William L. Masterton
`
`’
`
`Proiessor of Chemistry
`University of Connecticut
`
`Storrs. Connecticut
`Emil I. Slowinski
`
`Professor of Chemistry
`Macalester College
`St. Paul. Minnesota
`
`Conrad L. Stanitski
`
`Professor of Chemistry
`Randolph-Macon College
`Ashlcmd. Virginia
`
`'
`
`* Chemical
`Principles
`
`SAUNDERS COLLEGE PUBLISHING
`Philadelphia
`
`* SAUNDERS GOLDEN SUNBURST SERIES
`
`Page 1 of 19
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`Astraleneca Exhibit 2026
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`Chemical Principles
`
`ISBN 03-057804-3
`
`©1981 by Saunders College Publishing. Copyright I966, I969, I973. and I977 by W.B. Saunders
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`Page 2 of 19
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`

`
`‘I7
`
`ACIDS AND BASES
`
`In previous chapters we have referred from time to time to water solutions uf
`acids and bases. Such solutions are among the most useful of laboratory reagents. They
`are also common in the home. Vinegar. orange juice, and battery Fluid are, to varying
`degrees. acidic. Oven and drain cleaners, whitewash (like Tom Sawyer used). and
`antacid tablets are basic.
`This chapter examines the properties ofacidic and basic water solutions. We st-(1.1
`with working definitions of the terms acid and base. To begin. we take an acid to be ;.
`substance which. when added to water, produces hydrogen ions (ll+). Bases produce
`/1)’dI'0.\’id(' ions (OH') in water solution.
`All acidic water solutions have certain properties in common. They evolve hydi-Q.
`gen gas. Hz. with zinc or magnesium. They react with compounds containing the C03 -
`ion to form CO, (Fig. 17. I ). They affect the color of certain organic dyes, called indi-
`cators. F‘or example. litmus turns red in acidic solution. These properties of acidic
`solutions are due to the presence of H‘ ions.
`Water solutions of bases also have identifying properties. They feel slippery and
`
`Figure 17.1 The H" ions present in a water solution ol an acid ieacl
`with calcium carbonate to term caihon (lioxirle: CaCO_~.(s)
`I 2 H'(aq)"
`Ca“ (aq) + CO,(g) 1 H90 (tube a). The DH" |0l'l$ present in a solution
`of a ease precipitate Mg(OH), liom a solution containing Mg" i0"35
`Mg”(aq) l 2 OH‘ (aql -4 Mg[0H),(s) (lube h)
`
`Page 3 of 19
`
`

`
`tum litmus blue. In the presence of Mg“ ions. they precipitate magnesium hydroxide,
`Mg(OH)2. The characteristic properties of‘ basic solutions are caused by OH’ ions.
`
`401
`ACIDS I\l‘.l"‘
`B/‘« 1’
`
`17.1 WATER DISSOCIATION; ACIDIC, NEUTRAL, AND
`BASIC SOLUTIONS
`
`‘the acidic and basic properties of aqueous solutions are dependent upon an equi-
`librium that involves the solvent. water. Water, when pure or as a solvent. tends to
`dissociate to some extent into hydrogen ions and hydroxide ions:
`
`H2O(l) T—‘ H"(aq) +- OH‘(aq)
`
`(l7.l)
`
`The forward reaction proceeds only slightly before equilibrium is reached. Only a
`small fraction of the total number of water molecules is dissociated.
`Applying the general rules from Chapter I3 for equilibrium systems. we can write
`the equilibrium constant expression for Reaction l7.l as
`
`K = [H ‘] >< [OH 1
`°
`[H20]
`
`In aqueous solutions, the concentration ot‘H-,0 is very high, typically about 55 M, and
`is essentially constant. Hence the term [H20] can be combined with K‘. to give a new
`constant. K... called the dissociation constant of water:
`
`Kc X [H20] = K... = [H*] X [OH ]= l.0 ><
`
`I0 ”
`
`(17.2)
`
`In the equation for K,.
`lll:.Ol is nlwHy.<. utttilteil
`
`At 25°C. K... is Lil X I()‘”. This small value reflects the slight dissociation of water
`into I1“ and OH“ ions. In any aqueous solution or in pure water. the product of [H*]
`times [()H‘] at 25°C is always 1.0 X l()‘'‘'.
`We can readily calculate [H*] and [OH'j in pure water. From Equation l7.l, we
`see that equal amounts of these two ions form when water dissociates. Applying Equa-
`tion 17.2 to pure water:
`
`[14+] = [on-J; [H*] x [on-1 = [I-l*]'~’ = 1.0 x 10 M
`
`[H*] = l.0 X l0‘7 M = [OH‘l
`
`Any aqueous solution in which [l-1"] equals [OH’] is called a Il(’ll!l'(l/ solution. It has
`a [H‘] equal to 1.0 X l0“’ M at 25°C.
`Ordinarily, the concentrations of H‘ and OH’ in a solution are not equal. Note
`from Equation l7.2 that as the concentration ol'one ion goes up, that ofthe other must
`go down so that the product [H*] x [OH'] stays constant.
