Concentrations in solutions

Type of physical science: Chemistry

Field of study: Chemical reactions

The vast majority of chemical reactions, both those occurring naturally and those performed by researchers, occur in solution. In order to appreciate these reactions, it is necessary to know the forms taken by chemicals in solution and to be aware of the way that scientists describe solutions and solution concentrations.

Overview

The chemical and physical properties of many substances in water solutions are largely dependent upon their dissociation into electrically charged particles called ions. The majority of chemical reactions, both naturally occurring and those guided by scientists, occur in water solutions. Biological fluids are, for the most part, water solutions, so most of biochemistry is solution chemistry. Most reactions done under the guidance of a researcher are done in water solution, for convenience in handling and because of the greater speed of the reaction that is available in solution compared to reactions between solids or pure liquids. There are important chemical reactions that are carried out in nonaqueous solutions, but ions do not play a prominent role in most of them.

The material dissolved in water is called the solute, and the water itself is termed the solvent. Scientists have settled on several common ways to describe the concentration of solutions in terms of the amount of solute dissolved in a given amount of solvent. Some of the more common designations are: percentage of solute, parts per million or per billion, and molarity. Percentage (parts per hundred), parts per million, and parts per billion all are similar in that they are ratios of the amount of solute to the total amount of solution, with the amounts generally measured in grams. A part per billion is an extremely low concentration but one that often is used in measuring environmental contaminants. Molarity is a more technical unit, defined as the amount of chemical substance, measured in moles, dissolved per liter of solution.

This unit is very convenient for chemists because the amount of chemical substance relates directly to the amounts of materials that will react with one another. Thus, in determining how much of two solutions to mix to make a desired product, molarity is of direct application.

Substances that dissolve in water behave in a manner that allows them to be classified under two broad headings. Some materials, when they dissolve in water, cause the resulting solution to be electrically conductive. Others leave the water in its original, nonconductive condition. Materials of the first type, such as sodium chloride (table salt), are called electrolytes, and materials of the second type, such as sugar, are nonelectrolytes. The key to understanding this behavior is the realization that electrolytes, when they dissolve, undergo a dissociation process whereby they break up into their constituent ions. Nonelectrolytes, on the other hand, go into solution as neutral molecules.

Within the group of electrolytes, there is a variation in the extent to which water is made conductive. Some materials, such as table salt, greatly increase the conductivity of water, while others, such as vinegar, increase it to a significantly smaller degree. These materials are referred to as strong and weak electrolytes, respectively, and there is no clear line of demarcation between these classes of solutes. The degree of strength or weakness of an electrolyte is a description of the extent to which the material dissociates. Materials that dissociate completely form many ions in solution, and these solutions are good conductors of electricity, while those in which a large percentage of the molecules remain undissociated form fewer ions and are poorer conductors.

Electrolytes, when acted upon by water molecules, are broken up into positively and negatively charged particles called cations and anions, respectively. The force responsible for this dissociation is the attraction between the water molecules and the ions of the solute. Once separated, these ions draw water molecules in a cluster around themselves, which provides a shield to prevent the oppositely charged ions from being attracted back together. This clustering effect is called hydration. The strength of the force that binds the water molecules to the ions and gives them stability allows the formation of solutions of electrolytes. In cases in which this force is weak enough, the ions stay bound together in the form of an insoluble substance. Situations with an intermediate strength of this hydration force result in the cases of partial dissociation, and the solutes are referred to as weak electrolytes. In the case of nonelectrolytes, the solution process results in leaving the neutral solute molecules intact and surrounded by water molecules.

Naturally occurring water, even rainwater, contains dissolved materials. Groundwater sources leach ions from the minerals through which they flow, and it is the presence of such ions that causes water to be designated hard (with much dissolved material) or soft. Rainwater is not free of dissolved materials, as it dissolves both gases and small particles in its passage through the atmosphere. The fertilizing effect of rain, as well as the deleterious effect of acid rain, is a consequence of this process.

It is reasonable to expect that if two chemical species are to react with each other, they must come into contact. The parts of a solid in the interior of a piece are not able to come into contact with another reactant until the outer layers have been reacted and removed. Similarly, a poorly stirred liquid has parts in which such contact is not possible. The ability of a solvent to break a material into its constituent ions greatly increases the opportunity for contact between reacting species and thus allows for a much-increased reaction rate. Dissolving a cube of table salt 1 centimeter on a side produces about a million million million ions. This clearly represents an advantage over having just one cube of material in providing chances for collisions between reacting species.

There is a set of behaviors of solutions, called colligative properties, that are determined by the number of particles, ions, or molecules present in a given volume of solution.

