Acids and Base$ by Doris Koll Illinois Central Colleg
"Aluminum chloride is an aeid? I thought it was a salt." When a student has learned to recognize acids as compounds that contain ionizable hydrogen atoms, he may he perplexed to discover that some acids contain no hydrogen at all. The following equation shows a neutralization reaction of the usual type 3HCI + A1(OH13 AICI, + 3H20 (aeid) (base) (salt) (water) Most first-year chemistry students would have no trouble identifying this as an acid-base reaction, even without the labels. The acid provides H+ ions, the hase ~rovidesOH- ions, and they react form water pl& a salt. o n the other hand, it is doubtful that manv. beeinnine students would also recognize this equation as an acid-base reaction
-
-
AlCb + C1- A1CI6Yet the AICls in this reaction is acting as an acid, in the broader sense of the term (Lewis acid-base theory), and chloride ion is acting as a base. The fact that aluminum chloride can be hoth a salt and an acid should not be difficult to accept. After all, sodium bicarbonate is a salt that is hoth an acid and a base. (It is a weak acid because the hydrogen in NaHCOs is very slightly ionized in water solution, and it is a hase because i t can neutralize acids and causes litmus indicator to turn blue.) The acid, hase, and salt classifications are not mutuallv exclusive. There is really nothing wrong with dekning an acid in terms of hvdroeen ions. so lone as we are interested only in reactions tbaiocc;r in water solukon; hut sometimes thesituation demands that we broaden our definitions. Acids and bases (or alkalies) have long been known as chemical opposites capable of counteracting each other, hut just what constitutes an acid or a base has been a subject of controversy and changing opinion over the years. Earliest Ideas About Acids and Bases Vinegar wns the unly ncid known to the ancient Egyptians, Creeks. and Romans. It was made lw air oxidation of fermented fruit juice (wine), and the word acid literally meant "sour." Among the alkalies known to the ancients was potash (potassium carbonate) obtained from wood ashes (alkali comes from two words meaning "plant ashes"), soda (sodium carbonate) made by evaporation of alkaline waters, and lime (calcium oxide) made by roasting seashells. Caustic soda and caustic potash (sodium &d potassium hydroxides) were made by the action of lime on soda and potash. Later, during the middle ages, the alchemists learned to makeaqua fortis (nitric aeid), aqua regia ( a nitric-hydrochloric acid mixture), and oil of vitriol (sulfuric acid). ~ u r i o ~ t mid-160d's he Johann Rudolph Glauber did much experimentation with acids and alkalies. He recognized that salts were made up of two parts, one from an acid and one from a metal or its oxide (an alkali). and that acids differed in strength. In the followingdescription he likened neutralization to a battle
part has overcome and killed the other, neither a fiery liquor nor spiritus ocidus can be found in their dead bodies, but the same ha been made. as hoth were before and from which thev were derivec namely ordinary saltpeter (KNO,). Glauber is best known for his discoverv that hvdrochloric aci, rould In! made by reactinr sulfuric ncid wirh common sal (sndium chloride). Todav this isstill the easirst was to mak HCI. The residue from the reaction (sodium sulfate) becam known as "Glauber's salt." About the same time Otto Tachenius and Francois Sylvin tried to simplify the chemistry of life processes by reducin all chemical interactions within the living organism to acid hase reactions. They made much of the antagonism betweei these chemical oo~osites and the effervescence that occurrel .. when acids were poured onto materink such as cdrbonate! I achenius evrntuallv hecame rowinced rhnt all suhitanre were either acids or alkalies, and argued that all reactions wer acid-alkali neutralizations. Robert Boyle rejected the acid-alkali theory of chemistr) although he did recognize acids and alkalies as importan classes of substances. Boyle (1663) noted that acids, in addi tion to their sour taste, had exceptional solvent power, t h ability to color certain blue vegetable dyes red, and a precip itating action on dissolved sulfur. Alkalies, on the other hanc had a slippery feel and detergent properties, the ability t dissolve oils and sulfur, and the capacity to counteract acid and destroy their properties. Boyle's tests showed that som substances were neutral and did not classify either as acid or alkalies. Probably the earliest attempt at explaining the mechanisr of the neutralization reaction was that of Nicholas Lemer (1675), who described acids as having sharp spiky atoms which nroduced a pricking sensation on the skin, and alkalie as being made upbf rouid particles, which made them fet slippery or soapy. When acids and bases were mixed, he pic tured the sharp needles of the acids as penetrating the porou alkali globules, thus producing salts, which were neithe stinging nor slippery to the touch. It was quite an imaginativ - 3
Doris Kolb Illinois Central College East Peoria. Illinois 61635
"Acids and Bases" is part of a series of substantive reviews of chemical principles taught first in high school chemistry courses. Dr. Kolh received a BS degree from the University of Louisville and both MS and PhD degrees from The Ohio State University. She hes been employed as a chemist at the Standard Oil Company and as a television lecturer in a series " S ~ o t l i ~an h t Research." She has sewed on the staffs of Corning Community College and Bradley University. Since 1967, she has been Professor of Chemistry at Illinois Central College.
