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Acids and Base,, , , , 7.1 Introduction, nt classes of compounds in chemistry. These two are defined in man,, , Acids and bases are two very importa:, ways depending upon the type of scientific work involved. The most important ones and worthy used are, Arrhenius concept, Lowry-Bronsted concept, Lewis concept, Pearson concept etc. Some of these are, , discussed below., , , , /°7.2 Arrhenius Concept of Acids and Bases, The first modern definition of acids and bases was devised by Arrhenius. It isa hydrogen theory of acids and, Arrhenius with oswald in establishing the presence of ions in, , bases and it followed from the combined work of., aqueous solution. This theory proposed two specific classifications compounds, tumed as acid bases, When, , dissolved in an aqueous solution, certain ions were released into the solution., Arrhenius acid : Acids are defined as a species that releases hydrogen (H?*) ions into the aq. solution, In the reaction shown below, nitric acid (HNO;) disassociates into hydrogen (H*) and nitrate (NO3, , ions when dissolved in water. So, it is an Arrhenius acid., HNO3(aq) 22> H*(aq) + NO3 (aq), H,O _, or HNO3(aq) —27> H30* (aq) + NO3 (aq), Loss of proton or H* causes protonation of water, thereby yielding hydronium (H;O*) ion. In modem times., the use of H* is regarded as shorthand for H;O*, because it is now known that a bare proton does not exist 4, a free species in aqueous solution. Another example is shown below where hydrochloric acid is shown losing, , , , i, , , , a proton,, H,O cc A, , or HCI(aq) 22> H30* (ag) + CI" (aa), , 2. Arrhenius base : Bases are defined as a species which furnish hydroxide (OH™) ions into the, , solution. the. Lit (aq) + OH™~(aq), , , , , , LiOH(aq), , f Scanned with Oken Scanner

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Acids arm weve ee, Le’, , (Lit) and hydroxide (OH’) ions when, , s reaction lithium hydroxide (LiOH) dissociates into lithium i, ydroxide is shown losing a hydroxide, , thi:, In One more example can be sited where sodium hi, , dissolved in water., , (OH) ion,, : NaOH (aq) ake, Nat (aq) + OH™(aq), , | HT) +@, , 3. Limitation of Arrhenius Theory : The major limitation of the theory is that the Arrhenius definitions, , of acidity and basicity are restricted to aqueous solutions only. Under this definition, pure H,SO,, dissolved in toluene is not acidic. Similarly,molten NaOH and solutions of calcium amide in liquid, ammonia are not alkaline. ‘, , 4, Summary of Arrhenius Theory : Overall, to qualify as an Arrhenius acid, in the presence of water,, the chemical must either cause, directly or otherwise :, , 1. Anincrease in the aqueous H+ or H,O* concentration, or, , 2. Adecrease in the aqueous hydroxide concentration., , Conversely, to qualify as an Arrhenius base, in the presence of water, the chemical must either cause,, , directly or otherwise :, , 1. Adecrease in the aqueous H* concentration, or, , 2. An increase in the aqueous hydroxide concentration., , 5. Arthenius Acid Base Reaction : The universal aqueous acid-base definition and subsequent, reaction between the two leads to the formation of a water molecule from a proton and hydroxide ion, along with aa salt molecule. This is a neutralization reaction in which the acidic and basic properties of, , H+ and OH™ are neutralized. The acid-base neutralization reaction can be put as,, acid + base > salt + water, , The positive ion from a base and the negative ion from an acid form salt together - in other words, an, acid-base neutralization reaction is a double-replacement reaction. An example can be considered in which a, Neutralization reaction takes place between hydrochloric acid (HCI) and sodium hydroxide (NaOH) yielding, sodium chloride and water. The reaction is shown below,, , HCl(aq) + NaOH(aq) > NaCl + H,O, , , , -3 Bronsted—Lowry Concept, , ne Bronsted-Lowry definition was formulated in 1923, independently by Bronsted and Lowry. It is, based upon the ability of a compound, later termed as acid, to “donate” hydrogen ions (H*) to bases, which, ee them. It is important to note that the theory refers to the removal of a hydrogen ion (H+), More, ™Pportantly, Bronsted-Lowry acid-base behaviour is formally independent of any solvent making it more 5, , Useful than the Arrhenius model., , 1, . ;, Bronsted Lowry Acid : Those species which lose H* in any solvent are termed as acids. So, the, , Solvent may or may not be water. Using any solvent is the biggest advantage of the concept., , f Scanned with Oken Scanner

