1. 15 - 1 Arrhenius Definition An acid is a substance that increases the hydrogen (hydronium) concentration in a water solution.  HCl(aq) H + (aq) + Cl - (aq)  HCl(aq) + H 2 O(l) H 3 O + (aq) + Cl - (aq) Either equation is acceptable and H + (aq) or H 3 O + (aq) is a hydrated proton.

2. 15 - 2 H + is very strongly hydrated in water because of its small size and high positive charge density.  H + (aq) + H 2 O(l) H 3 O + (aq) Arrhenius definitions are limited to aqueous solutions.

3. 15 - 3 A base is a substance that increases the hydroxide ions in a water solution.  NaOH(aq) Na + (aq) + OH - (aq) Remember that Arrhenius definitions are limited to aqueous solutions.

4. 15 - 4 Bronsted-Lowry Theory (BLT) An acid is a molecule or ion that donates a proton.  HCl(aq) H + (aq) + Cl - (aq) A base is a molecule or ion that accepts a proton.  NH 3 (aq) + H 2 O(l) NH 4 + (aq) + OH - (aq)

5. 15 - 5 The BLT of an acid and a base are not limited to aqueous solutions. When working with BLT, it is common to use the terms conjugate acid and conjugate base.  NH 3 (aq) + H 2 O(l) NH 4 + (aq) + OH - (aq) base acid conjugate conjugate acid base

6. 15 - 6 H 2 SO 4 (aq) + H 2 O(l) H 3 O + (aq) + HSO 4 - (aq) acid base conjugate conjugate acid base HSO 4 - (aq) + H 2 O(l) H 3 O + (aq) + SO 4 2- (aq) acid base conjugate conjugate acid base HSO 4 - is called amphoteric or amphiprotic because it can act as either a BLT acid or a BLT base depending on its chemical environment.

7. 15 - 7 Here HSO 4 - is acting as the conjugate base of H 2 SO 4. H 2 SO 4 (aq) + H 2 O(l) H 3 O + (aq) + HSO 4 - (aq) Here HSO 4 - is acting as the conjugate acid of SO 4 2-. HSO 4 - (aq) + H 2 O(l) H 3 O + (aq) + SO 4 2- (aq) Each acid has one more proton than its conjugate base.

8. 15 - 8 Each base has one less proton than its conjugate acid. Two important points to remember:  The stronger the acid, the weaker its conjugate base.  The stronger the base, the weaker its conjugate acid.

9. 15 - 9 Acid-Base Reactions An acid-base reaction always proceeds toward the weaker acid and weaker base. HClO 4 (aq) + H 2 O(l) H 3 O + (aq) + ClO 4 - (aq) stronger stronger weaker weaker acid base

10. 15 - 10 When analyzing an acid-base reaction, remember that you can’t have it both ways. ClO 4 - is too weak a base to compete with the stronger base, H 2 O, to acquire the proton.  How do you know that ClO 4 - is such a weak base? Because HClO 4 is one of the six strong acids. When you have a strong acid such as HClO 4, 100% ionization is assumed.

11. 15 - 11 If a molecule wants to completely ionize, why would its anion want to undergo hydrolysis? Similarly, H 3 O + is too weak an acid to compete with the stronger acid, HClO 4, to donate a proton.

12. 15 - 12 Lewis Acids and Bases An acid is a substance that accepts an electron pair.  Al 3+ (aq) + 6H 2 O(l) Al(H 2 O) 6 3+ (aq) The Al 3+ cation has the empty orbitals 3s, 3p x, 3p y, 3p z, as well as the size to accommodate d-orbitals. Also, the Al 3+ has a large positive charge density resulting in its interaction with water molecules.

13. 15 - 13 The Al 3+ cation acts as a Lewis acid and the water with its two unshared pair of electrons acts as a Lewis base (an electron pair donor). The hydrated Al 3+ cation, Al(H 2 O) 6 3+, can now behave as an Arrhenius acid or a Bronsted-Lowry acid. Al(H 2 O) 6 3+ (aq) H + (aq) + Al(H 2 O) 5 (OH) 2+ (aq)

14. 15 - 14 A base is a substance that donates an electron pair.  Zn 2+ (aq) + 4OH - (aq) Zn(OH) 4 2- (aq) The Lewis definition of acids and bases expands the number of species that can be acids.

