1. The Chemistry of Acids and Bases

2. Acids pH below 7 turns litmus paper red taste sour
reacts with metals to produce H2(g)
generally starts with a hydrogen ion
[H+] > [OH-]
HCl

3. generally contains a hydroxide ion
Bases
pH greater than 7
turns litmus paper blue
taste bitter
feel slippery
generally contains a hydroxide ion
[H+] < [OH-]
NaOH

4. Both Acids and Bases
an electrolyte

5. Acidic, Basic, and Neutral Solutions
Type of Solution
pH Ranges
[H+] versus [OH-]
Example
Acidic
Below 7
[H+] > [OH-]
Orange Juice
Battery Acid
Your Stomach
Neutral
Equals EXACTLY 7
[H+] = [OH-]
Distilled
Water
Basic
Above 7
[H+] < [OH-]
Bleach
Sea Water
Blood

6. Indicators
Indicators are compounds that have one color in acidic solutions and another in basic.

7. Litmus Paper
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

8. Paper or plastic strips that contain combinations of indicators estimate the pH of a solution by simply dipping a piece of pH paper into it and comparing the resulting color with standards printed on the container
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

9. Acid Nomenclature Flowchart

10. BINARY ACIDS
HBr (aq)
Hydrobromic Acid

11. Naming Ternary Acids TERNARY ACIDS SO42- H2SO4 H2SO4(aq) Sulfuric acid
POLYATOMIC
IONS
PURE
FORMS
TERNARY ACIDS
SO42-
H2SO4
H2SO4(aq)
Sulfuric acid
SO32-
H2SO3
H2SO3(aq)
Sulfurous Acid

12. Naming Bases polyatomic Sodium hydroxide Calcium hydroxide
Use the same rules as for ions (name the cation, then name the anion)
NaOH 
Ca(OH)2 
KOH 
Sodium hydroxide
Calcium hydroxide
Potassium hydroxide

13. Some Common Acids and Bases
and their Household Uses.

14. What are Acids and Bases?
There are two common definitions to
describe acids and bases:
Arrhenius acids and bases
Bronsted-Lowry acids and bases
These are basically the same although they state different things.

15. Definitions for Acids & Bases
Arrhenius
Brønsted-Lowry
Definition for Acids
Definition for Bases
Key Examples
a proton producer in an aqueous solution
a proton donor
a hydroxide producer in an aqueous solution
a proton acceptor
Acid – HCl
Base - NaOH
Acid – HCl
Base – NH3
H+ = proton

16. Arrhenius Acids and Bases Definitions
acids in water produce hydronium ions, (H3O+, H+)
HNO3(aq) H+(aq) + NO3-
2. Arrhenius Base
bases in water produce hydroxide ions, (OH-)
KOH(s) K+(aq) + OH-(aq)

17. Bronsted-Lowry Definitions
Acids are proton (H+) donors.
Bases are proton (H+) acceptors. 
HCl + H2O  Cl– + H3O+
acid
conjugate base
base
conjugate acid
Courtesy Christy Johannesson

18. Bronsted-Lowry Come in Pairs General equation
HA(aq) + H2O(l) A-(aq) + H3O+(aq)
Acid + Base Conjugate base + Conjugate acid
This is an equilibrium.
B(aq) + H2O(l) BH+(aq) + OH-(aq)
Base + Acid Conjugate acid +Conjugate base

19. What to Focus On?
Arrhenius was the most restrictive definition. This definition required:
the solutions to be aqueous and
a base to contain a hydroxide (OH-) ion.
Bronsted-Lowry’s definition is the most commonly used. It is helpful to remember:
acids tend to “lose“ an H+ ion, while
bases tend to “gain“ an H+ ion.
Under this definition, ammonia (NH3) is considered a base even though it is NOT an Arrhenius base.