Colligative properties are not dependent on the chemical identity of the particles, just on their concentration. As a result of the presence of a solute, a solution has a lower vapor pressure, lower freezing point, and higher boiling point than the solvent itself. These are three colligative properties; the other is the existence of an osmotic pressure when two solutions of differing concentration are separated by a membrane that will allow the passage of solvent molecules but not solute particles. Such a membrane is called a semipermeable membrane. As a result of these effects, water solutions evaporate more slowly than water itself, and water in a cooling system protected by an antifreeze will not freeze until colder than water itself, or boil as readily as water itself. Cell membranes are semipermeable, and the osmotic effect is important in biological systems as a means of regulating cellular activity.

Applications

Solutions are an integral part of everyday life in ways both obvious and subtle. The water we drink, the landscape around us, the paint we put on walls, the cook's preparations, and most of our bodies consist of water solutions. In a host of important instances, the solutes in these solutions have been dissociated, and the behavior of the solution is a response to the presence of ions.

Water passing over the surface and through the ground changes the earth not only by physical actions but through the process of solution. Rain, as it passes through the atmosphere, becomes slightly acidic because of the reactions between water and oxides of carbon, sulfur, and nitrogen present there. Caves are carved out of the earth by the dissolving action of water on the minerals present. The stalagmites and stalactites present in caves are formed when the solvent, water, later evaporates from the solutions and deposits the minerals. This is also one source of the minerals in water supplies that is spoken of in terms of hardness. Other sources are minerals that have been put on the ground as fertilizers or that have entered the water system in water runoff from populated areas. In some circumstances, the level of dissolved material becomes too high, a circumstance made evident by the large amount of soap scum present when this water is used for cleaning. When this occurs, people resort to water-softening equipment, which normally comes in two varieties. A method that is practical for large volumes of water relies on the process of ion exchange, through which one ion, normally sodium, is substituted for the calcium ion that is the major source of hardness. Another process uses electrical energy to force water to flow through a semipermeable membrane in the direction opposite to its natural flow. This movement leaves the ions behind and results in softened water. Even in areas where the water is not overly hard, the cleaning materials are formulated with materials called surfactants, many of which are electrolytes that react with the components causing water hardness and allow for more effective cleaning.

Dissolved minerals are a necessary part of a healthy diet, and at times people take mineral supplements to ensure that they have enough iron, calcium, iodine, zinc, phosphorus, or others. It should be noted that some of these are the same elements added in plant foods to simulate the proper growth of plants. The biochemical importance of these being in a soluble form may be realized by considering the medical use of the barium sulfate cocktail to make the digestive tract opaque during X-ray analysis. This compound of barium is highly insoluble and does not enter the biological reactions in the body, and thus can be tolerated although barium ions in solution are highly toxic to humans.

Other medically important solutions are those given in the form of intravenous injections. The goal of these procedures is to adjust and maintain the correct electrolyte balance in the body fluids. A correct level is necessary to keep the ion concentration in the fluids surrounding cells at the same level as the intracellular fluid. If an imbalance occurs, the existence of an osmotic pressure across the semipermeable cell membrane would either drive water into the cell until it bursts or draw water out of the cell, causing it to shrivel. Maintaining the correct balance is therefore of utmost importance in any medical treatment.

A particular medical application of the osmotic effect deserves special mention. The function of a correctly operating kidney is to remove materials that are dissolved in the bloodstream whose presence could lead to toxic reactions. Patients with poorly functioning kidneys are sometimes treated by the process of dialysis. During the process of dialysis, the patient's blood is circulated outside the body and through a machine equipped with a semipermeable membrane, which allows the unwanted materials to be separated from the blood, which is then returned to the patient.

The colligative properties of solutions, other than the osmotic effect, are also seen in many household activities. The presence of an electrolyte dissolved in water causes the freezing point of that water to be lowered several degrees. This is used to advantage when salt is put on the icy sidewalk; the water no longer freezes at that cold temperature. The same effect is put to good use when homemade ice cream is made. In order to freeze the solution that is to become ice cream, the can of solution is placed in a mixture of ice, water, and salt. The presence of the salt causes the water and ice to make a slush, the temperature of which is about -18 degrees Celsius, many degrees lower than water's normal freezing point of 0 degrees Celsius.

The exact opposite effect, the raising of water's boiling point by the presence of salt, is often used when foods are cooked by boiling. For example, pasta tends to lose its firmness when it stands too long in water. Adding salt to the cooking water increases the boiling point and allows the water to attain a higher temperature, at which the chemical reactions that occur in cooking are speeded up. This means that the cooking can be done in a shorter time and that the pasta retains some of its firmness. Many other recipes call for a pinch of salt to be added to the cooking water for exactly the same purpose.

