Module 2.2: Water and Hydrogen Bonds - Biology

Module 2.2: Water and Hydrogen Bonds - Biology

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learning objective

  1. Explain the molecular structure of Water.

  2. Identify hydrogen bond donors and acceptors.

  3. Predict the strength of hydrogen bonds based on geometry.

  4. Predict the solubility of compounds in water.

Polar Bonds and Molecules

A bond is considered to be polar if there is a significant difference in the electronegativities of the participating atoms. Electronegativity is a measure of an atom's attraction for electrons; a more electronegative atom will pull some electron density from other bonded atoms that are more electropositive. The following table gives the electronegativities of atoms that are common in biochemistry. These values can be used to estimate the partial charge on atoms in molecules. For example, since the electronegativity of hydrogen is smaller than C, S, N, and O, any bonds between hydrogen and these atoms will result in a partial positive charge on the hydrogen and a negative charge on the other atom. The larger the difference in electronegativity, the larger the difference in partial charges.


The dipole moment, (mu), is defined by the following equation:

where (q) is the charge on the atom and r is the distance to the center of mass.

[mu=sum space_{Allspace atoms} space qr onumber]

A polar molecule will have an overall net dipole moment. It is possible for a non-polar molecule to have polar bonds. For example, carbon dioxide (O=C=O) contains two polar bonds, but the dipole moment of one bond cancels the other, leading to no net dipole and therefore a non-polar molecule.

The ammonia molecule (NH3) has three identical polar N-H bonds that are equally spaced around the nitrogen atom. Ammonia has a net dipole moment of 1.4 D, similar to that of water (1.85 D).

Based on this information, do you think ammonia is a planer molecule or not? Briefly justify your response.


What orientations of the individual dipole would give rise to a net dipole of zero?

Ammonia is not planer. If it were planer than the dipole moments associated with each N-H bond would cancel, giving a net dipole of zero.

Structure of Water

  1. Oxygen has the following electronic configuration: (1s^22s^22p^4).
  2. In water, the 2s and the three 2p orbitals form four (sp^3) hybrid orbitals.
  3. These orbitals are tetrahedral in their orientation, however, the ideal bond angle of (109^{circ}) is distorted to (104.5^{circ}).
  4. The orbitals are populated such that two orbitals are filled and two contain one electron each.
  5. The filled orbitals cannot form bonds and are called lone pairs of electrons.
  6. The half-filled orbitals participate in the formation of a sigma bond between oxygen and hydrogen.
  7. "Bent" water molecule generates a permanent dipole moment, making water a polar solvent.

learn by doing

Water and Ice

Drag the model with your mouse to rotate it in 3D space. Use your right mouse button or "control" click on the Jmol to bring up a menu of options for manipulating the model.



1. Identify the following atoms based on the colors from the figures above


focus on a single water molecule.


Red: Oxygen; White: Hydrogen

2. Which set of atoms on different molecules will have a greater distance between them? Hydrogen to Hydrogen or Hydrogen to Oxygen?


check the actual distance


Hydrogen to Hydrogen (The same positive partial charges of the hydrogen atoms causes them to slightly repel each other. The opposite partial charges of the oxygen and hydrogen atoms cause them to attract each other.)

3. Describe the relative orientation of the hydrogen atoms (white) with respect to the oxygen atoms (red). How does this orientation differ between water and ice?

What is the physical basis for this orientation?


It may be helpful to rotate the molecules to see the relative positions of the atoms.


The hydrogen atoms are generally found to be close in space to the oxygens. The distances between the hydrogen and the oxygen are more uniform in ice than in water. The hydrogens preferentially orient in this manner because they have a partial positive change and the oxygen has a partial negative charge.

4. Compare liquid water and ice, which seems to have the lower density, and why?


view the physical property of ice that causes this phenomena.


Ice has the lower density (that is why it floats) due to hexagonal channels in the ice. The hexagonal channels form to optimize hydrogen bonding in solid water.

a. H-bonds are stable because of electron sharing across the bond (i.e. a weak covalent bond) and an electrostatic attraction between:

  • Electropositive hydrogen, attached to an electronegative atom is the hydrogen bond donor (i.e. NH)
  • Electronegative hydrogen bond acceptor (e.g. the lone pairs of oxygen in the case of water, or C=O group of an amide)

b. Typical length: 1.8 Å (from hydrogen to oxygen, 2.7 Å from hydrogen to nitrogen)

c. Typical angle: (180^{circ}) (pm) (20^{circ})

d. Typical energy: 20 kJ/mole.

e. Number of hydrogen bonds depend on temperature, 4/molecule at 0(^{circ}C).

Biochemical Significance of Hydrogen Bonds:

  1. In ice, the hydrogen bonds cause the formation of cavities in the ice, lowering the density of the solid.
  2. In liquid water, the hydrogen bonds persist, and are transiently formed on a time scale of ~nano seconds, generating small short-lived clusters of "ice" in liquid water.
  3. Hydrogen bonds are present over a wide temperature range.
  4. The hydrogen bonds in water allow water to absorb heat by breaking the hydrogen bonds without a large increase in temperature, giving water a high heat capacity.