Volume 55, Number 7, July 1978 / 45
theory for its dav, and it was simple enough for anyone to understand. Lavolsier's Oxygen Theory
When Antoine Lavoisier named the gaseous element oxygen (1777), he took the name from two Greek words that meant "acid former." When sulfur or phosphorus was burned in oxygen, the products dissolved in water to form acids, so he concluded that oxvaen was the "acidifvina orinciple" in acid materials. His idea'that oxygen was the eiementcommon to all acids was a c c e ~ t e dby manv well into the nineteenth century. In contradiction to Lavoisier's oxygen theory of acids, Claude Louis Berthollet (1789) showed that prussic acid (HCN) did not contain oxygen; hut it had such weak acid properties that most people concluded simply that it was not a true acid. Humphry Davy proved Lavoisier's error more convincingly with muriatic acid (HCI), a very strong acid helieved to contain the oxide of the element "murium." Davy showed that the acid contained only hydrogen and one other element, which he called chlorine. Muriatic acid contained no oxygen; therefore, the "acidifying principle" could not he oxygen. Davy suggested that i t might be hydrogen, hut there were many hydrogen compounds that were not acids. Lieblg's Hydrogen Theory
In the early years of the 19th century the idea of polyhasic acids was unknown. When Thomas Graham (1833) studied the acids of phosphorus, he decided that ortho-, meta-, and pyro-phosphoric acids (H3P04, HPOa, and HaP207) were different because of the varying amounts of water they contained, hut he also noted that they varied in the number of units of base they could neutralize. Shortly thereafter Justus von Liebig did a similar study with some organic acids. He found, for example, that cyanic, tartaric, and citricacidsseemed to be capable of reacting with one, two, and three units of hase, respectively. He was soon thinking in terms of monobasic, dibasic, and tribasic acids. Liehig (1838) revived the idea, previously suggested by Daw. that hvdromn was the essential inaredient in acids. He . .. ncsuked all acids to he"cornpounds of h;drogen in u,hirh the hvdrorcn rould he easilv rt:daced t ~ vmetnls." The numher oisucK replaceable hydrogen atoms determined whether the acid was monobasic or dolvbasic. Bases continued to he regarded simply as any substances that could neutralize acids, forming salts. Dualistic Theory of Berzellus
Following the development of the battery by Alessandro Volta (1800). chemists beean t o use this new device t o decompose all kinds of suhst&ces. Many compounds upon being subiected to an electric current would break down to yield new materials a t the two electrodes. Jiins Jacob ~ e r z & u sand William Hisinaer (1803) found that when salt solutions were sul~jrctedro rlucrrolysis, hases were found at the negative pols and nrids ar rhe vositive ode. Thev interoreted this to mean that acids and hases must carry opposite electrical charges. Berzelius concluded that acid-base reactions were simolv the result of electrical attractions, and he extended this iiek r o mhrr kinds of chemiral rfwtions as well. His duolibric theory (1812) explained all chemical interactions in terms of neutralization of opposite electrical charges. Arrhenius Theory of Ionization
Svante August Arrhenius, during his study of electrochemistry, observed that solutions of salts, acids, and hases were the only liquids that would conduct an electric current. He suggested (1884) that when these compounds dissolved in water they dissociated into charged particles, which he called "ions." Michael Faraday had introduced that term earlier as a name for the particlesdischarged at the electrodes of electrolytic cells, hut Faraday had assumed that the ions did not form until the current started to flow. Arrhenius 460 1 Journal of Chemical Education
theorized that the ions formed as soon as the compounds were dissolved in water. According to the Arrhenius theory acids are compounds that produce hydrogen ions in water solution
-
HCI H+ + C1and bases are substances that provide hydroxide ions in water solution. NaOH Na+ +OHWhen an acid and a hase neutralize eath other, the products are a salt and water.
