1. Which items do you think are ACIDS & which are BASES?
Make a T chart and list the ones you think are acids on one side & the other side label & list the bases

2. NaOH H2O LiOH H3PO3 H3PO4 KOH HCl H2SO4
From the list of compounds, on your T chart continue to list the ones you think are ACIDS under the acid & bases under BASES
HCl H2SO4
NaOH H2O
H3PO4 KOH
LiOH H3PO3

3. Name each of the chemical compounds from the previous slide on your T chart the way you know how, at this time.
How do you identify if a compound is an acid when looking at the chemical formula?
How do you identify if a compound is a base when looking at the chemical formula?

4. Naming Acids
Rule 1: when the name of the anion ends in –ide, the acid name begins with the prefix hydro-. The stem of the anion has the suffix –ic and is followed by the word acid.
Ex: HCl – anion is chloride; therefore
Hydrochloric acid
Practice:
H2S
Hydrosulfuric acid

5. Naming Acids
Rule 2: when the anion name ends in –ite, the acid name is the stem of the anion with the suffix –ous, followed by the word acid. Ex: H2SO3 – anion is sulfite; therefore Sulfurous acid Practice: HClO2 Chlorous acid

6. Naming Acids
Rule 3: When the anion name ends in –ate, the acid name is the stem of the anion with the suffix –ic, followed by the word acid. Ex: HNO3 – anion is nitrate; therefore Nitric acid Practice: H2SO4 Sulfuric acid

7. Naming Bases
Name bases exactly as you would ionic compounds. Do not add the word base. Ex: NaOH – cation is sodium; anion is hydroxide Sodium Hydroxide

8. Practice Problems #1 Formula Anion Anion name Acid Name 1. HF F- is
fluoride
hydrofluoric acid
2. HCl
Cl-  is 
3. HBr
Br-  is 
4. H2SO3
SO32- is
5. H2SO4
SO42- is
6. HNO3
NO3- is 
7. HNO2
NO2-1 is 
8. H2S
S2- is 
9. H2CO3
CO32- is 
10. H3PO4
PO43- is 
11. HClO
ClO¯ is 
12. HClO2
ClO2¯ is 
13. HClO3
ClO3¯ is 
14. HClO4
ClO4¯ is 
15. HIO3
IO3¯ is 

9. Check your neighbors work:
Formula
Anion
Anion name
Acid Name
1. HF
F-  is 
fluoride
hydrofluoric acid
2. HCl
Cl-  is 
chloride
hydrochloric acid
3. HBr
Br-  is 
bromide
hydrobromic acid
4. H2SO3
SO32- is
iodide
hydriodic acid
5. H2SO4
SO42- is
sulfide
hydrosulfuric acid
6. HNO3
NO3- is 
nitrate
nitric acid
7. HNO2
NO2-1 is 
acetate
acetic acid
8. H2S
S2- is 
sulfate
sulfuric acid
9. H2CO3
CO32- is 
carbonate
carbonic acid
10. H3PO4
PO43- is 
phosphate
phosphoric acid
11. HClO
ClO¯ is 
hypochlorite
hypochlorous acid
12. HClO2
ClO2¯ is 
chlorite
chlorous acid
13. HClO3
ClO3¯ is 
chlorate
chloric acid
14. HClO4
ClO4¯ is 
perchlorate
perchloric acid
15. HIO3
IO3¯ is 
iodate
iodic acid
16. HC2H3O2
NO2¯ is 
nitrite
nitrous acid

10. Lemons and limes are examples of foods that contain acidic solutions.
5.3 Acids, Bases, and Salts
Lemons and limes are examples of foods
that contain acidic solutions.

11. Unit 5.3 Outline 5.3.1 Acids and Bases – naming & properties
Reactions of Acids and Bases
Salts
Electrolytes and Nonelectrolytes
Introduction to pH
Neutralization
Writing Net Ionic Equations

12. Learning Objectives 5.3.1 Acids and Bases
Compare acid & base examples, name acids using naming rules specific for acids
Compare the definitions of acids and bases, including
Arrhenius, Brønsted-Lowry, and Lewis acids/bases.
5.3.2 Reactions of Acids and Bases
Describe the general reactions of acids and bases.
5.3.3 Salts & Naming Acids/Bases
Explain how a salt is formed and predict the formula
of a salt given an acid and base precursor. Explain how to name acids and bases.