The 4 possible hydrogen bonds formed with a water molecule in ice. The number of hydrogen bonds formed/molecule in liquid water is less than four, and decreases as the temperature increases. At room temperature each water molecule forms on average approximately 3 hydrogen bonds.

learn by doing

1. Which of the following sets of atoms will not form a strong hydrogen bond?

a. C-H.....O-C

b. O-H....N

c. N-H....O

d. O-H....O


The donor hydrogen should be attached to a strongly electronegative atom.


a. (the donor hydrogen is not attached to an electronegative atom.)

2. Which of the following hydrogen bonds would be the strongest (lowest in energy)?

a. An H to O distance of 1.0 A

b. An H to O distance of 2.0 A

c. An H to O distance of 2.5 A


The optimal distance between donor and acceptor depend on the balance between van der Waals forces and electrostatic forces.


b. (this is very close to the optimal distance for the electrostatic interaction in hydrogen bonds.)

3. Which of the following hydrogen bonds would be the weakest?

a. An N-H...O angel of 150 degree

b. An N-H...O angel of 170 degree

c. An N-H...O angel of 180 degree


Hydrogen bonds are partially covalent.


a. ( hydrogen bonds are also partially covalent and optimal overlap of the shared electrons occurs when they are linear.)

Solvation and Solubility

Hydrophobic compounds do not contain polar atoms and therefore cannot interact with water via hydrogen bonding. Consequently, solvated hydrophobic compounds cause the formation of an ordered shell of hydrogen bonded water molecules. Removal of the solvated hydrophobic compound will release these water molecules, increasing the entropy of the solvent. This favorable increase in the entropy of the solvent drives the hydrophobic molecules from the aqueous phase. The hydrophobic effect is responsible for the spontaneous formation of a number of important biological structures, such as:

  • Proteins
  • cell membranes
  • Interaction of small molecules with larger proteins, such as substrates with enzymes.

Hydrophillic compounds contain polar atoms, such as nitrogen or oxygen. Consequently, they can form hydrogen bonds with water. The formation of hydrogen bonds is energetically favorable, thus hydrophillic compounds readily dissolve in water.

Ionic compounds are readily solvated by water. There are two factors that favor a dissolved solution of ions over the crystalline form.

  1. Increase in entropy of the ions. A crystal is highly ordered, with low entropy. Dissolved ions are dispersed throughout the solution, a high entropy state.
  2. Electrostatic shielding. The force between two charged particles is: The force depends on the distance between the two charges and the dielectric constant (D) of the media. A high dielectric constant, such as that found in water, is important because the forces between charges are attenuated or reduced. Making it less favorable to have an electrostatic interaction between the positive and negative ions.

[F=frac{1}{4 pi varepsilon_{o}} frac{q_{1} q_{2}}{D r^{2}} quad varepsilon_{o}=8.854 imes 10^{-12} C^{2} / N m^{2} onumber]

The dielectric constant is proportional to the dipole moment of the solvent, as the dipole moment increases, D, increases, as shown in the following table. A large dipole moment means that the solvent molecules can interact favorably with charged solute molecules.

CompoundDielectric ConstantDipole Moment

Amphipathic (or amphiphilic) compounds are both polar (or charged) and nonpolar. An example is a fatty acid, which has a charged carboxylate (red) and a non-polar hydrocarbon chain (yellow). These can form micelles if the nonpolar part is sufficiently large. Micelles are aggregates of amphipathic molecules that sequester the nonpolar part on the inside, much like the inside of an orange. Micelles will form spontaneous driven by the hydrophobic effect.

Review Quiz


1. The partial negative charge at one end of a water molecule is attracted to the partial positive charge of another water molecule. This attraction is called:

a. a covalent bond.

b. a hydration shell.

c. a hydrogen bond.

d. a hydrophobic bond.

e. an ionic bond.


c. (The statement is a fairly complete definition of the hydrogen bond; Hydration shell refers to organization of water molecules around a non-polar group; "Hydrophobic Bond" is not a real term, don't confuse it with a "hydrogen bond"; Partial charges cannot be involved in ionic bonds, only full charges can form ionic bonds.)

2. A hydrogen bond that is linear is more stable than one that is bent.




True (A linear geometry is more stable due to the partial covalent nature of the hydrogen bond.)

3. On average, there are ____ hydrogen bonds per molecule in water at room temperature and _____ hydrogen bonds per molecule in ice.


Look carefully at the Jmol images and static picture of hydrogen bonds in ice and water.


3; 4

4. Which of the following alcohols would be most soluble in water?

a. methanol ((CH3OH))

b. ethanol ((CH_3CH_2CH_2OH))

c. butanol ((CH_3CH_2CH_2CH_2OH))

d. octanol ((CH_3(CH_2)_6CH_2OH))

e. phenol (benzyl-OH)


All of these compound have the same polar group (-OH). In what way do they differ?


a. (The alcohol with the least amount of non-polar character would be the most soluble.)