`It‘ the concentration of
`one ion is known. then that ofthe other can be calculated using Equation l7.2 (Example
`l7. l l.
`
`Example 17.1
`
`An aqueous solution has a [H*] of 2.0 x IO“ M. What is its [0H“]‘.’
`
`From liquation l7.2:
`
`Solution
`
`[on ]=
`
`Kw = Le x lt)‘“
`[H*J
`2.0 x 10-‘
`
`= 5.0 x lU‘“ M
`
`Page 4 of 19
`
`

`
`402
`CHAPTER 17
`
`
`
`Exercise
`X 10"‘ M.
`
`What is [H*] in a solution in which [OH'] is 2.5 X 10*’ M‘? Answer: 4.0
`
`
`
`An aqueous solution, like that described in Example 17.] where [H*] is grcmm.
`than [OH’], is said to be acidic. An aqueous solution in which [()H’] is greater than
`[l-{*] is ba.\'i(.‘. Therefore,
`
`if [H‘‘] > 1.0 x 10*’ M, [OH“] < 1.0 X 10"’ M. solution is acidic
`
`if [OH‘] > 1.0 X 10” M. [I-l“] < 1.0 x 10*’ M, solution is basic
`
`Table 17.1 indicates some possible combinations of concentrations of these ions. so|u_
`tions 1 through 4 are decreasingly acidic, 5 is neutral and 6 through 9 are progressively
`more basic.
`
`17.2
`
`pH
`
`As we have seen, the acidity or basicity of a solution can be described by giving
`its 11* ion concentration. Sorensen, in 1909, proposed an alternative method of accom-
`plishing this purpose. He suggested using a term called pH, defined as
`
`pH = -Iogto [H*] = log... I/[W]
`
`(17.3;
`
`Thus we have
`
`[W] = 10-4 M; pH = —tog... (104) = —(—4) = 4
`
`[H+] = lo-‘7
`
`= _]0gm (l0—7) =
`
`= 7
`
`[1-1*] = 10“"M; pH = —log..(l0'"’) = —(-10) = 10
`
`A scientist will usually Most aqueous solutions have hydrogen ion concentrations between 1 M and 10"“ M.
`"em '°f*?‘ 3°'#"°'t‘h"" By Equation 17.3, such solutions have pl-Ts lying between 0 and 14. In this case, it is
`ems 0 '5 "
`fa er perhaps more convenient to express acidity in terms of pH rather than [H*]. This
`than its [H‘]
`_
`_
`avoids using small fractions or negative exponents.
`Looking at the pH values in Table 17.1, we see that [H*] and pH are inversely
`
`TABLE 17.1 RELATIONS BETWEEN [H‘], [OH‘], and pH IN
`AQUEOUS SOLUTIONS
`
`SOLUTION
`
`No.1
`
`10"
`
`10"“
`
`Page 5 of 19
`
`10*“
`
`No. 2
`
`No. 3 No.4
`
`No. S
`
`No. 6
`
`10”’
`
`10"’
`
`now
`
`10"“
`
`10''
`
`10"“
`
`10'’
`
`10"
`
`10*
`
`I01“
`
`104'"
`
`to-'
`
`in--'~'
`
`

`
`403
`ACIDS AND
`BASES
`
`j. K
`
`
`
`Lemonjuice
`Wine
`Vinegar
`Tomato juice
`Beer
`Chccse
`
`TABLE 17.2 pH OF SOME COMMON LIQUIDS
`
`4.8-8.4
`Urine. human
`2.2-2.4
`6.3-6.6
`Cow’s milk
`2.8-3.8
`6.5-7.5
`Saliva, human
`3.0
`6.5-8.0
`Drinking water
`3.5
`7.3-7.5
`Blood. human
`4-5
`8.3
`Seawater
`4.8-6.4
` J
`
`related. As [l-1*] decreases from I0" to 10*‘ M, pH increases from 2 to 4. ln general,
`the higher the pH, the less acidic (more basic) the solution. A solution of pH 4 has a lower
`concentration of H* and a higher concentration of OH‘ than does a solution of pH 2.
`Notice also that when pH increases by one unit, [I-1"] decreases by a factor of 10. A
`solution of pH 3 has a hydrogen ion concentration one tenth that of a solution of pH 2
`and ten times that of a solution of pH 4.
`Previously. we used concentration of H* or OH‘ to differentiate acidic. neutral,
`and basic solutions. pH can also be used for this purpose:
`
`if pH < 7.0, solution is acidic
`if pH = 7.0, solution is neutral
`if pH > 7.0, solution is basic
`
`Table l7.2 shows the pH of some common solutions. Example 17.2 shows how Equa-
`tion l7.3 can be used to calculate pH from [I-1*], or vice versa.
`
`Calculate
`Example 17.2
`a. the pH ofa solution in which [H"] = 4.0 X l0‘° M.
`b.
`the [H*] of a beer with a pH of 4.7.