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Neeeceucied ie, , 2. Bronsted Lowry Base : Those species which accept H* in any solvent are termed as bases,, , 3. Conjugate Acid-base : The removal of a hydrogen ion from an acid produces its conjugate base, Similarly, the reception of a proton by a base produces its conjugate acid. The general formula for, acid-base reactions according to the Bronsted-Lowry definition is :, , HA + B> BH*+ + A~, , where HA represents the acid and A“ represents the conjugate base of HA. Similarly, B represents the dase, BH+ represents the conjugate acid of B. It can be generalised as below,, acid] + base2 = conjugate basel + conjugate acid2., , As an example, a Bronsted-Lowry model for the dissociation of hydrochloric acid (HC!) in aqueous Solution, , can be shown as below,, HCl + H,O > H,O+ + Cl”, , The removal of H* from the HCI produces chloride ion, Cl”, the conjugate base of the acid. The addition o, H* to the H,O, which is acting as a base, forms the hydronium ion, HjO*, the conjugate acid of the base, Another example can be considered in which the reaction of ammonia, a base, with acetic acid (CH;COOH}, is shown and that too in the absence of water., , CH;COOH + NH, — NH4 + CH,;COO, The removal of H* from acetic acid forms its conjugate base, the acetate ion- CH,;COO“. The addition of the, H* to the ammonia forms its conjugate acid, the ammonium ion NHj. An acid-base reaction is, thus, the, , removal of a hydrogen ion from the acid and its addition to the base. Unlike the previous definition, the, Bronsted-Lowry definition does not refer to the formation of salt and solvent, but instead to the formation of, conjugate acids and conjugate bases, produced by the transfer of a proton from the acid to the base. In this, , approach, acids and bases are fundamentally different in behaviour from salts, which are seen as electrolytes., The concept of neutralization is thus absent., , 4. Amphoteric Species : Those species which can act both as a Lowry Bronsted acid and a base are, termed as amphoteric compounds. As an example, water can be considered. Water is an amphoteric, species. So, it can act as both an acid and a base. The Bronsted-Lowry model explains this, showing, the dissociation of water into low concentrations of hydronium and hydroxide ions :, , H,0 + H,O = H,0+ + OH™, , Here, one molecule of water acts as an acid, donating an H+, , second molecule of water acts as a base, accepting the H+ ion and forming the conjugate acid H,O+. Another, , example of amphoteric compound is ammonia (NH) whose dissociation is shown below,, NH, + H,O = NHj + OH, , and forming the conjugate base, OH™, and@, , Ammonia acts as a base, accepting an H* to form the ammonium ion-NH4, , The same NH; molecule functions.as a Lowry Bronsted acid as well,, , when it loses a roton,, NH; —— NH5 + H+ r, , f Scanned with Oken Scanner

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So, far a conjugate acid base pair,, , Kx K, = K,, , Asan example, HF and F- are conjugate acid-base pairs. So,, , (Ka)ur x (Kp),- = K,, , FO w, , List of conjugate acid-base pair, , , , , , , , , , , , , , , , , , , , , , Acid Conjugate base Base Conjugate acia, HCl cae H,O H,0®, HNO, NO} “OH? H,O, CH,COOH CH,COO® ROH ROH?, H,SO, HsO? NH@ NH,, HSO4 so2° NH, NH®, , , , , , 6. Limitation of Bronsted-Lowry Concept : The Bronsted-Lowry model calls hydrogen-containing, substances like HCI as acids. Thus, some substances, which many chemists considered to be acids,, such as SO, or BCI, are excluded from this classification due to lack of hydrogen., , , , 7.4 Lewis Concept of Acid Base, , The hydrogen atom based definitions of acid-base of Arrhenius and Bronsted-Lowry was completely, replaced by the Lewis definition devised by Gilbert N. Lewis in 1923. The theory was proposed in the same, year as Bronsted and Lowry's concept. Instead of defining acid-base reactions in terms of protons or other, bonded substances, the Lewis defines a base or an acid in the terms of an electron pair. In this system, an, acid does not exchange atoms with a base, but combines with it., , Lewis Acid : An acid is a compound that can accept electron pair and it is termed as Lewis acid. Lewis, acids are diverse. Simplest are those that react directly with the Lewis base. Examples of Lewis acids based, on the general definition of electron pair acceptor include the proton (H*) and acidic compounds, such as, NHj and H,O+, metal cations, such as Lit and Mg?