7.3: Bonding of Hydrogen and Water (2023)

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    learning goals
    • Name three special properties of water that make it unusual for a molecule of its size, and explain how they result from hydrogen bonding.
    • Explain what hydrogen bonding means and what molecular structural features it gives rise to.
    • Describe the "structure" of liquid water.
    • Draw structural examples of hydrogen bonding in three different small H molecules.2Ö.
    • Describe the role of hydrogen bonding in proteins and DNA.

    Most chemistry students quickly learn to relate the structure of a molecule to its general properties. Therefore, we generally expect small molecules to form gases or liquids, and large molecules to exist as solids under normal conditions. And then we come to the H2Oh, and we were shocked to discover that many of the predictions are wildly wrong and that water (and therefore life itself) shouldn't even exist on our planet! In this section, we'll learn why this tiny combination of three nuclei and ten electrons has special properties that make it unique among the more than 15 million chemical species we know of today.

    In water, each hydrogen nucleus is covalently bonded to the central oxygen atom by a shared pair of electrons. in H2Or, only two of the six electrons in oxygen's outer shell are used for this purpose, leaving four electrons arranged in two nonbonding pairs. The four electron pairs surrounding oxygen tend to be as far apart as possible to minimize repulsions between these negatively charged clouds. This would normally result in a tetrahedral geometry, where the angle between the electron pairs (and thus the H-O-Hconnection angle) is 109.5°. However, as the two non-bonding pairs get closer to the oxygen atom, they exert more repulsion against the two covalently bonded pairs, effectively pulling the two hydrogen atoms closer together. The result is a distorted tetrahedral arrangement where the HOH angle is 104.5°.

    7.3: Bonding of Hydrogen and Water (2)

    The large dipole moment of water leads to the formation of hydrogen bonds.

    Vaya2The molecule is electrically neutral, but the positive and negative charges are not evenly distributed. This is illustrated by the color grading in the schematic representation here. The (negative) electronic charge is concentrated at the oxygen end of the molecule, due in part to the non-bonding electrons (solid blue circles) and the high nuclear charge of oxygen, which exert stronger attractive forces on the electrons. . This charge transfer forms aelectric dipole, represented by the arrow below; You can think of this dipole as the electrical "image" of a water molecule.

    7.3: Bonding of Hydrogen and Water (3)

    Opposite charges attract, so it's not surprising that the negative end of a water molecule tends to line up so that it's next to the positive end of another nearby molecule. the strength of itdipole-dipole attractionit is smaller than a normal chemical bond and is therefore completely dominated by the thermal motions common in the gas phase. However, when H.2With the O molecules clumped together in the liquid, these attractive forces exert a very noticeable effect that (somewhat misleadingly)hydrogen bond. And at temperatures low enough to quench the disruptive effects of thermal motion, water freezes into ice in which hydrogen bonds form a rigid, stable network.

    7.3: Bonding of Hydrogen and Water (4)

    Note that the hydrogen bond (represented by the green dashed line) is slightly longer than the covalent OH bond. this is alsomuch weaker, about 23 kJ mol–1compared to the O-H covalent binding strength of 492 kJ mol–1.

    Forty-one water anomalies”, some of them quite esoteric.

    It has long been known that water has many physical properties that distinguish it from other small molecules of comparable mass. Although chemists refer to them as the "anomalous" properties of water, they are not mysterious; all are entirely predictable consequences of the way the size and nuclear charge of the oxygen atom distort the electronic charge clouds of the atoms of other elements when chemically bound to oxygen.

    (Video) 7.3 - Periodicity, bonding and structure - part 2

    7.3: Bonding of Hydrogen and Water (5)

    Boiling point

    The most obvious peculiarity of water is itsvery high boiling pointfor such a light molecule. liquid methane CH4(molecular weight 16) boils at -161°C. As you can see in this diagram, extrapolating the boiling points of the various compounds from group 16 hydrogen to H2He suggests that this substance must be a gas under normal conditions.

    7.3: Bonding of Hydrogen and Water (6)

    surface tension

    Compared to most other liquids, water also has a highsurface tension. Have you ever seen an insect walk across the surface of a pond? EITHERwater walkerIt makes use of the fact that the surface of the water acts like an elastic film that does not deform when a small weight is placed on it. (If you're careful, you can also "float" a small paper clip or steel clip on the surface of water in a glass.)surface tensionof water A molecule in the middle of a liquid experiences attractive forces in all directions on neighboring molecules, but since the average is zero, there is no net force on the molecule. For a molecule that isNoSuperficially, the situation is very different; it experiences lateral and downward forces only, and this produces the effect of a stretched membrane.

