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Biology LibreTexts

SS1_2018_Lecture_01

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  • Biology

    Biology is the scientific study of life. Studying biology is an opportunity to ask exciting questions about the world that surrounds us. It is an opportunity to dig into some of humanity's deepest questions about our origins, our planet's history, and our connections to other living beings (big and small/extant or extinct). It is also an opportunity to dive into a world of practical problem solving and to think hard about possible solutions for improving health care, maintaining sustainable food supplies, and producing renewable energy technologies.

    Studying biology helps us understand issues and address everyday problems. For instance, you can better understand how what you eat and the amount you exercise influence your health when you understand the biochemical reactions that describe how the food (matter) is transformed, how it and your body store energy, and how this energy can be transferred from the food to your muscles. Deciding whether or not to buy products labeled with terms like "antimicrobial" or "probiotic" can be easier if you understand what the microbes, which live in, on, and around us, do. Understanding the biochemical principles that describe the changes that happen to eggs as they cook can help us understand how similar physical processes may be central to the cellular stress response and some diseases. Your eye color can be better appreciated with an understanding of the genetic and biochemical mechanisms that link genetic information to physical traits.

    Studying biology even helps us understand things that are "out of this world." For instance, understanding the requirements for life can help us look for life in places like Mars or deep in Earth’s crust. When we understand how to properly “rewire” cellular decision-making networks, we may finally be able to regenerate functional limbs or organs from someone’s own tissue, or reprogram diseased tissues back to health. There are many exciting opportunities. The key point is that mastering a few basic principles helps you understand and think more deeply about a wide array of topics. Keep this notion in mind throughout the course.

    Biology: an interdisciplinary science

    Questions in biology span size scales in excess of ten orders of magnitude, from the atomic makeup and chemical behavior of individual molecules to planetary-scale systems of interacting ecologies. Whatever the scale of interest, to develop a deep and functional understanding of biology, we must first appreciate biological concepts. This involves integrating important ideas and tools from across the spectrum of science, including chemistry, physics, and mathematics. Biology is truly an interdisciplinary science.

    The potential application of knowledge is broad

    Some people may think studying biology is only about medicine—however, it can lead to or influence many different careers. Biology has applications that are both vast and wide-ranging. Applications include treating (human or other animal) patients, improving agricultural practices, developing new building materials, writing new energy policies, remedying global climate change, creating new works of art—the list goes on and on. For the curious, biology has plenty of unexplored mysteries.

    As you study biology, appreciate its exciting questions and topics and be open-minded. Even though course topics may not always seem related at first, they likely are. Being open-minded helps you discover and appreciate the connections between the course’s topics and your interests. Discovering how seemingly different topics interrelate can give you a deeper appreciation for the things you enjoy and maybe even spark a new passion.

    BIS2A—from molecules to cells

    BIS2A focuses on the cell, one of the most fundamental units of life. Cells can be as simple as the disease-causing bacterium Mycoplasma genitalium, whose genome encodes just 525 genes (only 382 of which are essential for life), or as complex as a cell belonging to the multicellular plant Oryza sativa (rice), whose genome likely encodes ~51,000 genes. However, in spite of this diversity, all cells share some fundamental properties. In BIS2A, we explore basic problems that must be dealt with by all cells. We study the building blocks of cells, some of their key biochemical properties, how biological information is encoded and expressed in genetic material, and how all this combines to make a living system. We will also discuss some of the ways in which living systems exchange matter, energy, and information with their environment (including other living things). We focus primarily on core principles that are common to all life on Earth, and due to biology's large breadth, we put these ideas into a variety of contexts throughout the quarter.

     

    Evolution and Natural Selection

    Brief overview

    Evolution and natural selection are core concepts in biology that are typically invoked to help explain the diversity of and relationships between life on Earth, both extant and extinct. Fortunately, in BIS2A, you need to understand and use only a few core ideas related to evolution and natural selection. We describe these below. You will expand your understanding and add details to these core concepts in BIS2B and BIS2C.  

    The first idea you need to grasp is that evolution can be simply defined as the development/change of something over time. In the automotive industry, the shapes and features of cars can be said to evolve (change in time). In fashion, it can be said that style evolves. In biology, life and, in particular, reproducing populations of organisms with different traits evolve.  

    The second thing to understand is that natural selection is a process by which nature filters organisms in a population. What is the filter? Here it becomes a little more complicated (but only a little). The simplest explanation is that the selective filter is just a combination of all living and nonliving factors in an environment, which influence how successfully an organism can reproduce. The factors that influence the ability of an organism to reproduce are known as selective pressures. A small but important complication is that these factors are not the same everywhere; they change in time and by location. Thus, the selective pressures that create the filter are constantly changing (sometimes rapidly, sometimes slowly), and organisms in the same reproducing population could experience different pressures at different times and in different locations. 

    The theory of evolution by natural selection puts these two ideas together; it stipulates that change in biology happens over time and that the variation in a population is constantly subjected to selection based on how differences in traits influence reproduction. But what are these characteristics or traits? What traits/features/functions can be subject to selection? The short answer is: just about anything associated with an organism for which variation exists in a population and for which this variation leads to a differential likelihood of generating offspring will probably be subject to filtering by natural selection. We also call these traits heritable phenotypes. Organisms in a population that have phenotypes, which enable them to pass the selective filter more efficiently than others, are said to have a selective advantage and/or greater fitness.  

