Video Transcript
In this video, we will learn how to describe and explain hydrogen bonding and the effect it has on the physical properties of molecules.
In chemistry, we talk a lot about the attractions and attractive forces that happen on a very small scale. Some attractive forces are what we call intermolecular forces. Intermolecular forces are forces that occur between different molecules. The prefix inter- means between, much like how an international issue is an issue between two or more countries. The reason that water forms droplets made up of many molecules instead of, say, dispersing into individual molecules is because of the intermolecular forces holding those water molecules together.
The opposite of an intermolecular force is an intramolecular force. Intramolecular forces occur within one molecule. For example, the covalent bonds that hold together the oxygen atom and the hydrogen atoms of a water molecule are intramolecular forces. They’re the forces within the molecule that hold it together.
Another category of force is electrostatic forces. Electrostatic forces are forces that occur because of the electrical charge of the particles involved. Since we’re talking about attractive forces, we’ll focus on the electrostatic attraction between a positively charged particle and a negatively charged particle. But it is worth noting that an electrostatic repulsion can occur when the charges are the same. Electrostatic attractions can occur on both the intramolecular and the intermolecular level. For example, an ionic bond between a positive and a negative ion is an intramolecular force that is also an electrostatic force.
One example of an electrostatic intermolecular force occurs between polar molecules. Polar molecules, like the hydrogen bromide drawn here, have a slightly positive end and a slightly negative end. There’s an electrostatic intermolecular force that attracts the partially positive end of one polar molecule to the partially negative end of another molecule. Since polar molecules are also called dipoles, we call this attraction a dipole–dipole attraction. The topic of this video is hydrogen bonding. Hydrogen bonding is an intermolecular force and also an electrostatic force that occurs in a very specific situation. Now that we’ve learned about attractive forces in general, let’s take a specific look at hydrogen bonding.
There are two key conditions necessary to form a hydrogen bond. The first condition is that there must be a strongly electronegative atom. Fluorine with an electronegativity of 4.0, oxygen with an electronegativity of 3.5, and nitrogen with an electronegativity of 3.0 all satisfy this condition. It is worth noting that in certain exceptional molecules, other atoms such as sulfur and chlorine can form hydrogen bonds as well. Chlorine has a high electronegativity value of 3.2 but will only form hydrogen bonds in special cases. While the precise reason for this is beyond the scope of this video, it has to do with the fact that chlorine’s larger radius spreads its charge out over a larger space.
The second key condition to form a hydrogen bond is that the electronegative atom needs to be bonded to a hydrogen atom. Hydrogen will share its lone electron with the other atom in a covalent bond. Two atoms next to one another that satisfy these two conditions can form a hydrogen bond between the hydrogen atom of one molecule and the electronegative atom of the other molecule. In a solution where many such molecules are placed together, each molecule can form one or more hydrogen bonds with its surrounding molecules.
It’s worth noting that the electronegative atom and the hydrogen atom do not have to be the only atoms present in the molecule. For example, alcohols like methanol satisfy both conditions. They have an electronegative atom, oxygen, bonded to a hydrogen atom. So a hydrogen bond can form between the hydrogen atom of one molecule and the oxygen atom of an adjacent molecule.
It’s also worth noting that the molecules do not have to all be the same for hydrogen bonds to form. For example, if we mixed together solutions of the two compounds drawn here, we would see hydrogen bonds between the electronegative atom of one type of molecule and the hydrogen atoms of the other type of molecule. So we know what is required to form a hydrogen bond. But the question remains, why does the hydrogen bond form? What is so special about this particular arrangement? To answer this question, we need to take a look at the electrons involved. The covalent bond between chlorine and hydrogen represents a pair of shared electrons. Since fluorine is more electronegative than hydrogen, the shared electrons will be drawn closer to the fluorine atom and farther away from the hydrogen atom.
