Acid -Base Theories

2. Acid-Base Theories

3. Arrhenius Theory     Svante Arrenhius (1857-1927) Acid: Substance that produces H+ in water.Base:  Substance that produces OH-1 in water. HCl(aq) H+ + Cl- produces H+ in water NH3(aq)  NH4+ + OH- produces OH- in water Although NH3 does not contain OH-, hydroxide ions form when added to water. Arrhenius acid and base neutralize each other to produce salt and water: HCl(aq) + NaOH(aq) NaCl(aq) + H2O)(l) H+(aq) + OH-(aq) H2O(l)

4. Bronsted/Lowry Theory     Johannes Bronsted (1879-1947)Thomas Lowry (1874-1936) Acid:  Substance that can donate proton (H+).Base: Substance that can accept proton (must contain lone pair of electrons). HCl + NH3 NH4+ + Cl- Acid base CA CB Acids may be cations, neutral molecules, or anions, while bases may be anions or neutral molecules.  Just as a reduction must always accompany an oxidation, a proton donor (acid) must accompany a proton acceptor (base).  Once an acid transfers its proton it becomes the conjugate base (CB) and once a base accepts the proton it becomes the conjugate acid (CA).  Since protons are always transferred in the Arrenhius concept, all Arrhenius acid/base reactions are also Bronsted-Lowry acid/base reactions. But if water is not involved (HCl & NH3), the reaction can be explained by Bronsted/Lowry concept and not Arrenhius.

5. Solvent system concept of acids and bases: Acid= cation of solvent via autodissociation Base = anion of the solvent by autodissociation. Solutes that increase the concentration of the cation of the solvent are considered acids and soultes that increase the concentration of the anion are considered bases The solvent must be able to behave as both an acid and a base (amphoteric) 2 H2O  OH- + H3O+ H2SO4 + H2O  H3O+ + HSO4- H2SO4is an acid 2BrF3  BrF2+ + BrF4- SbF5 + BrF3  BrF2+ + SbF6- SbF5 is an acid KF + BrF3  BrF4- + K+ KF is a base

6. Lewis Theory     Gilbert Lewis (1875-1946) Acid:Substance that can accept a pair of electrons from another atom to form a new bond.Base:Substance that can donate a pair of electrons to another atom to form a new bond. The product of Lewis acid-base reaction referred to as adduct. The proton itself can act as Lewis acid. Lewis expands acid/base reactions to include many substances without H in formula. F3B + NH3 F3B:NH3

7. Chapter 6 p166

8. Chapter 6 p169

9. Which theories can explain the following? HI + H2O      H3O+ + I- HI  +  NH3 NH4+ + I- I2 + NH3 NH3I+ + I- I2  +  Cl- ICl + I- X:- + Y+ Y:X  

10. Lewis Concept The lone pair in the HOMO of the ammonia molecule combines with the empty LUMO of the BF3, which has very large, empty orbital lobes on boron, to form the adduct. The B-F bonds in the product are bent away from the ammonia into a nearly tetrahedral geometry around the boron

12. Boron trifluoride-diethyl ether adduct Lone pair on the oxygen of the diethyl ether are attached to the boron. The result is that one of the lone pairs bonds to the boron, changing the geometry around B from planar to tetrahedral. As a result, BF3, with a boiling point of -99.9 oC, and diethyl ether, with a boiling point of 34.5 oC, form an adduct of 125 – 126 oC. Lewis acid-base adducts involving metal ions are called coordination compounds. Ag+ + 2 :NH3 [H3N:Ag:NH3]+

13. Frontier orbitals and acid-base reactions

14. HOMO-LUMO interactions In most acid-base reactions, a HOMO-LUMO combination form new HOMO and LUMO of the product. Orbitals whose shapes allow significant overlap and whose energies are similar form useful bonding and antibonding orbitals. On the other hand, if the orbital combinations have no useful overlap, no net bonding is possible and they can not form acid-base product. Even when the orbital shapes match, several reactions may be possible, depending on the relative energies. A single species can act as an oxidant, an acid, a base or a reductant, depending on the other reactant

15. HOMO-LUMO interactions • 2H2O + Ca  Ca2+ + 2OH- + H2 • water as oxidant 2. nH2O + Cl- [Cl(H2O)n]- water as acid 3. 6H2O + Mg2+ [Mg(H2O)6]2+ water as base 4. 2H2O + 2F2 4F- + 4H+ + O2 water as reductant A base has an electron pair in a HOMO of suitable symmetry to interact with the LUMO of the acid. The better the energy match between the base’s HOMO and the acid’s LUMO, the stronger the interaction.

