1. Chapter 9 Acids and Bases
CHE 101
Sleevi

2. Properties of Acids Taste sour Produce H+ in water (H3O+)
React with metals to produce H2 (g)
pH <7
Dissolve ionic compounds insoluble in water (e.g., CaCO3)

3. Properties of Bases Solutions containing bases feel soapy or slippery
pH >7
React readily with acids
React with fats and oils and convert them to smaller, soluble molecules
Used as part of cleaning solutions
Often contain hydroxide ion

4. Naming Acids Binary acids: Acids containing polyatomic ions:
Use prefix hydro- and derivative of element name with suffix -ic
HCl = hydrochloric
Acids containing polyatomic ions:
Ion ends in –ate, acid name ends in –ic
Ion ends in –ite, acid name ends in –ous
H2SO4 = sulfuric acid
H2SO3 = sulfurous acid

5. Naming Acids
H2Se
HClO3
H3PO4
HNO2
HBr
HC2H3O2

6. Definition of Acids & Bases Arrhenius Theory
A substance that forms H+ in solution
HCl(aq)  H+(aq) + Cl-(aq)
Base:
A substance that forms OH- in solution
KOH(aq)  K+(aq) + OH-(aq)

7. Arrhenius Acids
Write dissociation equations for the following acids to illustrate their behavior as Arrhenius Acids HBr  HCN  HClO4 

8. Arrhenius Bases
Write dissociation equations for bases the Arrhenius bases. ID those that are not Arrhenius bases LiOH C2H5NH2 Ca(OH)2 NH3

9. Definition of Acids & Bases BrǾnsted-Lowry Theory
Acid (H-A)
donates a proton to another substance
Base (B:)
accepts a proton from another substance
must contain a lone pair of electrons that can be used to form new bond to the proton
NH3, H2O, OH-, Cl-

10. Proton Transfer The Reaction of a Brønsted–Lowry Acid with a Brønsted–Lowry Base
This e− pair forms
a new bond to H+
This e− pair
stays on A
gain of H+ H A + B
A − + H B+
acid
base
loss of H+

11. Definition of Acids & Bases BrǾnsted-Lowry Theory
When a species gains a proton (H+), it gains a +1 charge
When a species loses a proton (H+), it effectively gains a -1 charge

12. Definition of Acids & Bases BrǾnsted-Lowry Theory
Conjugate Acid:
species formed when base gains a proton
Conjugate Base:
species formed when acid loses a proton
H—A(aq) + B:(l)  H—B+ + A-
acid base conjugate conjugate
acid base

13. Conjugate Acids and Bases
H-Br + H2O Br- + H3O+
Conjugate acid-base pairs
HBr and Br-
H2O and H3O+
Equation must be mass and charge balanced!

14. Identifying Acids, Bases and Conjugates
HNO3 (aq) + H2O (l)  H3O+ (aq) + NO3- (aq)
NH3 (aq) + H2O (l)  OH- (aq) + NH4+ (aq)
SO32- (aq) + H2O (l)  HSO3- (aq) + OH- (aq)

15. Identifying Acids, Bases and Conjugates
Write balanced equations:
HClO3 (aq) + NH3 (aq) 
H2SO4 (aq) + H2O (l) 

16. Types of Acids monoprotic acid diprotic acid triprotic acid
gives up only one proton per molecule when dissolved (HCl)
diprotic acid
gives up two protons per molecule when dissolved (H2SO4)
triprotic acid
gives up three protons per molecule when dissolved (H3PO4)

17. 1 mole of HCl reacts with 1 mole of NaOH
Acid Behavior
HCl + NaOH  NaCl + H2O
HCl is monoprotic acid
1 mole of HCl reacts with 1 mole of NaOH

18. Acid Behavior H2SO4 + 2NaOH  Na2SO4 + 2H2O H2SO4 is a diprotic acid
1 mole of H2SO4 reacts with
2 moles of NaOH

19. Amphoteric Compounds Compounds that can be either an acid or base
Contain both H and lone pair
Examples:
H2O
NH3
HSO3-

20. The pH Scale

21. Analyzing Acids and Bases
Determining pH range of an acid or base:
pH meter
acid-base indicator
universal indicator
red, orange, green, blue, purple
litmus paper
blue  pink, pink  blue
phenolphthalein
(acid/neutral  colorless; base  pink)