    It is important to reiterate that while the phenotypes carried by individual organisms may be subject to selection, the process of evolution by natural selection both requires and acts on phenotypic variation within populations. If neither variation nor populations in which that variation can reside exist, there is no opportunity or need for selection. Everything is and stays the same.

    Common misconceptions and a course specific note

    Finally, we draw your attention to a critical point and common misconception among beginning students in biology. This misconception can arise when, for the sake of discussion, we decide to anthropomorphize nature by giving it an intellect. For example, we may try to build an example for evolution by natural selection by proposing that a surplus of a particular food exists in an environment and there is an organism close by that is starving. It would be correct to reason that if the organism could eat that food that this might give it a selective advantage over other organisms that cannot. If later we find an example of organisms that have the capability to eat that surplus food, it might be tempting to say that nature evolved to solve the problem the surplus food. The process of evolution by natural selection, however, happens randomly and without direction. That is, nature does NOT identify “problems” that are limiting fitness. Nature does NOT identify features that would make an organism more successful and then start creating diverse solutions that meet this need. The generation of variation is not guided. Variation happens and natural selection filters what works best. The observation that an organism exists that can eat the surplus food is not a reflection of nature actively solving a problem, but rather, a reflection of whatever processes that led to phenotypic variation in an ancestral population that createdamong many other variantsa phenotype that increased fitness (possibly because the ancestral organisms were able to eat the surplus food).  

    This point of the preceding paragraph is particularly important to understand in the context of BIS2A because of the way we will be utilizing the Design Challenge to understand biology. While the Design Challenge is intended to help focus our attention on functions under selection and their relationship to determining fitness, it can be easyif we aren’t attentiveto lapse into language that would suggest that nature purposefully designs solutions to solve specific problems. Always remember that we are looking retrospectively at what nature has selected and that we are attempting to understand why a specific phenotype may have been selected over many other possibilities. In doing so, we will be inferring or hypothesizing to the best of our ability (which is sometimes wrong) a sensible reason to explain why a phenotype might have provided a selective advantage. We are NOT saying that the phenotype evolved TO provide a specific selective advantage. The distinction between these two ideas may be subtle, but it is critical!

    Note: possible discussion

    What physical traits can you think of that give a selective advantage to certain species? Under what conditions would this trait grant those advantages? Under what conditions might that trait be a selective disadvantage?

    Note: possible discussion

    The great varieties of domesticated dog breeds from which we can choose for companionship are also the result of a process of evolution by selection. Likewise, the development of many very different looking cropscabbage, brussel sprouts, kohlrabi, kale, broccoli and caulifloweris also the result of evolution by selection. However, in these two cases the selection or filtering process is referred to artificial selection rather than natural selection. Discuss how artificial and natural selection are similar and different? 

    Note: possible discussion

    How do environmental and political factors influence manufacturing processes such as automobile design? Fashion? Etc. What aspects are similar to the evolution of an organism, and what aspects are different?

    Note: possible discussion

    A related but slightly different misconception about evolution by natural selection is that this process leads to the creation of the most efficient solutions to problems. What is the problem with this notion?

     

    General Approach to Biomolecule Types in BIS2A

    Before you start

    If necessary please review the Design Challenge module to review the Design Challenge rubric.

    Some context and motivation

    In BIS2A, we are concerned primarily with developing a functional understanding of a biological cell. In the context of a design problem, we might say that we want to solve the problem of building a cell. If we break this big task down into smaller problems, or alternatively, ask what types of things do we need to understand in order to do this, it would be reasonable to conclude that understanding what the cell is made of would be important. That said, it isn't sufficient to appreciate WHAT the cell is made of. We also need to understand the PROPERTIES of the materials that make up the cell. This requires us to dig into a little bit of chemistrythe science of the "stuff" (matter) that makes up the world we know.  

    This prospect of talking about molecular chemistry and thermodynamics makes some students of biology apprehensive. Hopefully, however, we will show that many of the vast number of biological processes that we care about arise directly from the chemical properties of the "stuff" that makes up life and that developing a functional understanding of some basic chemical concepts can be tremendously useful in thinking about how to solve problems in medicine, energy, and environment by attacking them at their core. 

    Importance of chemical composition

    As a student in BIS2A, you will be asked to classify macromolecules into groups by looking at their chemical composition and, based on this composition, also infer some of the properties they might have. For example, carbohydrates typically have multiple hydroxyl groups. Hydroxyl groups are polar functional groups capable of forming hydrogen bonds. Therefore, some of the biologically relevant properties of various carbohydrates can be understood at some level by a balance between how they may tend to form hydrogen bonds with water, themselves or other molecules. 

    Linking structure to function

    Each macromolecule plays a specific role in the overall functioning of a cell. The chemical properties and structure of a macromolecule will be directly related to its function. For example, the structure of a phospholipid can be broken down into two groups, a hydrophilic head group and a hydrophobic tail group. Each of these groups plays a role in not only the assembly of the cell membrane but also in the selectivity of substances that can/cannot cross the membrane.