5. Hydrophobic interactions

a. are responsible for the surface tension of water

b. are stronger than covalent bonds

c. can hold two ions together

d. are the major force forming lipid (hydrocarbon) bilayers


Hydrophobic compounds do not contain polar atoms and therefore cannot interact with water via hydrogen bonding.



Module 2.2: Water and Hydrogen Bonds - Biology

1. Water molecules are polar and hydrogen bonds form between them.

2. Hydrogen bonding and dipolarity explain the cohesive, adhesive, thermal and solvent properties of water.

  • hydrogen bonds between polar water molecules cause them to cohere
  • allowing for transpiration in plants moving water against gravity
  • surface tension between cohering water molecules
  • allowing for animals such as water striders to walk over the surface of ponds even though they are denser than water
  • hydrogen bonds between polar water molecules and any other charged or ionic substance cause them to cohere
  • allowing for transport in an aqueous environment
  • hydrogen bonds between polar water molecules cause water to resist change
  • high specific heat (energy required to change water temperature)
  • high heat of vaporization (energy required to boil water)
  • high heat of fusion (loss of energy required to freeze water)
  • thus, water produces a stable environment for aquatic organisms
  • the polarity of water attracts, or dissolves, any other polar or charged particles by forming hydrogen bonds with them
  • proteins, glucose, or ions, such as sodium or calcium are all soluble
  • cytoplasm is primarily water, providing a polar medium in which other polar or charged molecules dissolve
  • many enzymes are globular proteins that are water soluble so they dissolve in cytoplasm where they control metabolic reactions

3. Substances can be hydrophilic or hydrophobic.

Applications and skills:

Application: Comparison of the thermal properties of water with those of methane.

Application: Use of water as a coolant in sweat.

  • hydrogen bonds between polar water molecules cause water to resist change
  • high heat of vaporization (energy required to change liquid water to vapor) because hydrogen bonds must be broken
  • thus, evaporation of water from plant leaves (transpiration) or from human skin (sweat) removes heat, acting as a coolant

Application: Modes of transport of glucose, amino acids, cholesterol, fats, oxygen and sodium chloride in blood in relation to their solubility in water.

  • the polarity of water attracts, or dissolves, any other polar or charged particles by forming hydrogen bonds with them
  • proteins, glucose, or ions, such as sodium or calcium are all soluble

Students should know at least one example of a benefit to living organisms of each property of water.

Transparency of water and maximum density at 4°C do not need to be included.

Comparison of the thermal properties of water and methane assists in the understanding of the significance of hydrogen bonding in water.

2.2 Water

In this section, you will investigate the following questions:

  • How does the molecular structure of water result in unique properties of water that are critical to maintaining life?
  • What are the role of acids, bases, and buffers in dynamic homeostasis?

Connection for AP ® Courses

Covalent bonds form between atoms when they share electrons to fill their valence electron shells. When the sharing of electrons between atoms is equal, such as O2 (oxygen) or CH4 (methane), the covalent bond is said to be nonpolar. However, when electrons are shared, but not equally due to differences in electronegativity (the tendency to attract electrons), the covalent bond is said to be polar. H2O (water) is an example of a polar molecule. Because oxygen is more electronegative than hydrogen, the electrons are drawn toward oxygen and away from the hydrogen atoms consequently, the oxygen atom acquires a slight negative charge and each hydrogen atoms acquires a slightly positive charge. It is important to remember that the electrons are still shared, just not equally.

Water’s polarity allows for the formation of hydrogen bonds between adjacent water molecules, resulting in many unique properties that are critical to maintaining life. For example, water is an excellent solvent because hydrogen bonds allow ions and other polar molecules to dissolve in water. Water’s hydrogen bonds also contribute to its high heat capacity and high heat of vaporization, resulting in greater temperature stability. Hydrogen bond formation makes ice less dense as a solid than as a liquid, insulating aquatic environments. Water’s cohesive and adhesive properties are seen as it rises inside capillary tubes or travels up a large tree from roots to leaves. The pH or hydrogen ion concentration of a solution is highly regulated to help organisms maintain homeostasis for example, as will be explored in later chapters, the enzymes that catalyze most chemical reactions in cells are pH specific. Thus, the properties of water are connected to the biochemical and physical processes performed by living organisms. Life on Earth would be very different if these properties were altered—if life could exist at all.

The information presented and the examples highlighted in this section support concepts and Learning Objectives outlined in Big Idea 2 of the AP ® Biology Curriculum. The Learning Objectives listed in the Curriculum Framework provide a transparent foundation for the AP ® Biology course, an inquiry-based laboratory experience, instructional activities, and AP ® Exam questions. A Learning Objective merges required content with one or more of the seven Science Practices.