20. Examples HCl(aq) + KOH(s) KCl(aq) + H2O(l)
3 Ca(OH)2(aq) + 2 H3PO4(aq)  Ca3(PO4)2(s) + H2O(l)
F-(aq) + H2O(l)  HF(aq) + OH-(aq)
HCO3-(aq) + H2O(l)  CO32-(aq) + H3O+(aq)
NH4+(aq) + CO32-(aq)  NH3(aq) + HCO3-(aq)
acid
base
base
acid
base
acid
conjugate
acid
conjugate
base
acid
base
conjugate
base
conjugate
acid
acid
base
conjugate
base
conjugate
acid

21. Remember Electrolytes?
Ionic
Covalent
C6H12O6
Na+
NaCl
Cl-

22. Acids and bases are both strong or weak electrolytes (conduct electricity)
• Electrolytes = dissociate (break apart into ions) when dissolved
• Strong = completely Weak = partially Non = not at all
Weak
Strong H+
HC2H3O2
C2H3O21- H+
H-Cl
Cl-
Only a few Ions
Lots of Ions

23. Strong Electrolytes
WORD DESCRIPTION
Completely breaks apart into its ions
Are good conductors of electricity
Will produce a bright light bulb
Examples of Acids and Bases that are Strong Electrolytes
Strong Acids Strong Bases
H2SO4 NaOH
HCl Ba(OH)2
Notice that all of the ions are separated or dissociated.
The title is hyperlinked to a NCSSM animation of HCl dissolving in water.

24. Weak Electrolytes
WORD DESCRIPTION
Partially breaks apart into its ions
Are poor conductors of electricity
Will produce a dim light bulb
Examples of Acids and Bases that are Weak Electrolytes
Weak Acid
HC2H3O2 (Vinegar)
Weak Base
NH3 (Ammonia)
Notice that only some of the ions are separated or dissociated.
The title is hyperlinked to a NCSSM animation showing HA (a weak acid) dissolving in water.

25. What makes a strong acid or a strong base?
Strong electrolytes make strong acids and bases
Strong Acids
HCl - hydrochloric acid
HBr - hydrobromic acid
HI - hydroiodic acid
HNO3 - nitric acid
H2SO4 - sulfuric acid
HClO4 - perchloric acid
Strong Bases
The hydroxides of the
Group I and Group II
LiOH - lithium hydroxide
NaOH - sodium hydroxide
KOH - potassium hydroxide
*Ca(OH)2 - calcium hydroxide
*Sr(OH)2 - strontium hydroxide
*Ba(OH)2 - barium hydroxide
The title is hyperlinked to a NCSSM animation of HCl dissolving in water.

26. pH Concept

27. pH Scale Pouvoir hydrogéne (hydrogen power)
Is a scale to measure the acidity of a sample, Range: 0 -14 1 14
Highly acidic
Very basic (not acidic)
neutral 7
Neutral = 7.0
Bases 7-14
Acids 0-7

28. Relationships between pH, [H+], and [OH-]
Click on the picture to get to the animation.
As pH increases…
The [H+] (increases or decreases).
The [OH-] (increases or decreases).
The solution becomes more (acidic or basic).

29. Relationships between pH, [H+], and [OH-]
What happens as pH decreases?
As pH decreases…
The [H+] (increases or decreases).
The [OH-] (increases or decreases).
The solution becomes more (acidic or basic).

30. The pH Scale The value of pH is unitless.
Solutions with a pH less than 7 are acidic and solutions greater than 7 are basic.
If a solution is equal to 7 it is neutral.
Here is a typical pH scale.

31. pH of Common Substances

32. pH is a Logarithmic Scale
Logarithm –The number of times a base must be multiplied by itself to reach a given number
# of multiples
Base
# you’re trying to reach

33. pH Calculations Given Solving for Formula to Use [H+] pH
pH = - log[H+]
[OH-]
pOH
pOH = - log[OH-]
[H+] is the concentration of H+ ions, in mol/L.