    7.3: Bonding of Hydrogen and Water (7) 7.3: Bonding of Hydrogen and Water (8)

    The distinction between surface and deep level molecules is particularly pronounced in H2Or, due to strong hydrogen bonding forces. The difference between the forces that a molecule is subjected to at the surface and those in the volume of the liquid results in the surface tension of the liquid. This drawing highlights two H2The molecules, one on the surface and one in most of the liquid. The molecule on the surface is attracted to its neighbors below and on both sides, but there are no points of attraction pointing to the solid 180° angle above the surface. As a result, a molecule on the surface tends to be attracted to the interior of the liquid. But since there must always be some surface area, the overall effect is to minimize the surface area of ​​a liquid.

    The geometric figure with the smallest surface area/volume ratio is theKugel, so very small amounts of liquid tend to form spherical droplets. As the drops grow, they deform their weight in the typical teardrop shape.

    7.3: Bonding of Hydrogen and Water (9)

    ice floats on water

    The energetically most favorable configuration of H2The molecule is one in which each molecule is linked to four neighboring molecules through hydrogen bonds. Due to the thermal motions described above, this ideal is never achieved in liquid, but when water turns to ice, the molecules are stored in exactly this type of arrangement in the ice crystal. This arrangement requires the molecules to be a bit further apart than they otherwise would be; Consequently, ice, in which hydrogen bonding is maximal, has a more open structure and therefore a lower density than water.

    7.3: Bonding of Hydrogen and Water (10)7.3: Bonding of Hydrogen and Water (11)

    Here are three-dimensional views of a typical local structure of water (left) and ice (right). Note the greater openness of the ice structure required to ensure the strongest degree of hydrogen bonding in a regular extended lattice. The most crowded and messy arrangement in liquid water can only be sustained with the greatest amount of heat energy available above freezing.

    (Video) Intermolecular Forces for H2O (Water)

    When the ice melts, the increased thermal motion destroys much of the hydrogen bonding structure, allowing the molecules to stick to each other more tightly. Water is therefore one of the few substances whose solid form has a lower freezing point density than the liquid. However, localized groups of hydrogen bonds still remain; They break down and reform all the time as thermal movements push and push against the individual molecules. As the temperature of the water rises above freezing, the size and lifetime of these clumps decrease, causing the density of the water to increase.

    7.3: Bonding of Hydrogen and Water (12)

    At higher temperatures, another effect common to all materials begins to dominate: the amplitude of thermal motion increases with increasing temperature. This increase in agglomeration causes the average distance between the molecules to increase, reducing the density of the liquid; This is normal thermal expansion.

    Since the two competing effects (hydrogen bonding at low temperatures and thermal expansion at higher temperatures) lead to a decrease in density, it follows that there must be a temperature at which the density of water passes through a maximum. This temperature is 4 °C; This is the temperature of the water found at the bottom of an ice-covered lake, where the densest of all water has driven the coldest water up to the surface.

    liquid water structure

    The nature of liquid water and how H2How O molecules are organized and how they interact are questions that have interested chemists for many years. There is probably no liquid that has been studied more extensively, and there is now a vast literature on the subject. The following facts are well documented:

    • H2O molecules are attracted to each other through the special type of dipole-dipole interaction known as hydrogen bonding.
    • a group of hydrogen bonds in which four H2Those located at the corners of an imaginary tetrahedron are a particularly favorable configuration (low potential energy),but...
    • Molecules undergo rapid thermal motion on a picosecond (10 to 10) time scale.–12b) so that the lifetime of a particular cluster configuration is temporarily short.

    Various techniques have been used to study the microscopic structure of water, such as infrared absorption, neutron scattering and nuclear magnetic resonance. The information obtained from these experiments and theoretical calculations led to the development of some twenty "models" that attempt to explain the structure and behavior of water. More recently, computer simulations of various kinds have been used to investigate how well these models can predict the observed physical properties of water.

    This work led to a gradual refinement of our views on the structure of liquid water, but did not produce a definitive answer. There are several reasons for this, but the main one is that the very concept of "structure" (and "clumps" of water) depends on both the time period and the volume considered. Therefore, the following questions remain open:

    • How are members of a "group" distinguished from neighboring molecules that are not in that group?
    • Since individual hydrogen bonds are continually broken and reformed on the picosecond time scale, do clumps of water have significant existence for long periods of time? In other words, the groups are transient, while "structure" implies a more permanent molecular arrangement. Can we then legitimately use the term "group" to describe the structure of water?
    • The possible locations of neighboring molecules around a given H2They are limited by energetic and geometric considerations, resulting in some degree of "structure" within each small volume element. However, it is not clear to what extent these structures interact when the size of the volume element is increased. And as mentioned above, how long do these structures persist for more than a few picoseconds?