     

    The Structure of an atom

    An atom is the smallest unit of matter that retains all of the chemical properties of an element. Elements are forms of matter with specific chemical and physical properties that cannot be broken down into smaller substances by ordinary chemical reactions. 

    The chemistry discussed in BIS2A requires us to use a model for an atom. While there are more sophisticated models, the atomic model used in this course makes the simplifying assumption that the standard atom is composed of three subatomic particles, the proton, the neutron, and the electron. Protons and neutrons have a mass of approximately one atomic mass unit (a.m.u.). One atomic mass unit is approximately 1.660538921 x 10-27kg—roughly 1/12 of the mass of a carbon atom (see table below for more precise value). The mass of an electron is 0.000548597 a.m.u. or 9.1 x 10-31kg. Neutrons and protons reside at the center of the atom in a region call the nucleus while the electrons orbit around the nucleus in zones called orbitals, as illustrated below. The only exception to this description is the hydrogen (H) atom, which is composed of one proton and one electron with no neutrons. An atom is assigned an atomic number based on the number of protons in the nucleus. Neutral carbon (C), for instance has six neutrons, six protons, and six electrons. It has an atomic number of six and a mass of slightly more than 12 a.m.u.

    Table 1. Charge, mass, and location of subatomic particles

    Protons, neutrons, and electrons
      Charge Mass (a.m.u.) Mass (kg) Location
    Proton +1 ~1 1.6726 x 10-27 nucleus
    Neutron 0 ~1 1.6749 x 10-27 nucleus
    Electron –1 ~0 9.1094 x 10-31 orbitals

    Table 1 reports the charge and location of three subatomic particles—the neutron, proton, and electron. Atomic mass unit = a.m.u. (a.k.a. dalton [Da])—this is defined as approximately one twelfth of the mass of a neutral carbon atom or 1.660538921 x 10−27 kg. This is roughly the mass of a proton or neutron.

     

     

    Figure 2. Elements, such as helium depicted here, are made up of atoms. Atoms are made up of protons and neutrons located within the nucleus and electrons surrounding the nucleus in regions called orbitals. (Note: This figure depicts a Bohr model for an atom—we could use a new open source figure that depicts a more modern model for orbitals. If anyone finds one please forward it.)
    Source:(https://commons.wikimedia.org/wiki/F...um_atom_QM.svg)
    By User: Yzmo (Own work) [GFDL (http://www.gnu.org/copyleft/fdl.html) or CC-BY-SA-3.0 (http://creativecommons.org/licenses/by-sa/3.0/)], via Wikimedia Commons

    Relative sizes and distribution of elements

    The typical atom has a radius of one to two angstroms (Å). 1Å = 1 x 10-10m. The typical nucleus has a radius of 1 x 10-5Å or 10,000 smaller than the radius of the whole atom. By analogy, a typical large exercise ball has a radius of 0.85m. If this were an atom, the nucleus would have a radius about 1/2 to 1/10 of your thinnest hair. All of that extra volume is occupied by the electrons in regions called orbitals. For an ideal atom, orbitals are probabilistically defined regions in space around the nucleus in which an electron can be expected to be found.  

    For additional basic information on atomic structure click here
    For additional basic information on orbitals here.

    Video clips

    For a review of atomic structure check out this Youtube video: atomic structure.

    The properties of living and nonliving materials are determined to a large degree by the composition and organization of their constituent elements. Five elements are common to all living organisms: Oxygen (O), Carbon (C), Hydrogen (H), Phosphorous (P), and Nitrogen (N). Other elements like Sulfur (S), Calcium (Ca), Chloride (Cl), Sodium (Na), Iron (Fe), Cobalt (Co), Magnesium, Potassium (K), and several other trace elements are also necessary for life, but are typically found in far less abundance than the "top five" noted above. As a consequence, life's chemistry—and by extension the chemistry of relevance in BIS2A—largely focuses on common arrangements of and reactions between the "top five" core atoms of biology.

     

    Figure 3. A table illustrating the abundance of elements in the human body. A pie chart illustrating the relationships in abundance between the four most common elements. 
    Credit: Data from Wikipedia (http://en.wikipedia.org/wiki/Abundan...mical_elements); chart created by Marc T. Facciotti

     

     

     

    The Periodic Table

    The different elements are organized and displayed in the periodic table. Devised by Russian chemist Dmitri Mendeleev (1834–1907) in 1869, the table groups elements that, due to some commonalities of their atomic structure, share certain chemical properties. The atomic structure of elements is responsible for their physical properties including whether they exist as gases, solids, or liquids under specific conditions and and their chemical reactivity, a term that refers to their ability to combine and to chemically bond with each other and other elements.

    In the periodic table, shown below, the elements are organized and displayed according to their atomic number and are arranged in a series of rows and columns based on shared chemical and physical properties. In addition to providing the atomic number for each element, the periodic table also displays the element’s atomic mass. Looking at carbon, for example, its symbol (C) and name appear, as well as its atomic number of six (in the upper right-hand corner indicating the number of protons in the neutral nucleus) and its atomic mass of 12.11 (sum of the mass of electrons, protons, and neutrons).