34. Logarithms Use your calculator!
If you have a log button, you’re all set.
Each calculator can have its
own method for entering logs.
If you don’t know what to do
your calculator manual
should give examples.
E -2
9 - 43

35. Logarithms If your calculator has a ln button - Don’t use it.
Its for taking natural logs.
This is different than base 10.
E -2
9 - 44

36. Calculating pH
If [H+] is written in scientific notation and has a coefficient of 1, then the pH of the solution equals the absolute value of the exponent
Ex x 10-4 M
pH = 4.0

37. Calculating pH Problem 1: If [H+] = 3.40 x 10-5 M, what is the pH?
Given Unknown Equation
[H+] = 3.40 x 10-5 M pH pH = - log[H+]
Solve:
pH = -log (3.40 x 10-5)
pH = 4.47

38. Calculating pH Problem 2: If [H+] = 1 X 10-10, what is the pH?
Given Unknown Equation
[H+] = 1 X pH pH = - log[H+]
Solve:
pH = - log 1 X 10-10
pH = - (- 10)
pH = 10

39. Calculating pH Problem 3: If [H+] = 1.8 X 10-5, what is the pH?
Given Unknown Equation
[H+] = 1.8 X pH pH = - log[H+]
Solve:
pH = - log 1.8 X 10-5
pH = - (- 4.74)
pH = 4.74

40. Calculating pOH Solve: Problem 1: pOH = -log (2.3 x 10-12) pOH = 11.6
If [OH-] = 2.30 x M, what is the pOH?
Given Unknown Equation
[OH-] = 2.30 x M pOH pOH = - log[OH-]
Solve:
pOH = -log (2.3 x 10-12)
pOH = 11.6

41. Calculating pOH
If [OH-] is written in scientific notation and has a coefficient of 1, then the pOH of the solution equals the absolute value of the exponent
Problem 2:
If [OH-] = 1.0 x 10-9 M, what is the pH?
pOH = 9.0

42. What’s in a glass of water?
distilled

43. Distilled H2O at the Molecular Level
What’s in a glass of distilled water?
Water Molecules (H2O)
Hydronium Ions (H3O+)
Hydroxide Ions (OH-)
What’s happens in the glass of water?
H2O + H2O ⇆ H3O+ + OH-
This is called the self-ionization of water.

44. pH + pOH = 14 Water Water ionizes- falls apart into ions.
H2O ® H+ + OH-.
Only a small amount.
[H+ ] = [OH-] = 1 x 10-7M
A neutral solution.
In water Kw = [H+ ] x [OH-] = 1 x 10-14
Kw is called the ion product constant.
pH + pOH = 14
Amphoteric
a molecule or ion that can react as an acid as well as a base
Ex: H2O, NH3

45. Calculating pOH from pH
Problem 1:
If the pH is 3.25, what is the pOH?
Given
pH = 3.25
Unknown
pOH ?
Equation
pH + pOH = 14
Substitute and solve :
pOH = 14
(- 3.25) +pOH = 14 (- 3.25)
pOH = 10.8

46. Calculating pH from pOH
Problem 2:
What is the pH of a solution if [OH-] = 4.0 x M?
Given
[OH-] = 4.0 x M
Unknown
pH?
Equation
pH + pOH = 14
Step 1: Find pOH
pOH = -log [OH]
pOH= -log[4.0 x ] = 10.4
Step 2: Calculate pH
pH + pOH= 14; pH = 14 – 10.4
pH = 3.6

47. Looking at the Math
Given
Solving for
Formula to Use pH
[H+]
pOH
[OH-]

48. Calculating [H+] from pH
If the pH of Coke is 3.12, [H+] = ??? [
[H+] = = 7.6 x 10-4 M
*** to find antilog on your calculator, look for “Shift” or “2nd function” and then the log button
Known
pH = 3.12
Unknown
[H+] ?
Analysis
[H+] = 10 -pH
Substitute and solve :

49. Calculating [H+] from pH
The pH of an unknown solution is What is its [H+]?
Known
pH = 6.00
Unknown
[H+] ?
Analysis
[H+] = 10 -pH
Substitute and solve :
[H+] = 1x M

50. Calculating [H+] from pH
A solution has a pH of What is the Molarity of hydrogen ions in the solution?
Known
pH = 8.5
Unknown
[H+] ?
Analysis
[H+] = 10 -pH
Substitute and solve :
[H+] =
3.16 X M

51. Acid-Base Reactions or Neutralization Reactions
acid + base  water + salt
HBr(aq) + NaOH(aq) 
H2SO4(aq) + KOH(aq) 
H3PO4(aq) + Ba(OH)2(aq) 
H2O + NaBr
H2O + K2SO4
H2O + Ba3(PO4)2
* Double replacement reactions