1. Inorganic chemistry I ( Chem-212 ) Chapter three Acid -Base definitions By Mengstu Etay 2011 Ec 1

2. Specific objectives  At the end of this lesson you will able to To describe the Arrhenius, Bronsted-Lowery and Lewis definition of acids and bases. To identify conjugate acid-base pairs. Compare and contrast HSAB. 2

3. Acid – base definition 3

4. ARRHENIUS DEFINITION Arrhenius acid as a substance that yield hydrogen ions (H+) in aqueous solution. Hydrogen chloric acid Hydrogen Chloride ion ion Arrhenius base as a substance that yield hydroxide ions (OH - ) in aqueous solution. 4

5. Cont… The ions accompanying the hydrogen and hydroxide ions form a salt, so the overall Arrhenius acid-base reaction can be written. 5

6. The Arrhenius definition has been subjected to many objections including the following: 1.He defined an acid or a base in terms of hydrogen or hydroxyl compounds only. 2. It offered no explanation for the acid property of various substances such as AlCl 3, NH 4 NO 3. 3. A bare proton cannot exist in solution. 4. This excludes non-aqueous solvents. 6

7. Brønsted–Lowry definition of acids and bases An acid is a substance that donates a proton. A base is any substance that accepts a proton. An extension of the acids and bases with the concept of the conjugate acid-base pair, – When a Brønsted-Lowry acid donates a hydrogen ion, a conjugate base is formed. – When a Brønsted-Lowry base accepts a hydrogen ion, a conjugate acid is formed. 7

8. The conjugate base of the pair has one fewer H and one more negative charge than the acid. The conjugate acid of the pair has one more H and one less negative charge than the base. A Bronsted-Lowry acid-base reaction occurs when an acid and a base react to form their conjugate base and conjugate acid, respectively. acid 1 + base 2 base 1 + acid 2

9. 9 This definition also introduced the concept of conjugate acids and bases, differing only in the presence or absence of a proton, and described all reactions as occurring between a stronger acid and base to form a weaker acid and base: Reactions in nonaqueous solvents having ionizable hydrogens parallel those in water. An example of such a solvent is liquid ammonia, in which NH 4 C1 and NaNH 2 react as the acid NH 4 + and the base NH 2 - to form NH 3, which is both a conjugate base and a conjugate acid:

10. examples 10

11. Table 1 Conjugate Pairs in Some Acid-Base Reactions base 1 acid 2 + acid 1 base 2 + conjugate pair reaction 4 H 2 PO 4 - OH - + reaction 5 H 2 SO 4 N2H5+N2H5+ + reaction 6 HPO 4 2 - SO 3 2 - + reaction 1HFH2OH2O+ F-F- H3O+H3O+ + reaction 3 NH 4 + CO 3 2 - + reaction 2HCOOH CN - + HCOO - HCN+ NH 3 HCO 3 - + HPO 4 2 - H2OH2O+ HSO 4 - N 2 H 6 2+ + PO 4 3 - HSO 3 - + Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

12. SAMPLE PROBLEM 18.4 Identifying conjugate acid-base pairs PROBLEM: The following reactions are important environmental processes. Identify the conjugate acid-base pairs. (a) H 2 PO 4 - ( aq ) + CO 3 2 - ( aq ) HPO 4 2 - ( aq ) + HCO 3 - ( aq ) (b) H 2 O( l ) + SO 3 2 - ( aq ) OH - ( aq ) + HSO 3 - ( aq ) SOLUTION: PLAN: Identify proton donors (acids) and proton acceptors (bases). (a) H 2 PO 4 - ( aq ) + CO 3 2 - ( aq ) HPO 4 2 - ( aq ) + HCO 3 - ( aq ) proton donor proton acceptor proton donor conjugate pair 1 conjugate pair 2 (b) H 2 O( l ) + SO 3 2 - ( aq ) OH - ( aq ) + HSO 3 - ( aq ) conjugate pair 2 conjugate pair 1 proton donor proton acceptor proton donor

13. several conclusions regarding acids and bases according to the Brønsted–Lowry theory. The stronger an acid is, the weaker its conjugate will be as a base. The stronger a base is, the weaker its conjugate will be as an acid. A stronger acid reacts to displace a weaker acid. A stronger base reacts to displace a weaker base. 13

14. The Advantages of the Brønsted-Lowry definition over Arrhenius definition  It expands the list of acids to include positive and negative ions, as well as neutral molecules.  It expands the list of bases to include any molecule/ion with at least one pair of nonbonding valence electrons.  It expands to non-aqueous solvents.  It links acids and bases into conjugate acid-base pairs.  It relates strengths of an acid with its conjugate base. 14

