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1 WHAT ARE SOLUTIONS? WHY ARE THEY IMPORTANT?

1 WHAT ARE SOLUTIONS? WHY ARE THEY IMPORTANT?

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Thus, it is clear that many of the liquids that people come in contact with are actually

solutions.

There are many examples of important and useful solutions. Sugar, for instance,

must be dissolved before it can be used by the body for food. Candy, which is sugar

with added flavoring, dissolves in the mouth to form a solution of sugar with saliva.

The coolant in an automobile’s cooling system is a solution of antifreeze in water.

This solution freezes at a much lower temperature than pure water, thus preventing

the liquid from freezing and cracking the engine block. A solution containing the

simple sugar glucose and some other substances may be injected directly into the

veins of an ill or injured person who cannot take food throught the mouth. Chemists

use many different kinds of solutions that undergo chemical reactions with other

kinds of chemicals. By measuring how much of solution is required to complete a

reaction in a procedure called titration, the chemist can tell how much of a

particular kind of chemical is in a solution.

O2(g)



CO 2(g)



O2(aq)



CO2(aq)

HCO3



Cl2

Ca2+



Figure 7.1 Tap water is actually a solution that contains, in small quantities, many chemical

species besides H2O.



Reactions in Solution

One of the most important properties of solutions is their ability to allow chemical species to come into close contact so that they can react. For example, if

perfectly dry crystals of calcium chloride, CaCl2, were mixed with dry crystals of

sodium fluoride, NaF, a chemical reaction would not occur. However, if each is

dissolved in separate solutions which are then mixed, a precipitation reaction occurs,

CaCl2(aq) + 2NaF(aq) → CaF2(s) + 2NaCl(aq)



(7.1.1)



in which calcium chloride and sodium fluoride in aqueous solution (aq) react to

produce calcium fluoride solid (s) and a solution of sodium chloride. This reaction

occurs because in solution the Ca 2+ ions (from dissolved CaCl2) and the F- ions (from

dissolved NaF) move around and easily come together to form CaF2. The calcium

fluoride product does not stay in solution but forms a precipitate; it is insoluble.

In other cases, solutions enable chemical reactions to occur that result in materials being dissolved. Some of these reactions are important in geology. Consider

limestone, which is made of calcium carbonate, CaCO3. Limestone does not react

with dry CO2 gas, nor is it soluble in pure water. However, when water containing

dissolved CO2 contacts limestone, a chemical reaction occurs:



© 2001 CRC Press LLC



CaCO3(s) + CO2(g) + H2O → Ca 2+(aq) + 2HCO3-(aq)



(7.1.2)



The calcium ion and the bicarbonate ion, HCO3-, remain dissolved in water; this

solution dissolves the limestone, leaving a cave or hole in the limestone formation.

In some regions, such as parts of southern Missouri, this has occurred to such an

extent that the whole area is underlain by limestone caves and potholes; these are

called karst regions.



Solutions in Living Systems

For living things the most important function of solutions is to carry molecules

and ions to and from cells. Body fluids consist of complex solutions. Digestion is

largely a process of breaking down complex, insoluble food molecules to simple,

soluble molecules that may be carried by the blood to the body cells, which need

them for energy and production of more cell material. On the return trip the blood

carries waste products, such as carbon dioxide, which are eliminated from the body.



Solutions in the Environment

Solutions are of utmost importance in the environment Figure 7.2). Solutions

transport environmental chemical species in the aquatic environment and are crucial

participants in geochemical processes. Dissolution in rainwater is the most common

process by which atmospheric pollutants are removed from air. Acid rain is a

solution of strong mineral acids in water. Many important environmental chemical

processes occur in solution and at the interface of solutions with solids and gases.



Figure 7.2 Many environmental chemical phenomena involve solutions.



© 2001 CRC Press LLC



Pollutant pesticides and hazardous waste chemicals are transported in solution as

surface water or groundwater. Many hazardous waste chemicals are dissolved in

solution; often the large amount of water in which they are dissolved makes their

treatment relatively more difficult and expensive.



