Term Paper on Water: Top 5 Papers | Biological Solvents | Biochemistry

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Term Paper on Water

Term Paper Contents:

1. Term Paper on the Introduction to Water
2. Term Paper on Water and the Hydrogen Bond
3. Term Paper on Water and Temperature
4. Term Paper on Water as a Solvent
5. Term Paper on the Ionization of Water

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Term Paper # 1. Introduction to Water:

Life on this planet began in water, and today, wherever liquid water is found, life is also present. There are one-celled organisms that eke out their entire existence in no more water than can cling to a grain of sand. Some kinds of algae are found only on the melting undersurfaces of polar ice floes.

Certain bacteria can tolerate the near-boiling water of hot springs. In the desert, plants race through an entire life cycle-seed to flower to seed—following a single rainfall. In the jungle, the water cupped in the leaves of a tropical plant forms a microcosm in which a myriad of small organisms are born, spawn, and die.

Observations of the polar ice caps on Mars and of its “canals,” first seen more than 100 years ago, kindled hope that life might be present there also. However, no liquid water was detected in the soil or atmosphere at the site of the first spacecraft landings.

Indeed, it was discovered that the air surrounding Mars is so dry and cold that liquid water could not exist on the planet. Nor was life discovered. On earth, by contrast, water is a common liquid. Three- quarters of the earth’s surface is covered by water. In fact, if the earth’s surface were absolutely smooth, all of it would be 2.5 kilometers under water.

Water makes up 50 to 95 percent of the weight of any functioning living system. If you weigh 60 kilograms, about 50 kilograms of that is water. But “common” is not the same as “ordinary.” Water is not in the least an ordinary liquid. Compared with other liquids it is, in fact, quite extraordinary. If it were not, it is highly probable that life could never have evolved.

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Term Paper # 2. Water and the Hydrogen Bond:

In order to understand why water is so extraordinary and how, as a consequence, it can play its unique and crucial role in relation to living systems we have to look again at its molecular structure. Each water molecule is made up of two atoms of hydrogen and one of oxygen held together by two covalent bonds.

The water molecule as a whole is neutral in charge, having an equal number of electrons and protons. However, the molecule is electrically asymmetric-that is, polar. This polarity occurs because oxygen is more electronegative than hydrogen. The paired electrons in the outer orbitals spend more time around the oxygen nucleus than they do around the hydrogen nuclei.

As a consequence, the region near the oxygen nucleus has two weakly negative zones, and each of the regions near the hydrogen nuclei has a weakly positive zone. Thus, the water molecule, in terms of its polarity, is four-cornered, with two positively charged “corners” and two negatively charged ones.

When any of these charged regions comes close to the oppositely charged region of another water molecule, the force of the attraction forms a bond between them, which is known as a hydrogen bond. Hydrogen bonds are not found only in water. A hydrogen bond can form between any hydrogen atom that is covalently bonded to an electronegative atom-usually oxygen or nitrogen-and the oxygen or nitrogen atom of another molecule.

In water, hydrogen bonds form between the negative “corners” of one water molecule and the positive “corners” of another. As a consequence, every water molecule can establish hydrogen bonds with four other water molecules. Liquid water is made up of water molecules bound together in this way.

Any single hydrogen bond is relatively weak and has an exceedingly short lifetime; on an average, each such bond lasts about 1/100,000,000,000th of a second. But, as one is broken, another is made. All together, the hydrogen bonds have considerable strength, making water both liquid and stable under ordinary conditions of pressure and temperature.

Now let us look at some of the consequences of these attractions among water molecules, especially as they affect living systems:

1. Surface Tension:

Look at water dripping from a faucet. Each drop clings to the rim and dangles for a moment by a thread of water; then just as the tug of gravity breaks it loose, its outer surface is drawn taut, to form a sphere as the drop falls free. Gently place a needle or a razor blade flat on the surface of the water in a glass. Although the metal is denser than water, it floats.

Look at a pond in spring or summer; you will see water striders and other insects walking on its surface almost as if it were solid. These phenomena are all the result of surface tension. Surface tension is the result of the cohesion or clinging together, of the water molecules. (Cohesion is, by definition, the holding together of like substances. Adhesion is the holding together of unlike substances.)

