`
`"4-H55
`
`GENERAL CHEMISTRY
`
`ATKINS
`
`Uxfmd (.:'un'm'\If\
`
`S(IIENTl_Fl(‘.
`s\MERl(IAN
`
`.
`
`_
`Distributed by W H. Frt-exmm
`
`Page 1 of 8
`
`SENJU EXHIBIT 2072
`
`LUPIN v. SENJU
`
`IPR2015—0l100
`
`
`
`C1J‘(’(‘l' imugw
`
`K/.‘?i. Kmp.
`
`Library of Congress Cataloging-in-Publication Data
`
`I‘. W. (Pctcr Williimn).
`Atkins,
`Ucucral ciiciiiislry.
`
`]‘.»)-'40-
`
`Includes inclcx.
`
`I. Title.
`1. Chemistry.
`Q_1')3l.2.A75
`1989
`540
`ISBN ()-7167-I940-I
`
`88-3()58()
`
`(Ir)})_yiig}ii
`
`198!) I’. W. Alkins.
`
`No part of this book may be rcprod1u.'Cd by;1ny lT1(’,(_’i'l‘dl1-
`i(dl. phulogruphic. or elm Ironic proitcss, or in any fun Ill
`01' 2: pliotograpliic 1'<:cot'diug‘. 1101' may it be stored in a
`retrieval system. Lrzmsmiursd, or otherwise copied for
`[)Lli)iiL or pm.-nu: usc. wilhuul wriLu:n pcrmis.~ion from
`the publisher.
`l’rinLcd in the United States of America
`
`Scientific Al]lCI‘iC£ll1 Books is a subsidiary of Sciciitific
`American. lnc. UisLi'ibuL(:d by W. H‘ Freeman and (‘.om—
`pzinv, 4| Madison Avenue, New York, New York 10010
`I2:s456789uK1’765432In O6
`
`Page 2 of 8
`
`
`
`Therefore,
`
`0. l
`
`mol Csfigg
`
`.\«Iolality =
`
`‘ 1.92 mol C5ll,3/kg CHf4(:5lI5
`EXERCISE Calculate the molality of a solution of toluene in benzene.
`given that the mole fraction of toluene is 0.150.
`
`A
`
`[A-nszuer: 2.26 m]
`
`SOLUBILITY
`
`In this chapter we are focusing on aqueous solutions because they are
`so important, but many of our remarks apply equally to nonaqueous
`solutions. In the following discussion. remember that substances which
`dissolve to give solutions of ions that conduct electricity (e.g., sodituii
`Chloride and acetic acid) are called electrolytes (Section 2.7); those giv-
`ing solutions that do not conduct electricity because the solute remains
`molecular (e.g., glucose and ethanol) are nonelectrolytes.
`
`11.3 SATURATION AND SOLUBILITY
`
`If we add 20 g of sucroseucane sugar—to 100 mL of water at room
`temperature, all the sucrose dissolves. However, if we add 200 g. most
`dissolves but some does not (Fig.
`1 1.4). When the solvent has dissolved
`all the solute it can and some undissolved solute remains, the solution is
`said to be “saturated."
`
`If we could follow a single sucrose molecule
`The definition ofsolubility.
`in a saturated solution, we might find that at some instant it is part of
`the surface layer ofa sucrose crystal (Fig. 11.5). Shortly after, the mole-
`cule might be found in solution. Still later,
`it might be buried more
`deeply in a crystal, under many layers of molecules that had settled on
`
`
`
`FIGURE 11 .4 When a little su-
`crose is shaken with 100 ml. of
`water, it all dissolves (left). How-
`ever, when a large amount (more
`than 200 g) is added, some undis-
`solved sucrose remains (right).
`
`400 Page 3 of 8
`
`CHAPTER 11
`
`THE PROPERTIES OF
`
`SOLUTIONS
`
`
`
`top of it. There it would remain until it became exposed again and was
`able to return to the solution. In other words, a saturated solution is
`another example of dynamic equilibrium (see Section 10.4), in which a
`forward process and its reverse occur at equal rates. In this case, the
`solute continues to dissolve, and it does so at a rate that exactly matches
`the rate of the reverse process, the return of solute from the solution.
`This suggests the following definition:
`
`A saturated solution is a solution in which the dissolved and undis-
`
`solved solute are in dynamic equilibrium.