16. Hydrogen bonding The lowest orbital is distinctly bonding, with all three component orbitals contributing and no nodes between the atoms. The middle (HOMO) orbitals is essentially nobonding, with nodes through each of the nuclei. The highest energy orbital (LUMO) is antibonding, with nodes between each pair of atoms 4 nodes 3 nodes 2 nodes

17. For unsymmetrical H-bonding B + HA  BHA, the pattern is similar

19. Electronic spectra Charge transfer: the transition transfers an electron from an orbital that is primarily of donor composition to one that is primarily of acceptor composition I2Donor  [I2]-[Donor]+

21. Hard and Soft Acid and Bases • Ag F(s) + H2O  Ag+(ag) + F-(aq) Ksp = 205 • Ag Cl(s) + H2O  Ag+(ag) + Cl-(aq) Ksp = 1.8 x 10-10 • Ag Br(s) + H2O  Ag+(ag) + Br-(aq) Ksp = 5.2 x 10-13 • Ag I(s) + H2O  Ag+(ag) + I-(aq) Ksp = 8.3 x 10-17 • Solvation of the ions is a factor in these reactions, with fluoride ion being much strongly solvated than the other anions. • Related to HSAB in which iodide is much softer (more polarizable) than the others and interacts more strongly with silver ions, a soft cation. The result is a more covalent bond.

22. Colors AgI yellow AgBr slightly yellow AgCl and AgF white Color depends on the difference in energy between occupied and unoccupied orbitals. A large difference results in absorption in the ultraviolet region of the spectrum; a smaller difference in energy levels moves the absorption into the visible region. Compounds absorbing violet appear to be yellow; as the absorption moves toward lower energy, the color shifts and become more intense. Black indicates very broad and very strong absorption. Color and low solubility typically go with soft-soft interactions; colorless compounds and high solubility generally go with hard-hard interactions.

23. Color and low solubility typically go with soft-soft interactions; colorless compounds and high solubility generally go with hard-hard interactions, although some hard-hard combination have low solubilities. LiBr> LiCl > LiI > LiF The solubilities show a strong hard-hard interaction in LiF that overcomes the solvation of water, but the weaker hard-soft interactions of the other halides are not strong enough to prevent solvation and these halides are more soluble than LiF.

24. Fajan's Rules (Polarization) Polarization will be increased by:- 1. High charge and small size of the cation Ionic potential ?Z+/r+ (= polarizing power) 2. High charge and large size of the anion The polarizability of an anion is related to the deformability of its electron cloud (i.e. its "softness") 3. An incomplete valence shell electron configuration noble gas configuration of the cation    better shielding less polarizing power i.e. charge factor in (1) should be effectivenuclear charge e.g. Hg2+ (r+ = 102 pm) is more polarising than Ca2+ (r+ = 100 pm)

25. Four rules can be summarized: • For a given cation, covalent character increases with increase in size of the anion. • For a given anion, covalent character increases with decrease in size of the cation. • Covalent character increase with increasing charge on either ion. • Covalent character is greater for cations with nonnoble gas electronic configuration. Q1. Ag2S is much less soluble than Ag2O A1. Rule 1: S2- is much larger than O2- Q2. Fe(OH)3 is much less soluble than Fe(OH)2 A2. Rule 3: Fe3+ has a larger charge than Fe2+

26. These rules are helpful in predicting behavior of specific cation-anion interaction, but not enough • Li series does not fit • Solubility MgCO3 > CaCO3 >SrCO3 >BaCO3 • Rule 2 predicts the reverse of the order. The difference lies in the aquation of the metal ions. Mg2+ (small with higher charge density) attracts water molecules much more strongly than the others, with Ba2+ (large with smaller charge density) the least strongly solvated.

27. Ahrland, Chatt and Davies: Class (a) ions: Most metals Class (b) ions: Cu2+, Pd2+, Ag+, Pt2+, Au+, Hg22+, Hg2+, Tl+, Tl3+, Pb2+, and heavier transition metal ions

28. The class (b) ions form halides whose solubility is in the order F- > Cl- > Br- > I-. The solubility of Class (a) halide is in the reverse order. The calss (b) metals ions also have a larger enthalpy of reaction with P donor than with N donor, again the reverse order of the Class (a) metal ion recations. Class (b) – having d electrons available for  bonding Tl(III) show stronger Class (b) character than Tl(I) because Tl(I) has two 6s electrons that screen the 5d electrons and keep them form being fully available for  bonding

29. Pearson’s Principle: Hard Lewis acids prefer to bind to hard Lewis bases; soft Lewis acids prefer to bind to soft Lewis bases Class (a)– hard acids Class (b)– soft acids

31. Hard and Soft Acids and Bases (HSAB) Let A be a Lewis acid, and B a base Measure log K for the reaction A + B  AB If for B = halide, the order of log K is I– > Br– > Cl– > F– then A is called a soft acid If for B = halide, the order of log K is I– < Br– < Cl– < F– then A is called a hard acid

