In this explainer, we will learn how to describe and explain hydrogen bonding and the effect it has on the physical properties of molecules.
Hydrogen bonds are dipole–dipole interactions that exist between covalently bonded hydrogen atoms and electron lone pairs on strongly electronegative atoms. Hydrogen bonds usually form between the partially positively charged () end of one molecule and the partially negatively charged () end of a second molecule.
Definition: Hydrogen Bond
Hydrogen bonds are dipole–dipole interactions that exist between covalently bonded hydrogen atoms and electron lone pairs on strongly electronegative elements.
Hydrogen bonds are a special class of dipole–dipole interaction that exists between molecules that have atoms with some of the highest possible electronegativity values. The highly electronegative atoms attract a significant amount of electron density from covalent bonds and they make molecules extremely polar and extremely effective at generating strong intermolecular forces.
Hydrogen bonds are usually only formed by molecules that contain at least one fluorine, oxygen, or nitrogen atom. The following table shows the strengths of different intermolecular forces of attraction, but it is important to realize that the data is both approximate and not comprehensive.
|Type of intermolecular force||Occurs Between||Involves||Approximate Strengths of Attraction (kJ/mol)|
|Dispersion||Nonpolar molecules||Temporary dipoles||0.5–2.5|
|Dipole–dipole||Polar molecules||Permanent dipoles||2.0–12.5|
|Hydrogen bonding||Polar molecules||Permanent dipoles between and , , or||15.0–35.0|
Water molecules have one highly electronegative oxygen atom that is covalently bonded to two hydrogen atoms. The oxygen atom withdraws a significant amount of electron density from the two covalent bonds, and this makes asymmetrical water molecules highly polar. The hydrogen atoms have a partial positive electrostatic charge, and the oxygen atoms have a partial negative electrostatic charge.
Each water molecule can form up to four hydrogen bonds, because they have two partially positively charged hydrogen atoms and two electron lone pairs. The following figure uses thin dotted lines for hydrogen bonds (labeled 1) to show how one water molecule can make up to four hydrogen bonds with other adjacent water molecules.
Example 1: Stating How Many Hydrogen Bonds Can Be Formed by a Single Water Molecule
What is the maximum number of hydrogen bonds that can be formed by one molecule of water?
Hydrogen bonds are strong intermolecular interactions that can form between neighboring molecules. Hydrogen bonds are formed between the electron lone pair of one molecule and the covalently bonded hydrogen atom of a second molecule. The covalently bonded hydrogen atom must be bonded to some highly electronegative atom such as fluorine, oxygen, or nitrogen. Each one of the covalently bonded hydrogen atoms can form one hydrogen bond with one lone pair of electrons.
Water molecules contain two hydrogen atoms that are covalently bonded to a single highly electronegative oxygen atom. The oxygen atom has a total of six outer-shell electrons, and four of these electrons essentially remain on one side of the water molecules. The electrons are arranged as two lone pairs of electrons that can each accept a single hydrogen bond. Each one of the lone pairs and the two covalently bonded hydrogen atoms can form a single hydrogen bond either with other water molecules or another type of complementary polar molecule. This statement can be used to determine that a single water molecule can form a total of four hydrogen bonds.
Hydrogen bonds tend to be both longer and weaker than covalent bonds. The bond length is between 0.95 angstroms and 1.00 angstrom and the hydrogen bond between two water molecules can be two or three times as long.
The following table compares data for most different types of covalent and hydrogen bonds. You will notice that hydrogen bonds tend to be two or three times longer than covalent bonds and that hydrogen bonds tend to be at least five times weaker than covalent bonds.
|Bond Length||Bond Strength (kJ/mol)|
|Covalent Bond||1.0–1.5||150–1 000|
Example 2: Identifying the Bond Enthalpies of Hydrogen Bonds
The table below shows the bond enthalpies for hydrogen bonds and covalent bonds between pairs of atoms. Which column corresponds to the bond enthalpies of hydrogen bonds?
|Pair of Atoms||Bond Enthalpies (kJ/mol)|
Hydrogen bonds are strong intermolecular interactions that can form between neighboring molecules. Hydrogen bonds generally only form between molecules that contain at least one highly electronegative atom that is covalently bonded to a hydrogen atom. The highly electronegative atom is almost always a nitrogen, oxygen, or fluorine atom. The covalently bonded hydrogen atom ends up having a partial positive charge () and the electronegative atom ends up having a partial negative charge ().