Big Idea 2 Biological systems utilize free energy and molecular building blocks to grow, to reproduce, and to maintain dynamic homeostasis.
Enduring Understanding 2.A Growth, reproduction and maintenance of living systems require free energy and matter.
Essential Knowledge 2.A.3 Organisms must exchange matter with the environment to grow, reproduce and maintain organization.
Science Practice 4.1 The student can justify the selection of the kind of data needed to answer a particular scientific question.
Learning Objective 2.8 The student is able to justify the selection of data regarding the types of molecules that an animal, plant, or bacterium will take up as necessary building blocks and excrete as waste products.

Teacher Support

Discuss with students why scientists use the criteria of the presence of liquid water to determine if an environment or planet can support life. More information on this topic is available at this site.

Have students create visual representations with annotations to explain how water’s molecular structure and the resulting polarity results in its unique properties. Have the students describe how these properties are vital to life processes.

The Science Practice Challenge Questions contain additional test questions for this section that will help you prepare for the AP exam. These questions address the following standards:
[APLO 2.8] [APLO 2.23]

Why do scientists spend time looking for water on other planets? Why is water so important? It is because water is essential to life as we know it. Water is one of the more abundant molecules and the one most critical to life on Earth. Approximately 60–70 percent of the human body is made up of water. Without it, life as we know it simply would not exist.

The polarity of the water molecule and its resulting hydrogen bonding make water a unique substance with special properties that are intimately tied to the processes of life. Life originally evolved in a watery environment, and most of an organism’s cellular chemistry and metabolism occur inside the watery contents of the cell’s cytoplasm. Special properties of water are its high heat capacity and heat of vaporization, its ability to dissolve polar molecules, its cohesive and adhesive properties, and its dissociation into ions that leads to the generation of pH. Understanding these characteristics of water helps to elucidate its importance in maintaining life.

Water’s Polarity

One of water’s important properties is that it is composed of polar molecules: the hydrogen and oxygen within water molecules (H2O) form polar covalent bonds. While there is no net charge to a water molecule, the polarity of water creates a slightly positive charge on hydrogen and a slightly negative charge on oxygen, contributing to water’s properties of attraction. Water’s charges are generated because oxygen is more electronegative than hydrogen, making it more likely that a shared electron would be found near the oxygen nucleus than the hydrogen nucleus, thus generating the partial negative charge near the oxygen.

As a result of water’s polarity, each water molecule attracts other water molecules because of the opposite charges between water molecules, forming hydrogen bonds. Water also attracts or is attracted to other polar molecules and ions. A polar substance that interacts readily with or dissolves in water is referred to as hydrophilic (hydro- = “water” -philic = “loving”). In contrast, non-polar molecules such as oils and fats do not interact well with water, as shown in Figure 2.14 and separate from it rather than dissolve in it, as we see in salad dressings containing oil and vinegar (an acidic water solution). These nonpolar compounds are called hydrophobic (hydro- = “water” -phobic = “fearing”).

Water’s States: Gas, Liquid, and Solid

The formation of hydrogen bonds is an important quality of the liquid water that is crucial to life as we know it. As water molecules make hydrogen bonds with each other, water takes on some unique chemical characteristics compared to other liquids and, since living things have a high water content, understanding these chemical features is key to understanding life. In liquid water, hydrogen bonds are constantly formed and broken as the water molecules slide past each other. The breaking of these bonds is caused by the motion (kinetic energy) of the water molecules due to the heat contained in the system. When the heat is raised as water is boiled, the higher kinetic energy of the water molecules causes the hydrogen bonds to break completely and allows water molecules to escape into the air as gas (steam or water vapor). On the other hand, when the temperature of water is reduced and water freezes, the water molecules form a crystalline structure maintained by hydrogen bonding (there is not enough energy to break the hydrogen bonds) that makes ice less dense than liquid water, a phenomenon not seen in the solidification of other liquids.

Water’s lower density in its solid form is due to the way hydrogen bonds are oriented as it freezes: the water molecules are pushed farther apart compared to liquid water. With most other liquids, solidification when the temperature drops includes the lowering of kinetic energy between molecules, allowing them to pack even more tightly than in liquid form and giving the solid a greater density than the liquid.

The lower density of ice, illustrated and pictured in Figure 2.15, an anomaly, causes it to float at the surface of liquid water, such as in an iceberg or in the ice cubes in a glass of ice water. In lakes and ponds, ice will form on the surface of the water creating an insulating barrier that protects the animals and plant life in the pond from freezing. Without this layer of insulating ice, plants and animals living in the pond would freeze in the solid block of ice and could not survive. The detrimental effect of freezing on living organisms is caused by the expansion of ice relative to liquid water. The ice crystals that form upon freezing rupture the delicate membranes essential for the function of living cells, irreversibly damaging them. Cells can only survive freezing if the water in them is temporarily replaced by another liquid like glycerol.

Link to Learning

Click here to see a 3-D animation of the structure of an ice lattice.