--
HCI + NaOH NaCI + HzO (H+ + C1- + Naf + OH- Na+ + C1- + H20) HN03 + KOH KNO, + Hz0 (H+ + Nos- + K+ + OH- K+ + Nos- + H20) In generalized form, the Arrhenius acid-base reaction might he written H+ + OH- HzO the hydrogen ions from the acid combining with hydroxide ions from the hase to produce water. Perhaps it should he mentioned that H+ cations are unique in having no electrons. Their small size makes their charge density very high. The H+ ion is a proton, and free protons do not really exist in water solution. Although often written as H+ for the sake of simplicity, the hydrogen ions in acid solutions are actually attached to molecules of the solvent. One way to indicate this in equations is by witing H+(aq) to signify the "aqueous" solvation of the ion. Another way is by using the hydronium ion, H30+, which is a proton attached to a molecule of water.
The hydronium ion emphasizes the role of the solvent in reactions of acids that occur in water solution.
-
H30+ + OH- 2H20 Since liquid water molecules exist in hydrogen bonded clusters, with about four HzO units in an average group a t room temperature, it may he that the actual formula for the aquated hydrogen ion is closer to Hg04+ (although it is rarely written this way). In any case, that fact does not lessen the usefulness of the hydronium ion, which remains the simplest and most practical way to indicate the covalent attachment of H+ ions to water. Franklin's Ammonia Solvent System
The Arrhenius theory was successful in explaining acid-base reactions in water solution, but there were similar reactions that occurred in non-aqu&us solvents that the Arrhenius definitions did not include. Edward Franklin (1905) proposed that in liquid ammonia the reaction of ammonium chloride with sodium amide was an acid-base reaction, even though there were no H+ or OH- ions present and no water was formed.
-
NH4CI+ NaNH2 NaCI + 2NH3 Franklin suggested that in liquid ammonia NH4+ acted as an acid and NH2- as a base, with neutralization represented by the equation NH4++ NHz-- 2NH3 This was analogous to the reaction of H30+ and OH- ions in water. Brpnsted's Proton Theory
In 1923 Thomas Martin Lowy in England and Johannes Nicholas Br$nsted in Denmark independently arrived a t a
more generalized concept of the acid-base phenomenon. They saw the neutralization reaction as the transfer of a hydrogen ion (a proton) from an acid to a base.