    In the 1950s, liquid water was thought to consist of a mixture of hydrogen bonding groups (H2Ö)nortein whichnorteit can have a variety of values, but no evidence for the existence of such aggregates has ever been found. The current view, supported by computer modeling and spectroscopy, is that, on a very short time scale, water is more like a "gel" made up of a single giant clump held together by hydrogen bonds. in a 10–12-10–9Time scale, rotations, and other thermal motions cause individual hydrogen bonds to break and reform into new configurations, inducing constantly changing local discontinuities whose extent and influence depend on temperature and pressure.


    esLike all solids, it has a well-defined structure; Each water molecule is surrounded by four neighboring H's.2You. two of which are hydrogen bonds with the central oxygen atom H2O molecule, and each of the two hydrogens is similarly bonded to an adjacent H2Ö.

    7.3: Bonding of Hydrogen and Water (13)

    (Video) Orgo 7.3 - alkene hydration

    Ice forms crystals with a hexagonal lattice structure which, when fully developed, would tend to form hexagonal prisms very similar to those sometimes seen in quartz. This happens from time to time, and anyone who has done a lot of winter mountaineering has probably seen prisms of needle-shaped ice crystals floating in the air. Under most conditions, however, the snowflake crystals we see flatten out into the beautiful, commonly observed hexagonal fractal-like structures.


    7.3: Bonding of Hydrogen and Water (14)Vaya2The O molecules, which make up the top and bottom surfaces of the prism, are very densely packed and connected (via hydrogen bonds) to the molecules inside. In contrast, the molecules that make up the sides of the prism, and particularly those at the hexagonal corners, are much more exposed than atmospheric H.2The O molecules, which come into contact with most places on the surface of the crystal, bind very loosely and migrate along it until they can form hydrogen bonds at these corners, thus becoming part of the solid and then throughout him. structure to expand these six directions. This process continues as new extensions take on a hexagonal structure.

    Why is the ice soft?

    At temperatures as low as 200 K, the surface of ice is very messy and looks like water. As the temperature approaches freezing, this area of ​​clutter spreads further below the surface and acts as a lubricant.7.3: Bonding of Hydrogen and Water (15)

    The illustration is from an article in the April 7, 2008 issue of C&EN honoring physical chemist Gabor Somorjai, a pioneer in modern methods for studying surfaces.

    "Pure water

    To a chemist, the term "pure" only has meaning in the context of a specific application or process. The distilled or deionized water we use in the laboratory contains dissolved atmospheric gases and occasionally some silica, but their small amounts and relative inertness make these impurities insignificant for most purposes. When water of the highest possible purity is required for certain types of demanding measurements, it is typically filtered, deionized, and triple distilled under vacuum. But this "chemically pure" water is also a mixture of isotopes: there are two stable isotopes of hydrogen (H1eh2, the latter often denoted by D) and oxygen (Osixteeny-oh18) giving rise to combinations like H2o18, HDOsixteenetc., all easily identifiable in the infrared spectrum of water vapor. Furthermore, the two hydrogen atoms in water contain protons whose magnetic moments can be parallel or antiparallel, creatingHuerta-miPro-water, respectively. Both forms are usually present in ao/pratio of 3:1.

    The amount ofrare isotopes of oxygen and hydrogen in waterit varies so much from place to place that it is now possible to determine the age and origin of a given water sample with some precision. These differences are reflected in the H and O isotopic profiles of the organisms. Therefore, isotopic analysis of human hair can be a useful tool for law enforcement and anthropological research.

    More about hydrogen bonds

    7.3: Bonding of Hydrogen and Water (16)Hydrogen bonds form when the electron cloud of a hydrogen atom attached to one of the more electronegative atoms is distorted by that atom, leaving a partial positive charge on the hydrogen. Due to the very small size of the hydrogen atom, the density of this partial charge is great enough to allow it to interact with the lone pair of electrons of a neighboring electronegative atom. Although hydrogen bonding is commonly described as a form of dipole-dipole attraction, it is now clear that it also involves some degree of electron sharing (between the nonbonding outer electrons and hydrogen), thus these bonds have a covalent character. . .

    Hydrogen bonds are longer than ordinary covalent bonds, and they are also weaker. Experimental evidence for hydrogen bonding usually comes from X-ray diffraction studies in solids, which show smaller than normal distances between hydrogen and other atoms.

    hydrogen bonds in small molecules

    The following examples demonstrate some of the broad applications of hydrogen bonding in molecules.