*, often as their aquo or ether complexes, trigonal planar, , species, such as BF; and carbocations H,C* , pentahalides of phosphorus, arsenic, and antimony etc. Some, electron rich m-systems, such as enones and tetracyanoethylene also function as a Lewis acid., , The most studied examples of such Lewis acids are the boron trihalides and organoboranes, but other, compounds also exhibit this behavior. The examples below show the acidic character of BF., , (@) BR+F > BFy, (b) BF, + OMe, > BF;OMe,, Both BF4 and BF,OMe, are adducts of boron trifluoride. In many cases, the adducts violate the octet rule,, Such as the triiodide anion :, Ibt+lo &, The variation of the colors of iodine solutions reflects the ability of F to form adducts with the Lewis acid I., , Lewis Base : Lewis defines a base to be a compound that can donate an electron pair and it is referred to as, 8 Lewis base, Typical Lewis bases are conventional amines such as ammonia and alkyl amines. Other, , f Scanned with Oken Scanner

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Inorganic Chemistry.), 1-178 a":, , , , common Lewis bases include pyridine and its derivatives. Some of the other examples of Lewis bases an, , amines of the formula NH3_,R, where R = alkyl or aryl, Related to these are pyridine and its detivatives, Phosphines of the formula PR3_,A,, where R = alkyl, A = aryl, compounds of oO, S, Se and Te in, oxidation state 2, including water, ethers, ketones are some other examples of Lewis bases. The Mog, common Lewis bases are anions. Some examples are H~ and F”, other lone-pair-containing species, such, as H,O, NH, HO-, and CH3,complex anions, such as sulfate, electron rich m-system such as ethyne, , ethene, and benzene etc,, , 7.4.1 Application of Lewis Bases, , Nearly all electron pair donors that form compounds by binding transition elements can be viewed as, collections of the Lewis bases-or ligands. Thus a large application of Lewis bases is to modify the activity ang, selectivity of metal catalysts. Chiral Lewis bases thus confer chirality on a catalyst, enabling asymmetric, catalysis, which is useful for the production of pharmaceuticals. Lewis bases are most commonly used as, ligands in co-ordination chemistry. Many Lewis bases are “multidentate,” that is they can form several, bonds to the Lewis acid. These multidentate Lewis bases are called chelating agents., , 7.4.2 Lewis Acid Base Reaction, , In reactions between Lewis acids and bases, there is the formation of an adduct. A simple case of adduct, formation is shown here between NH, and H+ ion., , H, , t, , — + [q]}* ——> H—N-H, , |, H H, , As Lewis acid-base reactions can be viewed as the formation of adducts, one more example of adduct, formation can be sited,, , H* +OH™ = H,O0, , In some cases, the Lewis acids are capable of binding two Lewis bases, a famous example being the, formation of hexafluorosilicate by SiF,., , SiF, + 2F~— SiF2Another interesting example is the formation of adducts of borane. Monomeric BH, does not exis, appreciably, so the adduct of borane is generated by degradation of diborane : °, , BH, + 2H--» 2BHq, , Many metal complexes serve as Lewis aci on id Li, cids, but usual! imi i al ewls, a le i ee ly ly after eliminating a more wea kly bound Le i, , [Mg(H.0),]?* + 6NH3— [Mg(NH),]2* + 6 HO, , The proton (H*) is one of the stron i, The gest but is also one i a, ignore the fact that a proton is heavily solvated which i ee nee, , Sean ; Is to complication. Typi < a Lewis, : ction is in the Friedel-Crafts alkylation reaction. The Lewis acid pl. spe ceminieots™ ting &", electrophilic,carbocation as shown below, sia oat oie We OPS, , RCI + AIC, > Rt + AIC, , f Scanned with Oken Scanner