Industrial Uses of Solutions

Solutions are used throughout industry. Many chemical reactions that are part of

the manufacture of important industrial chemicals occur in solutions, and the

chemical processes that take place in solution—solution chemistry—are very

important in the chemical industry. Natural brines are solutions that contain a lot of

dissolved materials and that occur underground and in some lakes (saline lakes).

Some natural brines are important sources of valuable chemicals. Commercially

important chemicals that are recovered from brines include borax, a compound

containing boron and oxygen, which is used as an antiseptic, in making ceramics,

and in some cleaning formulations; bromine salts; and potassium salts, including

those used in potassium fertilizer. Solutions of detergents are used for cleaning;

some dyes are applied as solutions. Ammonia fertilizer may be added to the soil as a

solution of NH3 in water.

Many organic materials do not dissolve significantly in water. Such substances

are usually soluble in organic solvents, such as benzene and carbon tetrachloride.

When used to dissolve organic substances, organic liquids are commonly called

solvents. Some aspects of solvents are discussed in the following section.



7.2 SOLVENTS

Water is the solvent for most of the solutions discussed in this chapter. However,

as noted above, many other liquids are also used as solvents. Other than water, most

solvents are organic (carbon-containing) liquids. Some of the more important

organic solvents are shown in Table 7.1.

There are many uses for solvents. One of the most important of these is their role

as media in which chemical reactions may occur. In the chemical industry, solvents

are employed for purification, separation, and physical processing. Solvents are also

used for cleaners; one important example is the use of organic solvents to dissolve

grease and oil from metal parts after they have been fabricated. (In the past, though

less so now, solvents used for parts cleaning were allowed to evaporate to the

atmosphere, a major source of atmospheric organic pollutants.) The chemicals that

make up synthetic fibers, such as rayon, are dissolved in solvents, then forced under

very high pressure through small holes in a special die to make individual filaments

of the fiber. One of the most important uses for solvents is in coatings, which include

paint, printing inks, lacquers, and antirust formulations. In order to apply these

coatings, it is necessary to dissolve them in a solvent (vehicle) so that they may be

spread around on the surface to be coated. The vehicle is a volatile liquid, one that

evaporates quickly to form a vapor; when it evaporates, it leaves the coating behind

as a thin layer.



© 2001 CRC Press LLC



Fire and toxicity are major hazards associated with the use of many solvents.

Some organic solvents, such as benzene, are even more of a fire hazard than

gasoline. Both benzene and carbon tetrachloride are toxic and can damage the body

in cases of excess exposure. Benzene is suspected of causing leukemia, and worker

exposure to this solvent is now carefully regulated. The toxicity hazard of solvents

arises from absorption through the skin and inhalation through the lungs. One

solvent, dimethyl sulfoxide, is relatively harmless by itself but has the property of

carrying toxic solutes through the skin and into the body. Exposure to solvent vapor

is limited by occupational health regulations which include a threshold limiting

value (TLV). This is the measure of solvent vapor concentration in the atmosphere

considered safe for exposure to healthy humans over a normal 40-hour work week.

Table 7.1 Important Organic Solvents



Solvent



Solvent Use (may have many

other nonsolvent uses)



Approximate Annual

U.S. Production,

millions of kilograms



Acetone



Solvent for spinning cellulose acetate

fibers and for spreading paints and

other protective coatings.



1,000



Benzene



Dissolves grease and other organic

compounds.



6,500



Perchloroethylene



Best solvent for dry cleaning, also

used for degreasing metals and

extraction of fats.