`
`'
`
`Although we cannot follow a single molecule in a saturated solution,
`we can show experimentally that the equilibrium is dynamic and not
`static. One way to do so is to add solid silver iodide, containing some
`iodine-131 in place of the usual iodine-I27, to a saturated solution of
`silver iodide. Iodine-13] is radioactive and can be detected with Geiger
`counters and other radioactivity-detection devices. After a time the so-
`lution becomes radioactive, but the total mass of dissolved solid re-
`
`mains unchanged. This shows that some 1‘ ions have dissolved and
`others have come out of solution, even though the solution was already
`saturated.
`
`A saturated solution represents the limit of a solute’s ability to dis-
`solve in a given quantity of solvent. It is therefore a natural measure of
`the solute’s “s01ubility" S:
`
`The solubility of a substance in a solvent is the concentration of the
`saturated solution.
`
`The solubilities of some substances are given in Table 1 1.3. They de-
`pend on the solvent, the temperature, and, for gases. the pressure.
`
`TABLE 11.3 The solubilities of some substances
`
`Solubility,
`g solute/100 g solvent,
`in water at
`
`Compound
`
`NH3
`
`NH.,NO3
`
`CaCl2
`
`Car,
`
`CuSO4 - 51-I20
`
`I
`
`I-{Cl
`
`AgF
`
`AgCl
`
`0°C
`
`89.5
`
`l 18
`
`59.5
`
`1.7 x 10""
`
`31.6
`
`82.3
`
`100°C or as
`
`specified
`
`Other
`
`solvents
`
`Organic solvents
`
`Alcohol, ammonia
`
`Alcohol
`
`7.4
`
`87]
`
`159
`
`203.3
`
`56.! at 60°C
`
`Alcohol, benzene
`
`6 x 10"
`
`8 x 10*‘ at 30°C
`
`182
`
`7 x 10-5
`
`205
`
`2 x 10-3
`
`FIGURE 11.5 The solute in a
`saturated solution is in dynamic
`equilibrium with the undissolved
`solute. If we could follow a single
`solute particle (the red circle), we
`would sometimes find it in solu-
`tion ancl sometimes in the solute.
`
`Page 4 of 8
`11.3 SATURATION AND SOLUBILITY
`
`401
`
` .3
`
`
`
`
`
`FIGURE 11.6 This Chile saltpe«
`tcr has survived in the arid region
`where it is mined in Chile because
`
`there is too little groundwater to
`dissolve it and wash it away.
`
`The dependence of solubility on the solute. Some substances are soluble
`in water, others sparingly (slightly) soluble. and others almost inso1u—‘
`ble. We can know which behavior to expect by referring to the “solubil-
`ity rules," which were given in Table 3.1. We used t.he rules in Chapter
`3 to choose reagents for precipitation reactions; they are also of help in
`understanding the behavior of sortie everyday substances and the
`properties of minerals. Because of the solubility of most nitrates, for
`instance, they are rarely found in mineral deposits, for they are usually
`carried away bv the water that trickles through the ground. An excep-
`tion is the large deposit of sodium nitrate in the arid coastal region of
`Chile, where groundwater is absent. This “Chile saltpeter" (Fig.
`l 1.6) '
`was the main source of nitrates for fertilizers and explosives until the
`Haber process for ammonia was developed at the start of this century.
`The low solubility of most phosphates is an advantage for skeletons,
`since bone consists largely of calcium phosphate (much of the rest is the
`protein collagen). However, this insolubility is inconvenient for agricul-
`ture. since it means that phosphorus, which is essential to the function
`of biological cells, is slow to circulate through the ecos_vstem. One of
`chemistry’s achievements has been the development of manufacturing
`processes to speed phosphates on their way as fertilizers. The phos-
`phates and hydrogen phosphates used for fertilizers are obtained from
`phosphate rocks (Fig. 11.7), principally the apatites—hvdroxyapatite,
`C1a5(P()..)3(.)I-I, and Iluorapatite, Ca5(l’(),,);4F—»b_v treating them with
`concentrated sulfuric acid:
`
`(.Iag,('P().,)3Ol'1(5) + 5HgSO.,(_aq) —-9
`3H;tP().,(aq) + 5CaS()s,(s) + l-120(1)
`
`The phosphate rocks themselves were once alive, for thev are the .
`crushed and compressed remains of the skeletons of prehistoric ani-
`mals. Calcium hydrogen phosphate (Cal-lP(),,) is more soluble than
`calcium phosphate and is included in commercial phosphate fertilizers.
`‘lust as hydrogen phosphates are more soluble than phosphates. so
`hydrogen carbonates (bicarbonates, I-lC();f) are more soluble than
`carbonates.
