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

2. Chapter 16 Acids and Bases

5. CaCO3 + 2 HCl CaCl2 + CO2 + H2O
Properties of Acids
Sour taste
Turn blue litmus paper red
Change color of vegetable dyes (red cabbage juice)
React with “active” metals
Like Al, Zn, Fe, but not Cu, Ag or Au
Zn + 2 HCl ZnCl2 + H2
Corrosive
React with carbonates, producing CO2
Marble, baking soda, chalk
CaCO3 + 2 HCl CaCl2 + CO2 + H2O
React with bases to form ionic salts, and often water 2

6. Properties of Bases Also known as alkalis Bitter Taste Feel slippery
Change color of vegetable dyes
Different color than acid
Turn red litmus blue
React with acids to form ionic salts, and often water
Neutralization 3

7. 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 4

8. Arrhenius Theory (cont.)
Definition only good in water solution
Definition does not explain why ammonia solutions turn litmus blue
Basic without OH- ions

9. Brønsted-Lowry Theory
H+ transfer reaction
Since H+ is a proton, also known as proton transfer reactions
Acids are proton donors, bases are proton acceptors
In the reaction, a proton from the acid molecule is transferred to the base molecule
Products are called the conjugate acid and conjugate base 5

10. Brønsted-Lowry Theory (cont.)

11. Brønsted-Lowry 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 6

12. Example #1: Write the conjugate base for the acid H3PO4
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
H3PO4  H+ + H2PO4- 7

13. Self- check p 490
Which of the following represent conjugate acid base pairs?
A. H2O, H3O+
B. OH-, HNO3
C. H2SO4, SO42-
D. HC2H3O2, C2H3O2-

14. 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 + H2O  A- + H3O+
A-1 is the conjugate base, H3O+ is the conjugate acid 8

15. Brønsted-Lowry Theory (cont.)
H3O+ 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.
H2O + :B  OH- + H-B+
:B is acting here as a base.

16. Strength of Acids & Bases
The stronger the acid, the more willing it is to donate H+ 9

17. Strength of Acids & Bases (cont.)
Strong bases will react completely with water to form hydroxide:
CO3-2 + H2O HCO3- + OH-
Only small fraction of weak base molecules pull H+ off water:
HCO3- + H2O H2CO3 + OH-

18. 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 H2SO4, only the first H is strong; transferring the second H is usually weak
H2SO4 + H2O  H3O+ + HSO4-
HSO4- + H2O  H3O+ + SO4-2 10

19. Water As an Acid and a Base
Amphoteric substances can act as either an acid or a base.
Water as an acid, NH3 + H2O  NH4+ + OH-
Water as a base, HCl + H2O  H3O+ + Cl-
Water can even react with itself:
H2O + H2O  H3O + + OH- 11

20. Autoionization of Water
Water is an extremely weak electrolyte.
Therefore there must be a few ions present
H2O + H2O  H3O+ + OH- 12

21. Acid Nomenclature Acids Examples: Compounds that form H+ in water.
Formulas usually begin with ‘H’.
In order to be an acid instead of a gas, binary acids must be aqueous (dissolved in water)
Ternary acids are ALL aqueous
Examples:
HCl (aq) – hydrochloric acid
HNO3 – nitric acid
H2SO4 – sulfuric acid

22. Acid nomenclature
If anion ending is –ide (Binary compound), the acid name is hydro(stem)ic acid
If ternary compounds-
-Ate ending: (stem)ic acid
-ite ending: (stem)ous acid

23. Acid Nomenclature Flowchart

24. Solved examples HBr – 2 elements-ide, hydrobromic acid
H2CO3- 3 elements- ate, carbonic acid
H2SO3- 3 elements- ite, sulfurous acid
Hydrofluoric acid: 2 elements= HF
Sulfuric acid: 3 elements, –ic= -ate, H2SO4
Nitrous acid: 3 elements, -ous= -ite, HNO2

25. Now your turn!
HI (aq)
HCl
H2SO3
HNO3
HIO4

26. Hydrobromic acid
Nitrous acid
Carbonic acid
Phosphoric acid

27. 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-] =
[H+] 13

28. Ion product of water

29. Example #2
Determine the [H+] and [OH-] in a
10.0 M H+ solution 14

30. Example #2 (cont.)
Determine the given information and the information you need to find
Given [H+] = 10.0 M, find [OH-]

31. Example #2 (cont.) Given [H+] = 10.0 M = 1.00 x 101 M Kw = 1.0 x 10-1415

32. Self check p 497 Calculate [H+] in a solution in which [OH-] =
2.0X 10-2 M. Is this solution acidic, neutral or basic?

33. pH & pOH
The acidity/basicity of a solution is often expressed as pH or pOH.
pH = -log[H3O+] pOH = -log[OH-]
pHwater = -log[10-7] = 7 = pOHwater
[H+] = 10-pH [OH-] = 10-pOH 16

34. pH scales

35. pH & pOH (cont.)
pH < 7 is acidic; 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

36. pH of Common Substances

37. Example #3 Calculate the pH of a solution with a [OH-] = 1.0 x 10-6 M17

38. Example #3 (cont.)
Find the concentration of [H+]

39. Example #3 (cont.)
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 18

40. Example #4
Calculate the pH and pOH of a solution with a [OH-] = 1.0 x 10-3 M
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 to get the other value
pH = – 3.0 = 11.0 19

41. Solving concentration from pH or pOH
Calculate the [OH-] of a solution with a pH of 7.41
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 = – 7.41 = 6.59 20

42. 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 10x
[OH-] = = 2.6 x 10-7 M

43. Self check p 499
Calculate the pH value for each of the following solutions at 25°C.
A. a solution in which [H+] =1.0 X 10-9M.
B. a solution in which [OH-] =1.0 X 10-6M.
P501- A sample of rain in an area with severe air pollution has a pH of 3.5. What is the pOH of this rainwater?

44. 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+] 21

45. Example #6
Calculate the pH of a 0.10 M HNO3 solution. 22

46. Example #6 (cont.) Determine the [H+] from the acid concentration
HNO3  H+ + NO3-
0.10 M HNO3 = 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

47. Self check p503
The pH of rainwater in a polluted area was measured to be 3.5. What is the [H+] in this rainwater?
The pOH of a liquid drain cleaner was found to be What is the [OH-] for this cleaner?
P505- Calculate the pH of a solution of 5.0 X 10-3 M HCl.

48. 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 23

49. 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.

50. Buffered Solutions (cont.)