Van der Waals forces and other non-covalent interactions

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 General 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.