The only liquid with a surface tension greater than that of water is mercury. Atoms of mercury are so greatly attracted to one another that they tend not to adhere to anything else.

Water, because of its negative and positive charges, also adheres strongly to any other charged molecules and to charged surfaces. The “wetting” capacity of water-that is, its ability to coat a surface-reflects its polar structure, as does its cohesiveness.

2. Capillary Action and Imbibition:

If you hold two dry glass slides together and dip one corner in water, the combination of cohesion and adhesion will cause water to spread upward between the two slides. This is capillary action. Capillary action similarly causes water to rise in very fine glass tubes, to creep up a piece of blotting paper, or to move slowly through the micro-pores of the soil and so become available to the roots of plants.

Imbibition (“drinking up”) is the capillary movement of water molecules due to adhesion into substances such as wood or gelatin, which swell as a result. The pressures developed by imbibition can be astonishingly great. It is said that stone for the ancient Egyptian pyramids was quarried by driving wooden pegs into holes drilled in the rock face and then soaking the pegs with water.

The swelling of the wood created a force great enough to break the stone slab free. Seeds imbibe water as they begin to germinate, swelling and bursting their seed coats.

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Term Paper # 3. Water and Temperature:

Specific Heat of Water:

The amount of heat a given amount of a substance requires for a given increase in temperature is its specific heat (also called heat capacity). One calorie is defined as the amount of heat that will raise the temperature of 1 gram (1 milliliter or 1 cubic centimeter) of water 1°C. The specific heat of water is about twice the specific heat of oil or alcohol; that is, approximately 0.5 calorie will raise the temperature of 1 gram of oil or alcohol 1°C.

It is four times the specific heat of air or aluminum and 10 times that of iron. Only liquid ammonia has a higher specific heat. In other words, it requires a high input of energy to raise the temperature of water. The high specific heat of water is also a consequence of hydrogen bonding.

Heat is a form of energy-the kinetic energy, or energy of movement, of molecules. In order for the kinetic energy of water molecules to increase sufficiently for the temperature to rise 1°C, it is necessary first to rupture a number of the hydrogen bonds holding the molecules together.

The reason that it takes more heat to warm a gram of water to a given temperature than to warm a gram of almost any other liquid is that so much of the heat energy added to the water is used in breaking the hydrogen bonds between the water molecules. Only a relatively small amount of heat energy is therefore available to increase molecular movement.

Note that heat and temperature are not the same. A lake may have a lower temperature than does a bird flying over it, but the lake contains more heat because it comprises many more molecules than the bird; therefore it has more molecular motion. Temperature is measured in degrees and reflects the average kinetic energy of the molecules. Heat is measured in calories and reflects both molecular movement and the mass and number of the molecules present.

What does the high specific heat of water mean in biological terms? It means that for a given rate of heat input, the temperature of water will rise more slowly than the temperature of almost any other material.

Conversely, the temperature will drop more slowly as heat is removed. Because so much heat input or heat loss is required to raise or lower the temperature of water, organisms that live in the oceans or large bodies of fresh water live in an environment where the temperature is relatively constant.

Also, the high water content of terrestrial plants and animals helps them to maintain a relatively constant internal temperature. This constancy of temperature is important because, as we shall see, biologically important chemical reactions take place only within a narrow temperature range.

Heat of Vaporization:

Vaporization or evaporation, as it is more commonly called-is the change from a liquid to a gas. Water has a high heat of vaporization. It takes more than 500 calories to change a gram of liquid water into vapor, fifty times as much as for ether and almost twice as much as for ammonia.

Hydrogen bonding is also responsible for water’s high heat of vaporization. Vaporization comes about because some of the more rapidly moving molecules of a liquid break loose from the surface and enter the air. The hot the liquid, the more rapid the movement of its molecules and, hence, the more rapid the rate of evaporation.