125



Stoddard solvent



Mixture of alkanes and aromatic hydrocarbons containing from 7 to 12 C

atoms per molecule used as a solvent

for organic materials



Toluene



Dissolves grease and other organic

materials; substitute for benzene, but

not so toxic



Trichloroethylene



Vapor degreasing of metal parts; solvent

for greases, oils, fats, waxes, and tars;

fabric cleaner; waterless dying; ingredient of formulations of adhesives, lubricants, paints, varnishes, paint strippers



17



2,700



90



7.3 WATER—A UNIQUE SOLVENT

The remainder of this chapter deals with water as a solvent. Water is such an

important compound that all of Chapter 11 is spent discussing it in detail. Here, just

those properties of water that relate directly to its characteristics as a solvent are

summarized.



© 2001 CRC Press LLC



At room temperature H2O is a colorless, tasteless, odorless liquid. It boils at

100°C (212°F) and freezes at 0°C (32°F). Water by itself is a very stable compound;

it is very difficult to break up by heating. However, as explained in Section 8.6,

when electrically conducting ions are present in water, a current may be passed

through the water, causing it to break up into hydrogen gas and oxygen gas.

Water is an excellent solvent for a variety of materials; these include many ionic

compounds (acids, bases, salts). Some gases dissolve well in water, particularly

those that react with it chemically. Sugars and many other biologically important

compounds are also soluble in water. However, greases and oils generally are not

soluble in water but dissolve in organic solvents instead.

Some of water’s solvent properties can best be understood by considering the

structure and bonding of the water molecule:

(+)

105˚



H



H



O

(-)



The water molecule is made up of two hydrogen atoms bonded to an oxygen atom.

The three atoms are not in a straight line, but form an angle of 105°.

Because of water’s bent structure and the fact that the oxygen atom attracts the

negative electrons more strongly than do the hydrogen atoms, the water molecule

behaves like a body having opposite electrical charges at either end or pole. Such a

body is called a dipole. Due to the fact that it has opposite charges at opposite ends,

the water dipole may be attracted to either positively or negatively charged ions.

Recall that NaCl dissolves in water to form positive Na+ ions and negative Cl- ions

in solution. The positive sodium ions are surrounded by water molecules with their

negative ends pointed at the ions, and the chloride ions are surrounded by water

molecules with their positive ends pointing at the negative ions, as shown in Figure

7.3. This kind of attraction for ions is the reason why water dissolves many ionic

compounds and salts that do not dissolve in other liquids. Some noteworthy

examples are sodium chloride in the ocean; waste salts in urine; calcium bicarbonate,

which is very important in lakes and in geological processes; and widely used

industrial acids (such as HNO3, HCl, and H2SO4).



Na+



Cl-



Figure 7.3 Polar water molecules surrounding Na+ ion (left) and Cl- ion (right).



© 2001 CRC Press LLC



In addition to being a polar molecule, the water molecule has another important

property which gives it many of its special characteristics: the ability to form hydrogen bonds. Hydrogen bonds are a special type of bond that can form between the

hydrogen in one water molecule and the oxygen in another water molecule. This

bonding takes place because the oxygen has a partly negative charge and the hydrogen, a partly positive charge. Hydrogen bonds, shown in Figure 7.4 as dashed lines,

hold the water molecules together in large groups.

Hydrogen bonds also help to hold some solute molecules or ions in solution.

This happens when hydrogen bonds form between the water molecules and

hydrogen or oxygen atoms on the solute molecule (see Figure 7.4). Hydrogen

bonding is one of the main reasons that some proteins can be put in water solution or

held suspended in water as extremely small particles called colloidal particles (see

Section 7.9).

O



Hydrogen bonds between

water molecules



H

O

H



H

H



O



H



O



Hydrogen bonds

between water and

solute molecules



Solute

molecule



H



N

H



O

H

Figure 7.4 Hydrogen bonding between water molecules and between water molecules and a solute

molecule in solution.