13. Arrhenius Acid: An acid solution contains
Arrhenius Acids
Arrhenius Acid: An acid solution contains
an excess of H+ ions.
Common Properties of Acids
1. Sour taste
2. Turns litmus paper pink
3. Reacts with:
Metals to produce H2 gas
Bases to yield water and a salt
Carbonates to give carbon dioxide

14. Arrhenius Bases: A basic solution contains
an excess of OH– ions.
Common Properties of Bases
1. Bitter/caustic taste
2. Turns litmus paper blue
3. Slippery, soapy texture
4. Neutralizes acids

15. Brønsted-Lowry Acids and Bases
A Brønsted-Lowry acid is a proton (H+) donor.
A Brønsted-Lowry base is a proton (H+) acceptor.
HCl (g) + H2O (l) H3O+ (aq) + Cl- (aq)
Acid
Base
Hydronium ion (hydrated H+)
Conjugate Acid/Base Pairs: Two species that differ
from each other by the presence of one hydrogen ion.
HCO3- (aq) + H2O (l) H2CO3 (aq) + OH- (aq)
Acid
Base
Conjugate
conjugate pair 1
conjugate pair 2

16. Brønsted-Lowry Acids and Bases
Identify the base in the following reaction
and the compound’s conjugate acid.
H2PO4- (aq) + H2O (l) HPO42- (aq) + H3O+ (aq)
Base
Conjugate Acid
a. H2PO4- (aq)
b. H2O (l)
c. HPO42- (aq)
d. H3O+ (aq)
a. H2PO4- (aq)
b. H2O (l)
c. HPO42- (aq)
d. H3O+ (aq)

17. Lewis Acid: electron pair acceptor. Lewis Base: electron pair donor.
Lewis Acid-Bases
Lewis Acid: electron pair acceptor.
Lewis Base: electron pair donor.

18. Summary of the Acid/Base Theories
The theory that best explains the reaction
of interest is used.

19. Reactions of Acids 1. Reactivity with Metals
Acids react with any metals above
hydrogen in the activity series.
2 HCl (aq) + Mg (s) MgCl2 (s) + H2 (g)
In general:
acid + metal salt + hydrogen
2. Reactivity with Bases
Also called a neutralization reaction.
2 HCl (aq) + Ca(OH)2 (aq) CaCl2 (aq) + 2 H2O (l)
In general:
acid + base salt + water

20. Reactions of Acids 3. Reactivity with Metal Oxides
2 HCl (aq) + Na2O (s) NaCl (aq) + H2O (l)
In general:
acid + metal oxide salt + water
(base)
4. Reactivity with Metal Carbonates
2 HCl (aq) + Na2CO3 (aq) NaCl (aq) + H2O (l) + CO2 (g)
In general:
acid + carbonate salt + water + carbon dioxide
(base)

21. Base Reactions Bases can be amphoteric
(act as either Brönsted acids or bases)
As a base:
Zn(OH)2 (aq) + 2 HBr (aq) ZnBr2 (aq) + 2 H2O (l)
As an acid:
Zn(OH)2 (aq) + 2 NaOH (aq) Na2Zn(OH)4 (aq)
NaOH and KOH can also react with metals.
2 NaOH (aq) + 2 Al (s) + 6 H2O (l)
2 NaAl(OH)4 (aq) + 3 H2 (g)
In general:
base + metal + water salt + hydrogen

22. Salts Salts: products from acid-base reactions.
HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l)
Sodium chloride
(table salt)
Salts are ionic compounds.
Salts contain a cation (a metal or ammonium ion)
derived from the base and an anion
(excluding oxide or hydroxide ions) derived from the acid.
Salts are generally crystalline compounds with
high melting and boiling points.

23. Salts Practice Which salt forms in the reaction of
aluminum oxide and hydrobromic acid?
Al2O3 (s) + 6 HBr (aq) AlBr3 (aq) + 3 H2O (l)
a. AlBr
b. AlBr3
c. Al2Br
d. Al2Br3

24. Learning Objectives 5.3.4 Electrolytes and Nonelectrolytes
Describe properties, ionization, dissociation, and strength
of electrolytes and compare them to nonelectrolytes.
5.3.5 Introduction to pH
Calculate the pH of a solution from the
hydrogen ion concentration.
5.3.6 Neutralization
Describe a neutralization reaction and
do calculations involving titrations.

25. Electrolytes and Nonelectrolytes
Electrolytes: compounds that conduct electricity
when dissolved in water.
Nonelectrolytes: substances that do not conduct
electricity when dissolved in water.
Comparing Solution Conductivity
(Distilled water)
(Sugar solution)
(Salt solution)

26. Electrolytes and Nonelectrolytes
Ion movement causes conduction of electricity in water.
3 classes of compounds, acids, bases, and salts
are electrolytes because they produce ions
in water when they dissolve.