  1. Red and white balls represent oxygen and hydrogen, respectively, loose arrangement of molecules results in low density of ice
  2. Red and white balls represent oxygen and hydrogen respectively, tightly packed arrangement of molecules results in a low density of ice
  3. Red and white balls represent hydrogen and oxygen, respectively, loose arrangement of molecules results in low density of ice
  4. Red and white balls represent oxygen and hydrogen, respectively, tightly packed arrangement of molecules results in high density of ice

Water’s High Heat Capacity

Water’s high heat capacity is a property caused by hydrogen bonding among water molecules. Water has the highest specific heat capacity of any liquids. Specific heat is defined as the amount of heat one gram of a substance must absorb or lose to change its temperature by one degree Celsius. For water, this amount is one calorie . It therefore takes water a long time to heat and long time to cool. In fact, the specific heat capacity of water is about five times more than that of sand. This explains why the land cools faster than the sea. Due to its high heat capacity, water is used by warm blooded animals to more evenly disperse heat in their bodies: it acts in a similar manner to a car’s cooling system, transporting heat from warm places to cool places, causing the body to maintain a more even temperature.

Water’s Heat of Vaporization

Water also has a high heat of vaporization , the amount of energy required to change one gram of a liquid substance to a gas. A considerable amount of heat energy (586 cal) is required to accomplish this change in water. This process occurs on the surface of water. As liquid water heats up, hydrogen bonding makes it difficult to separate the liquid water molecules from each other, which is required for it to enter its gaseous phase (steam). As a result, water acts as a heat sink or heat reservoir and requires much more heat to boil than does a liquid such as ethanol, whose hydrogen bonding with other ethanol molecules is weaker than water’s hydrogen bonding. Eventually, as water reaches its boiling point of 100° Celsius (212° Fahrenheit), the heat is able to break the hydrogen bonds between the water molecules, and the kinetic energy (motion) between the water molecules allows them to escape from the liquid as a gas. Even when below its boiling point, water’s individual molecules acquire enough energy from other water molecules such that some surface water molecules can escape and vaporize: this process is known as evaporation .

The fact that hydrogen bonds need to be broken for water to evaporate means that a substantial amount of energy is used in the process. As the water evaporates, energy is taken up by the process, cooling the environment where the evaporation is taking place. In many living organisms, including in humans, the evaporation of sweat, which is 90 percent water, allows the organism to cool so that homeostasis of body temperature can be maintained.

Water’s Solvent Properties

Since water is a polar molecule with slightly positive and slightly negative charges, ions and polar molecules can readily dissolve in it. Therefore, water is referred to as a solvent , a substance capable of dissolving other polar molecules and ionic compounds. The charges associated with these molecules will form hydrogen bonds with water, surrounding the particle with water molecules. This is referred to as a sphere of hydration , or a hydration shell, as illustrated in Figure 2.16 and serves to keep the particles separated or dispersed in the water.

When ionic compounds are added to water, the individual ions react with the polar regions of the water molecules and their ionic bonds are disrupted in the process of dissociation . Dissociation occurs when atoms or groups of atoms break off from molecules and form ions. Consider table salt (NaCl, or sodium chloride): when NaCl crystals are added to water, the molecules of NaCl dissociate into Na + and Cl – ions, and spheres of hydration form around the ions, illustrated in Figure 2.16. The positively charged sodium ion is surrounded by the partially negative charge of the water molecule’s oxygen. The negatively charged chloride ion is surrounded by the partially positive charge of the hydrogen on the water molecule.

Water’s Cohesive and Adhesive Properties

Have you ever filled a glass of water to the very top and then slowly added a few more drops? Before it overflows, the water forms a dome-like shape above the rim of the glass. This water can stay above the glass because of the property of cohesion . In cohesion, water molecules are attracted to each other (because of hydrogen bonding), keeping the molecules together at the liquid-gas (water-air) interface, although there is no more room in the glass.

Cohesion allows for the development of surface tension , the capacity of a substance to withstand being ruptured when placed under tension or stress. This is also why water forms droplets when placed on a dry surface rather than being flattened out by gravity. When a small scrap of paper is placed onto the droplet of water, the paper floats on top of the water droplet even though paper is denser (heavier) than the water. Cohesion and surface tension keep the hydrogen bonds of water molecules intact and support the item floating on the top. It’s even possible to “float” a needle on top of a glass of water if it is placed gently without breaking the surface tension, as shown in Figure 2.17.

These cohesive forces are related to water’s property of adhesion , or the attraction between water molecules and other molecules. This attraction is sometimes stronger than water’s cohesive forces, especially when the water is exposed to charged surfaces such as those found on the inside of thin glass tubes known as capillary tubes. Adhesion is observed when water “climbs” up the tube placed in a glass of water: notice that the water appears to be higher on the sides of the tube than in the middle. This is because the water molecules are attracted to the charged glass walls of the capillary more than they are to each other and therefore adhere to it. This type of adhesion is called capillary action , and is illustrated in Figure 2.18.

Why are cohesive and adhesive forces important for life? Cohesive and adhesive forces are important for the transport of water from the roots to the leaves in plants. These forces create a “pull” on the water column. This pull results from the tendency of water molecules being evaporated on the surface of the plant to stay connected to water molecules below them, and so they are pulled along. Plants use this natural phenomenon to help transport water from their roots to their leaves. Without these properties of water, plants would be unable to receive the water and the dissolved minerals they require. In another example, insects such as the water strider, shown in Figure 2.19, use the surface tension of water to stay afloat on the surface layer of water and even mate there.