7.4 THE SOLUTION PROCESS AND SOLUBILITY

Very little happens to simple molecules, such as N2 and O2, when they dissolve

in water. They mingle with the water molecules and occupy spaces that open up

between water molecules to accomodate the N2 and O2 molecules. If the water is

heated, some of the gases are driven out of solution. This may be observed as the

small bubbles that appear in heated water just before it boils. A fish can extract some

of the oxygen in water by “breathing” through its gills; just 6 or 7 parts of oxygen in

a million parts of water is all that fish require. Water saturated with air at 25°C

contains about 8 parts per million oxygen. Chapter 12 discusses how only a small

amount of an oxygen-consuming substance can use up this tiny portion of oxygen in

water and cause the fish to suffocate and die.

Although N2 and O2 dissolve in water in the simple form of their molecules, the

situation is much different when hydrogen chloride gas, HCl, dissolves in water. The

hydrogen chloride molecule consists of a hydrogen atom bonded to a chlorine atom



© 2001 CRC Press LLC



H



2O



with a covalent bond. (Recall that covalent bonds are formed by sharing electrons

between atoms.) Water can absorb large amounts of hydrogen chloride: 100 grams

of water at 0°C will dissolve 82 g of this gas. When HCl dissolves in water (Figure

7.5), the solution is not simply hydrogen chloride molecules mixed with water

molecules. The water has a strong effect upon the HCl molecule, breaking it into two

parts, with the 2 electrons in the chemical bond staying with the chlorine. This forms

a positively charged hydrogen ion, H+, and a negatively charged chloride ion, Cl-.



H 2O

Solution



H–Cl(g)



H+(aq) + Cl-(aq)



Figure 7.5 HCl dissolving in water. Water breaks apart a hydrogen chloride molecule to form a

hydrogen ion, H+ , and a chloride ion, Cl-.



In water solution, the chloride ion is surrounded by the positive ends of the water

molecules, which are attracted to the negatively charged Cl- ion. This kind of

attraction of water molecules for a negative ion has already been shown in Figure

7.3. The H+ ion from the HCl molecule does not remain in water as an isolated ion;

it attaches to an unshared electron pair on a water molecule. This water molecule,

with its extra hydrogen ion and extra positive charge, becomes a different ion with a

formula of H3O+; it is called a hydronium ion. Although a hydrogen ion in solution

is indicated as H+ for simplicity, it is really present as part of a hydronium ion or

larger ion aggregates (H5O2+, H7O3+).

H



H

H



+



O



+



H+



H



H2O



H



O



+



H



H3O+

(Hydronium ion)



+



Figure 7.6 A hydrogen ion, H , bonds to a water molecule, H 2O, to produce a hydronium ion,

H3O+ . Bonding to additional water molecules may form larger aggregates, H5O2+ , H7O3+ .



The solution of hydrogen chloride in water illustrates a case in which a neutral

molecule dissolves and forms electrically charged ions in water. While this happens

with other substances dissolved in water, the hydrogen ion resulting when substances

like HCl dissolve in water is particularly important because it results in the

formation of a solution of acid. So rather than calling this a solution of hydrogen

chloride, it is called a hydrochloric acid solution.



7.5 SOLUTION CONCENTRATIONS

In describing a solution it is necessary to do so qualitatively, that is, to specify

what the solvent is and what the solutes are. For example, it was just seen that HCl

gas is the solute placed in water to form hydrochloric acid. It is also necessary to



© 2001 CRC Press LLC



know what happens to the material when it dissolves. For instance, one needs to

know that hydrogen chloride molecules dissolved in water form H+ ions and Clions. In many cases it is necessary to have quantitative information about a solution,

its concentration. The concentration of a solution is the amount of solute material

dissolved in a particular amount of solution, or by a particular amount of solvent.

Solution concentration may be expressed in a number of different ways. Solutions

used in technical applications, such as for cleaning, are often made up of a specified

number of grams of solute per 100 mL (milliliters) of solvent added. On the other

hand, a person involved in crop spraying may mix the required solution by adding

several pounds of pesticide to a specified number of barrels of water.