22. pH pH = - log [H+] [H+] = 1 x 10-pH
Each unit on pH scale is a factor of 10 different from the next lower or higher number

23. Examples

24. Acid and Base Strength Strong acid, strong base
fully dissociated in water
NaOH
HCl

25. Acid and Base Strength Weak acid, weak base
partially dissociated in water
HC2H3O2, NH4+, H2CO3, citric acid, NH3

26. Acid and Base Strength Strong acid forms weak conjugate base
Strong base forms weak conjugate acid
Concentration of acids and bases measured in molarity (moles/L)
Dilute vs weak

27. Acid and Base Strength Predicting the Direction of Equilibrium
When the stronger acid and base are the reactants
on the left side, the reaction readily occurs and
the reaction proceeds to the right. H A + B: A- + H B+
stronger
acid
stronger
base
weaker
base
weaker
acid
A larger forward arrow means that products are
favored.

28. Acid and Base Strength Predicting the Direction of Equilibrium
If an acid–base reaction would form the stronger
acid and base, equilibrium favors the reactants
and little product forms. H A + B: A- + H B+
weaker
acid
weaker
base
stronger
base
stronger
acid
A larger reverse arrow means that reactants are
favored.

29. Predict Direction of Equilibrium
HCN(g) +
−OH(aq)
−CN(aq) +
H2O(l)
See Table Relative Strength of Acids and
Their Conjugate Bases

30. Equilibrium and Acid Dissociation Constants
For the reaction where an acid (HA) dissolves in
water,
HA(g) H2O(l)
H3O+(aq) A:- (aq)
the following equilibrium constant can be written:
[H3O+][A:-]
K =
[HA][H2O]

31. Equilibrium and Acid Dissociation Constants
Multiplying both sides by [H2O] forms a new constant,
called the acid dissociation constant, Ka.
[H3O+][ A:- ]
Ka = K[H2O] =
[HA]
acid dissociation
constant
The stronger the acid, the larger the value of Ka.
Equilibrium, though, favors formation of the weaker acid—that is, the acid with the smaller value of Ka.

32. Equilibrium and Acid Dissociation Constants

33. Dissociation of Water Water can behave as both a Brønsted–Lowry acid
and a Brønsted–Lowry base. Thus, two water molecules can react together in an acid–base reaction:
loss of H+ + H H O − H O H + H O H + H O H
acid
base
conjugate
base
conjugate
acid
gain of H+

34. Dissociation of Water From the reaction of two water molecules, the
following equilibrium constant expression can be
written:
[H3O+][−OH]
K =
[H2O]2
Multiplying both sides by [H2O]2 yields Kw, the
ion-product constant for water.
Kw = [H3O+][−OH]
ion-product
constant

35. Dissociation of Water Experimentally it can be shown that
[H3O+] = [−OH] = 1.0 x 10−7 M at 25 oC
Kw = [H3O+] [−OH]
Kw = (1.0 x 10−7) x (1.0 x 10−7)
Kw = x 10−14
Kw is a constant, 1.0 x 10−14, for all aqueous
solutions at 25 oC.

36. Dissociation of Water To calculate [−OH] when [H3O+] is known:
To calculate [H3O+] when
[−OH] is known:
Kw = [H3O+][−OH]
Kw = [H3O+][−OH] Kw Kw
[−OH] =
[H3O+] =
[H3O+]
[−OH]
1 x 10−14
1 x 10−14
[−OH] =
[H3O+] =
[H3O+]
[−OH]

37. Dissociation of Water
If the [H3O+] in a cup of coffee is 1.0 x 10−5 M, then
the [−OH] can be calculated as follows: Kw
1 x 10−14
[−OH] = = =
1.0 x 10−9 M
[H3O+]
1 x 10−5
In this cup of coffee, therefore, [H3O+] > [–OH], and
the solution is acidic overall.

