1.2: Chemistry Concepts

Introduction to Biomolecules for General Biology

Some context and motivation

Many general biology courses are concerned with getting students to start learning how living systems "work". To accomplish this we need to understand the "stuff" biological systems. We want to know what "stuff" organisms are made of and how the properties of this "stuff" contribute to organismal forms and to their behaviors. This requires us to dig into a little chemistry — the science of the "stuff" that makes up the world we know. By doing so, we can show that many of the biological processes we care about arise directly from the chemical properties of the "stuff" that makes up living systems. Developing a functional understanding of some basic chemical concepts can, therefore, be tremendously useful in thinking about how to solve problems in medicine, energy, and environment by attacking them at their molecular levels. For these reasons, we begin a study of biomolecules and their chemistry.

Importance of chemical composition

Most instructors of general biology will want you to learn how to classify major classes of biological macromolecules into groups by looking at their chemical composition and, based on this composition, to also infer some properties they might have. For example, you will learn to recognize the molecular structures of carbohydrates. You'll learn to do so by spotting chemical "parts" on carbohydrates called hydroxyl groups and appreciate that those chemical "parts" can form special types of interactions called "hydrogen bonds". You'll discover how hydrogen bonds give carbohydrates special properties that govern how they interact with water and other molecules and how these interactions give different properties to carbohydrates like simple sugars and starches. This will shed light on the functional roles these molecules play in living systems, how they are built, and how they can be broken down. We will also look at other types of common macromolecules in biology in the same way.

Linking molecular structure to function

In biology, structure is often said to "encode" function. That is, the shape and physical properties of something governs what it can do. We'll want to make this general idea a bit more specific for you and begin to help you form the related conceptual ideas that will help you explain how the chemical structure and properties of biomolecules can contribute to their functions.

For example, biological membranes have the biological function of defining the boundaries between the inside and outside of a cell or between the inside and outside of compartments within a cell. Membranes also serve as selective barriers that allow certain things to pass into and out of a cell while excluding others substances from crossing the barrier. These are very important functions indeed! The chemical structures of membranes determine these functional properties. You will learn, for instance, that most membranes are composed of two layers of molecules called phospholipids and that phospholipids can be schematically broken down into two chemically distinct subregions, a so-called hydrophilic head group and a hydrophobic tail group. Each of these chemical subregions and the chemical "parts" they are made of, contribute directly to determining the structure and permeability of cell membranes.

Likewise, you will also learn about important biological molecules called enzymes. Enzymes help catalyze many of the chemical reactions in an organism, including how they process the food they eat, how an organism senses and interacts with its environment, and how it replicates (among many other functions). We will learn that the three dimensional structures of enzymes encode their functions and that their structures are determined to a large degree by the number, type, and order of chemicals called amino acids. The properties of these amino acids (and thus how they behave in an enzyme) are directly related to their own chemical "parts".

In sum, while structure "encodes" biological function, structure is itself usually determined by chemistry. Knowing a little bit of chemistry can be powerful in biology.

Objectives and strategy

On these notes, we set out to learn a little bit of useful chemistry for biology.

• First, we want to learn some basic vocabulary, symbols and models that help us communicate about chemistry in biology.
• Second, we want to learn a few basic ideas about what the "stuff-of-life" is made of and some basic properties of that stuff that help us understand biological phenomena.
• Third, we want to learn to recognize some of the most common chemical structures in biology and something about their properties.
• Fourth, we want to understand how some of the chemistries of the "stuff-of-life" can be altered and how this might change the function of a biomolecule.
• Fifth, we want to develop a vocabulary and learn key ideas that help us talk about how things change at the molecular level in biology.

If you've already taken some chemistry, some of the ideas will be familiar to you. If you haven't taken chemistry - or it's been a long time since you did - don't worry. In general biology, these topics aren't typically covered as theoretically or in as much detail as they might be in a formal chemistry class. Here, we want to extract the most important ideas/concepts that help us function as general biologists. The core ideas and concepts aren't too complicated. However, there will likely be a lot of new vocabulary to learn and a good amount of practice required to get comfortable thinking about and building mental models of things that happen at the molecular scale. Success comes from practice. So, practice early and often. Strive to learn the new vocabulary quickly and take advantage of practice opportunities to think about molecules. That'll make it much easier to then think about and add new concepts as we go move forward in the course.

The Periodic Table

The periodic table is a tool for organizing information about the different elements found in nature. Devised by Russian chemist Dmitri Mendeleev (1834–1907) in 1869, the table groups elements that, because of 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 their chemical reactivity, a term that refers to their ability to combine and to bond chemically 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. Besides 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, and its atomic number of six (in the upper right-hand corner showing 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 associate with one another through bonds. It is reasonable to expect — and true empirically — that different atoms will exhibit different physical properties, including abilities to interact/bond with other atoms. We describe one such property, the tendency of an atom to attract electrons, by the chemical concept and term, electronegativity. While chemists have developed several methods for measuring electronegativity, Linus Pauling created the one most commonly taught to biologists.

