Brønsted Acidity Brønsted Theory of Acids and Bases Acid Substance which donates a proton Base Accepts proton from another substance HF+H O

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1 Acids and Bases Bronsted Acidity Periodic trends Lewis Acidity Systematics of acids and bases Description of different acids and bases Proton transfer Electron pair sharing Acceptor and donator Description of acidity Hard and soft Thermodynamic 5-1

2 Brønsted Acidity Brønsted Theory of Acids and Bases Acid Substance which donates a proton Base Accepts proton from another substance HF+H 2 O H 3 O + + F - NH 3 + HCl NH 4+ + Cl - H 2 O + HCl H 3 O + + Cl - NH 3 + H 2 O NH 4+ + OH - Remainder of acid is base Complete reaction is proton exchange between sets Extent of exchange based on strength Amphiprotic substance Acts as both Brønsted acid and base 5-2

3 Acid Strengths Each acid has conjugate base HSO 4- +OH - SO H 2 O HSO 4- is acid, SO 4 2- is base * In equilibrium in solution Strong acids tend towards completion HCl + H 2 O <--> H 3 O + + Cl - Strong acids tend to have weak conjugate bases Weak acid forms only slightly ionized species CH 3 COOH + H 2 O <--> CH 3 COO - + H 3 O + Stronger conjugate base Relative strengths of acids can be compared Can be rated relative to water Some salts can have acid-base properties NH 4 Cl 5-3

4 Relative Strengths of Acids and Bases Acid Strength Conjugate Acid HClO 4 ClO - 4 H 2 SO 4 SO 2-4 HCl Cl - H 3 O + H 2 O H 2 SO 3 HSO - 3 HF F - HC 2 H 3 O 2 C 2 H 3 O - 2 HSO - 3 SO 2-3 H 2 S HS - NH + 4 NH 3 HCO 3 CO 2-3 H 2 O OH - HS - S 2- OH O 2- H 2 H - Conjugate Base Base Strength 5-4

5 Dissociation Constants Equilibrium expression for the behavior of acid HA + H 2 O <--> A - + H 3 O + Water concentration is constant K = [A ][H 3 O + ] [HA][H 2 O] pk a =-logk a Can also be measured for base K a = K[H 2 O] = [A ][H 3 O + ] [HA] Constants are characteristic of the particular acid or base Strengths are measured by K a or K b K b = [HB+ ][OH ] [B] 5-5

6 Acid Constants at 25 C 5-6

7 Polyprotic acid Acid that can donate more than 1 proton H 2 SO 4, H 3 PO 4, ethylenediaminetetraacetic acid H 2 CO 3 +H 2 O HCO 3- +H 3 O + K a1 = [HCO 3- ][H 3 O + ]/[H 2 CO 3 ]=3.5E-7 HCO 3- +H 2 O CO H 3 O + K a2 = [CO 3 2- ][H 3 O + ]/[HCO 3- ]=5E-11 * K a2 is smaller than K a1 ➋ Larger negative charge on counter ion after 1 st deprotonation Data can be used for speciation calculation Determine which ion is dominate at a given ph α = i x [ i] [ i x ] 5-7

8 Phosphoric Acid speciation 5-8

9 Dissociation Constants for Acids at 25 C Acid Formula K a Acetic HC 2 H 3 O 2 1.8E-5 Hydrocyanic HCN 7.2E-10 Carbonic H 2 CO 3 3.5E-7 HCO - 3 5E-11 Nitrous HNO 2 4.5E-4 Hydrosulfuric H 2 S 1E-7 HS - 1E-14 Phosphoric H 3 PO 4 7.5E-3 H 2 PO E-8 HPO E-13 Oxalic H 2 C 2 O 4 5.9E-2 HC 2 O E-5 5-9

10 Dissociation Constants for Bases at 25 C Base Formula K b Ammonia NH 3 1.8E-5 Pyridine C 5 H 5 N 2.3E-9 Methylamine CH 3 NH 2 4.4E-4 Protonation of amine group 5-10