Science Practice Connection for AP® Courses

During a process called transpiration, water evaporates through a plant’s leaves. Water in the ground travels up from the roots to the leaves. Based on water’s molecular properties, create a visual representation (e.g., diagrams or models) with annotations to explain how water travels up a 300-ft. California redwood tree. What other unique properties of water are attributed to its molecular structure, and how are these properties important to life?

Teacher Support

This activity is an application of Learning Objectives 2.8 and Science Practice 4.1 and Learning Objectives 2.9 and Science Practices 1.1 and 1.4 because you are modeling the relationship between water’s molecular structure and its unique properties that are essential to maintaining life, including capillary action.

PH, Buffers, Acids, and Bases

The pH of a solution indicates its acidity or basicity.

litmus or pH paper, filter paper that has been treated with a natural water-soluble dye so it can be used as a pH indicator, to test how much acid (acidity) or base (basicity) exists in a solution. You might have even used some to test whether the water in a swimming pool is properly treated. In both cases, the pH test measures the concentration of hydrogen ions in a given solution.

Hydrogen ions are spontaneously generated in pure water by the dissociation (ionization) of a small percentage of water molecules into equal numbers of hydrogen (H + ) ions and hydroxide (OH - ) ions. While the hydroxide ions are kept in solution by their hydrogen bonding with other water molecules, the hydrogen ions, consisting of naked protons, are immediately attracted to un-ionized water molecules, forming hydronium ions (H30 + ). Still, by convention, scientists refer to hydrogen ions and their concentration as if they were free in this state in liquid water.

The concentration of hydrogen ions dissociating from pure water is 1 × 10 -7 moles H + ions per liter of water. Moles (mol) are a way to express the amount of a substance (which can be atoms, molecules, ions, etc), with one mole being equal to 6.02 × 10 23 particles of the substance. Therefore, 1 mole of water is equal to 6.02 × 10 23 water molecules. The pH is calculated as the negative of the base 10 logarithm of this concentration. The log10 of 1 × 10 -7 is -7.0, and the negative of this number (indicated by the “p” of “pH”) yields a pH of 7.0, which is also known as neutral pH. The pH inside of human cells and blood are examples of two areas of the body where near-neutral pH is maintained.

Non-neutral pH readings result from dissolving acids or bases in water. Using the negative logarithm to generate positive integers, high concentrations of hydrogen ions yield a low pH number, whereas low levels of hydrogen ions result in a high pH. An acid is a substance that increases the concentration of hydrogen ions (H + ) in a solution, usually by having one of its hydrogen atoms dissociate. A base provides either hydroxide ions (OH – ) or other negatively charged ions that combine with hydrogen ions, reducing their concentration in the solution and thereby raising the pH. In cases where the base releases hydroxide ions, these ions bind to free hydrogen ions, generating new water molecules.

The stronger the acid, the more readily it donates H + . For example, hydrochloric acid (HCl) completely dissociates into hydrogen and chloride ions and is highly acidic, whereas the acids in tomato juice or vinegar do not completely dissociate and are considered weak acids. Conversely, strong bases are those substances that readily donate OH – or take up hydrogen ions. Sodium hydroxide (NaOH) and many household cleaners are highly alkaline and give up OH – rapidly when placed in water, thereby raising the pH. An example of a weak basic solution is seawater, which has a pH near 8.0, close enough to neutral pH that marine organisms adapted to this saline environment are able to thrive in it.

The pH scale is, as previously mentioned, an inverse logarithm and ranges from 0 to 14 (Figure 2.20). Anything below 7.0 (ranging from 0.0 to 6.9) is acidic, and anything above 7.0 (from 7.1 to 14.0) is alkaline. Extremes in pH in either direction from 7.0 are usually considered inhospitable to life. The pH inside cells (6.8) and the pH in the blood (7.4) are both very close to neutral. However, the environment in the stomach is highly acidic, with a pH of 1 to 2. So how do the cells of the stomach survive in such an acidic environment? How do they homeostatically maintain the near neutral pH inside them? The answer is that they cannot do it and are constantly dying. New stomach cells are constantly produced to replace dead ones, which are digested by the stomach acids. It is estimated that the lining of the human stomach is completely replaced every seven to ten days.

Link to Learning

Watch this video for a straightforward explanation of pH and its logarithmic scale.

  1. Diabetic ketoacidosis decreases the normal pH (8.35-8.45) to a lower value.
  2. Diabetic ketoacidosis increases normal pH level of blood disrupting biological processes.
  3. Diabetic ketoacidosis keeps pH level of blood constant which disrupts biological processes.
  4. Diabetic ketoacidosis decreases normal pH (7.35-7.45) to a lower value.

