Richard A. Pacer Purdue University Fort Wayne Campus Fort Wayne, Indiana 46805
I
I
Conjugate Acid-Base and Redox Theory
Students in general chemistry courses are usually exposed to oxidation-reduction shortly after studying acids and bases. The Bronsted-Lowry concept is presented in treatment of the latter and the student acquires some feeling for relative acidlhase strength in terms of transfer of a proton between members of a conjugate acid-base pair. Unfortunately, this approach is usually not carried over when redox reactions are studied, although the analogy of oxidation-reduction reactions to acid-base reactions was first presented by Hazlehurstl as far back as 1940. While some general chemistry texts, such as Sisler, VanderWerf, and DavidsonZ do present the conjugate reductant-oxidant concept, usage of this terminology is not very widespread. Consideration of relative oxidizing/reducing strength would be greatly facilitated if this topic were presented in terms of transfer of one or more electrons between memhers of a conjugate redox pair. The approaches used to illustrate relative acid-base strength can be paralleled quite well when relative oxidizing-reducing strength is considered, as will be shown. Tables of Relative Acid/Base and Oxidant/Reductant Strength Relative acid-base strength is usually introduced by listing, in tahle form, several acids and their conjugate acid bases in order of decreasing strenath . of the con.iuaate (Table 1). The stronger an acid, the greater its tendency to lose a proton and transform to its conjugate hase. If an acid is strong, the position of equilibrium lies very extensively to the right; this indicates that the conjugate hase has little tendency to accept a proton and form its conjugate acid. Thus the stronger an acid, the weaker its conjugate hase and vice-versa. Absolute strength is difficult to measure, and it is customary to use a second acid-base couple, H3O+/HzO, as a point of reference. Thus ~
describes the transfer of a proton between conjugate acidhase pairs. The magnitude of the equilibrium constant for the reaction indicates the strength of the acid in question, relative to HsO+
or, for dilute aqueous solutions
Looking a t the same couple from the point of view of the conjugate base H,PO,-
+ H20 ZZ H,PO, + OH-
KD indicates the strength of the conjugate hase relative to OH- ions. The product of Ka and Kb is a constant, Kw, 178
/ Journalof Chemical Education
Table 1. Acids and Bases Conjugate Base
Conjugate Acid
HClO. inneasing strength of eonjugate acid
a
* =
HCI H30'
HtPOl H2P01-
-s
HPO? ~ncrea%es) H20
= =
NHs
+ CI0.+ CI-
H*
H*
+
H*
;H:+
H,0
H"0'HPOP
+ PolJ+ OH+ NHI-
H*
H*
increasing strength "1
increaaesl
Table 2. Oxidants and Reductants -
EO,,, -3046V -0.37 -0.136 0.m C0.361 tO.987 +0.5355 t1.842
iceductant Conjugsb&dunant Li
Ti**
increasing Sn strengthof HI eonjugate Va+
;\-
+
Cd*
3
e
Conjugate Oxidant c-
e2~
2 e 40 * ra 2r2 e a
a
c-
t Li*
+ + + + + +
Siix
1
t3.045 V t0.37
Tia* SnZ* increasing W.136 strengthof om 2Ht VO'* + 2 H * conjupate -0.361 Pd** oxidant -0.987 -0.5'865 Cola
-1.862
equal to the product of the molar concentrations of the conjugate acid and hase forms of the reference, Hz0.
Thus the stronger one memher of an acid-base couple is, the weaker the other memher is. Applying the same approach to the relative strengths of oxidants and reductants, electron transfer between memM+ + e-, is compared to bers of a redox couple, M transfer between the reference couple, H2/H+
--
Thus a tahle of oxidants and reductants may he prepared, indicating relative strength in terms of tendency to counle as accent or donate an electron. with the H?/H+ -, reference (Table 2). The reference counle.. (Hz/H+), is arhitrarilv assianed a . -, standard potential oi 0.000 . . . volts. Species with a *eater tendency to lose electrons than Hz have positive oxidation potentials, which become increasingly positive as reducing strength (tendency to lose electrons) increases. Similarly, species with a greater tendency to gain electrons. than H+ have positive reduction potentials, increasingly positive as oxidizing strength (tendency to gain electrons) increases. The stronger a reductant, the greater its tendency to lose electrons; consequently, its conjugate oxidant has little tendency to gain electrons. Thus, the stronger one memher of an acid-base or redox couple is, the weaker is its conjugate. Mathematically, this was indicated for acids and bases by the fact that the product of the acidic and basic ionization constants for members of a conjugate acid-base pair is a constant K.K,
=
constant = K . (for water as solvent)
'Hazlehurst, J. CHEM. EDUC. 17,466 (1940). 2Sisler, VanderWerf, and Davidson, "College Ch&istry," Ed.) MacMillan, 1967, pp 447-8.