    7.3: Bonding of Hydrogen and Water (17) Ammonia(mp -78, bp -33°C) is a liquid and solid hydrogen bond.
    7.3: Bonding of Hydrogen and Water (18) Hydrogen bonds are responsible for this.AmmoniaIt is remarkably high solubility in water.
    7.3: Bonding of Hydrogen and Water (19) Large amountorganic acids (carboxylic)form dimers linked by hydrogen bonds in the solid state.
    7.3: Bonding of Hydrogen and Water (20) Here, the hydrogen bond acceptor is the π electron cloud of a benzene ring. This type of interaction is important in maintaining the shape of proteins.
    7.3: Bonding of Hydrogen and Water (21)

    hydrogen fluoride(mp -92, bp 33°C) is another common substance that is strongly hydrogen bonded in its condensed phases.

    7.3: Bonding of Hydrogen and Water (22) Obifluoruro(for which no proper Lewis structure can be written) can be viewed as a complex ion held together by the strongest known hydrogen bond: about 155 kJ mol–1.
    7.3: Bonding of Hydrogen and Water (23) 7.3: Bonding of Hydrogen and Water (24)"Slow as molasses in winter!"Multiple hydroxyl groups offer many opportunities for the formation of hydrogen bonds and lead to high viscosities of substances such as, for example,glycerinmisugar syrups.
    (Video) Types of hydrogen bonds with examples | Intermolecular and Intramolecular Bonding - Dr K

    Hydrogen bonds in biopolymers

    Hydrogen bonding plays an essential role in bio-based natural polymers in two ways:

    • hydrogen bonds between adjacent polymer chains (intermolecular bonding);
    • hydrogen bonds between different parts of the same chain (intramolecular bonding;
    • Hydrogen bonding of water molecules to -OH groups on the polymer chain ("bound water"), helping to maintain the polymer's shape.

    The following examples are representative of various types of biopolymers.


    7.3: Bonding of Hydrogen and Water (25)celluloseit is a linear glucose polymer (see above) containing from 300 to more than 10,000 units, depending on the source. As the most important structural component of plants (along with lignin in trees), cellulose is the most abundant organic substance on Earth. The role of hydrogen bonds is to join individual molecules to build sheets as shown here. These sheets are then stacked in a staggered arrangement, held together by van der Waals forces. Additional hydrogen bonds from adjacent stacks bring them together into a stronger, more rigid structure.


    These polymers of the amino acids R-CH(NH2)COOHs depend on intramolecular hydrogen bonding to maintain their shape (secondary and tertiary structure), which is essential for their important role as biological catalysts (enzymes). The hydrogen-bonded water molecules incorporated into the protein are also important for its structural integrity.

    Primary hydrogen bonding in proteins occurs between the -NH groups of "amino" residues and the -C=O groups of "acid" residues. These interactions result in two main types of secondary structure, which are related to the arrangement of the polymeric amino acid chain:

    7.3: Bonding of Hydrogen and Water (26)

    alpha helix


    7.3: Bonding of Hydrogen and Water (27)

    beta sheet

    Although carbon is not normally considered to be particularly electronegative, CH----X hydrogen bonds are now also known to be important in proteins.

    DNA (deoxyribonucleic acid)

    Who you are depends entirely on the hydrogen bond!As you probably know, DNA is the most famous of the biopolymers, as it plays a central role in defining the structure and function of all living organisms. Each strand of DNA consists of a sequence of four differentnucleotide monomerswhich consists of adeoxyribose sugar, phosphate groups, it's abase nitrogenadacommonly identified by the letters A, T, C, and G. DNA itself consists of two of thempolynucleotide chainscoiled around a common axis in a configuration similar to the alpha helix of the protein shown above. The sugar and phosphate backbones are on the outside, so the nucleotide bases are on the inside, facing each other. The two strands are held together by hydrogen bonds, which connect a nitrogen atom on one nucleotide on one strand with a nitrogen or oxygen atom on the opposite nucleotide on the other strand.

    (Video) Hydrogen bonds drawing and explanation

    7.3: Bonding of Hydrogen and Water (28)

    Efficient hydrogen bonding within this configuration can only occur between the A-T and C-G pairs, so these two complementary pairs form the "alphabet" that encodes the genetic information that is transcribed when new protein molecules are built. Hydrogen-bonded water molecules to the outer parts of the DNA helix help stabilize it.


    1. Notes for IB Biology chapter 7.3
    (Cheryl Hickman)
    2. ALEKS: Identifying hydrogen bonding interactions between molecules
    (Roxi Hulet)
    3. 7.3 - Periodicity, bonding and structure - part 1
    (A Level Chemistry Exam Support)
    4. Why is water polar? Why does water have a bent shape?
    5. 7.3 - Solids
    (K. Pluchino)
    6. Properties of Water


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