A description of how Pauling electronegativity can be calculated is typically beyond the scope of introductory biology. It is important to know, however, that electronegativity values have been experimentally and/or theoretically determined for nearly all elements in the periodic table. The values are unitless. The larger the electronegativity value, the greater the tendency an atom has to attract electrons. Using this scale, one can quantitatively compare the electronegativity of different atoms. 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 most general biology courses 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 general biology is to provide a chemical basis for explaining the bonds that form between the commonly occurring elements in biological systems and to explain some 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 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 partial negative charge around the O atom (because of the higher tendency of the electrons to associate with the O atom). This also results in a slight partial positive charge around the H atom (because of the decrease in the probability of finding an electron around the hydrogen atom). Since the electrons distribute unevenly between the two atoms AND, by consequence, the electric charge also distributes unevenly, we describe this interaction or bond as polar. There are two poles in effect: the more negative pole near the oxygen and the more 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 the relationship between electronegativity and some physical properties used to organize the elements into the table. Certain trends are plain. 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).

You can find more information on electronegativity in the Chemistry LibreTexts.

The main use of the concept of electronegativity in general biology will therefore be to provide a conceptual grounding for discussing the different types of chemical interactions that occur between atoms in nature. We will focus primarily on covalent bonds, and several non-covalent interactions called ionic bonds, hydrogen bonds, and Van der Waals forces.

Covalent bonds and non-covalent molecular interactions

In general biology, we focus primarily on covalent bonds, ionic bonds, hydrogen bonds and Van der Waals forces. 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, in Chemistry we learn 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. 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 Bruce Blaus (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, 12 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, 18 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 Commons—Modified by Marc T. Facciotti

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Possible NB Discussion Point

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 seemingly behave differently in air and water? What is the significance of this observation to biology?

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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. It 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 non-uniform 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 in biology 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 linguistic 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 then intended 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 usually not be asked to assign polarity based on this guesstimate in a general biology course (though you may want to make sure with your individual instructor). 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. However, it is also true that a molecule may be nonpolar but still have some polar bonds. The typical example given to illustrate this case is that of carbon dioxide (CO₂) - this molecule is shown in the figure above. While the CO₂ molecule has two polar C-O bonds, these diploes are equal in magnitude and point in opposite directions. When the bond dipoles in CO₂ are added together to determine the molecule's dipole they cancel one another. This leads a molecule with no dipole even though it has individual bonds that are polar.

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)

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 general biology

In general biololgy, we are usually 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.

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Possible NB Discussion Point

Imagine that you were able to shrink yourself down to the size of an atom and see things like electrons and protons. Describe what you would see if you were standing on Carbon 1 in the molecule below and looking in different directions towards the bound oxygen, hydrogens or carbon 2. Compare and contrast what you expect to see along each bond.

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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 general biology, we typically 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." We will ask you to learn 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 many 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)

Dipoles, Van der Waals Forces, and Pi Interactions

In addition to ionic and hydrogen bonds, there are several other types of non-covalent molecular interactions that we encounter in biology. Key among these are dipole-dipole interactions, Van der Waals forces and pi interactions. In this section we briefly describe each of these interactions and some of their underlying basis. Developing a deep and comprehensive theoretical understanding of these interaction types requires a dive into more advanced chemistry. We don’t do that. Rather, we try to provide a more descriptive understanding of these phenomena that will hopefully be useful for interpreting common molecular interactions in biology. Recall that with respect to chemistry, our goals in general biology are relatively modest. We want students to recognize different chemical interactions between biomolecules, to appreciate that these interactions arise from the unique chemical properties of the elements that make up the molecules, and to appreciate how environmental and chemical factors can change these interactions. If you can identify obvious biological scenarios in which different interactions can take place, you’re doing great!

Dipole-Dipole Interactions

Dipole-dipole interactions are, as the name suggests, simply interactions between two dipoles. Recall how the differences in electronegativities between elements can explain the creation of polar covalent bonds. We describe these polar covalent bonds as permanent dipoles, “hard-coded” by the properties of the elements bonded together. The dipole-dipole interaction is an interaction between two permanent dipoles. If partial charges carrying the same (+ or -) sign interact (i.e. positive interacts with positive or negative interacts with negative) we say that the interacting molecules experience a repulsive dipole-dipole force which pushes the molecules away from one another. If partial charges of opposite sign (i.e. positive interacts with negative) interact, we say that the interacting molecules experience an attractive dipole-dipole force which attracts the molecules to one another.

