1. by Steven S. Zumdahl & Donald J. DeCoste University of Illinois Introductory Chemistry: A Foundation, 6 th Ed. Introductory Chemistry, 6 th Ed. Basic Chemistry, 6 th Ed.

2. Chapter 16 Acids and Bases

3. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 3 Properties of Acids Sour taste Change color of vegetable dyes React with “active” metals –Like Al, Zn, Fe, but not Cu, Ag or Au Zn + 2 HCl  ZnCl 2 + H 2 Corrosive React with carbonates, producing CO 2 –Marble, baking soda, chalk CaCO 3 + 2 HCl  CaCl 2 + CO 2 + H 2 O React with bases to form ionic salts, and often water

4. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 4 Properties of Bases Also known as alkalis Bitter Taste Feel slippery Change color of vegetable dyes –Different color than acid –Litmus = blue React with acids to form ionic salts, and often water –Neutralization

5. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 5 Arrhenius Theory Acids ionize in water to H + ions and anions Bases ionize in water to OH - ions and cations Neutralization reaction involves H + combining with OH - to make water H + ions are protons

6. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 6 Definition only good in water solution Definition does not explain why ammonia solutions turn litmus blue –Basic without OH - ions Arrhenius Theory (cont.)

7. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 7 Brønsted-Lowery Theory H + transfer reaction –Since H + is a proton, also known as proton transfer reactions In the reaction, a proton from the acid molecule is transferred to the base molecule Products are called the conjugate acid and conjugate base

8. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 8 Brønsted-Lowery Theory (cont.)

9. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 9 Brønsted-Lowery Theory (cont.) H-A + :B  A - + H-B + A - is the conjugate base, H-B + is the conjugate acid Conjugate acid-base pair is either the original acid and its conjugate base or the original base and its conjugate acid –H-A and A - are a conjugate acid-base pair –:B and H-B + are a conjugate acid-base pair

10. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 10 Example #1: Determine what species you will get if you remove 1 H +1 from the acid. –Conjugate base will have one more negative charge than the original acid H 3 PO 4  H + + H 2 PO 4 - Write the conjugate base for the acid H 3 PO 4

11. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 11 Brønsted-Lowery Theory (cont.) In this theory, instead of the acid, HA, dissociating into H + (aq) and A - (aq), the acid donates its H to a water molecule HA + H 2 O  A - + H 3 O + A -1 is the conjugate base, H 3 O + is the conjugate acid

12. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 12 Brønsted-Lowery Theory (cont.) H 3 O + is called the hydronium ion In this theory, substances that do not have OH - ions can act as a base if they can accept a H +1 from water. H 2 O + :B  OH - + H-B +

13. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 13 Strength of Acids & Bases The stronger the acid, the more willing it is to donate H +

14. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 14 Strength of Acids & Bases (cont.) Strong bases will react completely with water to form hydroxide: CO 3 -2 + H 2 O  HCO 3 - + OH - Only small fraction of weak base molecules pull H + off water: HCO 3 - + H 2 O  H 2 CO 3 + OH -

15. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 15 Multiprotic Acids Monoprotic acids have 1 acid H, diprotic 2, etc. –In oxyacids only the H on the O is acidic In strong multiprotic acids, like H 2 SO 4, only the first H is strong; transferring the second H is usually weak H 2 SO 4 + H 2 O  H 3 O + + HSO 4 - HSO 4 - + H 2 O  H 3 O + + SO 4 -2

16. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 16 Water As an Acid and a Base Amphoteric substances can act as either an acid or a base. –Water as an acid, NH 3 + H 2 O  NH 4 + + OH - –Water as a base, HCl + H 2 O  H 3 O + + Cl - Water can even react with itself: H 2 O + H 2 O  H 3 O + + OH -

17. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 17 Autoionization of Water Water is an extremely weak electrolyte. –Therefore there must be a few ions present H 2 O + H 2 O  H 3 O + + OH -

18. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 18 Autoionization of Water (cont.)

19. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 19 Acidic and Basic Solutions Acidic solutions have a larger [H + ] than [OH - ] Basic solutions have a larger [OH - ] than [H + ] Neutral solutions have [H + ]=[OH - ]= 1 x 10 -7 M [H + ] = 1 x 10 -14 [OH - ] [OH - ] = 1 x 10 -14 [H + ]

20. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 20 Example #2 Determine the [H + ] and [OH - ] in a 10.0 M H + solution

21. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 21 Example #2 (cont.) Determine the given information and the information you need to find –Given [H + ] = 10.0 M, find [OH - ]

22. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 22 Given [H + ] = 10.0 M = 1.00 x 10 1 M K w = 1.0 x 10 -14 Example #2 (cont.)

23. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 23 pH & pOH The acidity/basicity of a solution is often expressed as pH or pOH. pH = -log[H 3 O + ]pOH = -log[OH - ] –pH water = -log[10 -7 ] = 7 = pOH water [H + ] = 10 -pH [OH - ] = 10 -pOH

24. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 24 pH & pOH (cont.) pH 7 is basic, pH = 7 is neutral The lower the pH, the more acidic the solution; the higher the pH, the more basic the solution 1 pH unit corresponds to a factor of 10 difference in acidity pOH = 14 - pH

25. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 25 Example #3 Calculate the pH of a solution with a [OH - ] = 1.0 x 10 -6 M

26. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 26 Example #3 (cont.) Find the concentration of [H + ]

27. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 27 Enter the [H + ] concentration into your calculator and press the log key –log(1.0 x 10 -8 ) = -8.0 Change the sign to get the pH –pH = -(-8.0) = 8.0 Example #3 (cont.)

28. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 28 Enter the [H + ] or [OH - ] concentration into your calculator and press the log key log(1.0 x 10 -3 ) = -3.0 Change the sign to get the pOH pOH = -(-3) = 3.0 Subtract the calculated pH or pOH from 14.00 to get the other value pH = 14.00 – 3.0 = 11.0 Calculate the pH and pOH of a solution with a [OH - ] = 1.0 x 10 -3 M Example #4

29. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 29 If you want to calculate [OH - ] use pOH; if you want [H + ] use pH. It may be necessary to convert one to the other using 14 = [H + ] + [OH - ] pOH = 14.00 – 7.41 = 6.59 Calculate the [OH - ] of a solution with a pH of 7.41 Example #5

30. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 30 Example #5 (cont.) Enter the pH or pOH concentration into your calculator Change the sign of the pH or pOH -pOH = -(6.59) Press the button(s) on you calculator to take the inverse log or 10 x [OH - ] = 10 -6.59 = 2.6 x 10 -7 M

31. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 31 Calculating the pH of a Strong, Monoprotic Acid A strong acid will dissociate 100% HA  H + + A - Therefore the molarity of H + ions will be the same as the molarity of the acid Once the H + molarity is determined, the pH can be determined pH = -log[H + ]

32. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 32 Example #6 Calculate the pH of a 0.10 M HNO 3 solution.

33. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 33 Example #6 (cont.) Determine the [H + ] from the acid concentration HNO 3  H + + NO 3 - 0.10 M HNO 3 = 0.10 M H + Enter the [H + ] concentration into your calculator and press the log key log(0.10) = -1.00 Change the sign to get the pH pH = -(-1.00) = 1.00

34. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 34 Buffered Solutions Buffered solutions resist change in pH when an acid or base is added to it. Used when need to maintain a certain pH in the system –Blood

35. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 35 Buffered Solutions (cont.) A buffer solution contains a weak acid and its conjugate base. Buffers work by reacting with added H + or OH - ions so they do not accumulate and change the pH. Buffers will only work as long as there are sufficient weak acid and conjugate base molecules present.

36. Copyright © Houghton Mifflin Company. All rights reserved. 16 | 36 Buffered Solutions (cont.)