`
`If you have a ‘‘scientific‘’ calculator, you may be able to answer these ques-
`Solution
`lions instantly by pushing the proper button. In case your calculator does not do this,
`we will indicate briefly the procedure involved (see also Appendix 4 for a discussion of
`logarithms).
`
`Calculators are l|'\dQ9d
`"'9'P‘U| with PH Wob-
`‘ems
`
`X l0"’ M: 2.0 X 10”’ M (seawater is basic).
`
`a. pH = —log,.. [H*] = —log,., (4.0 X I0") = -(log,., 4.0 + log... I0")
`= —(0.60 — 3.00) = —(-2.40) = 2.40
`
`b. pH = 4.7 = —log,., [H‘]. This means that [H*] = I0“-', which can be rewritten
`as I0“ ><
`l0‘°. The antilog of 0.3 is 2; i.e.. 10“ = 2; that of 10"‘ is -5. Therefore.
`a pH of 4.7 corresponds to a [H*] of 2 x IO‘5 M.
`:_
`
`Exercise
`
`Using Table 17.2, obtain the [H*] of seawater;
`
`the [OH']. Answer: 5.0
`
`Several experimental methods can be used to determine the pH of an aqueous
`solution. Most often a so-called “pH meter“ is used (Figure 17.2). A more colorful
`(and time-consuming) way to determine pH is to use acid-base indicators. These sub-
`stances undergo a color change within a narrow pH range. perhaps I to 2 pH units
`(sec color plate 14, center of book). By using two or more of the indicators listed in
`Table l7.3, you can estimate quite accurately the pH of a solution. Suppose, for exam-
`ple, that a solution turns red when a drop of phenolphthalein is added. This means
`
`Page 6 of 19
`
`

`
` Figure 17.2 The simplest way to deter-
`
`mine the pH of a solution is to measure
`it with an electrical device called a pH
`meter. Care must be taken in making
`such a measurement; the electrodes are
`easily broken. (Courtesy ol the Arlhur H
`Thomas Co. and Coming Glass.)
`
`'
`
`that its pH is l0 or greater. If you find that this same solution gives a yellow color with
`alizarin yellow, its pH must be 10 or less. Putting these two observations together,
`the solution must have a pH of just about 10.
`"pH paper" is coated with a mixture of indicators. Strips of pH paper are used
`widely to test the pH of biological fluids, soil, ground water, and foods. Depending
`upon the indicators used, a test strip can measure pH over a wide or narrow range,
`
`17.3 STRONG AND WEAK ACIDS
`
`For a species to act as an acid, it must supply H* ions to water. The acid, which
`may be a molecule or ion. contains hydrogen atoms. The H’ ions are formed by the
`dissociation of the acid in water.
`Strong acids dissociate completely in water, forming H’“ ions and anions. A typical
`strong acid is HCI. It undergoes the following reaction upon addition to water:
`HCl(aq) —-> H*'(aq) + Cl'(aq)
`
`(l7.4)
`
`This reaction goes to completion. In a dilute water solution of hydrochloric acid there
`are no HCI molecules, only H* ions and Cl‘ ions. Consider, for example, a 0.10 M solu-
`
`TABLE 17.3 TYPICAL ACID-BASE INDICATORS
`Acm COLOR
`(Lowsn pH)
`
`pl-I lnrsxvm.
`
`BASE Coma
`(HIGHER pH)
`
`19.4
`crmnren 17
`
`Page 7 of 19
`
`INDICATOR
`
`Methyl violet
`Methyl yellow
`Methyl orange
`Methyl red
`Bromthymol blue
`Thymol blue
`Phenolphthalcin
`Alizarin yellow
`
`O\©r~J4=OO\CO
`
`5"?
`
`T‘i"°-B-F-—
`
`
`
`Emcee?—-H6000IIIII7;-5)°?°.°‘...
`
`--
`
`yellow
`red
`red
`red
`yellow
`yellow
`colorless
`yellow
`
`

`
`405
`tion of HCl. prepared by adding 0.10 mol of HCl to water to form one liter of solution.
`The concentration of HCl molecules is virtually zero. The concentrations of H* and Cl‘ we label the how 0 1
`are 0.l0 M. The pH of the solution is [.00. Any way you look at it, the HCl is com~ M HG,’ but perhalm.
`pletely dissociated into ions.
`0.1 M H‘. 0.1 M Ct
`In contrast, a weak acid is only partially dissociated in water. A typical weak acid ‘'‘'°“'‘‘ °‘’ 59"”
`is HF. When hydrogen fluoride is added to water, the following reversible reaction
`occurs:
`
`HF(aq) 3-‘ I-l"(aq) -l— F'(aq)
`
`(17.5)
`
`In a solution of hydrofluoric acid, there are HF molecules as well as H* and F‘ ions.
`In 0. IO M HF, prepared by adding 0. I0 mol of HF to enough water to give one liter of
`solution. more than 90% of the HF molecules remain undissociatcd. The concentra-
`tions of H" and F‘ ions are less than 0.0! M. The pH of the solution is slightly greater
`than 2.