Context

The understanding that some solutes spontaneously break apart into their constituent ions has changed in detail since its first formulation, but the broad features of the theory are still as they were put forward by Svante Arrhenius in his doctoral thesis in the early 1880's. Arrhenius' ideas were not well received by the chemists of that day, and there was dispute about accepting his thesis. It was quite a mental shift to accept the idea that very stable substances, such as sodium chloride, would break apart on their own. His work, however, came to the attention of Wilhelm Ostwald, an established chemist who championed his cause. It was through Ostwald's publishing efforts that the theory became widely known. As this happened and others realized its power in explaining the results of their own work, the theory of electrolytic dissociation took its place as one of the foundational theories of the newly developing branch of chemistry known as physical chemistry. The purpose of this branch of chemistry was to explain the physically measured properties of materials in terms of the chemical composition of the material.

The evidence for ionization that Arrhenius drew on to establish his theory can be placed in three categories: electrical evidence, chemical evidence, and colligative evidence.

Studies of the behavior of dissolved materials under the influence of electricity had resulted in the knowledge that some solutes greatly increased water's ability to conduct electricity. Further studies had identified the materials produced at the electrodes during the passage of a current and had connected these products with the solutes present; Michael Faraday had given a quantitative statement of this effect in his laws of electrolysis. The chemical evidence came in the form of the consistent color of solutions containing particular metal salts (copper solutions being blue); in the common properties of materials classed as acids or as bases; in the rapid reaction of water solutions compared with the slow reactions of the same materials as solids; and in the analysis of products formed during solution reactions, which appeared to have formed by the interchange of ions from the reactants. The colligative evidence came in the form of the measurements of freezing point lowering, boiling point raising, vapor pressure lowering, and osmotic pressures.

Prior to Arrhenius' hypothesis, the electrical evidence was explained by the assertion that the ionization was caused by the electrical energy present, and the details of the other effects were unexplained.

Current understanding of the dissociation process takes into account the importance of the role of the solvent molecules in providing the force needed to separate and keep separated the ions of an electrolyte. Indeed, these forces can be calculated, and the calculations confirmed by measurements. Much of this later refinement of Arrhenius' theory is the result of the work of Peter J. W. Debye and Erich Huckel.

Principal terms

DISSOCIATION: the process of a chemical compound breaking apart into its constituent charged particles (ions)

COLLIGATIVE PROPERTY: a change in the behavior of a solvent caused by adding a solute; the amount of the change is dependent only on the solute's concentration and not on its identity

ELECTROLYTE: a substance that, when dissolved in water, makes the water a better conductor of electricity

ION: a fragment of a chemical compound that has an electrical unbalance, causing it to be either positively changed (a cation) or negatively charged (an anion)

SOLUTE: the portion of a solution that changes state (solid or gas to liquid) when the solution is prepared

SOLVENT: the portion of a solution that does not change state when the solution is prepared

Bibliography

Ebbing, Darrell D., and Mark S. Wrighton. "Solutions." In GENERAL CHEMISTRY. Boston: Houghton Mifflin, 1990. A first-level college chemistry text that treats the topic of solutions in a quantitative way. The reader who finds the mathematics bothersome can read past the numerical portions and find an understanding of solutions as seen from the chemist's point of view.

Hammel, H. T. "Colligative Properties of a Solution." SCIENCE 192 (May, 1976): 748-755. A survey of colligative properties from both a descriptive and a quantitative point of view. Deals with the underlying, molecular-level cause of the effects.

Ihde, Aaron J. "Physical Chemistry I: Origins." In DEVELOPMENT OF MODERN CHEMISTRY. New York: Dover, 1964. In a classic survey of the historical development of chemistry, chapter 15 describes the origins of the subspecialty of physical chemistry. Includes a fine discussion of Arrhenius' theory, its origin and its development in the context of all that was happening in chemistry in those exciting years at the beginning of the twentieth century.

MacCarthy, Patrick. "Classification of Experimental Methods in Solution Chemistry." JOURNAL OF CHEMICAL EDUCATION 63 (1986): 339-343. Gives a brief description of several of the most common experimental techniques that are used to study reactions in solution and then proposes a method to compare the results available from the various methods.

Macintyre, Ferran. "Why the Sea Is Salt." SCIENTIFIC AMERICAN 223 (November, 1970): 104-115. Without a doubt, the most important ionic solution that exists outside the human body is the ocean. This article describes which ions are present, their concentrations, and their major sources and controlling cycles.

Snoeyink, Vernon L., and David Jenkins. WATER CHEMISTRY. New York: John Wiley & Sons, 1980. A first-level text applying principles of chemistry to the water environment. Most of the book is at a level that is available to general understanding, and it does well in setting the subject of electrolyte solutions in the context of both natural and polluted water.

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