But, whatever the temperature, so long as a liquid is exposed to air that is less than 100 percent saturated with the vapor of that liquid, evaporation will take place, right down to the last drop. In order for a water molecule to break loose from its fellow molecules-that is, to vaporize the hydrogen bonds have to be broken.

This requires heat energy. At water’s boiling point (100°C at a pressure of 1 atmosphere), 540 calories are needed to change 1 gram of liquid water into vapor. As a consequence, when water evaporates, as from the surface of your skin or a leaf, the escaping molecules carry a great deal of heat away with them.

Thus evaporation has a cooling effect. Evaporation from the surface of a land-dwelling plant or animal is one of the principal ways in which these organisms “unload” excess heat and so stabilise their temperatures.

Freezing: The Heat of Fusion:

The freezing point of water is 0°C. So is the melting point. To make the transition from a solid to a liquid, water requires 79.7 calories per gram. As water melts in an icebox or other closed, insulated container, it draws this much heat from the contents of the container. Conversely, as water freezes, it releases the same amount of heat into its surroundings. This 79.7 cal./gram is known as the heat of fusion.

In this way, ice and snow also serve as temperature stabilisers, particularly during the temperate zone transition periods of fall and spring. Moderation of sudden changes in temperature gives organisms’ time to make seasonal adjustments essential to survival. Water exhibits another peculiarity as it undergoes the transition from a liquid to a solid.

As the temperature of a liquid drops, its density tends to increase. This increase occurs because the individual molecules are moving more slowly and so the spaces between them decrease. The density of water also increases, until the temperature nears 4°C.

Then the water molecules come so close together that every one of them becomes hydrogen-bonded to four others, forming an open latticework. In the course of this bonding process, water expands again, so that water as a solid takes up more volume than water as a liquid.

This increase in volume has occasional disastrous effects on water pipes, but, on the whole, turns out to be so beneficial for life forms that it would seem to be almost miraculously devised. If water contracted as it froze, not only would ice cubes fail to tinkle in our glasses, but also lakes and ponds and other bodies of water would freeze from the bottom up.

Once ice began to accumulate on the bottom, it would tend not to melt, season after season, and eventually the pond would freeze solid and life in the pond would be destroyed. By’ contrast, a layer of ice tends to protect the life of the pond, keeping the liquid water beneath it at 0°C or above.

The presence of dissolved substances in water reduces the temperature at which water freezes, which is why salt is thrown on icy sidewalks and used in ice-cream freezers. The “hardening” process in several species of winter-hardy plants, by which they prepare themselves for cold weather, includes the breakdown of starch (which is insoluble) into simple sugars (which are soluble).

Freshwater fish, whose body fluids are salty compared to the pond or lake in which they live, do not freeze when the temperature of the water is at or near 0°C. However, logically speaking, saltwater fish, whose body fluids are less salty than the ocean water surrounding them, should freeze at the below-zero temperatures of Arctic water.

They do not, however, and animal physiologists investigating this phenomenon have discovered that at least one species, ghost fish, produces a complex protein named, appropriately, antifreeze protein. This protein, which is secreted into the bloodstream, appears, like the antifreeze in your car radiator, to prevent water molecules from crystallizing.

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Term Paper # 4. Water as a Solvent:

Many substances within living systems are found in solution. (A solution is a uniform mixture of the molecules of two or more substances; the substance present in the greatest amount-usually a liquid-is called the solvent, and the substances present in lesser amounts are called solutes.)

The polarity of the water molecules is responsible for water’s capacity as a solvent. The polar water molecules tend to separate substances such as NaCl into their constituent ions. Then, the water molecules cluster around and segregate the charged ions. Many of the molecules important in living systems-such as glucose (sugar)—also have areas of positive and negative charge.

(Such Polar Regions of partial charge arise, as you might expect, in the neighborhood of covalently bonded atoms of unequal electronegati­vity.) These molecules therefore attract water molecules and so dissolve in water.

Molecules that readily dissolve in water are often called hydrophilic (“water-loving”). Such molecules slip into aqueous solution easily because their partially charged regions attract water molecules and so compete with the attraction between the water molecules themselves.