Hydrogen bonds are established as the partially positive end of one molecule interacts with the partially negatively charged end of an adjacent molecule. The intermolecular interactions are strong, and hydrogen bonds tend to have bond enthalpy values of 15–35 kJ/mol.
Covalent bonds are a type of chemical bond and intramolecular interaction. Covalent bonds are formed as the valence electrons of one atom overlap with the valence electrons of a second atom. Covalent bonds are much stronger than intermolecular hydrogen bonds. Covalent bonds tend to have bond enthalpy values of 150–1 000 kJ/mol.
Column A shows values that are within the 15–35 kJ/mol range, but column B shows bond enthalpy values that lie outside this range. Column A shows bond enthalpies that could correspond to the bond enthalpies of hydrogen bonds, and column B shows bond enthalpies that cannot correspond to the bond enthalpies of hydrogen bonds. We can use these statements to determine that option A must be the correct answer to this question.
Hydrogen sulfide () and water () molecules both have two hydrogen atoms and one atom from the sixteenth group of the periodic table. Hydrogen sulfide and water molecules have similar V-shaped (angular) geometries, but they have different intermolecular interactions and different boiling points.
Water has a relatively high boiling point of , because each water molecule can form up to four hydrogen bonds. Hydrogen sulfide has a much lower boiling of , because hydrogen sulfide molecules cannot induce the formation of any hydrogen bonds whatsoever.
It takes a lot of energy to separate water molecules, because they are held together with strong hydrogen bonds. It takes much less energy to separate hydrogen sulfide molecules, because they do not have highly electronegative atoms that can induce the formation of any hydrogen bonds. Sulfur atoms have an electronegativity value that is 25% lower than the electronegativity value of oxygen. The following graphic shows the similarities between the structures of water and hydrogen sulfide molecules.
The strength of a hydrogen bond depends on the strength of induced electric dipole moments. The most electronegative atoms withdraw the most electron density from covalent bonds, and this produces more intensely asymmetric electric dipole moments. The most intensely asymmetric dipole moments generate the strongest hydrogen bonds. Fluorine atoms usually withdraw more electron density than nitrogen atoms, and this explains why hydrogen fluoride () molecules generally produce stronger hydrogen bonds than ammonia () molecules. The following figure shows how a single hydrogen bond can form between two adjacent hydrogen fluoride molecules and two adjacent ammonia molecules.
Example 3: Identifying the Correct Hydrogen Bonding Diagram
Which of the following diagrams correctly shows the hydrogen bonding between two molecules of ?
Hydrogen bonds are strong intermolecular forces that exist between the covalently bonded hydrogen atom of one molecule and the lone pair of electrons on an adjacent molecule. The hydrogen atom must be covalently bonded to a fluorine, nitrogen, or oxygen atom. Fluorine atoms have the highest electronegativity value of any element in the periodic table, and the bond is highly polar. The fluorine atoms have a partial negative () electrostatic charge, and the hydrogen atoms have a partial positive () electrostatic charge. Hydrogen bonds form between the covalently bonded hydrogen atom of one hydrogen fluoride molecule and the lone pair electrons of an adjacent fluorine atom.
The hydrogen bonds between adjacent hydrogen fluoride molecules can be depicted with relatively basic images that use and chemical symbols to represent the hydrogen and fluorine atoms and single straight lines to represent the covalent bonds. The partial positive and negative electrostatic charge values are represented with the and symbols. The hydrogen bonds are represented as thin dotted lines and never as single-sided arrows.
The question is asking us to determine which graphic correctly depicts the hydrogen bonding between two hydrogen fluoride molecules. The correct graphic would be drawn with the tag next to the hydrogen atom () symbol and with the tag next to the fluorine atom () symbol. The correct graphic would also have to have the hydrogen bond being placed between the section of one molecule and the of an adjacent molecule. Option A is the only image that has all of its symbols and bonds in the appropriate positions, and we can determine that option A has to be the correct answer for this question.