Since the hydrogen atom starts with just one electron, sharing that electron gives the hydrogen atom a particularly low electron density. As a result, it has a particularly strong partial positive charge. Conversely, fluorine side of the molecule has a strong partial negative charge due to the unequal sharing of electrons. The attraction between the positive end of one molecule and the negative end of an adjacent molecule in this special situation is called a hydrogen bond. Specifically, the positively charged hydrogen atom is attracted to a lone pair on the fluorine atom. We should also recognize that hydrogen bonds are about two times longer than the covalent bonds that hold together the molecule, although both types of bonds can vary in size.
In this example and others, we don’t always see these bonds drawn proportionally. But if we did, it would look something like this with the hydrogen bond approximately twice as long as the covalent bonds. We can summarize this situation with a definition. A hydrogen bond is an intermolecular electrostatic attraction between a hydrogen atom with a strong partial positive charge and a strongly electronegative atom on another molecule. Hydrogen bonding is a particularly strong intermolecular electrostatic attraction. It’s a strong attraction because the partial charges on either side of the hydrogen bond are also particularly strong.
Let’s take a closer look at the strength of hydrogen bonds compared to other forces. When sorting attractive forces by strength, the first comparison we can make is that intramolecular forces are stronger than intermolecular forces. It makes sense that the forces holding together a single molecule would be stronger than the forces holding together two separate molecules, as it should be easier to pull apart two separate molecules than to pull apart the constituent parts of a single molecule. Within these two categories, we can sort the individual forces by strength based on a simple principle.
Within each of these two categories, the greater the difference in charge, the stronger the attraction between the particles. Two of the strongest intramolecular forces are metallic bonds and ionic bonds. These bonds involve positive and negative ions or positive ions and electrons with full charges and therefore a very strong attraction. Descending in strength, we have polar covalent bonds that contain atoms with only partial charges due to the uneven sharing of electrons. Nonpolar covalent bonds are an even weaker attractive force because there’s no charge difference between the particles involved in the bond.
Looking at various intermolecular forces, we can sort by strength following the same pattern. Both hydrogen bonds and dipole–dipole attractions involve partial charges. However, the large electronegativity difference of a hydrogen bond makes the partial charges particularly strong, resulting in an attraction that’s stronger than an ordinary dipole–dipole attraction. A dipole-induced dipole attraction is weaker than a dipole–dipole attraction because instead of being between two particles with opposite partial charges, it is between one particle with a partial charge and one nonpolar molecule. The nonpolar molecule develops only a weak temporary area of charge in the presence of the other molecule.
The weakest intermolecular force, London dispersion forces, occur between any pair of molecules, even those without any difference in charge between them. More detail about these various forces can be found in other videos. Overall, we can see that a hydrogen bond is a strong intermolecular attraction. The presence of hydrogen bonding can affect the physical properties of a substance, most notably the boiling point. For example, water molecules form hydrogen bonds between the oxygen atom of one molecule and the hydrogen atom of a neighboring molecule. Hydrogen sulfide, on the other hand, has no hydrogen bonds, although there are the weaker dipole–dipole attractions.
When we boil a substance like water, we’re adding energy to the molecules until they become a gas. Since there’s a strong attractive force between a water molecule and the water molecules that surround it, it requires more energy for that water molecule to transition to the gaseous state. This energy requirement raises the boiling point of the substance. Conversely, when we boil hydrogen sulfide, a molecule that does not form hydrogen bonds, there are only the weaker dipole–dipole interactions to overcome. As a result, it requires less energy to transition to the gaseous state. And the boiling point will be lower than similar substances that do form hydrogen bonds. Overall, we can see that the presence of hydrogen bonds raises the boiling point of the substance.
The number of hydrogen bonds that a molecule forms also raises the boiling point. Each water molecule has two lone pairs on the oxygen atom and two hydrogen atoms. So each molecule can form a total of four hydrogen bonds. Hydrogen fluoride can also form hydrogen bonds. With one hydrogen bond at either end, each molecule forms two hydrogen bonds. Each bond forms between the hydrogen atom of one molecule and a lone pair on another. Although there are three total lone pairs on the fluorine atom, there are only enough hydrogen atoms to form a hydrogen bond with one lone pair per molecule. As a result, the number of hydrogen bonds is limited to two per molecule.