11 Buffers Weak acid or weak base with conjugate salt Acetate as example Acetic acid, CH 3 COONa CH 3 COOH + H 2 O <--> CH 3 COO - + H 3 O + large quantity huge quantity large quantity small quantity If acid is added hydronium reacts with acetate ion, forming undissociated acetic acid If base is added Hydroxide reacts with hydronium, acetic acid dissociates to removed hydronium ion 5-11

12 Buffer Solutions Buffers can be made over a large ph range Can be useful in controlling reactions and separations Buffer range Effective range of buffer Determined by pk a of acid or pk b of base HA + H 2 O <--> A - + H 3 O - Write as ph K a = [A ][H 3 O + ] [HA] [H 3 O + ] = K a[ha] [A ] 5-12

13 Buffer Solutions ph = pk a log [HA] [A ] The best buffer is when [HA]=[A - ] largest buffer range for the conditions ph = pk a - log1 For a buffer the range is determined by [HA]/[A - ] [HA]/[A - ] from 0.1 to 10 Buffer ph range = pk a ± 1 Higher buffer concentration increase durability 5-13

14 Buffers What is a good buffer system for maintaining a solution at ph 4? Want 0.1 M total in 1L total volume For acetic acid, pk a = = 4.75 log [CH 3 COOH] [CH 3 COO ] We have 0.1 M each Volume CH 3 COO - =1/6.62 L = 0.15 L Acid = =0.85 L log [CH 3 COOH] [CH 3 COO ] = 0.75 [CH 3 COOH] [CH 3 COO ] =

15 Calculations 1 L of 0.1 M acetic acid has ph = 2.87 What is the pk a for acetic acid CH 3 COOH + H 2 O <--> CH 3 COO - + H 3 O + [CH 3 COO - ] = [H 3 O + ] = pk a =4.73 K a = 10 (2*2.87) 2.87 = 1.84x

16 What is ph of 0.1 M NH 3 K b =1.8E-5 NH 3 + H 2 O <-->NH 4+ + OH - [NH 4+ ] = [OH - ] = x [NH 3 ] = x Calculations x E-5x -1.8E-6 = 0 x x = 1.8x10 5 [OH - ] = 1.33E-3 M, poh = 2.87, ph 14-pOH x = 1.8E 5 ± (1.8E 5)2 + 4*1.8E

17 Solvent effect Acidity can be affected by solvent Solvent is involved in reaction Solvents can be compared by autoprotolysis constant Compare with water Any acid stronger than H 3 O + will donate proton in water Both HBr and HI donate protons, which is stronger? Evaluate in acetic acid Can compare with K a of solvent Strong acids have pk a <0 5-17

18 Periodic trends Acidic proton Proton that can be donated Often ROH * Consider CH 3 COOH Aqua Acid Central atom with low oxidation state s, d, left p block, Acid on water coordinated to metal MOH 2 +H 2 O MOH +H 3 O + Reaction of water with metal ion * Common reaction * Environmentally important * Dependent upon metal ion oxidation state M z+ + H 2 O <--> MOH z-1+ + H + Constants are listed for many metal ion with different hydroxide amounts 5-18

19 Periodic Trends Hydroxo acid High oxidation state central metal ROH that is not carboxylic Si(OH) 4, phenolic group Oxoacid Right p block Proton on hydroxyl group with an oxo group on the same atom Formation of oxo groups on metals (H 2 O) 2 M HOM=O 3- +3H

20 Trends in aqua acid strength Increase with charge and Increase with decreasing ionic radius Charge density Electrostatic ζ = parameter IR + Deviation due to role of covalent bonding * Electron density on ligand orbitals enhances deprotonation z ( diameter H 2O 2 ) 5-20