15. LEWIS DEFINITION The Brønsted–Lowry definition of acids and bases focuses on the transfer of a proton between species. a base as an electron-pair donor and an acid as an electron-pair acceptor. A is Lewis acid and B is Lewis base; Coordinate covalent bond is formed between the A and B species 15

16. Examples of Lewis acid. a)A molecule with an incomplete octet of valence electrons can complete its octet by accepting an electron pair. (b) A metal cation can accept an electron pair supplied by the base (ligand) in a coordination compound. Eg. ammonia-silver adduct Ag + Lewis acid, NH 3 is Lewis base (Ligand) 16

17. 17

18. Lewis bases include the following types of species: 1. Anions that have an unshared pair of electrons (e.g., OH-, H-, F-, PO4 3- ). 2. Neutral molecules that have unshared pairs of electrons (e.g., NH3, H2O, R3N, ROH, PH3 ). 18 Lewis bases

19. Activities 19

20. Solutions (a) (b) Hg 2+ is the Lewis acid and CN - is the Lewis base. 20

21. exercise 21

22. SOLVENT SYSTEM CONCEPT The solvent system definition applies to any solvent that can dissociate into a cation and an anion (autodissociation), where the cation resulting from autodissociation of the solvent is the acid and the anion is the base. Solutes that increase the concentration of the cation of the solvent are considered acids and solutes that increase the concentration of the anion are considered bases. 22

23. 23 For example: The classic solvent system is water, which undergoes autodissociation By the solvent system definition, the cation, H 3 O + is' the acid and the anion, OH -, is the base. For example, in the reaction sulfuric acid increases the concentration of the hydronium ion and is an acid by any of the three definitions given.

24. 24 The solvent system approach can also be used with solvents that do not contain hydrogen. For example, BrF 3 also undergoes autodissociation: Solutes that increase the concentration of the acid, BrF 2 + are considered acids. For example, SbFs is an acid in BrF 3 : and solutes such as KF that increase the concentration of BrF 4 - are considered bases:

25. 25 Acid-base reactions in the solvent system concept are the reverse of autodissociation: The Arrhenius, Bronsted-Lowry, and solvent system neutralization reactions can be compared as follows:

26. Cont… Caution is needed in interpreting these reactions. For example, SOCl 2 and SO 3 2- react as acid and base in SO 2 solvent, with the reaction apparently occurring as It was at first believed that SOCl 2 dissociated and that the resulting SO 2+ reacted with SO 3 2- However, the reverse reactions should lead to the exchange of oxygen atoms between SO 2 and SOC1 2, but none is observed. The details of the SOCl 2 + SO 3 2- reaction are still uncertain, but may involve dissociation of only one chloride, as in 26

27. 27 Properties of Solvents

28. 28 There have been many attempts to categorize various metal ions and anions to predict reactivity, solubility, etc. R.G. Pearson (1963) categorized acids and bases as either hard or soft Theory of Hard and soft acids and bases

29. Much of the hard-soft distinction depends on polarizability, the degree to which a molecule or ion is easily distorted by interaction with other molecules or ions. Electrons in polarizable molecules can be attracted or repelled by charges on other molecules, forming slightly polar species that can then interact with the other molecules. 29

30.  Hard acids/bases are  small in size  nonpolarizable  The electrons held fairly tightly by the nucleus  Soft acids/bases are  large in  polarizable  Hard/Hard and Soft/Soft interactions are the most favorable 30

31. 31 Hard acids react preferentially with hard bases, and soft acids react preferentially with soft bases. Hard-Hard interactions are ionic and Soft-Soft interactions are covalent involving π-bonding as well. For example: Solubility of Lithium Halides LiBr> LiCl> LiI> LiF Li + is a hard ion LiF would be expected to be very ionic and soluble Very favorable hard-hard LiF interaction even overcomes solubility LiBr, LiCl are more soluble because of less favored interactions LiI is out of order because of poor I - solvation

32. 32  Soft Lewis bases: those in which the donar atoms high polarizability, low electronegativity and high reducing power are under soft bases.  Hard Lewis bases: those in which the donar atoms  low polarizability,  high electronegativity and  low reducing power are under hard bases.

33. 33 For example: Hard and. Soft Bases

34. 34  Soft acids are Lewis acids with o nearly full d-electrons, o comparatively larger in size, and o easily polarisable. o mostly heavy metal ions generally associated with low (or zero) oxidation state.  Hard acids are Lewis acids with no d-electrons (since the d-orbitals are either vacant or existent), small in size and not so easily polarisable. These are mostly light metal ions generally associated with high positive oxidation state.