38. Dissociation of Water

39. Chemical Equations and Solutions
Molecular equation
Each substance represented by its formula
HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)
2 Al (s) + 3 Cu(NO3)2 (aq)  2 Al(NO3)3 (aq) + 3 Cu (s)

40. Chemical Equations and Solutions
Total Ionic Equation
All soluble ionic substances represented by the ions they form in solution
Solids, liquids, gases and aqueous solutions of molecular compounds do not dissociate
HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)
=>H+ (aq) + Cl- (aq) + Na+ (aq) + OH- (aq) 
Na+ (aq) + Cl- (aq) + H2O (l)

41. Spectator Ions
Ions that appear on both sides of the chemical equation (not changed in the chemical reaction)

42. Net Ionic Equation
Contains only unionized or insoluble materials and ions that undergo changes in the reaction
All spectator ions are eliminated
H+ (aq) + Cl- (aq) + Na+ (aq) + OH- (aq) 
Na+ (aq) + Cl- (aq) + H2O (l)
H+ (aq) + OH- (aq)  H2O (l)

43. Net Ionic Equations
2 Al (s) + 3 Cu(NO3)2 (aq)  2 Al(NO3)3 (aq) + 3 Cu (s) 2 Al (s) +3Cu2+ (aq) + 6NO31-(aq)  2 Al3+ (aq) + 6NO31- (aq) + 3 Cu (s) 2 Al (s) +3 Cu2+ (aq)  2 Al3+ (aq) + 3 Cu (s) Net ionic equations must be mass and charge balanced

44. How to Write a Net Ionic Equation
Write balanced molecular equation
Write total ionic equation
Aqueous ionic compounds written as individual ions
Multiply by subscript and coefficient to balance mass and charge
Compounds that appear as solids, liquids, gases or aqueous solutions of molecular compounds are written in molecular form.
Write net ionic equation
Eliminate spectator ions
Include all solids, liquids, gases and non-spectator ions
Verify mass and charge balance

45. Analyzing Acids and Bases
Determine concentration of acid or base using neutralization reactions and titration
Equivalence point – the point at which the acid has exactly neutralized the base (neither is in excess)

46. Performing a Titration
Slowly add base from a buret to an acid in a receiving flask
Use phenolphthalein to indicate when endpoint is reached
Measure volumetric amount of base of known concentration
Calculate concentration of acid using solution stoichiometry

47. Performing a Titration

48. Calculating Unknown Concentration of Acid Solution
Write balanced equation for neutralization reaction
Titrate acid solution with known concentration of base solution (to phenolphthalein endpoint)
Determine accurate volume of base used in neutralization
Calculate concentration of acid solution using solution stoichiometry

49. Examples

50. Buffers
Solutions with ability to resist change in pH when acids or bases are added
Made up of weak acid and salt of its conjugate base
Amount of acid or base that can be added without significant pH change is called the buffer capacity
Common in the human body and nature

51. Change in pH with addition of Acid or Base a. No buffer b. With buffer

52. Common Buffers

53. Focus on the Human Body Buffers in the Blood
Normal blood pH is between 7.35 and 7.45.
The principle buffer in the blood is carbonic acid/
bicarbonate (H2CO3/HCO3−).
H2O
CO2(g) + H2O(l)
H2CO3(aq)
H3O+(aq) + HCO3−(aq)
CO2 is constantly produced by metabolic
processes in the body.
The amount of CO2 is related to the pH of the blood.

54. Salts
Ionic compound formed in neutralization reaction of acid and base
Contain anion of acid and cation of base
What salt forms from reaction of HNO3 + KOH?
What acid and base react to form CuSO4?

55. Reactions of Acids acid + metal  salt + H2 HCl + Zn 
acid + metal oxide  salt + H2O
HCl + NiO 
acid + base  salt + H2O
H2SO4 + NaOH 

56. Reactions of Acids acid + metal carbonate  salt + H2O + CO2
HCl + K2CO3 
acid + metal hydrogen carbonate
 salt + H2O + CO2
HC2H3O2 + NaHCO3 

57. Reactions of Acids
Write balanced molecular equations and net ionic equations for the reaction of nitric acid with:
water
calcium oxide
magnesium hydroxide
copper(II)carbonate
potassium bicarbonate
magnesium metal

58. Reactions of Bases
Write balanced equations for the reaction of KOH and the following acids: HCl HNO3 H2SO4

59. Formation of Salts
Write balanced equations for the formation of the following salts: LiCl Cu(NO3)2 MgCO3 Ca(C2H3O2)2