    Figure: The periodic table shows the atomic mass and atomic number of each element. The atomic number appears above the symbol for the element and the approximate atomic mass appears to the left.
    Source: By 2012rc (self-made using inkscape) [Public domain], via Wikimedia Commons Modified by Marc T. Facciotti - 2016

     

     

    Electronegativity

    Molecules are collections of atoms that are associated with one another through bonds. It is reasonable to expectand the case empiricallythat different atoms will exhibit different physical properties, including abilities to interact with other atoms. One such property, the tendency of an atom to attract electrons, is described by the chemical concept and term, electronegativity. While several methods for measuring electronegativity have been developed, the one most commonly taught to biologists is the one created by Linus Pauling.

    A description of how Pauling electronegativity can be calculated is beyond the scope of BIS2A. What is important to know, however, is that electronegativity values have been experimentally and/or theoretically determined for nearly all elements in the periodic table. The values are unitless and are reported relative to the standard reference, hydrogen, whose electronegativity is 2.20. The larger the electronegativity value, the greater tendency an atom has to attract electrons. Using this scale, the electronegativity of different atoms can be quantitatively compared. For instance, by using Table 1 below, you could report that oxygen atoms (O) are more electronegative than phosphorous atoms (P).

    Table 1. Pauling electronegativity values for select elements of relevance to BIS2A as well as elements at the two extremes (highest and lowest) of the electronegativity scale.

    Attribution: Marc T. Facciotti (original work)

    The utility of the Pauling electronegativity scale in BIS2A is to provide a chemical basis for explaining the types of bonds that form between the commonly occurring elements in biological systems and to explain some of the key interactions that we observe routinely. We develop our understanding of electronegativity-based arguments about bonds and molecular interactions by comparing the electronegativities of two atoms. Recall, the larger the electronegativity, the stronger the "pull" an atom exerts on nearby electrons.

    We can consider, for example, the common interaction between oxygen (O) and hydrogen (H). Let us assume that O and H are interacting (forming a bond) and write that interaction as O-H, where the dash between the letters represents the interaction between the two atoms. To understand this interaction better, we can compare the relative electronegativity of each atom. Examining the table above, we see that O has an electronegativity of 3.44, and H has an electronegativity of 2.20.

    Based on the concept of electronegativity as we now understand it, we can surmise that the oxygen (O) atom will tend to "pull" the electrons away from the hydrogen (H) when they are interacting. This will give rise to a slight but significant negative charge around the O atom (due to the higher tendency of the electrons to be associated with the O atom). This also results in a slight positive charge around the H atom (due to the decrease in the probability of finding an electron nearby). Since the electrons are not distributed evenly between the two atoms AND, by consequence, the electric charge is also not evenly distributed, we describe this interaction or bond as polar. There are two poles in effect: the negative pole near the oxygen and the positive pole near the hydrogen.

    To extend the utility of this concept, we can now ask how an interaction between oxygen (O) and hydrogen (H) differs from an interaction between sulfur (S) and hydrogen (H). That is, how does O-H differ from S-H? If we examine the table above, we see that the difference in electronegativity between O and H is 1.24 (3.44 - 2.20 = 1.24) and that the difference in electronegativity between S and H is 0.38 (2.58 – 2.20 = 0.38). We can therefore conclude that an O-H bond is more polar than an S-H bond. We will discuss the consequences of these differences in subsequent chapters.

    Figure 2. The periodic table with the electronegativities of each atom listed.

    Attribution: By DMacks (https://en.wikipedia.org/wiki/Electronegativity) [CC BY-SA 3.0 (http://creativecommons.org/licenses/by-sa/3.0)], via Wikimedia Commons

    An examination of the periodic table of the elements (Figure 2) illustrates that electronegativity is related to some of the physical properties used to organize the elements into the table. Certain trends are apparent. For instance, those atoms with the largest electronegativity tend to reside in the upper right hand corner of the periodic table, such as Fluorine (F), Oxygen (O) and Chlorine (Cl), while elements with the smallest electronegativity tend to be found at the other end of the table, in the lower left, such as Francium (Fr), Cesium (Cs) and Radium (Ra).

    More information on electronegativity can be found in the LibreTexts.

    The main use of the concept of electronegativity in BIS2A will therefore be to provide a conceptual grounding for discussing the different types of chemical bonds that occur between atoms in nature. We will focus primarily on three types of bonds: Ionic Bonds, Covalent Bonds and Hydrogen Bonds.

    Bond types

    In BIS2A, we focus primarily on three different bond types: ionic bonds, covalent bonds, and hydrogen bonds. We expect students to be able to recognize each different bond type in molecular models. In addition, for commonly seen bonds in biology, we expect student to provide a chemical explanation, rooted in ideas like electronegativity, for how these bonds contribute to the chemistry of biological molecules.