-
+
(in water) H30t OH- Hz0 + HzO (in ammonia) NHaC+ NH2NH8 + NH3 A proton (H+) is transferred from HsO+ to OH- in the first equation and from NH4+ to NH2- in the second. Although the basic idea of the proton theory was introduced almost simultaneously by Bransted and Lowry, its further development was mainly due to Bransted. For this reason the theory is sometimes credited to Br6nsted-Lowry, hut more often simply to Bransted. Bransted defined an acid as "a species having a tendency to lose a oroton" and a hase as "a soecies havine a tendencv to add 0n.a proton." In other words,'heconddere; acids to de proton donors and bases to be proton acceptors. Bransted acid-base reactions can occur in various solvents, or they can h a.~.o e nin the gas with no solvent Dresent at all. as in - ohase . the case of hydrogen chloride and ammonia gases. HCI + NH3 NHI+ + CIacid, basez acid2 base, HCI gives up a proton to become C1-, while NHs accepts the proton to become NHd+. The products, in turn, are also an acid and a base in the Bransted sense. The NH,+ ion is the "conjugate acid" of the hase NHa, and the CI- ion is the "conjugate base" of the acid HCI. Notice that the formulas for an acid and its conjugate base, or a base and its conjugate acid, differ from each other by a proton. A + B + H t acid base proton A and B in this equation represent any conjugate acid-base pair. Solvents can also function as Bdnsted acids or bases. In the following reactions notice that water is a hase (a proton acceotor) when it acts as a solvent for HCI, hut an acid (a proton donor) when i t reacts with NH3. HCI + H20 H30++ CIacid, basez acidz basel NH3 + HOH NH4+ + OHbase, acid2 acid, base* A molecule such as water, that can act as either an acid or a base, is said to he amphiprotic (capable of either accepting or losing a proton).
-
-
-
-
was general enough to include all known acid-base reactions. The "positive-negative" theory proposed by the Russian chemist M. Usanovich (1939) was an attempt to establish a more universal acid-base concept. Usanovich defined an acid as any substance capable of giving up cations (H+or other) or combining with anions, and a hase as any substance that could eive UD anions or comhine with cations. For examole.. SO? was an acid becausc it could combine with theanion 02' to form S O P - . T h ethrorv was criticized because it olaced too much stress on the importance of ions and the formation of salts. and because Usanovich extended i t to oxidation-reduction reactions as a special class of acid-base interactions, an idea that most chemists found unacceptable.
.
Eleclronlc Theory of Lewis
In 1923, the very same year when BrCsted and Lowry introduced the oroton theorv. a more fundamental theorv was proposed by dilhert ~ e w t o c ~ e w iSo s .different was the Lewis theory from all the other acid-base theories that had gone before, that most chemists paid little attention to i t a t first. It was widely ignored for about fifteen years. G. N. Lewis perceived the neutralization reaction as the formation of a coordinate covalent bond (both bondine electrons being supplied by the same atom).
+
-
-
-
-
Positive-Neoative Theorv
The proton theory of Br6nsted limited acid-base reactions to proton transfers, and the solvent-systems theory limited
-
A :B A:B Lewis saw the hase as a supplier of electrons and the acid as an electron acceptor. Ionization might also take place, hut the primary step in an acid-base reaction was the formation of a coordinate covalent bond. According to the Lewis theory, an acid was an electron pair acceptor, and a base was an electron pair donor. The Lewis theory can be used to explain the traditional acid-base reaction in water hut it also applies equally well to acid-base systems that involve neither solvents nor transferral of protons.
Solvent-Systems Theory
There were some cases which appeared t o be acid-base reactions but which did not involve any protons. Albert F. 0. Germann (1925) showed, for example, that aluminum chloride dissolved in ~ h o s a e n ebehaved as an acid. even though there were no protons the system. (solvent alone) COClz COCP + CIAlCl3 + COCI, COCli + AIC14The effect of the aluminum chloride was to increase the concentration of the solvent cation, COCl+. Germann concluded that it was the solvent that determined the acid-base system. The "solvent-systems" theory had its origin about twenty vears earlier in the ideas of Franklin reeardine the ammonia system. Germann mainly expanded Franklin's theory to include non-protonic solvents. The solvent-systems theory, in simplified form, defined an acid as a solute that increased the solvent cation concentration and a base as a solute that increased the solvent anion concentration. Although more general in scope than most other acid-base theories, it still had several important limitations. I t required that a solvent be present, and the solvent had to he slightly ionizable.
