between carbon and
hydrogen]. It is their very weakness that makes noncovalent bonds so essential, for it allows them to be
continually broken and re-formed in the dynamic molecular interplay that is
life. This interplay depends on rapid exchanges of molecular partners, which
could not occur if intermolecular forces were so strong as to lock the
molecules in conformation and in place.” p. 25, Christopher K. Mathews and K.
E. Van Holde, 1996, Biochemistry,
Second Edition, Benjamin/Cummings Publishing Company)
(g)
[weak chemical bonds,
noncovalent interactions,
noncovalent forces,
noncovalent bonds (Google Search)] [index]
(a)
Hydrogen bonds are both covalent-bond-like
and ionic-bond-like but nevertheless are very weak
(b)
Hydrogen bonds are a consequence of one atom in one molecule (or
different part of the same molecule) having too much charge (due to
participation in a polar covalent bond) and a second atom having too little
charge (ditto)
(c)
These polar covalent bonds give
atoms partial charges and partially charged atoms attract other oppositely
partially charged atoms
(d)
Example: O-H···O-H where the dotted line represents a
hydrogen bond between a hydrogen (italicized) and an oxygen atom (also
italicized)
(i)
the former (H) has a partial positive charge – a partial loss of electron to
the polar covalently bonded, not-italicized oxygen
to the left
(ii)
the latter (O) has a partial
negative charge – a partial gain of an electron from the polar covalently bonded, not italicized hydrogen to the right)
(e)
[hydrogen bonds (Google Search)] [hydrogen bonding, hydrogen bond (MicroDude)] [index]
(f)
See Figure 2.16, A hydrogen
bond
(g)
See Figure 3.1, Hydrogen
bonds between water molecules
CHEMICAL STRUCTURE