Figure \(\PageIndex{1}\): Dipole-dipole interactions. Permanent dipoles established through the covalent interaction between atoms A and B can interact via dipole-dipole interactions. Atom B is more electronegative than atom A and thus recruits electrons near it causing an imbalance in charge around the molecule (this is depicted by an oblong electron cloud around nuclei for atoms A and B) - more negative is red; more positive is blue. The imbalance of charge creates a dipole with partial negative charges (delta-) and partial positive charges (delta+). The figure depicts various ways in which two of these dipoles can interact with one another that, depending on orientation, lead to either attractive or repulsive dipole-dipole interactions.
Attribution: Marc T. Facciotti (original work)

The core idea underlying the formation of dipole-dipole interactions - the interaction between two permanent dipoles - should sound familiar. Recall that we describe a hydrogen bond as an electrostatic interaction between a partially charged hydrogen (the positive end of a bond dipole) with a partial negative charge from the negative end of a different bond dipole. While the hydrogen bond has some special properties not discussed in this text, you can think of it as a sub-type of attractive dipole-dipole interaction.

For a deeper dive into dipole-dipole interactions, see this LibreText Chemistry reading.

Van der Waals Forces

All molecules can experience Van der Waals forces, a type of molecular interaction found when molecules get very close together, typically at distances between 4-5 Angstroms. Just for reference, recall that 1 Angstrom = 10^-10 meters. Van der Waals forces are, yet again, based on the attraction or repulsion of electrical poles. However, unlike the dipole-dipole interactions discussed above that arise from the interaction between permanent dipoles in molecules, the Van der Waals forces arise from the spontaneous and/or induced transient polarization of molecules. The local polarization (i.e. polarization on a part of a molecule) may last only a short time as electrons dynamically redistribute. When two molecules are close together one of them may spontaneously form a transient dipole (or more accurately, multipole). In response, the second molecule may “sense” the partial charge nearby and react by adjusting its own charge distribution in response, thus becoming polar itself. We say that the first dipole/multipole induces the formation of the second. If the two molecules remain between 4 and 5 Angstroms long enough, this process can repeat, and even synchronize, leading to molecular attraction at very short distances. At distances closer than ~4 Angstroms, electron clouds can overlap and this creates repulsive interaction. While all molecules can engage in Van der Waals interactions, in introductory biology we usually introduce students to these interactions in a discussion of lipid membrane structures. As you will soon discover, Nature creates biological membranes by packing many lipid molecules together at distances that allow many simultaneous Van der Waals interactions to occur. Collectively, these many small and transient interactions between lipid molecules contribute to the stability of membrane structures.

Figure \(\PageIndex{1}\): Schematic of dynamic, induced dipoles involved in Van der Waals forces. Molecules composed of atoms A are depicted approaching one another near the top of the figure. The dynamics (change in time) of these molecules is depicted row-by-row as they change in through arbitrary jumps in time. This dynamic nature is critical to the formation of Van der Waals interactions. Attribution: Marc T. Facciotti (original work)

Pi Interactions

Pi interactions are a type of molecular interaction that biologists typically encounter in when discussing stabilizing interactions in nucleic acid and protein structures. In a course of general biology, you may also encounter pi interactions in a discussion of protein-DNA interactions. These types of interactions derive their name from the involvement of pi bonds, a specific type of covalent bond between two atoms in which neighboring electron orbitals are close enough to overlap. We’ll leave the underlying discussion of molecular orbital theory for your chemistry course and just say that we usually associate pi bonds with double or triple covalent bonds. In biology, these types of bonds occur in many kinds of molecules, particularly those with so called conjugated pi systems including aromatic ring structures like those seen in some amino acids, vitamins and cofactors, and nucleic acids.

Figure \(\PageIndex{1}\): Examples of molecules with conjugated pi systems. Each of the biomolecules has at least a portion of it that contains so called conjugated pi bonds that can engage in pi interactions with other molecules. With the exception of the retinol molecules the conjugated pi systems are present in planar ring structures. The retinol's conjugated pi system is in the linear hyrodrocarbon portion of the molecule. Arrows point to examples of double bonds. Yellow highlights the systems of delocalized electrons. Attribution: Marc T. Facciotti (original work)

The distribution of electrons within these pi systems can create regions of more negative and more positive charge and thus creating areas that may “attract” or “repel” other charged or partially charged molecules depending on their relative alignments to one another. Again, we can reserve a deeper discussion of pi systems for your upper division chemistry classes. For now, simply appreciate that - once again - the unique properties of the elements that make up molecules contribute to how electrons distribute within those molecules. The often uneven distributions of electrons about the molecule can create local positive and negative charges and when these regions of partial charges on different molecules (or different parts of molecules) come together in appropriate orientations, that electrostatic interactions (attractive and repulsive) can happen.

A TAKE-HOME POINT ON MOLECULAR INTERACTIONS

Hopefully, you appreciate a common theme to our discussion of non-covalent molecular interactions. Whether we consider ionic bonds, hydrogen bonds, dipole-dipole interactions, Van der Waals forces, or pi interactions, all share the feature of being interactions between full or partial electrostatic charges. The key differences between each of these types of interactions types have to do with how the charges arise on molecules (i.e. the atomic basis for the charge) and/or how the charges interact. This depends, of course, on the underlying unique chemical properties of each element and how they behave at the subatomic level with one another - a suitable topic for discussion in your chemistry class.