`
`As pointed out in Chapter 16, hydrogen fluoride behaves as a weak electrolyte
`in aqueous solution. A 0.10 M solution of HP has a small electrical conductivity due to
`the I-1* and P" ions present. in contrast, a 0.10 M solution of HCl is a strong electrolyte.
`with a conductivity more than ten times that of 0.10 M HF. In its colligative proper-
`ties, hydrogcn fluoride is intermediate between a nonelectrolyte such as glucose and a
`strong electrolyte such as HCl. The freezing point of 0.10 M HF is —0.2l°C as com-
`pared to —0.l9°C for 0.10 M glucose and —0.37"C for 0.10 M HCl.
`
`The Strong Acids
`
`There are very few strong acids. For our purposes, we need consider only the six
`species listed in Table l7.4. You should learn the names and molecular formulas of
`these six strong acids. They will be referred to again and again, in this and following
`chapters.
`
`o,;_
`
`TABLE 17.4 COMMON STRONG ACIDS
`
`STRONG ACID
`
`MOLECULAR FORMULA
`
`M()lJ€CULAR S'l‘l(U(.TUllE
`
`Hydrochloric acid
`
`Hydrobromic acid
`
`Hydriodic acid
`Nitric acid
`
`Sulfuric acid
`
`0
`
`0
`I
`H—O——S—()—H
`
`H,SO,,. HNO-_,. and HCl
`rank among the most
`important industrial
`chemicals
`
`Perchloric acid
`
`Page 8 of 19
`
`

`
`9.5CHAPTER 17
`
`Notice that all the substances that act as strong acids are molecular species when
`pure. in dilute water solution. they dissociate completely to form a H’" ion and an anion_
`The dissociation reactions; are similar to Reaction [724 for HCI. Examples include
`
`HN0-.i(aq) e> H*(aq) + NOa‘(aq)
`
`l-l2SO4(aq) A 1-l*(aq) + l-lSO.,‘{aq)
`
`(17.5;
`
`(173)
`
`Species Which Act As Weak Acids
`
`A wide variety of solutes behave as weak acids in water. For convenience, they
`can be classified into three categories: molecules, anions, and cations.
`
`MOLECULES CONTAINING AN IDNIZABLE PHOTON. There are literally
`thousands of molecular weak acids. A few of the more common ones are listed in
`Table 17.5. All these species contain a hydrogen atom bonded covalently to a non-
`metal atom in the molecule. When placed in water, the weak acid molecule partially
`dissociates, forming H+ ions and anions at low concentrations. The equations for dis.
`sociation are analogous to Reaction 17.5 for HP. Thus, for hypochlorous acid, HCIO,
`and acetic acid, I-lCg'H3O2, we have
`
`HClO(aq) r—‘ H+(aq'] + ClO‘{aq)
`
`HC2H302(aq) 7-‘ H+(aq) + C2H3O2—[3-Cl)
`
`(17.8)
`
`{[7-9)
`
`TABLE 17.5 SOME COMMON MOLECULAR WEAK ACIDS
`MOLECULAR
`Cone. H* IN
`STRUCI‘URE
`0.10 M SOLN.
`
`MOLECULAR
`Foamum
`
`pH
`(0.10 M Some.)
`
`WEAK
`Acm
`
`Phosphoric acid
`
`H3PO4
`
`Hydrofluoric acid
`Nitrous acid
`
`HF
`HNO2
`
`Acetic acid
`
`HC2H:,O.,,
`
`Carbonic acid
`
`HZCOS
`
`Hydrogen sulfide
`
`HZS
`
`Hypochlorous acid
`
`l-ICIO
`
`Hydrogen cyanide
`
`HCN
`
`0.024 M
`
`0.0080 M
`0.0065 M
`
`0.0013 M
`
`0.0002} M
`
`0.00010 M
`
`0.000056 M
`
`0.0000063 M
`
`Page 9 of 19
`
`

`
`The properties of 0.10 M solutions of HCIO or I-IC2H302 resemble those of 0.10 M
`HF. All three species are weak electrolytes with small electrical conductivities. Each
`has a pH greater than 1 (Table 17.5).
`Acid-base indicators, referred to previously, are weak acids. We might represent
`the dissociation of such an indicator by the equation:
`
`407
`ACIDS AW
`BASES
`
`Le Chélelior"-3 Princi-
`ple applies nicely horn
`
`Hlntaq) :—‘ H+(aq) + In”(aq)
`
`The formula Hln stands for the weak acid molecule, which has a color different from
`that ofihe in‘ onion. In the case of bromthymol blue (Table 17.3), the weak acid mole-
`cule is colored yellow while the anion is blue.
`The position ofthe above equilibrium is sensitive to the concentration of H’' ions.
`If [Ht] is “high,” the equilibrium lies far to the left and the principal species present
`is the Hln molecule. We see its color (yellow with bromthymol blue) when we look at
`a solution containing a few drops of the indicator. When [H*] is “low,” the equilib-
`rium shifts to the right, forming In‘ ions. The solution takes on the color of In’ (blue
`with brornthymol blue).