-
acid
base
acid-base complex
The Lewis electronic theory is a unifying concept that can be applied to all kinds of acid-base reactions. I t is also consistent with the fundamental orinciole that chemical interactions occur in order to form molecules of more stable electron configuration. ~
~
~
~
Current Definitions of "Acid" and "Base"
How one chooses t o define the terms acid and base todav is likely to depend upon the circumstances. Of all the acid-base theories that have been proposed over the past three hundred years, the three that have remained mostuseful are those of Arrhenius, Bransled, and Lewis. For those whose interest in acids and bases is limited strictly to water solutions, the Arrhenius definitions are usually sufficient for identifying acids and bases and explaining neutralization. The broader definitions of Br$nsted have the advantage that they can he applied to many non-aqueous systems as well as to water solutions, and his conjugate acid-base concept has proved to be extremelv useful. Brdnsted's Droton theorv also orovides a very simple way to explain neutralization. since most common acid-base reactions do involve Droton transfer. the Brdnsted definitions are the ones preferrkd by many chemists. 1t is the Lewis theow, however, that ismost inclusive in its definitions of acids andbases. a he three theories are compared in Table 1. Volume 55, Number 7, July 1978 / 461
Table 1. Comparison of Acid-Base Theories Lewis Anhenius Bnn~ted Proton ElechDnic Water-Ion Theory Theory Theory: Theory AcidDefi- Provider of H+ Proton donor in water nition: Bsss Provider of OH- Proton acceptor in water Deli-
Electron pair acceptor Eleclron pair donor
acid base acid-base complex In the vapor phase aluminum chloride has been shown to exist as a dimer, AlzC16, formed by the mutual attraction of two AIC13 molecules.
nition:
Neutrali- Formation of zation: water Equation: HC + OHH-0
Limit* lions:
Proton transfer
-
+
HA B +A
-
Coordinate covalent bond
-
formation
BH
Water solutions Proton transfer only reactions only
A
+ :B
A:B
Generalized theory
I t might be noted that BrQnsted significantly broadened the definition of the word base, although he did not change the definition of acid very much. Br6nsted acids are not greatly different from Arrhenius acids (both supply hydrogen ions), but BrQnsted bases extend far beyond the hydroxide bases ofhrhenius and are almost as broadly defined as Lewis bases. (An unshared electron pair is usually a proton acceptor.) The Lewis acid definition is much broader than either of the others. A proton, or H+ ion, is an electron pair acceptor, but i t is only one of many. Since water is by far the most common solvent, the terms acid and base are still generally taken to mean suppliers of hvdroeen and hvdroxide ions. For the sake of claritv. basic " species such as fluoride ion and pyridine are usually referred to as Lewis or Brdnsted bases. and uon-orotonic acids such e ~ e w iacids.' s as sozand ~ e c l s ' a r called
~~~-
Aluminum Chloride a s a Lewis Acid Neither the Arrhenius nor the Br$nsted theory classifies AICIBas an acid, but the Lewis theory does. It might be argued that aluminum chloride does hydrolyze to "produce hydrogen ions in water solution" in keeping with the Arrhenius definition of an acid, or that the hydrated aluminum ion is a proton donor, thereby qualifying as a Br6nsted acid; but AIC13 does not need water in order to function as an acid.
--
AICh + 6Hz0 AI(Hz0)c3++ 3C1+ H20 A1(H20)50H2++ HsO+ Al(H~0)6~+ An important criterion for determining whether a substance is an acid or a base is the way i t affects acid-base indicators. Aluminum chloride causes the indicator crystal violet to turn yellow, its acid color. When pyridine is added, the color changes to blne-violet as the point of neutralization is passed.
pyridine
acid-base complex
By carefully measuring thequantities used, a titrationof AICI:. aeainst thc base ~vridinrcan he carried out. This wmuld seem :t confirm that ~ i c 1 is 3 indeed an acid. The aluminum atom can accommodate eight electrons in its valence shell (the octet rule), but in AlC13 aluminum has onlysix valence electrons. There is one empty orbital that can hold another electron pair. :~
:CI:AI .. .. :c1:
-
unoccupied orbital
[t is this vacancy that makes AIC13 a Lewis acid, able to accept 3 uair of electrons. Anv substance that has an electron pair it :an share (such as chloride ion, for example) is a ~ e w i s base. 462 / Journal of Chemical Education
This dimerization can be looked upon as a double acid-base reaction, with each AIC13 molecule behaving simultaneously as a Lewis acid and a Lewis base. (The Al2Cl6molecule has the shape of a pair of tetrahedra with a common edge. The aluminum and central chlorine atoms lie in a common plane, with the outer chlorine atoms in a perpendicular plane.)