15. 15 - 15 Strength of Binary Acids The H – X bond strength is the most important factor to consider when determining acid strength in a group or family. Consider the following bond enthalpies:  H – F567 kJ mol -1  H – Cl431 kJ mol -1  H – Br366 kJ mol -1  H – I299 kJ mol -1

16. 15 - 16 The bond enthalpies from the previous slide indicate that the strength of the H – X bond decreases as the atomic radii of the halogen increases.  Longer bonds are generally weaker or less stable than shorter bonds. Similarly, H 2 S is a stronger acid than H 2 O, Ka(H 2 S) > Ka(H 2 O).

17. 15 - 17 The H – X bond polarity is the most important factor to consider when determining acid strength in a period or series.  Because electronegativity increases from left to right in a period, the acid strength also increases proceeding from left to right.  The acid strength increases from left to right in a period.  Ka(HF) > Ka(H 2 O) > Ka(NH 3 ) > Ka(CH 4 )

18. 15 - 18 Strength of Oxyacids When comparing oxyacids, there are additional factors to consider. – X – O – H │ │ As the electronegativity of element X increases, the stronger the acid. When the electronegativity of X increases, the polarizability of the O – H bond increases.

19. 15 - 19 As more O terminal atoms are added to the central atom, X, the more the electron density is pulled from the O – H bond. By adding more electronegative atoms to X, the acid strength is increased. Cl – O – H < Cl – O – H < Cl – O – H < O – Cl – O – H │ │ │ │ │

20. 15 - 20 By adding more oxygens (the second most electronegative element) to the central atom, Cl, the electron density shifts more towards the oxygens, making the O – H bond more polarizable. For oxoacids with the same number of O – H bonds and the same number of oxygen atoms, the acid strength will increase with an increase of electronegativity of the central atom.

21. 15 - 21 O – Cl – O – H > O – Br – O – H > O – I – O – H │ │ │ │ │ │

22. 15 - 22 Oxoacid Wrap Up For an oxoacid, the H atom that ionizes is bonded to an O atom which in turn is bonded to a nonmetal atom. The strength of any acid depends on how easily the O – H bond is broken.  One deciding factor is the oxidation of the central atom.  The higher the oxidation number the stronger the acid.

23. 15 - 23 To increase the ionization, the electron density surrounding the O atom which is bonded to the ionizable H, should be as low as possible. To decrease the electron density around the O atom:  Make the central atom more electronegative.  Add more O atoms to the central atom.

24. 15 - 24 A second deciding factor is the electronegativity of the central atom.  The more electronegative the central atom, the stronger the acid.

25. 15 - 25 Extent of Hydrolysis There are six strong acids There are six strong acids that completely ionize in water.  HCl, HBr, HI, HNO 3, HClO 4, H 2 SO 4 When representing the ionization of these acids, a single arrow is used and 100% ionization is assumed.  HClO 4 (aq) H + (aq) + ClO 4 - (aq) The K a of these acids is assumed to be infinite.

26. 15 - 26 Most acids are weak and only partially ionize in water.  HC 2 H 3 O 2 (aq) + 2H 2 O(l) H 3 O + (aq) + C 2 H 3 O 2 - (aq) Or alternatively  HC 2 H 3 O 2 (aq) H + (aq) + C 2 H 3 O 2 - (aq) The K a of these acids is small and can be looked up for each acid.

27. 15 - 27 Also note that with weak acids, a double arrow is used and a dynamic equilibrium results. The most commonly encountered bases The most commonly encountered bases that completely dissociate in water are:  LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH) 2, Sr(OH) 2, Ba(OH) 2 When representing the dissociation of these bases, a single arrow is used and 100% dissociation is assumed.  LiOH(aq) Li + (aq) + OH - (aq)

28. 15 - 28 The K b of these bases is assumed to be infinite. Most bases are weak and only partial ionization takes place.  NH 3 (aq) + H 2 O(l) NH 4 + (aq) + OH - (aq) The K b of these bases is small and can be looked up for each base. Note that in the case of a weak base, water must be explicitly written as a reactant unlike the case of a weak acid.