35. Examples of hard acids and soft acids 35

36. Example Is OH- or S 2- more likely to form insoluble salts with 3+ transition metal ions? Which is more likely to form insoluble salts with 2+ transition metal ions? Because OH - is hard and S 2- is soft, OH - is more likely to form insoluble salts with 3+ transition metal ions (hard) and S 2- is more likely to form insoluble salts with 2+ transition metal ions (borderline or soft). 36

37. Problem Predict the solubility (high or low) of silver fluoride, silver iodide, lithium fluoride and lithium iodide using the hard-soft acid/base approach. Identify each Lewis acid and Lewis base, and categorize each as hard or soft.

38. solution

39. Weak Acids and Their Ionization Constants The dissociation of a weak acid is not complete HA(aq) + H 2 O(l)  H 3 O + (aq) + A - (aq) The equilibrium constant for this reaction is the ionization constant of the acid, K a When the acid is stronger, its K a is larger We can calculate the concentrations of each species in equilibrium with the methods that we have seen in the chapter on chemical equilibrium

40. Weak Acids and Their Ionization Constants To calculate equilibrium concentrations, we make ​​the following approximations The concentration of H + (1.0 x 10 -7 M) before adding the acid to the pure water is negligible The amount of acid that dissociates is negligible, to a first approximation Once we find x (the quantity of acid that dissociates), we verify whether x is less than ~ 5% of the initial quantity of acid If yes, this is the value of x and we can calculate all of the concentrations If no, we need to solve for x, without making the second approximation

41. Weak Acids and Their Ionization Constants Example: Calculate the concentrations of H +, of A -, and of non-ionized HA in a solution of 0.20M HA. The value of K a for HA is 2.7 x 10 -4. Solution: [H + ] = x, [A - ] = x, and [HA] = 0.20 - x  0.20 Verify our approximation: Our approximation is acceptable, therefore [H + ] = 7.3 x 10 -3 M [A - ] = 7.3 x 10 -3 M [HA] = 0.20 - 7.3 x 10 -3 = 0.19 M

42. Weak Acids and Their Ionization Constants Example: What is the pH of a 0.122 M solution of a monoacid, HA, for which the value of K a = 5.7 x 10 -4 ? Solution: [H + ] = x, [A - ] = x, et [HA] = 0.122 - x  0.122 Verify our approximation: We cannot make the approximation that [HA]  0.122 M

43. Weak Acids and Their Ionization Constants Solution: [H + ] = x, [A - ] = x, et [HA] = 0.122 - x The second solution is not acceptable Thus, [H + ] = 8.06 x 10 -3 M, and the pH = -log(8.06 x 10 -3 ) = 2.09 or

44. Weak Acids and Their Ionization Constants Example: A 0.060 M of HA has a pH equal to 3.44. Calculate the value of K a for this acid. Solution: If pH = 3.44, [H + ] = 10 -3.44 = 3.63 x 10 -4 M [A - ] = [H + ] = 3.63 x 10 -4 M [HA] = 0.060 - 3.63 x 10 -4 = 0.0596 M

45. Percent Ionization The percent ionization is defined by For an acid that donates a single proton where [HA] o is the initial concentration of the acid The percent ionization decreases when the acid becomes more concentrated

46. Weak Bases and Their Ionization Constants We treat weak bases the same way we treat weak acids ex.; NH 3 (aq) + H 2 O(l)  NH 4 + (aq) + OH - (aq) K b is the ionization constant of a base

47. Weak Bases and Their Ionization Constants Example: Calculate the pH of a 0.26 M solution of methylamine (CH 3 NH 2 ). K b = 4.4 x 10 -4 for methylamine. CH 3 NH 2 (aq) + H 2 O(l)  CH 3 NH 3 + (aq) + OH - (aq) Solution: [CH 3 NH 3 + ] = x, [OH - ] = x, and [CH 3 NH 2 ] = 0.26 - x  0.26 Verify our approximation: Our approximation is acceptable, therefore pOH = -log(0.011) = 1.96, and therefore the pH = 14.00 - 1.96 = 12.04

48. The Relationship Between the Ionization Constants in Conjugate Acid-Base Pairs For a weak acid For its conjugate base The product of the two ionization constants gives K water

49. The Relationship Between the Ionization Constants in Conjugate Acid-Base Pairs The fact that K a K b = K water is not surprising because the sum of the two reactions is The stronger an acid becomes, the weaker its conjugate base becomes, and vice versa Example: For acetic acid, K a = 1.8 x 10 -5. What is the value of K b for the acetate anion (its conjugate base)? Solution:

50. The end!! 50