c1. .Al/
C1\M.'
c1
c1' c14 Relative Strength of Acids and Bases Chemists recognized the fact that all acids and bases were not of eoual strength lone before thev understood whv. Sulfuric was a much srrungrr acid than avctic, and sodium hgdroxide a stroneer IIRW than lime water. After Arrhcnius developed his iorh.ation theory, the differenrrs n,uld he v x ulained on the basisof extent of ionizntinn. In wntrr sdution strong acids and bases diisociat~completely intu ions, while weak arids and bases are only slightlv dis~oriated. Strong Acid: HX + HzO i.H30t + X- (almost l O W d ions) Weak Acid: HA + H20 t HsOt A- (few % ions) Such equilibria also exist in nonaqneous solvents. Sometimes the solvent can conceal differences in acid or base strength. Although perchloric acid (HCIOd is a much stronger acid than hydrochloric (HCI), both are completely ionized in water and appear to be about equal in strength. This is due to the levelling effect of the solvent. In water both acids are reduced to the strength of HsOi. HCIOd + H20 HsO+ + ClodHCI + H20 H,O+ + CIThe strongest acid that can exist in a given solvent is the conjugate acid of the solvent. (In liquid ammonia, for instance, all acids are reduced to NH4+, and acetic acid appears to be as strong as sulfuric.) The strongest hase that can exist in a given solution is the conjugate base of the solvent (OHin the case of water, or NH2- in liquid ammonia). The useful concept of conjugate acid-base pairs is, of course, part of the BrQnsted theory. One of the factors that determine acid or base strength is electronegativity (the attraction an atom has for a pair of shared electrons). A compound of the formula EOH might be either an acid or a hase, depending on the electronegativity of the element E. If it is low (as in NaOH), the compound acts .as a base; if i t is high (as in CIOH), the compound is an acid. :0:6;: .. H+ Naf :i):H-
-
-
(The electropositive sodium atom readily gives up its valence electron, but the chlorine atom clings tightly to the electron pair shared with oxygen.) Metals are generally low in electronegativity, so metal oxides and hydroxides are usually bases. Non-metals are hieher in electroneeativitv. so nonmetal oxides and hydroxibes are acids. ~ l e k e n t s - d medial f electronegativitv (such as antimony and tin) form amohoteric oxides, wkch can act either as acids or as bases. The oxidation state of an element also affects the acidity or basicity of its compounds. Again, this is easiest to illustrate
with hydroxy compounds. The higher the oxidation state of an element. the more acidic its hydroxy compound should be. orid. state of Cl +I
HOCl HOClO HOC102 HOCIOa
+2 +5
+I
very weak
acid
weak acid strong acid very strong acid
With metals, the lower the oxidation state, the more basic the oxide or hydroxide. oxid. state o f Tl
TlOH strong base TI(OHk weak base Although the oxy compounds of metals are usually bases, they can sometimes he acidic if the oxidation state of the metal is high enough. +1
+3
oxid. state o f Cr +2 +3 +6
CrO or Cr(OHh CrzO3 or Cr(OH)s CrOs or HzCrOa
basic amphoterie acidic
Metals that form acidic oxides are mainly found among the transition metals in the vanadium, chromium, and manganese families. The +4 oxy acids of titanium family elements are amphoteric. The following rule generally applies to the oxides and hydroxides of metal; its oxy compound will be:
If the oxidationstotr of the metal is: +I or +2 +3 or +4 +5, +6, or +7
basic amphoteric acidic A list showing the relative strengths of some common acids and their conjugate bases is given in Table 2. Notice that the stronger the acid is, the weaker its conjugate base, and the stronger the hase, the weaker its conjugate acid. The strength of an acid and that of its conjugate hase have a definite reciprocal relationship, and both are also related to the solvent. (In water, for example, K,Kb = K,, where K., Kb, and K , represent the dissociation constants of the acid, its conjugate hase, and water, respectively.) Pre~arine . . a similar list for Lewis acid and hases was fraught with problems. An experimental test for deciding which was the stronger of two Lewis acids (or bases) was displacement. (A strong acid displaces a weaker one from combination with a hase, and a strong base displaces a weaker one from combination with an acid.) I t was assumed that the order of strength should be the same no matter what reference acid or hase was being used, hut occasionally the results were contradictory. Order of base strength as measured against one acid was sometimes reversed when another acid was used. Efforts to explain this anomaly have resulted in a useful extension of the
Table 2.