So how can organisms whose bodies require a near-neutral pH ingest acidic and basic substances (a human drinking orange juice, for example) and survive? Buffers are the key. Buffers readily absorb excess H + or OH – , keeping the pH of the body carefully maintained in the narrow range required for survival. Maintaining a constant blood pH is critical to a person’s well-being. The buffer maintaining the pH of human blood involves carbonic acid (H2CO3), bicarbonate ion (HCO3 – ), and carbon dioxide (CO2). When bicarbonate ions combine with free hydrogen ions and become carbonic acid, hydrogen ions are removed, moderating pH changes. Similarly, as shown in Figure 2.21, excess carbonic acid can be converted to carbon dioxide gas and exhaled through the lungs. This prevents too many free hydrogen ions from building up in the blood and dangerously reducing the blood’s pH. Likewise, if too much OH – is introduced into the system, carbonic acid will combine with it to create bicarbonate, lowering the pH. Without this buffer system, the body’s pH would fluctuate enough to put survival in jeopardy.

Other examples of buffers are antacids used to combat excess stomach acid. Many of these over-the-counter medications work in the same way as blood buffers, usually with at least one ion capable of absorbing hydrogen and moderating pH, bringing relief to those that suffer “heartburn” after eating. The unique properties of water that contribute to this capacity to balance pH—as well as water’s other characteristics—are essential to sustaining life on Earth.

Water Stabilizes Temperature

The hydrogen bonds in water allow it to absorb and release heat energy more slowly than many other substances. Temperature is a measure of the motion (kinetic energy) of molecules. As the motion increases, energy is higher and thus temperature is higher. Water absorbs a great deal of energy before its temperature rises. Increased energy disrupts the hydrogen bonds between water molecules. Because these bonds can be created and disrupted rapidly, water absorbs an increase in energy and temperature changes only minimally. This means that water moderates temperature changes within organisms and in their environments. As energy input continues, the balance between hydrogen-bond formation and destruction swings toward the destruction side. More bonds are broken than are formed. This process results in the release of individual water molecules at the surface of the liquid (such as a body of water, the leaves of a plant, or the skin of an organism) in a process called evaporation. Evaporation of sweat, which is 90 percent water, allows for cooling of an organism, because breaking hydrogen bonds requires an input of energy and takes heat away from the body.

Conversely, as molecular motion decreases and temperatures drop, less energy is present to break the hydrogen bonds between water molecules. These bonds remain intact and begin to form a rigid, lattice-like structure (e.g., ice) (Figure 2.8 a). When frozen, ice is less dense than liquid water (the molecules are farther apart). This means that ice floats on the surface of a body of water (Figure 2.8 b). In lakes, ponds, and oceans, ice will form on the surface of the water, creating an insulating barrier to protect the animal and plant life beneath from freezing in the water. If this did not happen, plants and animals living in water would freeze in a block of ice and could not move freely, making life in cold temperatures difficult or impossible.

Figure 2.8 (a) The lattice structure of ice makes it less dense than the freely flowing molecules of liquid water. Ice’s lower density enables it to (b) float on water. (credit a: modification of work by Jane Whitney credit b: modification of work by Carlos Ponte)

The Folding of Proteins and Nucleic Acids Whether Hydrogen Bonding Contributes to Protein Stability is Controversial

That the hydrogen bond may stabilize protein conformation was first recognized when Pauling and Corey proposed the model of hydrogen-bonded structures of α helix and β sheets in 1951. 5 As shown in the vast number of protein structures deposited in the protein data bank, the hydrogen bond is ubiquitous in secondary and tertiary structures of proteins. Protein contains many functional groups capable of forming hydrogen bonds : backbone peptide amide and polar side chain groups (e.g., the hydroxyl group in serine, threonine, and tyrosine the amide group in asparagine and glutamine). Backbone peptide hydrogen bonds contribute to ∼70% of all intramolecular hydrogen bonds formed in proteins. 64

Despite the ubiquitous nature of the hydrogen bond in proteins, whether it stabilizes proteins is controversial. There is no question that the formation of the hydrogen bond itself is favorable – the energy of an amide-amide hydrogen bond in vacuum, estimated by quantum chemistry, is −25 kJ mol −1 . 65,66 The main concern is that in order to form protein-protein hydrogen bonds, the polar groups have to break their hydrogen bonds with water. It was believed that the protein-water hydrogen bond more or less cancels out the stabilization by the protein-protein hydrogen bond so that the contribution of the hydrogen bond to protein stability was expected to be small. 3,6 In his influential review, Dill elaborated the thermodynamics cycle of hydrogen bond formation as in Figure 6 . For example, formation of the N-methylacetamide (NMA) dimer is a popular model for peptide hydrogen bond. It was observed that the formation of the NMA dimer is disfavored in water (ΔG1=+13 kJ mol −1 ) but favored in nonaqueous solvent CCl4G5=−10 kJ mol −1 ). 67 On the other hand, because of the dehydration penalty, transfer of NMA from water to CCl4 is highly unfavorable, with ΔG4=+26 kJ mol −1 . 68 It follows that transfer of hydrogen-bonded groups from water to CCl4 is also unfavorable: ΔG2=+3 kJ mol −1 . 68 According to this analysis, the hydrogen bond is destabilizing, no matter whether it is exposed (ΔG1=+13 kJ mol −1 ) or buried (ΔG3=+16 kJ mol −1 ).

