• CLASSIFICATION BASED ON THEORIES/CONCEPTS • 1. ARRHENIUS CONCEPT OF ACID BASE • 2.BRONSTED LOWRY THEORY • 3. LEWIS THEORY • 4. LUX-FLOOD CONCEPT • 5. SOFT AND HARD ACID BASE THEORY ARRHENIUS CONCEPT
• ACID PRODUCES • Base produces HYDROGEN ION IN Hydroxyl ion in aqueous AQUEOUS SOLUTION solution + - • HCl+H2O→H3O + Cl Or HCl(aq) → H O (aq) + Cl- (aq) NaOH(aq) → OH-(aq) + Na+(aq) • 3 • HI>HBr>HCl>HF(decreasing strength of acidity) KOH> NaOH> LiOH (Fajan’s concept) SiH + H O → SiO + H • 4 2 2 2 CH4 +H2O → no reaction Te(OH)6 > Si(OH)6 > B(OH)3
HF, H2O, NH3, CH4 Acidity?? DRAWBACKS
• Unable to explain why NH3 which contains no OH- ions, is a • base and not an acid Arrange in order of increasing acidity: • Why a solution of FeCl3 is acidic 1.H3PO2, H3PO3, H3PO4 or why Na2S is alkaline • Limited to aqueous media only. 2. HClO , HClO , HClO , HOCl 4 3 2 Failed to explain why NaNH2 is 3. [Al(H O) ] +3 , [Fe(H O) ] +2 , [Co(H O) ] +2 alkaline and NH4Cl acidic in 2 6 2 6, 2 6 liquor ammonia • Failed to define inherent acid- base character(phenol/picric acid) Bronsted-Lowry Theory The Proton-donor-acceptor system
• Bronsted Acids • Bronsted Bases • Molecular: HCl → H+ + Cl- • Molecular: H2O + H⁺ → H3O⁺ Cationic: [Al(H O) ] ⁺³ • 2 6 • Cationic: ↓ [Al(H O) (OH)] ⁺² [Al(H O) (OH)]⁺² + H⁺(proton) 2 5 2 5 • Anionic : HCO ¯ → H+ + CO ¯ Anionic: 3 3 CN- + H+ → HCN 2- - CO3 + H⁺ → H CO3
- HCl + H2O ↔ H3O⁺ + Cl Acid 1 Base 2 Acid 2 Base 1 (Conjugate acid-base pair) Reactions in aqueous media Effect of charge on acidity Comparison of Hydracids
• The acid strength increases with increase in hydracid Ka value BDE(kJ/mol) the size of atoms HI> HBr> HCl> HF HF 7.2 x 10-4 570 H2Te> H2Se> H2S> H2O
HCl 1x 106 430 Gas phase acidity increases across a period and down HBr 1 x 109 370 a group in the p-block binary acids HI 3.3 x 109 300 Oxo and Hydroxo acids
• An oxo acid is one in which the Oxy M(O)x Value pKa Ka pka of X acidic proton is on a hydroxyl acid (OH)y =8-5x expt group with an oxo group attached to the same HOCl Cl(OH) 0 8- 3x 10-8 7.5 atom.Usually represented as (5 x 0) M(O)x(OH)y =8 • The acidity of these oxyacids increases HClO ClO(O 1 8- 1 x 10-2 2.0 significantly as the oxidation state of the 2 H) (5x1) central atom becomes larger =3 HOCl< HClO < HClO < HClO • 2 3 4 2 HClO3 Cl(O)2 2 8- 5x 10 -1.3 A hydroxo acid is one in which the acidic (OH) (5x2) proton is on a hydroxy group without a = -2 neighboring oxo group. Usually - + 3 represented as M-O-H → MO + H HClO4 Cl(O)3 3 8- 1x 10 -10 (OH) (5x3) Te(VI)OH > Si(IV)(OH) > B(III)(OH) = -7 6 4 3 Oxo acids of Phosphorus H⁺δ • H3PO2 (Monobasic with one hydroxyl and one | oxo group) • H― P(I) = Oδ¯² • H3PO3 (Dibasic with two | hydroxyl and one oxo δ¯O…H⁺δ group)
• H3PO4 (Tribasic, Three δ¯O…H⁺δ hydroxyl and one oxo | group) H…O δ¯―P(v) = Oδ¯² • H PO > H PO >H PO 3 2 3 3 3 4 | Strong acid weak acid δ¯O…H⁺δ Lewis concept of Acid and bases • Lewis acid: A chemical species that contains an empty orbital to accept a pair of electrons and form a coordinate covalent bond • Lewis base: A substance that has a filled orbital containing non-bonding electron pair which it can donate to form a dative bond with a Lewis acid.