`
`ANIONS CONTAINING AN IONIZABLE PROTON. Let us refer back for a
`moment to Equation 17.’? for the dissociation of I-I,SO.,. Notice that the anion formed,
`HSO.,‘, contains a hydrogen atom. In water, the H80; ion undergoes further dis-
`sociation. producing a H’' ion and a SO42’ ion. The dissociation reaction is reversible,
`so we classify I-130,,‘ as a weak acid:
`
`HSO.,‘(aq) = H"(aq) + S0.,“‘{aq)
`
`(17.10)
`
`In practice, very few anions give acidic solutions when added to water. The only other
`anion of this type that we need be concerned with is the H,PO_( ion (see Example
`17.3, p. 408).
`
`CATIONS. The ammonium ion, NH9‘, behaves as a weak acid in water because
`of the following reversible reaction:
`
`NH4*(aq) : H*(aq) + NH,-,.(aq)
`
`(l7.ll}
`
`The products are an ammonia molecule, NH,,, and a H’' ion which makes the solution
`acidic. Note that the behavior of the NH.,'* ion in water is very similar to that of the
`HF molecule (Eq. 17.5) or the I-130,‘ ion (Eq. 17.10). All three species contain hydro-
`gen atoms which are converted to 1-I+ ions when the weak acid dissociates.
`
`You may be surprised to learn that most metal cations, except those of Groups
`1 and 2, are weak acids. At first. it is not at all obvious how a cation such as Zn“ can
`make a water solution acidic. To understand how this is possible, we must realize
`that this cation and others like it are hydrated in water solution. When ZnCl2 or
`Zn(NO3)2 is added to water, the cation formed is Zn(H,O).,2+. Here, a Zn“ ion is bonded
`to four water molecules. This complex cation is slightly dissociated in water, accord-
`ing to the following reaction:
`
`Zn(Ha0}.i‘”‘(‘dC1) “—‘ Hllaql + ZI1(H20)s(0Hl*"(aCI)
`
`(17-12}
`
`The H‘‘ ion, which makes the solution acidic, comes from the ionization of one of the
`H20 molecules bonded to Zn“. The OI-I‘ ion formed at the same time remains bonded Free H‘ and OH ions
`to Zn“ and so does not directly affect the pH ofthe solution. (It does, however, affect "e‘”'”“"@ “H
`the charge of the complex cation, reducing it from +2 to +1.)
`
`63
`
`Page 10 of 19
`
`

`
`£3CHAPTER 1?
`
`Example 17.3
`water:
`a. H28
`
`Write dissociation equations for each of the following week acids in
`
`b. H,i=o,,-
`
`c. A1(H,o;,3+
`
`In each case, :1 proton (l-1+ ion) is produced to make the solution acidic. The
`Solution
`other species formed is the residue from the weak acid after removal of‘ the proton. A1]
`the reactions go to equilibrium, indicated by a double arrow.
`
`21.
`
`I-l.;S(aq') 7-‘ l-I"(aq) + HS’(aq): Compare Eqs. 17.5.
`
`l7.8, 17.9
`
`b. H2P04‘(aq) : H“”[aq) + HPO,,"(aq); compare Eq. 17.10
`
`l7.l2
`c. Al(H2O)3“[aq] = l-l*[aq} + Al(HgO),,(OH)2*(aCl)1compare Eq.
` ajj
`
`l-IS‘, HPO3’, Al(H20).~.(0i-1)“. Each of these
`Consider the following ions:
`Exercise
`can dissociate in water to produce a H" ion and another species. In each case, give the
`formula of the "other species" formed. Answer: S2‘; P04 ‘; Al(H2O),(OH)g*.
`
`It may seem strange.
`but
`in 0.1 M NaDH,
`conc. NaOH = 0!
`
`17.4 STRONG AND WEAK BASES
`
`For a species to act as a base, it must supply OH‘ ions to water. Bases, like acids,
`are classified as “strong” or “weak.” As with acids, there are only a few strong bases
`but a great many weak bases.
`
`Strong Bases
`
`A strong base dissociates completely in water to release 0H* ions. Sodium hy-
`droxide, NaOH, is the most common strong base. It dissolves readily in water to give
`a solution containing Na* and OH“ ions:
`
`NaOI-1(5) —> Na"(aqJ + OH‘(aq)
`
`(17.13)
`
`As with all strong bases, this reaction goes to completion. In a 0.10 M NaOH solution.
`prepared by dissolving 0.10 mol NaOH in enough water to give one liter of solution.
`the concentration of imdissociated NaOI-I is virtually zero. In this solution, the con-
`centrations of Na* and O1-1* are 0.10 M. The pH is 13.00.
`Strong bases are limited to:
`I. The hydroxides of the Group 1 metals (LiOH, NaOH, KOH, RbOH, CsOH)-
`2. The hydroxides of the Group 2 metals, Mg(OH)g, Ca(Ol-1),, S1*(0H)2, Ba(0H)2-
`With these compounds, two moles of OH‘ are produced for every mole of solid that
`dissociates.