Hydrogen fluoride and ammonia molecules have lower boiling points than water, because there are fewer hydrogen bonds between groups of hydrogen fluoride or ammonia molecules and more hydrogen bonds between groups of water molecules. It takes a moderate amount of heat energy to separate groups of ammonia or hydrogen fluoride molecules, because they are interlinked with a relatively low number of hydrogen bonds. It takes significantly more heat energy to break up groups of water molecules, because the water molecules are linked with a higher number of hydrogen bonds. The differences in hydrogen bonding capacity can be understood by considering the unequal number of lone pairs and partially positively charged hydrogen atoms in ammonia and hydrogen fluoride liquids.
Each ammonia molecule has three partially positively charged hydrogen atoms and a single lone pair of electrons. The ammonia molecules cannot create the same sort of hydrogen bond networks that are formed in liquid water. Each ammonia molecule will only be able to make an average of two hydrogen bonds with surrounding ammonia molecules. They tend to make one hydrogen bond with their single electron lone pair and a second hydrogen bond with one of their partially positively charged hydrogen atoms.
Each hydrogen fluoride molecule has one partially positively charged hydrogen atom and three lone pairs of electrons. The hydrogen fluoride molecules cannot create the same sort of hydrogen bond networks that are formed in liquid water. Each hydrogen fluoride molecule will only be able to make an average of two hydrogen bonds with surrounding hydrogen fluoride molecules. They tend to make one hydrogen bond with their single partially positively charged hydrogen atom and a second hydrogen bond with one of their electron lone pairs.
The fact that hydrogen fluoride is only able to make an average of two hydrogen bonds limits the forms that several molecules can make when joined by hydrogen bonds. Examples of potential straight-chain and closed-ring arrangements can be seen in the diagram below.
It was already stated that water has a higher boiling point than hydrogen sulfide, but it is interesting to note that hydrogen fluoride and ammonia also have higher boiling points than other comparable mono- and trihydride molecules. Boiling point data shows that hydride molecules almost always have higher boiling points if they contain highly electronegative atoms that can induce the formation of hydrogen bonds. The hydrogen bonds increase the force of attraction between neighboring molecules, and it takes more thermal energy to break them apart. Hydrogen fluoride () has a boiling point that is 105 degrees higher than the boiling point of hydrogen chloride (), and ammonia () has a boiling point that is 54 degrees higher than the boiling point of phosphine ().
Example 4: Understanding How Hydrogen Bonds Affect Boiling Points of Simple Molecular Compounds
The graph below shows the boiling points of hydrides of groups 14, 15, 16, and 17.
- Why are the boiling points of
and so much higher
than the other hydrides in their respective groups?
- The smaller size of the atoms means these molecules contain stronger covalent bonds.
- These molecules have higher vapor pressures compared with the other hydrides.
- These molecules can form network covalent structures, giving them higher boiling points compared with the other hydrides in the group.
- These molecules are able to hydrogen bond, giving them stronger intermolecular attractions compared with the other hydrides.
- The smaller size of these molecules means they can pack closer together.
- Why does the boiling point increase as you go down group 14?
- The elements become metallic in character and so experience metallic bonding, which is stronger than covalent bonding.
- The number of electrons in each atom increases, giving rise to greater van der Waals attractions between molecules.
- The bond strength between the group 14 atom and hydrogen atoms increases, meaning more energy is needed to separate them.
- The molecules are able to form stronger hydrogen bonds, increasing the boiling point.
- The reactivity of the molecules decreases, meaning more energy is required to change their state from liquid to gas.
Boiling points depend on the strength of intermolecular interactions between molecules. Materials have relatively high boiling points when there are strong intermolecular interactions between the constituent molecules. Materials have relatively low boiling points when there are weaker intermolecular interactions between the constituent molecules.
