1. Unit 13 Acids and Bases

2. D. Finding the pH of Solutions Self- ionization of water – the simple dissociation of water H 2 O H + + OH - Concentration of each ion in pure water: [H + ] = 1.0 x 10 -7 M + [OH - ] = 1.0 x 10 -7 M Where K w = 1.0 x 10 -14 Ion-product constant for water (K w ), Where K w = 1.0 x 10 -14 K w = [H + ] [OH - ] Acid [H + ] > [OH - ] Base [H + ] < [OH - ] Neutral [H + ] = [OH - ]

3. [OH - ]pOHpH[H + ] 1 x 10 -14 1401 x 10 0 1 x 10 -13 1311 x 10 -1 1 x 10 -12 1221 x 10 -2 1 x 10 -11 1131 x 10 -3 1 x 10 -10 10Increasing acidity41 x 10 -4 1 x 10 -9 951 x 10 -5 1 x 10 -8 861 x 10 -6 1 x 10 -7 7Neutral71 x 10 -7 1 x 10 -6 681 x 10 -8 1 x 10 -5 591 x 10 -9 1 x 10 -4 4Increasing basicity101 x 10 -10 1 x 10 -3 3111 x 10 -11 1 x 10 -2 2121 x 10 -12 1 x 10 -1 1131 x 10 -13 1 x 10 0 0141 x 10 -14

4. Calculating [H+] and [OH-] reverse the pH equation The pH of a solution is 7.52. Find the [H+] and [OH-] and determine whether it is acidic, basic, or neutral. –basic

5. Example 1. If the [H + ] in a solution is 1.0 x 10 -5 M, is the solution acidic, basic or neutral? 1.0 x 10 -5 M What is the concentration of the [OH - ]? Use the ion-product constant for water (K w ): K w = [H + ] [OH - ] 1.0 x 10 -14 = [1.0 x 10 -5 ] [OH - ] 1.0 x 10 -14 = [OH - ] 1.0 x 10 -5 1.0 x 10 -(14-5) pH 5 = acidic 1.0 x 10 -9 OH -

6. Examples 2. If the pH is 9, what is the concentration of the hydroxide ion? K w = [H + ] [OH - ] 1.0 x 10 -14 = [1.0 x 10 -9 ] [OH - ] 1.0 x 10 -5 = [OH - ] 3. If the pOH is 4, what is the concentration of the hydrogen ion? K w = [H + ] [OH - ] 1.0 x 10 -14 = [H + ] [1.0 x 10 -4 ] 1.0 x 10 -10 = [H + ] 14 = pH + pOH 14 = 9 + pOH 5 = pOH 14 = pH + pOH 14 = pH + 4 10 = pH

7. Example A solution has a pH of 4. Calculate the pOH, [H+] and [OH-]. Is it acidic, basic, or neutral? 14 = pH + pOH 14 = 4 + pOH 10 = pOH –acidic

8. Practice Problems: Classify each solution as acidic, basic or neutral. 1. [H + ] = 1.0 x 10 -10 2. [H + ] = 0.001 3. [OH - ] = 1.0 x 10 -7 4. [OH - ] = 1.0 x 10 -4 Basic pH 10 1.0 x 10 -3 acid pH 3 Neutral 14 – 4 = 10 base pH 10

9. Fill in the chart. [OH - ]pOHpH[H + ] 8 1x 10 -12 10 1 x 10 -3 5 1 × 10 -1 1.0 X 10 -8 1.0 X 10 -2 1.0 X 10 -4 1.0 X 10 -6 1.0 X 10 -10 1.0 X 10 -11 6 3 11 4 212 9 113 1.0 X 10 -5 1.0 X 10 -9 1.0 X 10 -13

10. pH = -log[H + ] E. pH Scale 0 7 INCREASING ACIDITY NEUTRAL INCREASING BASICITY 14 pouvoir hydrogène (Fr.) “hydrogen power” pH is the negative logarithm of the hydrogen ion concentration

11. E. pH Scale pH = -log[H + ] pOH = -log[OH - ] pH + pOH = 14

12. E. The pH Scale

13. E. pH Scale pH of Common Substances

14. F. Neutralization Chemical reaction between an acid and a base. Products are a salt (ionic compound) and water.

15. F. Neutralization ACID + BASE  SALT + WATER HCl + NaOH  NaCl + H 2 O HC 2 H 3 O 2 + NaOH  NaC 2 H 3 O 2 + H 2 O –Salts can be neutral, acidic, or basic. –Neutralization does not mean pH = 7. weak strong neutral basic

16. G. Titration TitrationTitration –Analytical method in which a standard solution is used to determine the concentration of an unknown solution. standard solution unknown solution

17. Equivalence point (endpoint)Equivalence point (endpoint) –Point at which equal amounts of H + and OH - have been added. –Determined by… indicator color change G. Titration dramatic change in pH

18. G. Titration moles H + = moles OH - M  V  n = M  V  n M:Molarity V:volume n:# of H + ions in the acid or OH - ions in the base

19. G. Titration 42.5 mL of 1.3M KOH are required to neutralize 50.0 mL of H 2 SO 4. Find the molarity of H 2 SO 4. H3O+H3O+ M = ? V = 50.0 mL n = 2 OH - M = 1.3M V = 42.5 mL n = 1 MV# = MV# M(50.0mL)(2)=(1.3M)(42.5mL)(1) M = 0.55M H 2 SO 4

20. Naming Acids Binary acids –Contains 2 different elements: H and another –Always has “hydro-” prefix –Root of other element’s name –Ending “-ic” –Examples: HI, H 2 S, HBr, HCl

21. Naming Acids Ternary Acids - Oxyacids –Contains 3 different elements: H, O, and another –No prefix –Name of polyatomic ion –Ending “–ic” for “-ate” and “–ous” for “- ite” 3 –Examples: HClO 3, H 3 PO 4, HNO 2

22. Practice H 2 SO 3 –Sulfurous acid HF –Hydrofluoric acid H 2 Se –Hydroselenic acid Perchloric acid –HClO 4 Carbonic acid –H 2 CO 3 Hydrobromic acid –HBr

23. Definitions of Acids and Bases Arrhenius –Most specific/exclusive definition –Created by Svante Arrhenius, Swedish –Acid : compound that creates H + in an aqueous solution –Base : compound that creates OH - in an aqueous solution –HNO 3  H + + NO 3 - –NaOH  Na + + OH -

24. Definitions of Acids and Bases Bronsted-Lowry –A bit more general than Arrhenius definition –Most commonly used definition –Created by two scientists around the same time (1923) –Acid: Molecule or ion that is a proton (H + ) donor –Base: Molecule or ion that is a proton (H + ) acceptor –HCl + H 2 O  H 3 O + + Cl - –NH 3 + H 2 O ↔ NH 4 + + OH -

25. Definitions of Acids and Bases Lewis –The most general definition –Defined by electrons and bonding instead of H + –Created by same scientist who electron-dot diagrams are named after –Acid: atom, ion, or molecule that accepts electron pair to form covalent bond –Base: atom, ion or molecule that donates and electron pair to form covalent bond –NH 3 + Ag +  [Ag(NH 3 ) 2 ] 1+ –BF 3 + F -  BF 4 -