    Ionic bonds

    Ionic bonds are electrostatic interactions formed between ions of opposite charges. For instance, most of us know that in sodium chloride (NaCl) positively charged sodium ions and negatively charged chloride ions associate via electrostatic (+ attracts -) interactions to make crystals of sodium chloride, or table salt, creating a crystalline molecule with zero net charge. The origins of these interactions may arise from the association of neutral atoms whose difference in electronegativities is sufficiently high. Take the example above. If we imagine that a neutral sodium atom and a neutral chlorine atom approach one another, it is possible that at close distances, due to the relatively large difference in electronegativity between the two atoms, that an electron from the neutral sodium atom is transferred to the neutral chlorine atom, resulting in a negatively charged chloride ion and a positively charged sodium ion. These ions can now interact via an ionic bond.

    Figure 1. The formation of an ionic bond between sodium and chlorine is depicted. In panel A, a sufficient difference in electronegativity between sodium and chlorine induces the transfer of an electron from the sodium to the chlorine, forming two ions, as illustrated in panel B. In panel C, the two ions associate via an electrostatic interaction. Attribution: By BruceBlaus (own work) [CC BY-SA 4.0 (http://creativecommons.org/licenses/by-sa/4.0)], via Wikimedia Commons

    This movement of electrons from one atom to another is referred to as electron transfer. In the example above, when sodium loses an electron, it now has 11 protons, 11 neutrons, and 10 electrons, leaving it with an overall charge of +1 (summing charges: 11 protons at +1 charge each and 10 electrons at -1 charge each = +1). Once charged, the sodium atom is referred to as a sodium ion. Likewise, based on its electronegativity, a neutral chlorine (Cl) atom tends to gain an electron to create an ion with 17 protons, 17 neutrons, and 18 electrons, giving it a net negative (–1) charge. It is now referred to as a chloride ion.

    We can interpret the electron transfer above using the concept of electronegativity. Begin by comparing the electronegativities of sodium and chlorine by examining the periodic table of elements below. We see that chlorine is located in the upper-right corner of the table, while sodium is in the upper left. Comparing the electronegativity values of chlorine and sodium directly, we see that the chlorine atom is more electronegative than is sodium. The difference in the electronegativity of chlorine (3.16) and sodium (0.93) is 2.23 (using the scale in the table below). Given that we know an electron transfer will take place between these two elements, we can conclude that differences in electronegativities of ~2.2 are large enough to cause an electron to transfer between two atoms and that interactions between such elements are likely through ionic bonds.

    Figure 2. The periodic table of the elements listing electronegativity values for each element. The elements sodium and chlorine are boxed with a teal boundary. Attribution: By DMacks (https://en.wikipedia.org/wiki/Electronegativity) [CC BY-SA 3.0 (http://creativecommons.org/licenses/by-sa/3.0)], via Wikimedia CommonsModified by Marc T. Facciotti

    Note: possible discussion

    The atoms in a 5 in. x 5 in. brick of table salt (NaCl) sitting on your kitchen counter are held together almost entirely by ionic bonds. Based on that observation, how would you characterize the strength of ionic bonds?

    Now consider that same brick of table salt after having been thrown into an average backyard swimming pool. After a couple of hours, the brick would be completely dissolved, and the sodium and chloride ions would be uniformly distributed throughout the pool. What might you conclude about the strength of ionic bonds from this observation?

    Propose a reason why NaCl's ionic bonds in air might be behaving differently than those in water? What is the significance of this to biology?

    For additional information:

    Check out the link from the Khan Academy on ionic bonds

    Covalent bonds

    We can also invoke the concept of electronegativity to help describe the interactions between atoms that have differences in electronegativity too small for the atoms to form an ionic bond. These types of interactions often result in a bond called a covalent bond. In these bonds, electrons are shared between two atoms—in contrast to an ionic interaction in which electrons remain on each atom of an ion or are transferred between species that have highly different electronegativities.

    We start exploring the covalent bond by looking at an example where the difference in electronegativity is zero. Consider a very common interaction in biology, the interaction between two carbon atoms. In this case, each atom has the same electronegativity, 2.55; the difference in electronegativity is therefore zero. If we build our mental model of this interaction using the concept of electronegativity, we realize that each carbon atom in the carbon-carbon pair has the same tendency to "pull" electrons to it. In this case, when a bond is formed, neither of the two carbon atoms will tend to "pull" (a good anthropomorphism) electrons from the other. They will "share" (another anthropomorphism) the electrons equally, instead.

    Aside: bounding example

    The two examples above—(1) the interaction of sodium and chlorine, and (2) the interaction between two carbon atoms—frame a discussion by "bounding," or asymptotic analysis (see earlier reading). We examined what happens to a physical system when considering two extremes. In this case, the extremes were in electronegativity differences between interacting atoms. The interaction of sodium and chlorine illustrated what happens when two atoms have a large difference in electronegativities, and the carbon-carbon example illustrated what happens when that difference is zero. Once we create those mental goal posts describing what happens at the extremes, it is then easier to imagine what might happen in between—in this case, what happens when the difference in electronegativity is between 0 and 2.2. We do that next.

    When the sharing of electrons between two covalently bonded atoms is nearly equal, we call these bonds nonpolar covalent bonds. If by contrast, the sharing of electrons is not equal between the two atoms (likely due to a difference in electronegativities between the atoms), we call these bonds polar covalent bonds.