32. Hard metal ions form their most stable complexes with Hard Bases Hard Bases: contain the smaller electronegative atoms, especially O, N, F and Cl. The bonding between a Hard Lewis Acid and a Hard Lewis Base is predominantlyionic

33. Soft metal ions form their most stable complexes with Soft Bases Soft Bases: contain the larger, more polarisable and less electronegative atoms, especially S, Se, P, C and As. The bonding between a Soft Lewis Acid and a Soft Lewis Base is predominantly covalent

34. Chapter 6 p183

35. Chapter 6 p184

36. The smaller drop in energy in the hard-hard case does not indicate small interaction. The hard-hard interaction depends on a longer range electrostatic force, and this interaction can be quite strong. Many comparisons of hard-hard and soft-soft interactions indicates that the hard-hard combination is stronger and is the primary driving force for the reaction.

37. Quantitative measure Absolute hardness  = (I –A)/2 Mulliken’s definition of electronegativity  = (I + A)/2 Softness  = 1/

38. Chapter 6 p189

39. Chapter 6 p189

40. Chapter 6 p191

41. Measurement of Acid-Base Interactions • Change in boiling points • Direct calorimetric methods or temperature dependent of equivalent constants can be used to measure enthalpies and entropies • Gas phase measurements of the formation of protonated species • IR • NMR • UV-vis

42. Thermodynamic measurements The enthalpy and entropy of ionization of a weak acid HA can be found by measuring (1) the enthalpy of reaction with NaOH, (2) the enthalpy of reaction of a strong acid (HCl) with NaOH and the equivalent constant for dissociation of the acid. (1) HA + OH-  A- + H2O H1o (2) H3O+ + OH- 2H2O H2o Ka (3) HA + H2O ⇋ H3O+ A- Ka H3o H3o = H1o - H2o S3o = S1o - S2o G3o = -RTlnKa = H3o - T S3o Ln Ka = - H3o/RT + S3o /R On a plot of Ka vs 1/T, the slope is - H3o/R and the intercept is S3o/R

43. Chapter 6 p193

44. Proton Affinity BH+(g) B(g) + H+ proton affinity = H A large proton affinity means it is difficult yo remove the hydrogen ion; this means that B is a strong base and BH+ is a weak acid . 1. The alkali metal hydroxide, which are of equal basicity in aqueous solution have gas phase basicities in the order LiOH < NaOH < KOH <CsOH. This order matches the increase in the electron-releasing ability of the cation in these hydroxides. 2. Pyridine and analine are stronger base the ammonia in the gas phase, but they are weaker than ammonia in aqueous solution., presumably because the interaction of the ammonium ion with water is more favorable than the interaction with pyridinium or anilinium ions,

45. In binary acids, such as the hydrogen halides, the strength of the acid is determined by the strength of the H–X bond. For a series such as the hydrogen halides, the strength of the H–X bond decreases as the size of X increases. In terms of acidic strength, HF < HCl < HBr < HI As we move from left to right across a row in the periodic table, there is less change in bond strength. In this case, what determines acid strength is the polarity of the H- X bond. The electronegativity of elements increases from left to right across a period in the periodic table. As the electronegativity of X increases, the polarity of the H- X bond increases, increasing acidity. Acidity of the second row hydrides varies as CH4 < NH3 < H2O < HF

46. Chapter 6 p195

47. Inductive effects An atom like fluorine which can pull the bonding pair away from the atom it is attached to is said to have a negative inductive effect. Most atoms that you will come across have a negative inductive effect when they are attached to a carbon atom, because they are mostly more electronegative than carbon. You will come across some groups of atoms which have a slight positive inductive effect - they "push" electrons towards the carbon they are attached to, making it slightly negative. Inductive effects are sometimes given symbols: -I (a negative inductive effect) and +I (a positive inductive effect).

48. Inductive effects (electron releasing and withdrawing) PF3 is a much weaker base than PH3 Base strength NMe3 > NHMe2 > NH2Me > NH3 But the acid strength BF3 < BCl3≤ BBr3 BF3 and BCl3 have significant  bonding that increase the electron density on B

49. Inductive Effect Effect Acid Base Examples weaken Electron withdrawing by inductive effect (-I) strengthen electron releasing by inductive effect (+I) weaken strengthen

50. Strength of Oxyacids Many bases contain the OH- ion, but O–H groups are found in acids as well. Whether an OH compound is a base or an acid depends on whether the OH groups are combined with a metal or a nonmetal. For instance, sodium is a metal; NaOH is a base. Chlorine is a nonmetal; HClO is an acid. Acids that contain one or more O–H bonds are called oxyacids. The strength of an oxyacid depends on the electronegativity of the central nonmetal to which the OH groups are bound and on the number of oxygen atoms bound to the central nonmetal atom. For a series of oxyacids with the same number of oxygen atoms, the acidity increases with the electronegativity of the nonmetal. The table below gives such a series and the corresponding Ka values.