(3rd
Thus the position of equilihrium lies principally to the left and we conclude that Pd will not liberate Hz from acids.
For oxidants and reductants, this inverse strength relationship is indicated by the fact that E", + EO,*n = 0 for the members of a redox counle. That is. the sum of the standard oxidation and reduction potentials for the reduced and oxldized forms.. resnectivelv. . - , of a redox counle (which indicates reducing and oxidizing strength, resp.) must equal zero. The thermodynamic relationship of the two systems can be shown as follows Eom
=
Leveling Effect
Just as the solvent plays a role in controlling acid or base strength, it similarly places a limit on the strength of the oxidant or reductant. In aqueous solutions, H20 stands ready to accept or donate a proton to level acids and bases to the strength of H 3 0 + and OH-. It similarly stands ready to undergo oxidation or reduction, and therebv level the streneth of the oxidant or reductant which may be present. Thus just as the NH2- ion is leveled to the strength of the conjigate hase of HZO,OHI - ion
RT nF InK,,
-
RT lnK,d E0,d = -
nF
NH,-
=
+
k.
g ( l n ~ , , In-
""
F Z
NH,
+ OH-
Na metal, when placed in water is similarly leveled in reducing strength to that of HZ .
=
0
Therefore, as reducing strength (tendency to undergo oxidation and thereby reduce something else) increases, oxidizing strength (tendency to undergo reduction and thereby oxidize something else) decreases. For example EDlr* = +1.842 V Co2+ Co3+ + e- E", = -1.842 V
+
+ H,O
An example of leveling to the strength of Oz is illustrated by the fact that Ag (11) is too strong an oxidant to exist in aqueous solution
+
E", EDlr" = (-1.842 V) (+1.842 V) = 0 The potentials show that Co3+ is a much better oxidant than H + , since it has a much greater tendency to gain an electron than H+, as indicated by its large, positive reduction potential. CoZ+, on the other hand, is a much poorer reductant than Hz, since its tendency to lose electrons, as indicated by its oxidation potential, is a large, negative value.
One important difference should he pointed out, however. Whereas reactions involving transfer of a proton are quite rapid, reactions involving electron transfer are often sluggish (i.e., show a higher kinetic stability than might be expected). Thus very strong oxidants and reductants can exist in aqueous solution for some time, without being leveled to Hz and 02, respectively. For example, M n O c ion can exist in aqueous solution as a stable species, even though it is capable of oxidizing HzO to 02, assuming unit activity for all species.
Position of Equilibrium
In an acid-base reaction, which consists of transfer of a proton from one acid-base couple to another, the position of equilibrium favors the formation of the weaker acid and weaker hase. Thus, the position of equilibrium lies to the right in the following reaction eoniugate pair
MnO,-
+ 8 ~ '+ 5e0,
Mn"
+ 4 ~ '+ 4e-
+ 4H,O
E",,n = 2H20 EO,,,
=
1
+ 1.51 V + 1.221 V
Autoprotolysis versus Autoredox
stronger acid
stranger base
conjugate pair
weaker base
weaker acid
and to the left in the following coniuzate pair I NH,
+
I
OH-
' NH,.
+
H,O
A second important difference between acid/ base and oxidant/reductant behavior concerns autoprotolysis and autoredox reactions. In autopmtolysis reactions, K,, is always small. Thus, in the reactions 2HP0,'2H20
ZZ
PO,'H30C
+ H,PO,-
+ OH-
conjugate pair weaker acid
weaker base stronger base stronger acid The same rule applies to redox reactions. Thus, in answer to the question, "Will palladium metal liberate hydrogen from acids?", we can write the redox reaction and look at the position of equilibrium conjugate pair i
weaker reduetant
weaker oxidant
the position of equilibrium lies to the left. However, with autoredox reactions, the position of equilibrium may lie either to the left or right, (K,, small or large) as shown below. to the left:
I
stronger oxidant
stronger reductant
to the right:
conjugate pair Volume 50,Number 3, March 7973
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179
Conclusion
A clear understanding of conjugate acid-base theory will be of great assistance in helping the student understand oxidation-reduction. However, while many similarities exist, there are valid differences which should be pointed
180
/ Journal of Chemical Edocafion
out in any comparison between proton-transfer and electron-transfer reactions. The author is grateful for the reviewer's comments, particularly the derivation of E 4 x EOred = 0 for members of a conjugate redox pair.
+
