Chemistry for Biology: Basics

Stephen T. Abedon, Ph.D.

These lecture notes were inspired by the Campbell Biology for Majors series. They were used when I had taught major’s biology at The Ohio State University (Bio 113 and Bio 114). I have slightly updated them but keep in mind that they have not been either used or updated in earnest since the early-mid 2000s.

These notes over the following terms: Chemical bonds, Chemical equilibria, Chemical history, Covalent bonds, Electronegativity, Electrons store energy, Energy, Essential elements, Form follows function, Hydrogen bonds, Ionic bonds, Metabolism, Polar covalent bond, Product, Reactant, Reversible reactions, Valence electrons, Weak bonds.

This document was last updated on Friday, March 02, 2012

(1) Introduction

(a)                  In this chapter we 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 going to be a complete overview everything that is chemistry and I am not a chemistry instructor, so there is only so much that can be taught J)

(c)                   This lecture considers:

(i)                    The stuff you are made of (elements)

(ii)                 Energy

(iii)               Chemical bonding

(iv)               That, in biology, form consistently follows function

(v)                 Chemical equilibrium

(vi)               For those of you with some interest, I will additionally discuss 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

(d)                  Note that a number of you have not completed a college chemistry sequence so will not have taken enough college chemistry to necessarily grasp all of the chemistry that is covered in an introductory biology course; as a consequence, don’t assume that you know enough chemistry to blow off various introductions to what chemistry is all about as found in your typical introductory biology textbook




(2) Essential elements

(a)                  The essential elements are what organisms are made of

(b)                  The most prevalent elements in your body (96%) are carbon, oxygen, hydrogen, and 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 up the bulk of the remaining 4% (Ca, P, K, S, Na, Cl, Mg); 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; these are many of the minerals found in your daily multivitamin and mineral supplements




(3) 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

(4) Electrons store energy

(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 that they are associated with possess more energy than electrons that are closer to the nucleus

(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)                  Less simplistically, electrons actually exist within specific energy levels or electron shells, which exist as probabilistic clouds surrounding atomic 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.)

(e)                   Quantum mechanics is (at least in part) the study of the discrete storage of energy by electrons

(f)                    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)

(5) Chemical bonds

(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 their 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 its atomic nucleus, the more energy it contains. This distance from an atomic nucleus can be locked into an electron, more or less, 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 at the same time it becomes associated with electrons that are now farther from atomic nuclei than they otherwise might be (in fact, otherwise were). Finally, note that all else held constant, an electron that is shared between two atoms that possess relatively equal electronegativity will be trapped at a further distance from the two atomic nuclei than an electron 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, i.e., oxygen) than an electron found between C and C, or even O and O.

(6) 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)                  Note that there is a tendency in representing organic molecules to not explicitly show C-H bonds

(c)                   In biology versus chemistry there is less of a tendency to 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)

(d)                  Single, double, and triple covalent bonds will represented as in the following examples: C-C, C-H, C=O, CºC, NºN (note that not all H's shown)

(7) Valence electrons

(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)

(c)                   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)

(d)                  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 consequently is referred to as the valence shell. The valence of an atom, however, is its 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. In the phosphate ion, however, phosphorus is actually exhibiting a valence of 5 since it forms a total of five bonds with four atoms 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.

(8) Covalent bonds

(a)                  Covalent bonds are, particularly in aqueous solutions, 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

(d)                 Covalent bonds 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

(9) Electronegativity

(a)                  The concept of electronegativity 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 its 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.

(10) Polar covalent bond

(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)                    FAQ: What is the definition of a nonpolar covalent bond? Covalent bonds, as opposed to ionic bonds, form 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) that occurs between two atoms of similar electronegativity are considered to be non-polar. With ionic bonds, electrons are not shared between the two contributing atoms. Polar covalent bonds lie somewhere between these two extremes.

(g)                  FAQ: How do you know that the bond is polar or nonpolar? Does it have to do with the element’s 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, etc.). 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 number (i.e., going up columns in the periodic table) the higher the electronegativity. Hydrogen turns out to be somewhat exceptional, possessing a much higher electronegativity 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=O, 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 nonpolar 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.

(11) Ionic bonds

(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)                    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., nonpolar covalent bonds) 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).