The concentrations of water pollutants frequently are given in units of

milligrams per liter (mg/L). Most of the chemicals that commonly pollute water are

harmful at such low levels that milligrams per liter of water is the most convenient

way of expressing their concentrations. For example, water containing more than

about one-third of a milligram of iron per liter of water can stain clothing and

bathroom fixtures. To get an idea of how small this quantity is, consider that a liter

of water (strictly speaking, at 4˚C) weighs 1 million milligrams. Water containing

one-third part per million of iron contains only 1 mg of iron in 3 million milligrams

(3 liters) of water. Because 1 liter of water weighs 1 million milligrams, 1 mg of a

solute dissolved in a liter is a part per million, abbreviated as ppm. The terms part

per million and milligrams per liter are both frequently used in reference to levels

of pollutants in water.

Some pollutants are so poisonous that their concentrations are given in micrograms per liter (µg/liter). A particle weighing a microgram is so small that it cannot

be seen with the naked eye. Since a microgram is a millionth of a gram, a liter of

water weighs 1 billion micrograms. So, 1 microgram per liter is 1 part per billion

(ppb). Sometimes it is necessary to think in terms of concentrations that are this low;

a good example of this involves the long-banned pesticide Endrin. It is so toxic that,

at a concentration of only two-thirds of a microgram per liter, it can kill half of the

fingerlings (young fish) in a body of water over a four-day period.

At the other end of the scale, it may be necessary to consider very high concentration; these are often given as percent by weight. As indicated in Figure 7.7, the

concentrations of commercial acids and bases are often expressed in this way. For

example, a solution of concentrated ammonia purchased for laboratory use is 28%

by mass NH 3; this means that out of 100 g of the ammonia solution, 28 g are NH 3

and 72 g are water. Commercial hydrochloric acid is about 36% by mass hydrogen

chloride, HCl: of 100 g of concentrated hydrochloric acid, 36 g are HCl and 64 g are

water.



Molar Concentration

The mass of a solute in solution often does not provide full information about

its effect. For example, a solution of either NaCl or sodium iodide (NaI) will remove

silver from a solution of silver nitrate (AgNO3) when the two solutions are mixed.

The two possible chemical reactions are,

AgNO3(aq) + NaCl(aq) → NaNO 3(aq) + AgCl(s)



© 2001 CRC Press LLC



(7.5.1)



AgNO3(aq) + NaI(aq) → NaNO 3(aq) + AgI(s)



(7.5.2)



in which solid precipitates of AgCl and AgI, respectively, come out of solution. The

molcular mass of NaCl is 58.5 and that of NaI is 149.9. By comparing the removal

of silver from solution by solutions of NaCl and NaI, both containing the same mass

of solid in a specific volume of solution, it can be seen that a given volume of NaCl

solution removes more silver. In chemical reactions such as these, it is the number of

ions or molecules that is important, not their masses in solution. It would be convenient to have a way of expressing concentrations in quantities that would be directly



NH3

28%

Concentrated



HCl

36%

Concentrated



HNO3

69%

Concentrated



Figure 7.7 Percentages of solutes in commercial solutions of ammonia, hydrochloric acid, and

nitric acid.



related to the number of molecules or ions in solution. Instead of expressing such

huge numbers, however, the chemist normally works with moles. (Recall from

Section 2.3 that a mole is the number of grams of a substance that is equal

numerically to the mass in amu of the smallest unit of the substance. For example,

since the molecular mass of NaCl is 58.5, a mole of NaCl weighs 58.5 g.) The molar

concentration of a solution is the number of moles of solute dissolved in a liter of

solution. Molar concentration is abbreviated with the letter M. A 1 M solution has 1

mole of solute dissolved in a liter of solution. To understand this concept better,

consider the following:

• The molecular mass of HCl is 36.5; a 1 M solution of HCl has 1 mole of

HCl (36.5 g) dissolved in a liter of solution.

• The molecular mass of NH3 is 17; a 1 M solution of NH3 has one mole of

NH3 (17 g) in a liter of solution.