H3N: + BF3 → H3N►BF3 Lewis base L. acid acid-base adduct BF₃ < BCl₃ < BBr₃ < BI₃ -----Lewis acid st increases → Group Characteristics of Lewis acids
• SiF4 > SiCl4 > SiBr4 > SiI4 Correlates with the decrease in electron withdrawing power of the halogen from F to I δ δ- + - Si ⁺ F4 + H2O ( H ....OH )= Si(OH)4 + 4HF ↓ SiO₂
SbF + 2HF → [SbF ] ¯ + [H F] ⁺ 5 6 2 Super acid Oxides of Group 16 family like Sulfur dioxide can act as Lewis acid by accepting an electron pair at S-atom, and can also act as a base by donating its lone pair either from sulfur or from oxygen
Bromine and Iodine molecules of Group 17 family act as mild Lewis acids Boric acid behaves as strong acid in presence of diols
• Boric acid (H3BO3) is a weak acid which does not act as a proton donor but acts as a Lewis acid, accepting a pair of electrons from hydroxyl ion. -10 ) B(OH)3 +H2O ↔ [B(OH)4]¯ + H⁺ (ka= 6x10
[B(OH)4]¯ + cis-diol → (anionic chelate complex) + H⁺ + 4H2O The proton can now be treated against strong alkali using phenolphthalein indicator - - - - • 1. Arrange the bases HS , F , I , NH2 in order of increasing proton affinity
2. Thermally most stable: PH4Cl, PH4Br, PH4I
3. Most acidic oxide: Ag2O, V2O5, CO Soft and Hard Acid Base Principle • Ralph G. Pearson (1963) proposed the principle, abbreviated as SHAB. The complex A:B is most stable when A and B are either both soft or both hard. The complex is least stable when one of the reactants is very hard and the other is very soft. A + B → A:B Lewis acid Lewis base adduct (Acceptor) (donor) Hard base: hard to oxidise Characteristics: High electronegativity/ low polarisibility/presence of filled orbitals Hard acids: Difficult or hard to polarise. Characteristics: high oxidation state/a small size/a noble gas electronic configuration + Methyl mercury cation CH3Hg is proton like and monopositive soft Lewis acid + + + + if K>1, B is soft BH + CH3Hg ↔ CH3HgB + H if K< 1, B is hard Indicate with reason the direction of the following reaction: i. CF H + CH F = CF + CH + - + - 3 3 4 4 (CF3 H ,CH3 ,F ) ii. H SO + FeS (s)= H S(g) + FeSO 2 4 2 4 Lux Flood Concept • The definition covers things which would become acids or bases if dissolved in water
CaO+ CO2 =CaCO3 CO is considered an anhydride f carbonic acid and CaO as base since it 2 would give Ca(OH) in water 2 An acid is an oxide ion acceptor and a base is an oxide ion donor The concept is generally used in high temperature anhydrous systems and in molten state Chemistry
CaO+SiO =CaSiO 3Na O +P O = 2Na PO 2 3 2 2 5 3 4
Base acid Other solvent system Ammonia as solvent Sulphur dioxide as solvent + - • H2O + H2O ↔ H3O + OH 2+ 2- • SO2 + SO2 ↔ SO + SO3
NH + NH ↔ NH ⁺ + NH ¯ SOCl + K SO = 2KCl + 2 SO 3 3 4 2 • 2 2 3 2 Acid Base salt solvent NH Cl (ammonium chloride) reacts 4 with KNH (potassium amide) 2 to give salt and solvent
NH Cl + KNH = KCl + 2NH 4 2 3 Acid Base Salt Solvent BiI₃ + KNH₂ → product? BiI₃ +KNH₂ = BiN + 3KI + 2NH₃
+ 2- NH₂¯ ↔ H + NH
2- 3- + NH ↔ N + H Drago-Wayland equation • Drago and Wayland proposed a quantitative system of acid-base parameters to account for Lewis acid-base reactivity by including electrostatic and covalent factors. They proposed a two parameter equation which predicts the formation of a very large number of adducts quite accurately.