How would the number of hydrogen bonds affect the boiling point? Water has more hydrogen bonds per molecule than hydrogen fluoride. With a stronger attraction between the molecules, more energy is required to transition the molecules to the gaseous state. As a result, the boiling point is higher. In comparison, hydrogen fluoride has fewer hydrogen bonds per molecule. So less energy is required to transition the molecules to the gaseous state. And its boiling point will be lower than water’s, although its boiling point will still be higher than the boiling point of similar molecules without hydrogen bonding at all.
We can observe the intensity of the effect of hydrogen bonding on boiling point by taking a look at this graph that shows the boiling points of various hydrides or combinations of hydrogen and another element. As represented by the key on the graph, each colored line represents a group on the periodic table. For example, the green line represents hydrides of group 16 elements. And as we move to the right in the graph, we move down in the group, from oxygen to sulfur to selenium to tellurium. If we look at the second, third, and fourth points along each line, we can see that as the molecular mass of the hydride increases, the boiling point of the hydride increases with it.
If we take, say, hydrogen sulfide and replace the sulfur with selenium or tellurium, the elements one and two periods below sulfur on the periodic table, the structure of the molecule will be very similar. But the molecular mass will be higher, and so will the boiling point. While the precise mechanisms of this trend are outside the scope of this video, we can briefly say that a compound with a higher molecular mass has stronger London dispersion forces between its molecules. Stronger attractive forces means a higher boiling point. However, if we look at the leftmost portion of the graph, we can see a few points that clearly deviate from this pattern. Of the hydrides included in this graph, these three are the ones that form hydrogen bonds.
As mentioned earlier, the electronegativity difference in these three bonds is large enough to allow for the hydrogen bond to form between molecules. Even though these three molecules are of a lower molecular mass than the other hydrides in the graph, the presence of hydrogen bonds raises their boiling point. The fact that CH4 or methane, which does not form hydrogen bonds, follows the trend based on mass supports the notion that these three abnormally high values above methane on the graph are due to the presence of hydrogen bonds between those molecules. In addition, water, which forms four hydrogen bonds per molecule, has the highest boiling point of any hydride on the graph.
Hydrogen fluoride and ammonia, on the other hand, have two hydrogen bonds per molecule. With fewer hydrogen bonds per molecule, there’s a weaker attractive force between any two given molecules. And therefore, the compound has a lower boiling point. Overall, we can quantitatively see in this graph that the presence of hydrogen bonds, as well as the number of hydrogen bonds formed per molecule, can increase the boiling point of a compound.
Hydrogen bonds also play an important role in biology. You may be familiar with a double helix shape of a DNA strand. What might not be clear based on the shape alone is that the double helix structure of DNA is held together by hydrogen bonds. Each rung of the ladder is really two separate bases sticking inward and joined in the middle by two or three hydrogen bonds. The number and alignment of these hydrogen bonds ensures that adenine is always opposite thymine and cytosine is always opposite guanine and vice versa. These four bases are often shortened to the first letter of their name, A, C, T, and G. The instructions in our genetic code are written with these four letters.
The hydrogen bonds between the two opposite sides of the helix are complementary to one another, allowing the two strands to be separated, copied, and reattached. DNA contains instructions for cells to make proteins by arranging amino acids in a series. These proteins can then carry out various functions in and around the cell. However, the function of a protein is dependent not just on the order of the amino acids in the chain, but also how this chain twists and folds depending on attractions and repulsions between different regions of the chain. The coils and sheets that form in a folded protein are in part due to the hydrogen bonds between different amino acids along the chain. Without hydrogen bonding, proteins would be unable to form their particular shape to complete their particular function. Hydrogen bonding is critical to the function of life on the smallest scales.
Let’s review the key points of the video. Hydrogen bonding is a strong intermolecular attractive force. Hydrogen bonding occurs between a strongly electronegative atom on one molecule and a hydrogen atom with a strong partial positive charge on another. This usually means that hydrogen bonds will form in compounds with FH, NH, or OH bonds, although there are other arrangements that form hydrogen bonds as well. Substances with hydrogen bonds have higher boiling points than other molecules of similar size. Lastly, hydrogen bonds are critical to the structure of DNA and proteins.