• The molecular mass of glucose, C6H12O6, is 180; a 1 M solution of

glucose has 1 mole of glucose (180 g) in a liter of solution.

Of course, solutions are not always exactly 1 M in concentration. However, it is

easy to perform calculations involving molar concentration using the relationships

expressed in the following equation:



© 2001 CRC Press LLC



M =



Moles solute

Mass solute, g

=

Volume solution, L

Molar mass solute × Volume solution, L



(7.5.3)



To see how this equation can be used, consider the following examples:

Example: What is the molar concentration of a solution that contains 2.00 moles of

HCl in 0.500 L of solution?

Answer:

M = 2.00 mol = 4.00 mol/L

0.500 L

Example: What is the mass of HCl, molar mass 36.5 g/mol in 2.75 L of a 0.800 M

solution?

Answer:

0.800 mol/L =



Mass solute

36.5 g/mol × 2.75 L



Mass solute = 0.800 mol/L x 36.5 g/mol x 2.75 L = 80.3 g

M =



Mass solute, g

Molar mass solute × Volume solution, L



Example: What is the molar concentration of a solution containing 87.6 g of HCl in

a total volume of 3.81 L of solution?

Answer:

M =



87.6 g

= 0.630 mol/L

36.5 g/mol × 3.81 L



To summarize the steps involved in making up a certain quantity of a solution

with a specified concentration, consider the task of making up 4.60 liters of 0.750 M

NaCl, which would involve the following steps:

Step 1. Calculate the mass of NaCl, molar mass 58.5 g/mol, required.

0.750 mol/L =



Mass solute

58.5 g/mol × 4.60 L



Mass solute = 0.750 mol/L × 58.5 g/mol × 4.60 L = 202 g

Step 2. Weigh out 202 g NaCl

Step 3. Add NaCl to a container with a mark at 4.60 L

Step 4. Add water and mix to a final volume of 4.60 L



© 2001 CRC Press LLC



Diluting Solutions

Often it is necessary to make a less concentrated solution from a more concentrated solution; this process is called dilution. This situation usually occurs in the

laboratory because it is more convenient to store more-concentrated solutions in

order to save shelf space. Also, laboratory acid solutions, such as those of

hydrochloric acid, sulfuric acid, and phosphoric acid, are almost always purchased as

highly concentrated solutions; these are more economical because there is more of

the active ingredient per bottle. For standard solutions to use in chemical analysis, it

is more accurate to weigh out a relatively large quantity of solute to make a

relatively concentrated solution, then dilute the solution quantitatively to prepare a

more dilute standard solution.

For example, the concentration of commercial concentrated hydrochloric acid is

12 M. A laboratory technician needs 1 liter of 1 M hydrochloric acid. How much of

the concentrated acid is required? The key to this problem is to realize that when a

volume of the concentrated acid is diluted with water, the total amount of solute

acid in the solution remains the same. The problem can then be solved by

considering the equation

M =



Moles solute

Volume solution



(7.5.3)



using subscripts “1” and “2” to indicate values before and after dilution,

respectively,

M1 =



(Moles solute)1



(7.5.4)



(Volume solution)1

M2 =



(Moles solute)2



(7.5.5)



(Volume solution)2

and setting “Moles solute” before and after dilution equal to each other:

M1 × (Volume solution) 1 = M2 × (Volume solution) 2



(7.5.6)



In the example cited above, the volume of HCl before dilution is to be calculated:

(Volume solution)1 = M2 × (Volume solution) 2

M1

(Volume solution)1 = 1 Mol/L × 1 L = 0.083 L

12 mol/L



(7.5.7)

(7.5.8)



The result of this calculation shows that 0.083 L, or 83 mL of 12 M HCl must be

taken to make 1.00 L of 12 M solution.

The same general approach used in solving these dilution problems can be used

when concentrations are expressed in units other than molar concentration. Concentrations of metals dissolved in water are frequently measured by atomic absorption



© 2001 CRC Press LLC



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