Lewis theory known as the principle of "hard and soft"acids and bases. Hard and Soft Acids and Bases
During the 1950's it became obvious that Lewis acids and bases should be divided into two categories: class a, or "hard," and class b, or "soft." The "hard" and "soft" labels were introduced by Ralph G. Pearson in 1963 and seem to have caught on in spite of their informality. In general, the strongest acid-base complexes are those formed from hard acids in combination with hard bases. Hard acids have acceptor atoms that are small and difficult to polarize, and hard hases are those that hold tightly to their valence electrons. Some examples are Hard Acids
AI3+, Mg2+,Caz+,H30+,BF8,and AlClr
F-,OH-, NH3,S O F , P O P ,
Hard Bases
amines, alcohols, and ethers Soft acids, on the other hand, bind mostly strongly to soft bases. Soft acids have large acceptor atoms that are readily polarizable, and soft bases have valence electrons that are easily shifted or pulled away. Some examples are Soft Acids
Hg2+,Ag+, TP+, Cdz+,Ip,Bra, BH2, and Inch
I-, CN-, SCN-, S2-, CO,
Soft Boses
organic phosphines, and benzene Soft bases often have loosely held d orbital electrons, and soft acids often have empty d orbitals available. Hardness and softness also seem to have some correlation with properties such as electronegativity, polarizahility, and oxidation potential. It is not possible to make ageneral listing of Lewis acids (01 hases) in order of decreasing strength. The best one can do is to try to list them according to decreasing hardness or softness hut there are inconsistencies even then.. deoendina . - on the parameters used. When a hard acid is used as reference, the strength of Lewis bases tends to decrease accordine- to the following list of donor atoms
-
F > O > N > Cl > Br > I > P > S (decreasinghardness) When the reference acid is soft, the sequence is essentiall) reversed. The situation is further complic&d by the fact that manv Lewis acids and bases are "borderline," neither clearlq hardnor really soft. The concept of hard and soft acids and bases helps to ex. plain why n~etalssuch as calcium and magnesium art: found in nature a? uxides and cnrhonates, while metals such ar copper and mercury often occur as sulfides. Calcium and magnesium ions are hard acids, and oxide and carbonate ions are hard hases. Copper and mercury ions are soft acids, and sulfide is a soft base. I t is also interesting to note that most ol the acids and bases found in biological systems are hard. Ir fact, many soft acids and bases are poisonous to living organisms. The hard-soft principle has been useful in correlatin~ various chemical facts, and it has helped tie together diversf ideas regarding chemical bonding.