`
`Cfl(0H}g(S] —> Ca"’+(aq) + 2 OI-I‘(aq}
`
`(17-14)
`
`Of the strong bases, only NaOH and KO!-I are commonly used in the chemistry
`laboratory. All compounds oflithium, rubidium, and cesium, including the hydroxides.
`are expensive. The Group 2 hydroxides, particularly Mg(OI-1);. have limited 501"‘
`bilities. Calcium hydroxide is often used in industry when a strong base is neflded
`and high solubility is not critical.
`
`Page 11 of 19
`
`

`
`Weak Bases
`
`Weak bases furnish OH‘ ions by a reversible reaction involving a water molecule.
`Perhaps the most common weak base is ammonia, NH3. It reacts with water as follows:
`
`N]-[,.,(aq) + H20 7-" NI-.l.,'*(aq) + OH‘[aq)
`
`(l7. l5)
`
`409
`ACIDS AND
`
`“SE3
`
`The forward reaction occurs to only a slight extent. In 21 (H0 M solution of NH3, pre-
`pared by adding 0.10 mol of NH3 to enough water to form a liter of solution, nearly
`99% of the NH; molecules remain unreacted. The concentration of NH3 is about
`0.099 M. In contrast, the concentrations ofNH.,+ and OH‘ are only about 0.001 M. The W”,-W mkes H G,,,,,._.
`pH of the solution is about 1 I.
`basic
`Most weak bases are anions. A typical exampte is the fluoride ion, F‘. It under-
`goes the following reversible reaction with water:
`
`F‘(aq) + H20 r—" I-lF(aq) + Ol-l‘(aq)
`
`(17.16)
`
`As with NI-I3, the forward reaction occurs only to a slight extent. In a solution in which
`the F‘ concentration is 0.10 M, the concentration of OH‘ at equilibrium is only about
`10‘“ M. As small as this is, it is enough to make the solution basic, with a pH of about
`8. A reaction similar to l7.l6 occurs with the anion of any weak acid. With the acetate
`ion, C_..;H302‘, the reaction is
`
`CgH;,O2"(aq} 1- [-120 «‘i HC,H3O2(aq} + OH‘(aq)
`
`(I7.I7)
`
`Looking at Equations 17.15 through 17.17, we see that they resemble each other in 3 very real sense a
`very closely. In each case, a weak base {N}-I3, F‘, C,H3O2‘) picks up a proton (H+ ion) We?“ “sf °|j'".”“"e5
`from a water molecule. The products are:
`Wm H’O W M5
`
`—-H. Weak acid (NH4+, HF, HC2H3Og).
`
`-an OH‘ ion, which makes the solution basic.
`
`The general reaction is
`
`weak base(aq) + H20 : weak acid(aq) + OH‘(aq)
`
`Since the reaction does not go to completion, relatively few OH‘ ions are formed.
`This is why we refer to species such as NH3, F‘, and C2H3O2‘ as weak bases.
`
`Example 17.4
`water:
`a. CN‘
`
`Write equations for the reactions of the following weak bases with
`
`b. (30,2-
`
`The equations are cntirely similar to 17.16 and 1?. I7. In each case. the weak
`Salution
`base picks up a proton (l-1+ ion) from an H30 molecule. This produces an OH‘ ion, which
`makes the solution basic.
`
`a. CN‘(-aq) + l"-I20 :—* t-lCN(aq) + Ol-l‘(aq)
`
`b. CO;,“(aq) + H20 T—‘ HCOg‘[aq) + OH'(aq)
`
`Write the formula for the weak acid formed by the reaction of each of the
`Exercise
`following weak bases with water: SO32‘; Cl-l,,N1-lg. Answer: HSO3‘; CH=,NH,," {analogous
`to NH4"].
`
`63
`
`Page 12 of 19
`
`

`
`1.1.”
`CHAPTER 1?
`
`17.5 ACID-BASE PROPERTIES OF SALT SOLUTIONS
`
`After completing Sections l7.3 and 17.4, you should be able to predict con-e
`Ctly
`that an aqueous solution of H] or H2804 is acidic while a solution of NaOH or
`NHB
`is basic. Solutions of NaNO2 or NH.,I might be more difficult for you to classify, These
`two compounds, and many others, such as NaCl, Z.n(N'O,,)2. and CLISO4 are s-a1¢S_
`A salt is an ionic compound containing a cation other than H", and an anion other than
`OH‘ or 02".
`
`In dilute water solution, a salt is completely dissociated. Hence, the acid-base
`properties of its solution are determined by those of the ions present. A particular ign
`may be neutral, having no effect on the [Hi] or [OH“] of water. Other ions are acidic.
`they increase [H"] to greater than Li) X l0‘7 M. Still other ions are basic, since they
`raise [OH*] beyond 1.0 X IO" M. Table 17.6 summarizes the acid-base properties of
`ions commonly present in aqueous solutions.