    In a polar covalent bond, the electrons are unequally shared by the atoms and are attracted to one nucleus more than to the other. Because of the unequal distribution of electrons between atoms in a polar covalent bond, a slightly positive (indicated by δ+) or slightly negative (indicated by δ–) charge develops at each pole of the bond. The slightly positive (δ+) charge will develop on the less electronegative atom, as electrons get pulled more towards the slightly more electronegative atom. A slightly negative (δ–) charge will develop on the more electronegative atom. Since there are two poles (the positive and negative poles), the bond is said to possess a dipole.

    Examples of nonpolar covalent and polar covalent bonds in biologically relevant molecules

    Nonpolar covalent bonds

    Molecular oxygen

    Molecular oxygen (O2) is made from an association between two atoms of oxygen. Since the two atoms share the same electronegativity, the bonds in molecular oxygen are nonpolar covalent.

    Methane

    Another example of a nonpolar covalent bond is the C-H bond found in the methane gas (CH4). Unlike the case of molecular oxygen where the two bonded atoms share the same electronegativity, carbon and hydrogen do not have the same electronegativity; C = 2.55 and H = 2.20—the difference in electronegativity is 0.35.

    Figure 3. Molecular line drawings of molecular oxygen, methane, and carbon dioxide. Attribution: Marc T. Facciotti (own work)

    Some of you may now be confused. If there is a difference in electronegativity between the two atoms, is the bond not by definition polar? The answer is both yes and no and depends on the definition of polar that the speaker/writer is using. Since this is an example of how taking shortcuts in the use of specific vocabulary can sometimes lead to confusion, we take a moment to discuss this here. See the mock exchange between a student and an instructor below for clarification:

    1. Instructor: "In biology, we often say that the C-H bond is nonpolar."

    2. Student: "But there is an electronegativity difference between C and H, so it would appear that C should have a slightly stronger tendency to attract electrons. This electronegativity difference should create a small, negative charge around the carbon and a small, positive charge around the hydrogen."

    3. Student: "Since there is a differential distribution of charge across the bond, it would seem that, by definition, this should be considered a polar bond."

    4. Instructor: "In fact, the bond does have some small polar character."

    5. Student: "So, then it's polar? I'm confused."

    6. Instructor: "It has some small amount of polar character, but it turns out that for most of the common chemistry that we will encounter that this small amount of polar character is insufficient to lead to "interesting" chemistry. So, while the bond is, strictly speaking, slightly polar, from a practical standpoint it is effectively nonpolar. We therefore call it nonpolar."

    7. Student: "That's needlessly confusing; how am I supposed to know when you mean strictly 100% nonpolar, slightly polar, or functionally polar when you use the same word to describe two of those three things?"

    8. Instructor: "Yup, it sucks. The fix is that I need to be as clear as I can when I talk with you about how I am using the term "polarity." I also need to inform you that you will find this shortcut (and others) used when you go out into the field, and I encourage you to start learning to recognize what is intended by the context of the conversation.

    A real-world analogy of this same problem might be the use of the word "newspaper". It can be used in a sentence to refer to the company that publishes some news, OR it can refer to the actual item that the company produces. In this case, the disambiguation is easily made by native English speakers, as they can determine the correct meaning from the context; non-native speakers may be more confused. Don't worry; as you see more examples of technical word use in science, you'll learn to read correct meanings from contexts too."

    Aside:

    How large should the difference in electronegativity be in order to create a bond that is "polar enough" that we decide to call it polar in biology? Of course, the exact value depends on a number of factors, but as a loose rule of thumb, we sometimes use a difference of 0.4 as a guesstimate.

    This extra information is purely for your information. You will not be asked to assign polarity based on this criteria in BIS2A. You should, however, appreciate the concept of how polarity can be determined by using the concept of electronegativity. You should also appreciate the functional consequences of polarity (more on this in other sections) and the nuances associated with these terms (such as those in the discussion above).

    Polar covalent bonds

    The polar covalent bond can be illustrated by examining the association between O and H in water (H2O). Oxygen has an electronegativity of 3.44, while hydrogen has an electronegativity of 2.20. The difference in electronegativity is 1.24. It turns out that this size of electronegativity difference is large enough that the dipole across the molecule contributes to chemical phenomenon of interest.

    This is a good point to mention another common source of student confusion regarding the use of the term polar. Water has polar bonds. This statement refers specifically to the individual O-H bonds. Each of these bonds has a dipole. However, students will also hear that water is a polar molecule. This is also true. This latter statement is referring to the fact that the sum of the two bond dipoles creates a dipole across the whole molecule. A molecule may be nonpolar but still have some polar bonds.

    Figure 4. A water molecule has two polar O-H bonds. Since the distribution of charge in the molecule is asymmetric (due to the number and relative orientations of the bond dipoles), the molecule is also polar. The element name and electronegativities are reported in the respective sphere. Attribution: Marc T. Facciotti (own work)

    For additional information, view this short video to see an animation of ionic and covalent bonding.