-∆HAB=EA.EB + CA.CB Here, ∆HAB is the enthalpy of the reaction A+B →AB in the gas phase or in an inert solvent. E and C are parameters calculated from experimental data. E is a measure of the capacity for electrostatic(ionic) interactions and C is a measure of the tendency to form covalent bonds. The subscripts refer to values assigned to the acid A and Base B. Ques: Using the Drago-Wayland equation, calculate ∆H for the reactions (CH ) N: and BF and of (CH ) N: and B(CH ) Discuss the values in terms of SHAB theory.3 3 3 3 3 3 3. (CH ) N: ; C = 11.54 kcal/mol, E =0.808kcal/mol; BF , C =1.62 kcal/mol, E =9.88 kcal/mol 3 3 B B 3 A A
B(CH ) : C = 1.70kcal/mol; E =6.14 kcal/mol 3 3 A A solution • For the reaction between boron trifluoride and trimethyl amine → F B:N(CH ) BF3 + :N(CH3)3 3 3 3 Acid Base Adduct Putting the given values in Drago wayland equation: -∆HAB= [(9.88x0.808)+(1.62x11.54)] kcal/mol or -∆HAB = (7.9804+ 18.6948)= 26.68 kcal/mol For the reaction between trimethyl boron and trimethyl amine, (CH ) B + :N(CH ) → (CH ) B :N(CH ) 3 3 3 3 3 3 3 3 Acid Base Adduct -∆HAB= [(6.14x0.808)+(1.70x11.54)] kcal/mol or -∆HAB = (4.96112+ 19.6180) kcal/mol = 24.58 kcal/mol Agreement is very good with Pearson’s SHAB principle which sows a better adduct formation between hard acid BF and hard base :N(CH ) 3 3 3. Conjugate acid base pair + pH of a solution=-log10[H ]
• pH + pOH=14 Equations relating to acid-base equilibria • pK =-log K -pKa -3 a a ; Ka=10 Calculate the pH of 10 (M) NaOH solution -pKb • PKb=-log Kb ; Kb=10 -8 -9 • PK =-log K = 14 Calculate the pH of 10 (M) HCl and 10 (M) w w NaOH K =K x K -8 -7 -7 -1 • Ans: [H⁺] = (10 +10 )= 10 (10 +1) w a b -7 =1.1 x 10 (M) • At neutral solution,[H+]= [OH-] -7 • For an acidic solution, [H+]>[OH-], pH<7 or, pH= -log1.1 + - (log 10 )= 7- log 1.1 • For a basic solution, [H+]<[OH-], pH>7 Or pH= 7- 0.414=6.9586 (Kw=ionic product of water=10 ¯14at 25 degree celsius Which of the following solutions will act as buffer? i. BaCl2 + Ba(OH)2 ii.CH3COOH +NaOH (2:1) pOHpH==pKpKa+log+log[salt]/[acid][salt]/[base] iii. CH3COOH + NaOH (1:2) b iv. CH3COONH4 v. H3BO3 + Na2B4O7.10H2O vi. NaH2PO4 + Na2HPO4 Simple solutions
For a neutral solution, pH=pOH And pH+pOH = pKw Therefore 2pH=12.30, or pH=6.15