`
`Neutral ions
`
`A neutral ion does not react with water to produce H’' or OH‘ ions. Hence, it does
`not affect the pH. There are relatively few neutral ions. We see from Table 17.6 that:
`—t}2e neutral unions are those derived from strtmg acids.
`——the neutral ccflions are those derived firom strong bases.
`A typical neutral anion is the chloride ion, produced by the dissociation of hydm-
`chloric acid:
`
`HCl(aq] —> H+(aq} + Cl‘{aq)
`
`Explain In your own
`words why the NO-,
`ion is neutral
`
`You will recall trom Section 17.3 that HCl is completely dissociated in water; it is a
`strong acid. This means that there is no tendency for the reverse of the above reaction
`to occur. Chloride ions. wherever they come from, do not combine with 11* ions. In
`particular, Cl’ ions from a salt such as NaCl do not pick up Hi‘ ions from water. As a
`result, CI‘ ions. and the other neutral anions listed in Table 17.6 do not change the
`[H*] or pH of water.
`A similar argument applies to cations like Na+, produced by the dissociation of
`strong bases such as Na0H. Dissociation is complete:
`
`NaOH(s) —> Na+(aq} + O]-{‘(aq)
`
`TABLE 17.6 ACID-BASE PROPERTIES OF SOME COMMON IONS IN
`WATER SOLUTION
`
`Neurnar.
`
`BASIC
`
`C2 H.-:02 _
`F-
`C032-
`3”’
`p0_I rt-
`
`Ci‘
`or
`1-
`
`No,-
`C10,.-
`son-
`
`Mgfli
`Ca“
`B32!
`
`HSO; , H2PO{
`
`Aliii
`N H."
`transition metal
`ions
`
`Page 13 of 19
`
`

`
`Hence, there is no tendency for the reverse reaction to occur. That is, Na" ions, regarcl—
`less of their source, do not combine with OH‘ ions in watet'. As a result, Na" and the
`other cations of the Group I and 2 metals are neutral.
`
`411
`ACIDS AWE
`BASES
`
`Basic Anions
`
`Recall from Section I7.4 that any anion derived from a weak acid acts as a weak
`base in water solution. There is a small army of such anions. Those listed in Table 17.6
`are typical examples. In contrast. there are no common basic cations.
`
`Acidic tons
`
`These include:
`
`——all cations except those ofthe metals in Groups 1 and 2. Typical examples of acidic
`cations are listed in Table 17.6.
`
`—the HS04‘ and I-IZPO4’ ions (see discussion on p. 407).
`
`The principlesjust discussed can be applied to predict the acid-base properties of
`salt solutions. Here, two ions are present, a cation and an anion. To decide whether
`the solution is acidic, basic, or neutral, we have to consider the properties of both
`ions. The way this is done is shown in Example 17.5.
`
`Using Table 17.6, describe each of the following 0.! M solutions as
`Example 17.5
`acidic, basic, or neutral:
`a.
`b.
`
`Cu Na3P0,,
`
`d.
`
`solution
`
`Solute
`ll
`
`'
`l<.ClO
`. NH,I
`. Na;,PO,
`. Zn(N03),,
`
`Cation
`
`K+ (neutral)
`NH: (acidic)
`Nat (neutral)
`Zn“ (acidic)
`
`Anion
`4
`C10 ‘ (neutral)
`1‘ (neutral)
`PO? (basic)
`N03‘ (neutral)
`
`Aqueous Solution
`
`neutral
`acidic
`basic
`acidic
`
`acid '7 nemraiflacidic
`basis I neutral
`—- basic
`
`Experiment confirms these predictions; see color plate 15 (center of book).
`m
`
`Write an equation to explain why 0.] M Na,.P0.. is basic. Answer: P0,,-“‘(aq)
`Exercise
`+ H,o .—_‘ HPO,.‘*”(aq) + on-(aq).
`
`Anions containing ionizable protons can show both acidic and basic character.
`The hydrogen carbonate ion, HCO3‘, is typical. It can dissociate via the reaction
`
`HCO,,‘(aq) '1: H*{aq) + CO32‘(aq}
`
`(17.18)
`
`This reaction, by itself, would tend to make the solution acidic. However, since the
`HCO3’ ion is the anion ofa weak acid, HECO3, it can also undergo the following reac-
`tion:
`
`HCO5-“(aq) + H20 ?—“ H2CO3(aq) + OH’(aq)
`
`(17.19)
`
`This reaction, which tends to make the solution basic, occurs to a greater extent than
`
`Page 14 of 19
`
`

`
`9.‘?
`nnnptes 1?
`
`17.13. This explains why a water solution of Nal-{CO3 (“bicarbonate of soda") is
`slightly basic, with a pH above 3.
`Most anions of this type behave like the HCO3” ion. Solution containing H33
`I-IP03“, and H,._BO;,‘ are all slightly basic. As previously mentioned, the only two acidic
`anions are H530; and I-l.;P0.,‘. With these two ions, a reaction similar to l7.l8 pl-e_
`dominates.