    The continuum of bonds between covalent and ionic

    The discussion of bond types above highlights that in nature you will see bonds on a continuum from completely nonpolar covalent to purely ionic, depending on the atoms that are interacting. As you proceed through your studies, you will further discover that in larger, multi-atom molecules, the localization of electrons around an atom is also influenced by multiple factors. For instance, other atoms that are also bonded nearby will exert an influence on the electron distribution around a nucleus in a way that is not easily accounted for by invoking simple arguments of pairwise comparisons of electronegativity. Local electrostatic fields produced by other non-bonded atoms may also have an influence. Reality is always more complicated than are our models. However, if the models allow us to reason and predict with "good enough" precision or to understand some key underlying concepts that can be extended later, they are quite useful.

    Key bonds in BIS2A

    In BIS2A, we are concerned with the chemical behavior of and bonds between atoms in biomolecules. Fortunately, biological systems are composed of a relatively small number of common elements (e.g., C, H, N, O, P, S, etc.) and some key ions (e.g., Na+, Cl-, Ca2+, K+, etc.). Start recognizing commonly occurring bonds and the chemical properties that we often see them showing. Some common bonds include C-C, C-O, C-H, N-H, C=O, C-N, P-O, O-H, S-H, and some variants. These will be discussed further in the context of functional groups. The task is not as daunting as it seems.

    Note Common Point of student confusion

    In this reading we have been talking about the polarity of bonds.  That is, we have been learning how to describe the polarity of a single bond joining two atoms (i.e. how are the electrons shared between two atoms distributed about the respective nuclei?). In biology we also sometimes talk about the polarity of a molecule. The polarity of a molecule is different than the polarity of a bond within the molecule. The latter is asking whether the whole molecule has a net dipole. The molecule's dipole can be roughly thought of as the sum of all of its bond dipoles.  For example, let us examine a molecule of CO2 depicted in the figure above.  If we ask whether one of the C=O bonds is polar we would conclude that it is since the oxygen is significantly more electronegative that the carbon to which it is covalently bonded. However, if we ask whether the molecule O=C=O is polar we would concluded that it is not. Why? Look at the figure of CO2 above. Each CO bond has a dipole. However, these two dipoles are pointed in directly opposite directions. If we add these two bond dipoles together to get the net dipole of the molecule we get nothing - the two bond dipoles "cancel" one another out.  By contrast, if we examine the structure of water above, we also see that each O-H bond has a dipole. In this case when we ask whether the molecule has a net dipole (done by adding the bond dipoles together) we see that the answer is yes. The sum of the the two bond dipoles still yields a net dipole moment. We therefore say that this molecule is polar.  We can do this same exercise for parts of molecules so long as we define what specific part we are looking at.  

     

    Hydrogen Bonds

    When hydrogen forms a polar covalent bond with an atom of higher electronegativity, the region around the hydrogen will have a fractional positive charge (termed δ+). When this fractional positive charge encounters a partial negative charge (termed δ-) from another electronegative atom to which the hydrogen is NOT bound, AND it is presented to that negative charge in a suitable orientation, a special kind of interaction called a hydrogen bond can form. While chemists are still debating the exact nature of the hydrogen bond, in BIS2A, we like to conceive of it as a weak electrostatic interaction between the δ+ of the hydrogen and the δ- charge on an electronegative atom. We call the molecule that contributes the partially charged hydrogen atom "the hydrogen bond donor" and the atom with the partial negative charge the "hydrogen bond acceptor." You will be asked to start learning to recognize common biological hydrogen bond donors and acceptors and to identify putative hydrogen bonds from models of molecular structures.  

    Hydrogen bonds are common in biology both within and between all types of biomolecules. Hydrogen bonds are also critical interactions between biomolecules and their solvent, water. It is common, as seen in the figure below, to represent hydrogen bonds in figures with dashed lines.

    Figure 1: Two water molecules are depicted forming a hydrogen bond (drawn as a dashed blue line). The water molecule on top "donates" a partially charged hydrogen while the water molecule on the bottom accepts that partial charge by presenting a complementary negatively charged oxygen atom. 

    Attribution: Marc T. Facciotti (original work)

     

    Functional groups

    A functional group is a specific group of atoms within a molecule that is responsible for a characteristic of that molecule. Many biologically active molecules contain one or more functional groups. In BIS2A, we will review the major functional groups found in biological molecules. These include the following: hydroxyl, methyl, carbonyl, carboxyl, amino, and phosphate (see Figure 1). 

    Figure 1. The functional groups shown here are found in many different biological molecules. "R" represents any other atom or extension of the molecule.  
    Attribution: Marc T. Facciotti (own work adapted from previous image of unknown source)

    A functional group may participate in a variety of chemical reactions. Some of the important functional groups in biological molecules are shown above: hydroxyl, methyl, carbonyl, carboxyl, amino, phosphate, and sulfhydryl (not shown). These groups play an important role in the formation of molecules like DNA, proteins, carbohydrates, and lipids. Functional groups can sometimes be classified as having polar or nonpolar properties depending on their atomic composition and organization. The term polar describes something that has a property that is not symmetric about it—it can have different poles (more or less of something at different places). In the case of bonds and molecules, the property we care about is usually the distribution of electrons and therefore electric charge between the atoms. In a nonpolar bond or molecule, electrons and charge will be relatively evenly distributed. In a polar bond or molecule, electrons will tend to be more concentrated in some areas than others. An example of a nonpolar group is the methane molecule (see discussion in Bond Types Chapter for more detail). Among the polar functional groups is the carboxyl group found in amino acids, some amino acid side chains, and the fatty acids that form triglycerides and phospholipids.