27. Electrolytes and Nonelectrolytes Practice
Which compound will not dissociate in water?
a. HCl
b. KBr
c. NaOH
d. CH3OH
a. HCl (acid, dissociates)
b. KBr (salt, dissociates)
c. NaOH (base, dissociates)
d. CH3OH (organic, does not dissociate)

28. Dissociation of Electrolytes
Salts dissociate into their respective cations
and anions when dissolved in water.
NaCl (s) Na+ (aq) + Cl- (aq)
Hydrated sodium (purple) and chloride (green) ions
The negative end of
the water dipole is
attracted to the
positive Na+ ion.
When NaCl dissolves in water,
each ion is surrounded by several water molecules.
The permanent dipoles in the water molecules cause
specific alignment around the ions.

29. Electrolyte Ionization
Ionization: process of ion formation in solution.
Ionization results from the chemical reaction
between a compound and water.
Acids ionize in water, producing the
hydronium ion (H3O+) and a counter anion.
HCl (g) + H2O (l) H3O+ (aq) + Cl- (aq)
H3PO4 (aq) + H2O (l) H2PO4- (aq) + H3O+ (aq)
Bases ionize in water, producing the
hydroxide ion (OH-) and a counter cation.
NH3 (aq) + H2O (l) OH- (aq) + NH4+ (aq)

30. Strong and Weak Electrolytes
Strong electrolytes: undergo complete ionization in water.
Example: HCl (strong acid)
Weak electrolytes: undergo incomplete ionization in water.
Example: CH3COOH (weak acid)
HCl (left) is 100% ionized.
CH3COOH exists primarily in the unionized form.

31. Strong and Weak Electrolytes
Double arrows indicate incomplete ionization
(typically weak electrolytes).
HF (aq) + H2O (l) F- (aq) + H3O+ (aq)

32. Salts Salts can dissociate into more than 2 ions,
depending upon the compound.
A 1 M solution of NaCl produces a total of 2 M of ions.
NaCl (s) Na+ (aq) + Cl- (aq) 1M 1M 1M
A 1 M solution of CaCl2 produces a total of 3 M of ions.
CaCl2 (s) Ca2+ (aq) + 2 Cl- (aq) 1M 1M 2M

33. Salts Practice What is the concentration of bromide ion in a
1.5 M solution of magnesium bromide?
a. 0.75M
b. 1.5M
c. 3.0M
d. 4.5M
For every one mole of MgBr2, 2 moles of Br- ionize.
MgBr2 (s) Mg2+ (aq) Br- (aq)
1.5 M
1.5M
3.0M

34. Autoionization of Water
Pure water auto(self) ionizes according to the reaction:
H2O (l) + H2O (l) H3O+ (aq) + OH– (aq)
Based on the reaction stoichiometry:
Concentration H3O+ = Concentration OH– = 1 x 10–7 M
[H3O+] x [OH–] = (1 x 10–7)2 = 1 x 10–14
When acid or base is present in water,
[H3O+] and [OH-] change.
In acidic solutions, [H3O+] > [OH–].
In basic solutions, [H3O+] < [OH–].

35. In pure water, [H3O+] = 1 x 10-7 M, so
Introduction to pH
pH = - log[H3O+]
In pure water, [H3O+] = 1 x 10-7 M, so
pH = - log(1 x 10-7) = 7
Neutral
High H3O+
Low OH-
Increasing acidity
Increasing basicity
Low H3O+
High OH-
The pH scale

36. pH

37. pH Practice By what factor is a solution of pH = 3 more acidic
than a solution with a pH = 5?
a. 2
b. 20
c. 200
d. 100
pH = 5 means [H3O+] = 1 x 10-5 M
pH = 3 means [H3O+] = 1 x 10-3 M
1 x 10-3 M
Factor =
= 100
1 x 10-5 M

38. pH Calculations pH = - log[H3O+] [H3O+] = 10-pH Generalizations
Exponent = pH
pH = 5
[H3O+] = 1 x 10-5 M
The pH is between
the exponent and
next lowest whole
number
pH = 4.7
If exactly 1
[H3O+] = 2 x 10-5 M
If a number between 1 and 10

39. pH Calculations Calculate the pH of a 0.015 M [H3O+] solution.
pH = - log(0.015) = 1.8
The number of decimal places for the answer must equal
the number of significant figures in the original number.

40. pH Calculations Practice
What is the pH of a M HCl solution? a
b. -2.0
c. 1.70
d. 1.7
pH = - log(0.020) = 1.7

41. What is the pH of a NaOH solution with a
pH Calculations Practice
What is the pH of a NaOH solution with a
[OH-] = 2.5 x M?
a. 3.4
b. 10.6
c. -3.4 d
pOH = - log [OH-]
pOH = - log (2.5 x 10-11)
= 10.6
pH + pOH = 14
pH = 14 – 10.6 = 3.4

42. Neutralization General Reaction acid + base salt + water Example
HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l)
Overall Ionic Equation:
H+(aq) + Cl- (aq) + Na+ (aq) + OH- (aq)
Na+ (aq) + Cl- (aq) + H2O (l)
All species are included; soluble compounds shown as ions.
Net Ionic Equation:
H+ (aq) + OH- (aq) H2O (l)
Spectator ions (green) are removed from both sides.