`[his difference is responsible for the behavior of hmrl
`water. water that contains dissolved calcium and magnesium salts. In
`particular, the difference accounts for the deposit of scale inside hot
`pipes and for the formation of a scum with soap in hard water. The .
`
`1*.&'.‘rsA'uu\n;...'.;.2.«
`
`
`
`FIG U RE 11. 1 Mining of phos-
`phate rock, the crushed remains
`of the skeletons of prehistoric ani-
`mals.
`
`402
`
`CHAPTER 11
`
`THE
`
`PROPERTIES OF SOLUTIONS-t
`
`Page 5 of 8
`
`
`
`London forces
`
`
`
`(C)
`
`FIGURE 1 1 .8 Like often dis«
`solves like. (:1) Intermolecular in-
`teractions help a polar solvent to
`dissolve other polar substances,
`(b) a hydrogen-bonding solvent to
`dissolve substances held together
`by hydrogen bonds, and (c) a sol-
`vent with strong London forces to
`dissolve nonpolar molecular sol-
`ids.
`
`
`
`\ Hydrogen
`bonding
`
`( '0)
`
`
`
`Dipole-dipole H
`interactions
`
`(3)
`
`behavior of hard water begins with the fact that rainwater contains
`dissolved carbon dioxide, and hence some carbonic acid from the reac-
`tion
`
`C0-Ag) + 1420(1) —> H2C0.~s(aq)
`
`As the water runs along and through the ground, the carbonic acid
`-reacts with the calcium carbonate of limestone or chalk and forms the
`
`more soluble hydrogen carbonate:
`
`CaCO3(s) + H2(2O3(aq) ——> Ca(HCO3)«_;(aq)
`
`These reactions are reversed when the water is heated in a kettle or
`furnace:
`
`4.
`
`\F
`
`:,.
`
`
`
`2HCO3‘(aq) L) CO32‘(aq) + C02(g) + 1420(1)
`
`The carbon dioxide is driven off, leaving carbonate ions in solution,
`-and the almost insoluble calcium carbonate is deposited as scale.
`
`.
`
`In many instances, the de-
`The dependence of solubility on the solvent.
`pendence of the solubility of a substance on the identity of the solvent
`can be summarized by the rule that “like dissolves like." That is, a polar
`liquid, such as water, is generally a much better solvent than a nonpolar
`one (such as benzene) for ionic and polar compounds. Conversely,
`nonpolar
`liquids.
`including benzene and the tetrachloroethylene
`<CQCl4) used for dry cleaning, are often better solvents for nonpolar
`'._ ‘compounds than for polar compounds (Fig. ll.8). The reason is that
`the energy of the solute molecules is similar in the solution to what it
`was in the original solid if the intermolecular forces in solution and
`solid are similar.
`If the principal cohesive forces in a solute are hydrogen bonds, the
`“like dissolves like” rule implies that it is more likely to dissolve in a
`hydrogen-bonding solvent than in others. Sucrose, for example, dis-
`solves readily in water but not in benzene. Similarly, if the principal
`Cohesive forces are London forces, the best solvent is likely to be one
`.held together by the same kind of forces. One example is carbon disul-
`fide, which is a far better solvent for sulfur than is water (Fig. 11.9),
`because solid sulfur is a molecular solid of S3 molecules held together
`by London forces.
`
`1.
`J - application of the principle oflike dissolving like. Soaps are the sodium
`1", salts of organic acids with long hydrocarbon chains, including sodium
`
`
`
`FIGURE 11.9 The molecular
`solid sulfur does not dissolve in
`water (left) but does dissolve in
`carbon disulfide (right), with
`which its molecules have strongly
`favorable London interactions.
`
`711.3
`
`Pageufizouf 81 o N
`
`AND SOLUBILITY
`
`403
`
`
`
`
`
`
`
`1 Sodium stearate
`
`
`
`2 Polyphosphate ion
`
`stearate (1); we shall denote them NaA, where HA is the organic acid.
`The anions have a polar group (called the “head group”) at one end of
`a long nonpolar group, the hydrocarbon chain. The anions (A‘) sink
`their nonpolar and thus hvdroghobic, or water-repelling, hydrocarbon
`tails into a blob of grease. Their hydroghilic, or water-attracting, head.
`groups remain on the surface of the grease blob, coating it with a skin
`of polar hydrogen—bonding groups (Fig. 11.10). The polar head
`groups enable the grease blob to dissolve in water and to be washed
`away.