    Nonpolar functional groups

    Methyl R-CH3

    The methyl group is the only nonpolar functional group in our class list above. The methyl group consists of a carbon atom bound to three hydrogen atoms. In this class, we will treat these C-H bonds as effectively nonpolar covalent bonds (more on this in the Bond Types chapter). This means that methyl groups are unable to form hydrogen bonds and will not interact with polar compounds such as water.

    Figure 2. The amino acid isoleucine is on the left, and cholesterol is on the right. Each has a methyl group circled in red. Attribution: created by Marc T. Facciotti (own work adapted from Erin Easlon)

    The methyl groups highlighted above are found in a variety of biologically relevant compounds. In some cases, the compound can have a methyl group but still be a polar compound overall due to the presence of other functional groups with polar properties (see the discussion on polar functional groups below).

    As we learn more about other functional groups, we will add to the list of nonpolar functional groups. Stay alert!

    Polar functional groups

    Hydroxyl R-OH

    A hydroxyl (alcohol group) is an -OH group covalently bonded another atom. In biological molecules the hydroxyl group is often (but not always) found bound to a carbon atom, as depicted below. The oxygen atom is much more electronegative than either the hydrogen or the carbon, which will cause the electrons in the covalent bonds to spend more time around the oxygen than around the C or H. Therefore, the O-H and O-C bonds in the hydroxyl group will be polar covalent bonds. Figure 3 depicts the partial charges, δ+ and δ-, that are associated with the hydroxyl group.

    Figure 3. The hydroxyl functional group shown here consists of an oxygen atom bound to a carbon atom and a hydrogen atom. These bonds are polar covalent, meaning the electron involved in forming the bonds is not shared equally between the C-O and O-H bonds. Attribution: created by Marc T. Facciotti (own work)

    Figure 4. The hydroxyl functional groups can form hydrogen bonds, shown as a dotted line. The hydrogen bond will form between the δ - of the oxygen atom and the δ + of the hydrogen atom. Dipoles are shown in blue arrows. Attribution: Marc T. Facciotti (original work)

    Hydroxyl groups are very common in biological molecules. Hydroxyl groups appear on carbohydrates (A), on some amino acids (B), and on nucleic acids (C). Can you find any hydroxyl groups in the phospholipid in (D)?

    Figure 5. Hydroxyl groups appear on carbohydrates (A, glucose), on some amino acids (B, Serine), and on nucleotides (C, adenosine triphosphate). D is a phospholipid. Attribution: created by Marc T. Facciotti (own work)

    Carboxyl R-COOH

    Carboxylic acid is a combination of a carbonyl group and a hydroxyl group attached to the same carbon, resulting in new characteristics. The carboxyl group can ionize, which means it can act as an acid and release the hydrogen atom from the hydroxyl group as a free proton (H+). This results in a delocalized negative charge on the remaining oxygen atoms. Carboxyl groups can switch back and forth between protonated (R-COOH) and deprotonated (R-COO-) states depending on the pH of the solution. 

    The carboxyl group is very versatile. In its protonated state, it can form hydrogen bonds with other polar compounds. In its deprotonated state, it can form ionic bonds with other positively charged compounds. This will have several biological consequences that will be explored more when we discuss enzymes.

    Can you identify all the carboxyl groups on the macromolecules shown above in Figure 5?

    Amino R-NH3

    The amino group consists of a nitrogen atom attached by single bonds to hydrogen atoms. An organic compound that contains an amino group is called an amine. Like oxygen, nitrogen is also more electronegative than both carbon and hydrogen, which results in the amino group displaying some polar character. 

    Amino groups can also act as bases, which means that the nitrogen atom can bond to a fourth hydrogen atom, as shown in Figure 6. Once this occurs, the nitrogen atom gains a positive charge and can now participate in ionic bonds.

    Figure 6. The amine functional group can exist in a deprotonated or protonated state. When protonated, the nitrogen atom is bound to three hydrogen atoms and has a positive charge. The deprotonated form of this group is neutral. Attribution: created by Erin Easlon (own work)

    Phosphate R-PO4-

    A phosphate group is a phosphorus atom covalently bound to four oxygen atoms and contains one P=O bond and three P-O bonds. The oxygen atoms are more electronegative than the phosphorous atom, resulting in polar covalent bonds. Therefore, these oxygen atoms are able to form hydrogen bonds with nearby hydrogen atoms that also have a δ+(hydrogen atoms bound to another electronegative atom). Phosphate groups also contain a negative charge and can participate in ionic bonds.

    Phosphate groups are common in nucleic acids and on phospholipids (the term "phospho" referring to the phosphate group on the lipid). In Figure 7 are images of a nucleotide, deoxyadenosine monphosphate (left), and a phosphoserine (right).

    Figure 7. A nucleotide, deoxyadenosine monphosphate, is on the left, and phosphoserine is on the right. Each has a phosphate group circled in red.  
    Attribution: created by Marc T. Facciotti (own work)