43. Titration Titration: experiment where the volume of one reagent
(titrant) required to react with a measured amount
of another reagent is measured.
Titrations allow the amount of an acid or base
present in a sample to be determined.
Indicators are used to signal the endpoint of a titration,
the point when enough titrant is added to react
with the acid/base present.
Burets deliver measured amounts of the titrant
into a solution of the unknown reagent.
Endpoint

44. Titration Calculations
If mL of M HCl is used to titrate mL
of NaOH, what is the molarity of the base?
Reaction
HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l)
Knowns mL of M HCl
25.00 mL NaOH
Solving for
Molarity of base
Calculate
mol HCl
1000 mL soln
1 mol NaOH
22.59 mL × ×
= mol NaOH
1 mol HCl
mol NaOH
L soln
Molarity NaOH =
= M NaOH

45. Titration Calculations
What is the molarity of a NaOH solution if mL of
base is required to titrate g of oxalic acid (H2C2O4)?
Reaction
H2C2O4 (aq) + 2 NaOH (aq) NaC2O4 (aq) + 2 H2O (l)
Knowns g H2C2O4
21.93 mL NaOH
Solving for Molarity of base
Calculate
1 mol H2C2O4
90.04 g H2C2O4
2 mol NaOH
0.243 g H2C2O4 × ×
= mol NaOH
1 mol H2C2O4
mol NaOH
L soln
Molarity NaOH =
= M NaOH
Always check reaction stoichiometry!

46. Titration Calculations Practice
What is the concentration of a nitric acid solution
if 10.0 mL of the solution is neutralized by
3.6 mL of 0.20 M NaOH?
Reaction
a M
b M
c M
d. 5.6 M
HNO3(aq)+NaOH(aq) NaNO3(aq)+ H2O(l)
Knowns mL, 0.20 M NaOH
10.00 mL HNO3
Solving for Molarity of acid
Calculate
0.20 mol NaOH
1000 mL NaOH
1 mol HNO3
3.60 mL NaOH × ×
= mol HNO3
1 mol NaOH
mol HNO3
L soln
Molarity HNO3 =
= M HNO3

47. Learning Objectives 5.3.7 Writing Net Ionic Equations
Write net ionic equations using the stated rules.

48. Rules for Writing Net Ionic Equations
1. Strong electrolytes are written as the corresponding
ions. Example: NaOH (aq) is written as Na+(aq) + OH-(aq)
2. Weak electrolytes and nonelectrolytes are written as
molecules. Example: CH3OH(aq), CH3COOH(aq), etc.
3. Solids and gases are written as their molecular forms.
4. The net ionic equation does not include spectator ions.
5. The net ionic equation must balance atoms and charge.

49. Net Ionic Equations Practice
Write the net ionic equation (NIE) for
the following reaction:
BaCl2 (aq) + Na2SO4 (aq) BaSO4 (s) + 2 NaCl (aq)
Total Ionic Equation
Ba2+ (aq) + 2 Cl- (aq) + 2 Na+ (aq) + SO42- (aq)
BaSO4 (s) + 2 Na+ (aq) + 2 Cl- (aq)
Spectator Ions
Net Ionic Equation
Ba2+ (aq) + SO42- (aq) BaSO4 (s)

50. Net Ionic Equations Practice
Write the net ionic equation (NIE) for
the following reaction:
Na2CO3 (aq) + 2 HCl (aq) CO2 (g) + 2 NaCl (aq) + H2O (l)
Total Ionic Equation
2 Na+ (aq) + CO32- (aq) + 2 H+ (aq) + 2 Cl- (aq)
2 Na+ (aq) + 2 Cl- (aq) + CO2 (g) + H2O (l)
Spectator Ions
Net Ionic Equation
2 H+ (aq) + CO32- (aq) CO2 (g) + H2O (l)

51. Net Ionic Equations Practice
What is the net ionic equation when
hydrobromic acid reacts with potassium hydroxide? a. b. c. d.
H+ (aq) + Br- (aq) HBr (l)
K+ (aq) + OH- (aq) KOH (s)
K+ (aq) + H- (aq) KH (s)
H+ (aq) + OH- (aq) H2O (l)
Equation
HBr (aq) + KOH (aq) KBr (aq) + H2O (l)
Total Ionic Equation
H+(aq)+Br-(aq)+K+(aq)+OH-(aq) K+(aq)+Br-(aq)+H2O(l)