`
`A problem with soaps is that they form a scum in hard water. The
`scum is the product ofa precipitation reaction that occurs because cal-
`cium salts are less soluble than Sodium salts:
`
`Ca2*(aq) + 2A”(aq) ————-> CaA2(5)
`
`One way of avoiding the problem is to use another precipitation reac-
`tion to remove the Ca“ ions from the water before the soap is used.
`This can be done by adding sodium carbonate (“washing soda”) to the
`water and precipitating calcium carbonate:
`
`Ca(HC03)2(aq) + Na2CO3(aq) _s CaCO3($) + 2NaHCO3(a.q)
`
`Another way to avoid soap scum is to add polyphosphate ions to the
`water as a component of the detergent. Polyphosphate ions (2) are
`formed when phosphates are heated, and they consist of chains and-
`rings of P04 groups. The first step in their formation is
`
`(ll)
`‘i’
`‘H’
`ll
`I
`HO—l!3—OH + HO—E|’—OH L) HO—l|’~O*l|’+OH + H20
`OH
`OH
`OH
`on
`
`Polyphosphate ions are big, and when they are added to hard water.
`they wrap around the calcium cations and hide them away from other’
`anions with which they would normally precipitate. This wrapping up
`of one ion by another is called seflestration of the ion (from the Latin
`word for “hiding away"), and the polyphosphates are called “sequester-
`ing agents."
`the
`Modern commercial detergents are mixtures of compounds,
`most important of which is the “surface-active agent,” or “surfactant.”
`Surfactant molecules are synthetic organic compounds that resemble
`the one shown below (3). Like the stearate ion, they have a hydrophilic
`
`
`
`3 A typical surfactant molecule
`
`CHAPTER 11
`
`THE
`
`PROPERTIES OF
`
`$OLUTlONS
`
`Page 7 of 8
`
`
`
`head group and a hydrophobic tail, and they act similarly. Detergents
`also contain polyphosphates to sequester calcium ions and to adjust the
`acidity. Other additives in the mixture “fluoresce" (absorb ultraviolet
`light and then give out visible light) to give the impression of greater
`cleanliness.
`
`Head group
`
`Hydrophobic tail
`
`11.4 THE EFFECT OF PRESSURE ON GAS SOLUBILITY
`
`We have noted that solubility depends on the pressure the solution
`experiences. The strongest dependence is shown by gases, which are
`more soluble at higher pressures (Fig. 1 1.1 1). A practical application of
`this phenomenon is the production of soft drinks and champagne. In
`each case carbon dioxide is dissolved in the liquid under pressure (in
`champagne, as a result of fermentation that continues in the sealed
`bottle). When the bottle is opened the pressure is released, the solubil-
`ity of the gas is greatly reduced, and the gas effervesces (bubbles out of
`solution) with a pop. A more serious consequence of the dependence of
`gas solubility on pressure is the additional nitrogen that dissolves in the
`blood of deep-sea divers. The dissolved nitrogen effervesces when the
`diver returns to the surface, resulting in the formation of numerous
`small bubbles in the bloodstream (Fig. 11.12). These bubbles can block
`the capillaries—the narrow vessels that distribute the blood--and
`starve the tissues ofoxygen, causing the painful condition known as the
`“bends,” which in serious cases can lead to death. The risk of the bends
`is reduced if helium is used instead of nitrogen to dilute the diver's
`oxygen supply, for helium is much less soluble than nitrogen.
`
`Henry’: law. The dependence of the solubility of a gas on its pressure
`was summarized in 1801 by the English chemist William Henry:
`
`Henry’s law: The solubility of a gas in a liquid is proportional to its
`partial pressure.
`
`_This law is normally written
`
`3 = It“ X P
`
`where P is the partial pressure of the gas, and kg, which is called Hemzis
`constant. depends on the gas. the solvent. and the temperature (see
`
`FIGURE 11.10 The hydropho-
`bic tail of a soap or surfactant
`molecule enters the blob of
`
`grease, leaving the hydrophilic
`polar head group on the surface.
`
`(mM)
`Solubility
`
`Pressure (atm)
`
`FIGURE 11.1 1 The variation
`
`of the solubilities of oxygen. ni-
`trogen, and helium with the pres-
`sure. Note that the solubility of
`each gas is doubled when the
`pressure is doubled.
`
`
`
`(a)
`
`(1'))
`
`FIGURE 11.12 The small bub-
`
`bles of air are responsible for the
`“bends." (a) Normal blood vessels;
`(b) catastrophic collapse as bub-
`bles of gas escape from solution
`in the blood plasma.
`
`11.4
`
`rPrage 8 e£& c r
`
`OF
`
`PRESSURE ON GAS
`
`SOLUBILITY
`
`405