Figure 6 . Thermodynamics cycle for the formation of a hydrogen bond.


The aim of this study was to gain a better understanding of the contribution of hydrogen bonds by tyrosine -OH groups to protein stability. The amino acid sequences of RNases Sa and Sa3 are 69% identical and each contains eight Tyr residues with seven at equivalent structural positions. We have measured the stability of the 16 tyrosine to phenylalanine mutants. For two equivalent mutants, the stability increases by 0.3 kcal/mol (RNase Sa Y30F) and 0.5 kcal/mol (RNase Sa3 Y33F) (1 kcal = 4.184 kJ). For all of the other mutants, the stability decreases with the greatest decrease being 3.6 kcal/mol for RNase Sa Y52F. Seven of the 16 tyrosine residues form intramolecular hydrogen bonds and the average decrease in stability for these is 2.0(±1.0) kcal/mol. For the nine tyrosine residues that do not form intramolecular hydrogen bonds, the average decrease in stability is 0.4(±0.6) kcal/mol. Thus, most tyrosine -OH groups contribute favorably to protein stability even if they do not form intramolecular hydrogen bonds. Generally, the stability changes for equivalent positions in the two proteins are remarkably similar. Crystal structures were determined for two of the tyrosine to phenylalanine mutants of RNase Sa: Y80F (1.2 Å), and Y86F (1.7 Å). The structures are very similar to that of wild-type RNase Sa, and the hydrogen bonding partners of the tyrosine residues always form intermolecular hydrogen bonds to water in the mutants. These results provide further evidence that the hydrogen bonding and van der Waals interactions of polar groups in the tightly packed interior of folded proteins are more favorable than similar interactions with water in the unfolded protein, and that polar group burial makes a substantial contribution to protein stability.

Module 2.2: Water and Hydrogen Bonds - Biology

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Hydrogen bonding is a special case of permanent dipole-permanent dipole bonding. It is important to be clear that although it is called "hydrogen-bonding" it really is an intermolecular force. It is alsovital that you refer to the hydrogen bonding as being between molecules and not within them.

It exists where one of the most electronegative elements (fluorine, oxygen or nitrogen) is bonded to hydrogen.

Hydrogen bonding causes stronger intermolecular forces than would otherwise be predicted.

This increases the boiling point of substances such as water.

To investigate the power of hydrogen bonding, look at the boiling points of the group VI hydrides.

There is a steady increase down the group from H2S downwards but water (H2O) has a much higher boiling point despite being higher in the group.

It also explains why water is one of the few substances for which the solid is less dense than the liquid. On freezing, the hydrogen bonding organises the molecules into a fairly open structure.

Dynamics of Hydrogen Bond Desolvation in Protein Folding

As proteins fold, a progressive structuring, immobilization and eventual exclusion of water surrounding backbone hydrogen bonds takes place. This process turns hydrogen bonds into major determinants of the folding pathway and compensates for the penalty of desolvation of the backbone polar groups. Taken as an average over all hydrogen bonds in a native fold, this extent of protection is found to be nearly ubiquitous. It is dynamically crucial, determining a constraint in the long-time limit behavior of coarse-grained ab initio simulations. Furthermore, an examination of one of the longest available (1 μs) all-atom simulations with explicit solvent reveals that this average extent of protection is a constant of motion for the folding trajectory. We propose how such a stabilization is best achieved by clustering five hydrophobes around the backbone hydrogen bonds, an arrangement that yields the optimal stabilization. Our results support and clarify the view that hydrophobic surface burial should be commensurate with hydrogen-bond formation and enable us to define a basic desolvation motif inherent to structure and folding dynamics.

Hydrogen Bonds

Ionic and covalent bonds between elements require energy to break. Iconic bonds are not as strong as covalent, which determines their behavior in biological systems. However, not all bonds are ionic or covalent bonds. Weaker bonds can also form between molecules. Two weak bonds that occur frequently are hydrogen bonds and van der Waals interactions. Without these two types of bonds, life as we know it would not exist. Hydrogen bonds provide many of the critical, life-sustaining properties of water and also stabilize the structures of proteins and DNA, the building block of cells.

When polar covalent bonds containing hydrogen form, the hydrogen in that bond has a slightly positive charge because hydrogen’s electron is pulled more strongly toward the other element and away from the hydrogen. Because the hydrogen is slightly positive, it will be attracted to neighboring negative charges. When this happens, a weak interaction occurs between the δ+ of the hydrogen from one molecule and the δ– charge on the more electronegative atoms of another molecule, usually oxygen or nitrogen, or within the same molecule. This interaction is called a hydrogen bond. This type of bond is common and occurs regularly between water molecules. Individual hydrogen bonds are weak and easily broken however, they occur in very large numbers in water and in organic polymers, creating a major force in combination. Hydrogen bonds are also responsible for zipping together the DNA double helix.


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