21 Mononuclear acids Electronegative elements upper right of periodic table Can be substituted F, NH 2 * Electron withdrawing or donating * Change strength of proton Pauling rules O p M(OH) q pka 8-5p q>1, pka increase by 5 Not accurate with H 2 CO 3 and H 2 SO 3 Oxoacids 5-21

22 Acidic oxide Hydrolysis of water Carbonic acid from carbon dioxide Basic oxide Proton transfer to oxide Oxide that reacts with acid CaO+2H 3 O + Ca 2+ +3H 2 O Correlated with properties Basic oxides are ionic (metallic) Acidic oxides are covalent (nonmetallic) Amphoteric oxide both acidic and basic properties Al 2 O 3 + 6H 3 O + +3H 2 O 2[Al(OH 2 ) 6 ] 3+ Al 2 O 3 + 2OH - +3H 2 O 2[Al(OH)] - Anhydrous Oxides 5-22

23 Polyoxo compound formation Formation of polymers or colloids Formed by loss of H 3 O + Protonation of O and departure as H 2 O 2[Al(OH 2 ) 6 ] 3+ [(H 2 O) 5 Al(OH)Al(OH 2 ) 5 ] 5+ +H 3 O + Polymerization of aqua ions and polyoxoanions 5-23

24 Lewis acid Electron pair donor and acceptor A :B A-B Bronsted acids are Lewis acids Bronsted bases are Lewis bases Wider range of substances are Lewis acids or bases 5-24

25 Lewis Acids Metal cation bonding to electron pair in coordination Completion of octet by accepting electron pair (CH 3 ) 3 B+NH 3 (CH 3 ) 3 BNH 3 (CH 3 ) 3 B is Lewis acid Octet accepting additional electron pair CO 2 +OH - CO 2 OH Expansion of valence shell to accept another electron pair SiF 4 +2F - SiF 2-6 SiX 4, AsX 3, PX 5 Closed shell molecule use one unoccupied antibonding MO to accommodate electron pair 5-25

26 Boron and Carbon groups BX 3 Vacant p orbitals can accept electron pair Geometry changes Complex stability for X F<Cl<Br * Order based on π bonding strength with B AlX 3 Dimeric in the gas phase 5-26

27 Si and Sn complexes Expansion of valence shells Increase in coordination SiF 4 +2HF SiF H + N and O group acids Heavier members of group SbF 5 +2HF SbF - 6 +H 2 F + Sulfur trioxide is a strong Lewis acid and weak base SO 3 +NH 3 SO 3 NH 3 * Higher polymer sulfur oxides ➋ H 2 S 2 O 7 I 2 and Br 2 Charge transfer bands * Electron transfer from base to acid Can accept electrons Lewis acid complexes 5-27

28 Lewis acid reactions Complex formation BF 3 NH 3 complex Net lowering of energy upon formation Electrons from base form bonding orbital Displacement reaction Displacement of one Lewis base by another Also called substitution reaction Double displacement can also occur A-B + A -B A-B +A -B 5-28

29 Hard and Soft Acids and Bases Description based on interactions Hard acid Bond in order I - <Br - <Cl - <F - Soft acids in reverse order Hard acids bind to hard bases Hard interactions electrostatic Soft interactions covalent 5-29

30 Thermodynamic acidity parameters Need to include electronic, structural, and steric effects Enthalpies of reaction include these parameters H (A-B)=E A E B +C A C B * E electrostatic * C covalent Describe electronic and covalent contribution to bonding Does not include solvent effects 5-30

31 Solvents Solvents can act as Lewis acids or Lewis bases Dissolution of solute in solvent Basic solvents Polar solvents * Water, ROH, ethers, amines, DMSO, dimethylformamide, acetonitrile Reactions are displacement Acidic solvent Hydrogen bonding * H 2 OHN 3 +H 3 O + NH H 2 O Solvent parameters Donor number * Negative of H (in kcal/mol) Acceptor number * Measure of solvent acidity 5-31