Relative Strength of Some Common Acids and Bases Conclusion
Conjugate Acid STRONG HCiO* A HCI
Increasing acidity
weak
Base CIOai
weak
61-
H2S04 HSOIHNO, NOaH30+ H,O HSOISO4*H3P04 H2P04i FHF HC2H302 C2H302C H2C03 HCOJNH4+ NHs H20 OH1 NHa NH2STRONG
In the light of modern acid-base theory, just what is ar acid-base reaction? I t is the familiar kind of neutralizatior that produces a salt and water, of course, but it is much morf than that. It is also the formation of esters, amides, ethers, anc anhvdrides. It~ is ~hvdrolvsis. ammonolvsis. and every other ~"~~ ~ kind of solvolysis. is acid d a t a ~ ~ s iand s , base catalysis, and chemical adsorotion. I t is the combination of dissolved ions to form covalent molecules or precipitated solids. Even hy. droeen hondine is a kind of acid-base phenomenon. " Transition metal complexes and other coordination com. nounds are a soecial class of acid-base adducts. The centra metal ions are Lewis acids, and the ligands surrounding then are Lewis bases. The ligands may be as simple as the watel ~
Increasing basicity
k
Volume 55, Number 7, July 1978 / 461
nolecules in a hydrated salt or as complicated as a polydentate :helating agent. All cations can act as Lewis acids, and all anions can act as Lewis bases. Many ions and molecules can act either as acids x as bases. All compounds can he looked upon as Lewis acid-base adducts. Methyl alcohol, for example, is derived From the acid CH3+ and the base OH-, or from the acid H+ m d the base CH30-. calcium carbonate is a combination of the acid C02 and the hase CaO, or the acid Ca2+and the hase C03=-. Even the simple HZmolecule is made up of the acid H+ and the hase H-. Then can we assume that all chemical h a ~ ~ e n i n gare s acid-hase reactions? What about redox reactions? when we we the I.ewis definitions. acid-hase neutralization dcrs sound a bit like oxidation-reduction. Reducing agents are electron donors, and so are bases. Oxidizing agents are electron acceptors, and so are acids. An oxidizing agent cannot act unless a reducing agent is present, and an acid cannot act unless a base is present. There are some interesting parallels, hut there are also important differences. Reducing agents donate varying numbers of electrons, while bases always donate electronpairs. An oxidizing agent may accept one or several electrons, while acids always accept paired electrons. A reducing agent gives up electrons (either partially or entirely) to an oxidizing agent, while a base only shares its electrons with an acid in a coordinate covalent bond. When redox reactions occur, there are always changes in oxidation state; but when acid-base reactions occur, there are usually no changes
A
164 1 Journal of Chemical Education
in oxidation state. The dissimilarities certainly seem to justify senarate categories for these two maior reaction types. keactions &volving odd-electronmolecules (s"ch as the combination of two free radicals) would also be outside the scope of the Lewis acid-base reaction. In other words. according to the Lewis electronic theory, acid-base reactions do not include oxidation-reduction (i.& electron transfer) or the reactions of odd-electron molecules, but they do include just about every other kind of chemical reaction, ?'he hmr hor r u o elecrn.nc t r con rhnre T h e o r i d c o n urcornrnodatr o pair
Their mutual attraction Leads to on interaction, And both ore neutralized by the affair.
References Alyea. H. N.. J. CHEM.EDUC., 18.206 (1941). Bell. 8. P.,"Acidsand Ba~e8:'Methuen. London, 1969. Bell, R. P.."The Proton in Chemistry? 2nd ed., Cornell Univ. Prm, Ithsea, N.Y., 1973. Drago, R.S.. and Mafviyoff, N. A . "Addaand Bm3,"D. C. Heath. Boston. 1968. Hali,N.F..J. CHEM. EDUC., 17.124 (1940). ~ ~ i N.,"vslenee ~ . c . and thestructureof Atoms andMolecules."TheChemieslCatalog Co.. New York, 1923. Levin,G. N., J. Franklinlnst., 226.293(19381. Ludor, W . F.,J. CHEM. EDUC.,25.555 (1948). Luder. W. F..and Zuffmfi.. S.."EledronicThwmofAcids and Basea:'Joho Wiev &Sam, . NewYork, 1946. P-n, R. G.,"Hardand SofiA~idsandBma,"Dowden,HuVhenmn &Ross.Stroudsbug, ~
Pa..
1913.
VanderWerf, C. A.,"Aeido. Bsws, and the Chemistn ofthe C~valentBond."ReinhoId. New York, 1961. Walden, P., "Salts, Aeida, and Bases: MeCraw-Hill. NnuYork. 1929.