Molecules such as fats that lack polar, or charged, regions tend to be very insoluble in water. The hydrogen bonding between the water molecules acts as a force to exclude the nonpolar molecules. As a result of this exclusion, nonpolar molecules tend to cluster together in water, just as droplets of fats tend to coalesce, as for example, on the surface of chicken soup.

Such molecules are said to be hydrophobic (“water-fearing”), and the clustering’s are known as hydrophobic interactions. These weak forces-hydrogen bonds and hydrophobic forces are play very important roles in shaping the architecture of large biomolecules and, as a consequence, in determining their properties.

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Term Paper # 5. Ionization of Water:

In liquid water, there is a slight tendency for a hydrogen atom to jump from the oxygen atom to which it is covalently bonded to the oxygen atom to which it is hydrogen-bonded. In this reaction, two ions are produced- the hydronium ion (H₃O⁺) and the hydroxide ion (OH^–). In any given volume of pure water, a small but constant number of water molecules will be ionized in this way. The number is constant because the tendency of water to ionize is exactly offset by the tendency of the ions to reunite; thus even as some molecules are ionizing, an equal number of others are forming, a state known as equilibrium.

Although the hydrogen ion does not exist in a separate form, by convention the ionization of water is expressed by the equation:

The arrows indicate that the reaction can go in either direction. The fact that the arrow pointing toward HOH is longer indicates that, at equilibrium, most of the H₂O is not ionized. As a consequence, in any sample of pure water, only a small fraction exists in ionized form.

In pure water, the number of H⁺ ions exactly equals the number of OH^– ions. This is necessarily the case since neither ion can be formed without the other when only H₂O molecules are present. A solution acquires the properties we recognise as acidic when the number of H⁺ ions exceeds the number of OH^– ions; conversely, a solution is alkaline (basic) when the number of OH^– ions exceeds the number of H⁺ ions.

The pH Scale:

Chemists define degrees of acidity by means of the pH scale. In the expression “pH,” the p stands for a negative power of 10, and the H stands for the concentration of hydrogen ions in moles per liter.

A mole is the amount of an element equivalent to its atomic weight expressed in grams, or the amount of a substance equivalent to its molecular weight expressed in grams.

Thus, a mole of atomic hydrogen (atomic weight 1) is 1 gram of hydrogen atoms, a mole of atomic oxygen (atomic weight 16) is 16 grams of oxygen atoms, and a mole of water is 18 grams of water molecules. The most interesting thing about the mole is that a mole- of any substance-contains the same number of particles as any other mole.

This number, known as Avogadro’s number, is 6.02 × 10²³. Thus, a mole of water molecules (18 grams) contains exactly the same number of molecules as a mole of hydrochloric acid molecules (36.5 grams). Use of the mole in specifying quantities of substances involved in chemical reactions makes it possible for us to consider comparable numbers of reacting particles.

A liter of pure water contains 1/10,000,000 mole of hydrogen ions. For convenience, this is written exponentially as 10^-7 mole per liter. In terms of the pH scale, it is referred to simply as pH 7. At pH 7, the concentrations of free H⁺ and OH^– are exactly the same, and thus pure water is said to be neutral. Any pH below 7 is acidic, and any pH above 7 is basic.

The lower the pH number, the higher the concentration of hydrogen ions. Thus pH 2 means 10^-2 mole of hydrogen ions per liter of water, or 1/100 mole per liter-which is, of course, a much larger figure than 1/10,000,000. The difference of one pH unit represents a tenfold difference in the concentration of hydrogen ions.

We can now define “acid” and “base” more exactly:

1. An acid is a substance that donates H⁺ ions to a solution, that is, a hydrogen-ion donor. A solution with a pH below 7 (with more than 10^–7 mole of H⁺ ions per liter) is acidic.

2. A base is a substance that decreases the number of H⁺ ions, that is, a hydrogen-ion acceptor. A solution with a pH above 7 (with less than 10^-7 mole of H⁺ ions per liter) is basic.

In short, a solution that contains more H⁺ ions than OH^– ions is acidic, and one that contains more OH^– ions than H⁺ ions is basic.