`
`17.6 ACID-BASE REACTIONS
`
`When an acidic water solution is mixed with a solution containing a base. a reac.
`tion occurs. The nature of this reaction, and the equation we write for it. depend upon
`whether the acid and base are strong or weak. In this section, we will look at several
`different types of acid-base reactions. All those considered have large equilibrium
`constants and. for all practical purposes, go to compieriori.
`
`Reactions of Strong Acids with Strong Bases
`
`Consider what happens when we add a solution of ‘:1 strong acid like HNO3 to a
`solution of a strong base such as NaOH. Both HNO3 and NaOH are completely dig-
`sociated into ions:
`
`solution of I-IND“: I-l'*. NO; ions
`solution ot”NaOH: Na*, OH‘ ions
`
`The ac-id—base reaction that occurs involves the H‘‘ ion of the l-INO3 solution and the
`OH‘ ion of the NaOH solution. The equation for the reaction is simply
`
`I-l”(aql + OH’taq}—-> H20
`
`(17.20)
`
`This is the net reaction that occurs when any strong acid reacts with (my .rrron.g bcrse.
`Note that we do not include in the equation “spectator ions“ such as Nat" or N03”,
`which do not take part in the reaction.
`The reaction between a strong acid and a strong base is called neutralization.
`Ifjust enough base is added to react with all the acid, the resulting solution is neutral.
`For example, if equal volumes of 0.10 M HNO3 and 0.10 M NaOH are used, the final
`solution will contain only Na* and N03" ions. Since both these ions are neutral, the
`solution will have a pH of 7.
`Robert Boyle of gas law fame was probably the first to recognize that when an acid reacts
`with a base, they neutralize each other’s properties. Nearly I50 years passed before the nature
`of this reaction was determined. The delay came because Lavoisier (1787) insisted that oxyge-_“
`was the fundamental component of all acids. In [Si l, Humphry Davy showed that hydro-:hi0l‘|C
`acid contained no oxygen. Shortly thereafter (I814), Ciay~Lussac concluded that it is the hyd1‘0'
`gen in acids that neutralizes bases.
`
`Reactions of Weak Acids with Strong Bases
`When a strong base such as NaOl-l is added to a weak acid, a reaction similar in
`many ways to 17.20 occurs. The equation for the reaction, however, is diffe1‘e11t- T“
`see why this is the case. let us consider the nature of the principal species present
`
`Page 15 of 19
`
`

`
`Lactic acid is found in sour milk (L. lacris. milk) and sore muscles.
`Example 13.4
`Its molecular formula is HC3H503 (MM 2 90.1"). To save space. "we will abbreviate this
`as HLac. A solution is prepared by adding 9.01 g (0.100 mol} of lactic acid. HLac, to
`enough water to give a liter of solution. In this solution. [I-["] is measured to be 3.6? X 10‘3
`M. Calculate 11,, for lactic acid.
`
`Solution
`
`For lactic acid, we have
`
`HLac(aq) i H"(aq) + l_.ac“[aC1); Kn =
`
`To calculate K", we need the equilibrium concentrations of 1-1+, Lac‘, and unriissot-icrred
`HLac. The [Ht] has been measured as 3.67 X 10*‘ M. From the dissociation equation
`we note that one mole of lactate ion. Lac‘, is produced with every mole of H‘' ion. There-
`fore. in the acid solution,
`
`[H*'] = [Lac-1 = 3.67 x we M
`
`To obtain the equilibrium concentration of HLac, we note that a mole of I-Il..ac must be
`consumed for every mole of H'* produced. Therefore. by subtracting [1-1*] from the original
`concentration of HLac. we obtain [HLac] at equilibrium:
`
`@C
`
`HAPTER 18
`
`The H‘ comes from
`the HLaC
`
`[1-lLac] = orig. conc. HLac ~ [I-l"]
`
`= 0.100 M — 0.00367 M = 0.096 M
`
`Substituting in the expression for K“:
`
`K3 : [H+[]HxL£IC.]ac—] _{3.67 x ions) 39367 X 10,3] _ M X 10“
`
`
`A 0.100 M HF solution has a [I-1*] of 8.0 X l0“’ M. Calculate the K,, of HF.
`Exercise
`Answer: 7.0 X 10”‘.
`
`18.3 SOME AF’PLlCATlONS OF Ka
`
`The dissociation constant of a weak acid can be used for several purposes. In this
`section, we will consider two of its more important uses:
`1. The determination of [I-1+] in a buffer solution. where we know or c.an cal~
`culatc both the concentration of the weak acid, HB, and that of its conjugate base, B‘.
`2. The determination of [H’'] in a solution prepared by dissolving the weak acid
`HB in water.
`
`Buffers; pH Control‘ Systems
`
`Some solutions have a pH which is resistant to change upon addition of 11* OT
`OH’ ions. Such solutions are called buffers. Blood is an example ofa buffer. Suppose
`we add I cm3 of 10 M I-1C1 or NaOH (0.010 mol 1-1*‘ or 0H’ ions) to a liter of