Hydrogen fluoride has a higher boiling point than other group seventeen hydrides, because hydrogen fluoride is the only group seventeen hydride that can form hydrogen bonds. Water and ammonia have higher boiling points than the other group fifteen and sixteen hydrides, because they are the only group fifteen and sixteen hydrides that can form hydrogen bonds. These statements can be used to determine that option D is the correct answer to this question.
Boiling points depend on the strength of intermolecular interactions between molecules. Some materials have relatively high boiling points because they are made up of molecules that can form strong intermolecular hydrogen bonds. Other materials have lower boiling points because they are made up of molecules that can only form weaker dispersion intermolecular interactions.
The group 14 tetrahydride compounds all have relatively low boiling points, because they are nonpolar molecules that are held together with nothing more than weak dispersion forces. Some of the group 14 tetrahydride compounds have higher boiling points than the other group 14 tetrahydride compounds because they are made up of nonpolar compounds that can form stronger dispersion interactions.
The strength of the dispersion force depends on the strength of the molecular dipole that can be induced in a molecule. Molecules can generate stronger electric dipole moments when they contain a greater number of electrons and a greater amount of negatively charged electron density that can be redistributed. The number of electrons increases as we move down group 14, and this explains why stannane () has a higher boiling point than germane () and why germane has a slightly higher boiling point than silane () and a much higher boiling point than methane (). This line of reasoning is summarized in option B, and we can conclude that option B must be the correct answer for this question.
Hydrogen bonds regulate the structure and functioning of some of the most important biological macromolecules, and this includes the -helical transmembrane proteins and double-helix strands of deoxyribonucleic acid (DNA). There are two or three hydrogen bonds between complementary nitrogenous base pairs in DNA and lots of hydrogen bonds between the backbone carbonyl and amide hydrogen atoms of most integral membrane proteins. Most DNA and -helical or -sheet protein structures would instantly lose their important shape and biological function(s) if they lost all or even some of their intramolecular hydrogen bonds. The following image shows how hydrogen bonds help to maintain the complex three-dimensional double-helix structure of DNA. The hydrogen bonds are depicted as thin dotted lines. There are two hydrogen bonds between the complementary thymine (T) and adenine (A) base pairs and three hydrogen bonds between the complementary cytosine (C) and guanine (G) base pairs.
Example 5: Stating How Many Hydrogen Bonds Link Guanine and Cytosine Base Pairs Together
The base pairing in DNA between molecules of guanine and cytosine is shown in the given structure. How many hydrogen bonds can be formed between molecules of guanine and cytosine?
Deoxyribonucleic acid (DNA) is a macromolecule that is composed of two polynucleotide chains. Each polynucleotide chain is made up of a sugar–phosphate backbone and different nitrogenous base pairs. The nitrogenous base pairs contain partially positively charged hydrogen atoms and highly electronegative atoms that have exposed electron lone pairs. The adenine base pair forms two hydrogen bonds with the complementary thymine base pair, and the guanine base pair forms three hydrogen bonds with the complementary cytosine base pair. The following figure shows how two and three hydrogen bonds are formed between complementary DNA base pairs. The first part of the figure shows how two hydrogen bonds form between the complementary adenine (A) and thymine (T) base pairs, and the second part shows how three hydrogen bonds form between the complementary guanine (G) and cytosine base pairs (C). The hydrogen bonds are depicted as red dotted lines.
This question is asking us to determine the number of base pairs between the guanine and cytosine complementary base pairs and not the number of bonds between the complementary adenine and thymine base pairs. The figure clarified that there are three hydrogen bonds between the guanine and cytosine base pairs, and we can conclude that three must be the correct answer for this question.
- Hydrogen bonds are dipole–dipole interactions that exist between covalently bonded hydrogen atoms and electron lone pairs on strongly electronegative elements like fluorine.
- Hydrogen bonds are stronger than dispersion and conventional dipole–dipole intermolecular interactions.
- Materials have relatively high melting points and boiling points when they are made up of molecules that can form hydrogen bonds.
- Hydrogen bonds give many polypeptides specific three-dimensional shapes that make them ideally suited for specific biological functions.
- There are three hydrogen bonds between complementary cytosine and guanine base pairs and two hydrogen bonds between complementary adenine and thymine base pairs.