Strong and Weak Acids and Bases:

Hydrochloric acid (HCI) is an example of a common acid. It is a strong acid, meaning that it tends to be almost completely ionized in aqueous solution into H⁺ and CI^– ions. Sodium hydroxide (NaOH) is a common strong base; in aqueous solution, it exists almost entirely as Na⁺ and OH^– ions. Weak acids and weak bases are those that ionize only slightly.

Most organic acids owe their acidic properties to a group of atoms called the carboxyl group, which includes one carbon, two oxygen and a hydrogen atom (symbolised as -COOH).

When an organic acid is dissolved in water, the carboxyl group dissociates to yield hydrogen ions:

Thus any compound containing carboxyl is a hydrogen donor, or an acid. It is a weak acid, however, because, as indicated by the arrows, the —COOH ionizes only slightly.

Compounds that contain the amino group (—NH₂) act as weak bases since the —NH₂ has a weak tendency to accept hydrogen ions (protons), thereby forming —NH ₃+:

Because of the strong tendency of H and OH^– to combine and the weak tendency of water to ionize, the concentration of H⁺ ions will always decrease as the concentration of OH^– increases, and vice versa.

If HCl is added to a solution containing NaOH, the following reaction will take place:

Na⁺ + OH- + H⁺ + CI^– → H₂O + Na⁺ + Cl^–

In other words, if an acid and a base of comparable strengths are added in equivalent amounts, the solution will once more be neutral.

Buffers:

Solutions more acidic than pH 1 or more basic than pH 14 are possible, but these are not included in the scale because they are almost never encountered in biological systems. In fact, almost all the chemistry of living things takes place at pH’s between 6 and 8.

Notable exceptions are the chemical processes in the stomach of humans and other animals, which take place at a pH of about 2. Human blood, for instance, maintains an almost constant pH of 7.4 despite the fact that it is the vehicle for a large number and variety of nutrients and other chemicals being delivered to the cells, as well as for the removal of wastes, many of which are acids or bases.

The maintenance of a constant pH-an example of homeostasis is important because the pH greatly influences the rate of chemical reactions. Organisms resist strong, sudden changes in the pH of blood and other fluids by means of buffers.

Buffers help maintain constant pH by their tendency to combine with H+ ions and thus remove them from solution as the H⁺ ion concentration begins to rise and to release them as it falls. Buffers are often combinations of H⁺ donor and H⁺ acceptor forms of weak acids or bases.

The capacity of a buffer system to resist changes in pH is greatest when the concentrations of H⁺ donor and H⁺ acceptor are equal. As the ratio changes in either direction, the buffer becomes less effective.

The major buffer system in the human bloodstream is the acid-base pair H₂CO₃ HCO₃^–.

The weak acid H₂CO₃ (carbonic acid) dissociates into H⁺ and bicarbonate ions as shown in the equation below:

The H₂CO₃ HCO₃^– buffer system resists the changes in pH that might result from the addition of small amounts of acid or base by “soaking up” the acid or base. For example, if a small amount of H⁺ in the form of an organic acid is added to the system, it combines with the H⁺ acceptor HCO₃^– to form H₂CO₃.

This reaction removes the added H⁺ and maintains the pH near its original value. If a small amount of OH^– is added, it combines with the H⁺ to form H₂O; more H₂CO₃ tends to ionize to replace the H⁺ as it is used.

Control of the pH of the blood is rendered even “tighter” by the fact that the H₂CO₃ is in equilibrium with dissolved carbon dioxide (CO₂) in the blood:

As the arrows indicate, the two reactions are in equilibrium and the equilibrium favours the formation of CO₂; in fact, the ratio is about 100 to 1 in favour of CO₂ formation.

Dissolved CO₂ in the blood is, in turn, in equilibrium with the CO₂ in the lungs. By changing your rate of breathing, you can change the HCO₃ concentration in the blood and thus adjust the pH of your internal fluids.

Obviously, if the blood should be flooded with a very large excess of acid or base, the buffer would fail, but normally it is able to adjust continuously and very rapidly to the constant small additions of acid or base that normally occur in body fluids.