(12) Weak bonds (noncovalent interactions, noncovalent forces, noncovalent bonds)

(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)

(13) Hydrogen bonds

(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)




(14) Form follows function

(a)                  Molecule shape is very important for the functioning of organisms, just as the shape of parts is very important for the proper functioning of a bicycle

(b)                  The shape and chemistry of a biomolecule together define its form, which in turn helps define its function; change the shape or chemistry of a molecule, even subtly, and typically you will make that molecule nonfunctional

(c)                   A wonderful example of the role of form in function in living things is the cell membrane

(d)                  Note that the word “conformation” is often used as a synonym for “shape” in biological systems, as in “protein conformation




(15) Metabolism

(a)                  The sum of all of the chemical reactions that go on in an organism represent that organism’s metabolism

(b)                  Metabolism can be differentiated into what is known as catabolism versus what is known as anabolism

(16) Reactant

(a)                  In all chemical reactions, those chemicals that represent the starting material are called reactants

(17) Product

(a)                  Similarly, those chemicals that result from the chemical reaction are called products

(18) Reversible reactions

(a)                  Many chemical reactions of biological importance are reversible such that the products of one reaction serve as the reactants of the same chemical reaction operating in reverse (reversible reactions)

(b)                  For example, note the following equation:

(i)                    3H2 +N2 <==> 2NH3 (pardon my non-standard arrow usage)

(ii)                 In this equation the rate of the forward reaction (read left to right) is proportional to rate of collision between molecular hydrogen and molecular nitrogen

(iii)               The rate of this collision is dependent on the concentrations of the two reactants

(iv)               The rate of the reverse reaction (read right to left) is an intrinsic property of ammonia (i.e., the decomposition of NH3)

(c)                   Another, very important example of a reversible reaction is the dissociation of water:  H2O <==> OH- + H+

(19)  Chemical equilibria      

(a)                  Dynamic equilibrium occurs when the rate of the forward reaction is exactly equal to the rate of the reverse reaction in a reversible reaction

(b)                  (dynamic equilibria is a subcategory of chemical equilibria; the only chemical equilibria we will be concerned with are dynamic equilibria; in a dynamic equilibrium the reactions continue to proceed but there is no net gain or net loss in either products or reactants; in a static equilibrium no reaction proceeds)

(c)                   For example, again see the last equation on page 34 (3H2 + N2 <==> 2NH3); in this chemical reaction the point of the dynamic chemical equilibrium between molecular hydrogen, molecular nitrogen, and ammonia occurs when the reactant concentration is sufficiently low than the rate of collision leading to ammonia formation exactly equals the intrinsic rate of ammonium reversion to its parent gasses (i.e., H2 and N2)

(d)                  Equilibrium represents maximum disorder (maximum entropy) (and minimal energy) and all processes move toward increased entropy (and minimal energy), unless energy is employed to counter this process locally; thus, a scientist says grace: “Thank-you to all those who increased in entropy – in particular the sun – so that we may have this food.”




(20) Chemical history

(a)                  The energy and matter that make up your bodies were formed during the cosmological expansion of the universe known as the big bang

(b)                  According to the inflationary-universe hypothesis, all of the energy in the universe (and then some) was created during a brief interval very early in the cosmic big-bang expansion

(c)                   All of the elements that make up your body condensed as the young universe cooled, first as hydrogen (and a little helium and lithium), then into stars

(d)                  To a cosmologist, metals are elements larger (in atomic number) than hydrogen and metals are formed within large stars via a process known as fusion

(e)                   Explosion of these large stars (type I supernovae) results in the fusion of many elements including transuranic (heavier than uranium) elements

(f)                    Subsequent fission (splitting of atoms) resulted in the formation of all known elements not formed within stars, as well as the radioactive decay (and loss) of all not-naturally-occurring elements (including those heavier than uranium)

(g)                  These various elements exist within diffuse clouds of hydrogen as interstellar dust

(h)                  Further star formation via the condensation of these clouds of dust leads to accretion of dust particles into planetesimals which, in turn, collide (accrete) into larger bodies (planets)

(i)                    Most scientists assume that any planet that possesses liquid water, organic molecules, and a reasonable supply of energy either already has, or instead has the potential to evolve life

(j)                    Important to the evolution of life is the existence of a reducing atmosphere, that is, very little molecular oxygen (i.e., O2), since molecular oxygen decomposes organic molecules and therefore is toxic to all life forms except those that have evolved mechanisms by which they are protected from the toxic effects of oxygen