Acids, Bases and Salts (Oh my!) Name:
Unit 11
REGENTS
March/April
Mon
Tue
3/28 C
4/4 D
Notes- Titration,
practice problems
Wed
Thu
3/22 D
3/23 A
BROOKS BBQ
Test
3/24 B
Lab – EGGcellent
Experiment
3/29 D
Preview due
Unit 11Electrolytes
3/30 A
HW #1 due
3/31 B
4/5 A
Notes- pH
4/6 B
HW#2 due
Lab- Titration
4/11 A
4/12 B
Begin Unit 12
Test
Lab- Prop of A/B
4/13 C
Fri
No school
4/1 C
Notes- Properties &
Reactions
4/7 C
4/14 D
4/8 D End of Quarter
HW#3 due
Review
4/15A
Assignments
#1
#2
#3
I. Electrolytes
II. Acids and Bases- definitions
A. Operational
B. Arrhenius
C. Indicators
III Reactions
A. Neutralization
B. Titration
C. Hydrolysis
IV. pH and [H+]
A. pH
B. pOH
Read p 594-602, 606 (skip
Bronsted Lowry)
w/s #1, 2, 3
p. 601 #4, 5
p. 630 #44, 46, 52, 53, 57, 73
V. Bronsted-Lowry
Read p 598-599
P 599 #2 a, b
Read p 617-620
w/s #4, 5
p. 633 # 93a, 95, 96
Read p 608-616
w/s #6, 7, 8
p. 616 #24
p. 632 #63, 64, 66, 88 a, *b
Unit 11 ThINKs
1. Behavior of many acids and bases can be explained by the Arrhenius theory. Arrhenius acids and
bases are electrolytes.
 given properties, identify substances as Arrhenius acids or Arrhenius bases
2. An electrolyte is a substance which, when dissolved in water, forms a solution capable of
conducting an electric current. The ability of a solution to conduct an electric current depends on the
concentration of ions.
3. Arrhenius acids yield H+ (aq), hydrogen ion as the only positive ion in an aqueous solution. The
hydrogen ion may also be written as H3O+ (aq), hydronium ion.
4. Arrhenius bases yield OH- (aq), hydroxide ion as the only negative ion in an aqueous solution.
5. In the process of neutralization, an Arrhenius acid and an Arrhenius base react to form a salt and
water.
 write simple neutralization reactions when given the reactants
6. Titration is a laboratory process in which a volume of solution of known concentration is used to
determine the concentration of another solution.
 calculate the concentration or volume of a solution, using titration data
7. There are alternate acid-base theories. One theory states that an acid is an H+ donor and a base is
an H+ acceptor.
8. The acidity or alkalinity of a solution can be measured by its pH value. The relative level of acidity
or alkalinity of a solution can be shown by using indicators.
 interpret changes in acid-base indicator color
 identify solutions as acid, base, or neutral based upon the pH
9. On the pH scale, each decrease of one unit of pH represents a tenfold increase in hydronium ion
concentration.
Unit 11 VOCABULARY
Electrolyte (nonelectrolyte)
Arrhenius Acid & Base
Bronsted Lowry Acid & Base
Indicator
Titration
Neutralization
pH
Unit 11 Acids, Bases and SaltsNotes
General Concepts and Review
Electrolytes

ionization, solubility

Acids (Arrhenius)

Bases (Arrhenius)
Electrolytes-
ex:
ex:
*Strong - Produce _______________
*Weak - Produce _______________
Write the ionization equations:
HCl(aq)
Ca(OH)2

Salts

pH
HC2H3O2
CH3COOH
NH4OH
I. Electrolytes- Substances that are able
to _______________ _______________
when they are _______________ in water
because they _______________ into
_______________.
Non Electrolytes- _______________
Examples:
Water
Sample Ionization Reactions:
NaCl (s) 
Li2SO4(s) 
Sugars
Alcohols
II. Arrhenius Definitions and properties
Acids- Produce _______________ (or
hydronium ions) in solution.
Example:
Acid Properties
1. pH _____
2. Change indicators
• Litmus- __________
• Phenolphthalein- __________
3. Electrolytes?
4. React with __________to form a
__________ and __________
5. Have __________taste, __________
feeling
6.
React with some __________ to form
H2 gas
Use RT “J”- metals…
Reactions with Acids (acids with metals and
neutralization)
Al +
HCl 
Mg +
H2SO4 
Ag +
HBr 
HBr +
HNO3 +
NaOH 
Ca(OH)2 
REMEMBER H+(aq) can be written as:
and we call these:
Bases-
Produce _______________ in
solution. Example:
Base Properties
1. pH _____
2. Change indicators
a. Litmus- __________
b. Phenolphthalein- __________
3. Electrolytes?
4. React with __________to form a
__________ and __________
5. Have __________taste, __________
feeling
6. Emulsify fats and oils
Reactions with Bases (neutralization)
H2CO3 + NH4OH 
H2SO4 +
Mg(OH)2 
Using RT “M”
Indicator/pH
1.2-2.8
red-yellow
2
7
9
12
Range =
Methyl
orange
Yellow
Bromthymol
blue
Yellow
Bromcresol
green
Blue
Thymol blue
Yellow
III. Reactions
Neutralization- adding ________ and ________ to make ________ and ________
 Equal number of ________ of H+ and OH HCl + NaOH 
 Net ionic reaction:
How many moles of HCl can be neutralized by 0.042 L of a 0.10 M NaOH?
Titration
Lab procedure used to find unknown concentration
Ma = concentration of H+ ions or [H+]
Ma and Mb
 1 M HCl =>
Mb = concentration of OH- ions or [OH-]
Ma = _____ M H+
 0.1 M CH3COOH=>
Ma = _____ M H+
 0.1 M H2SO4 =>
Ma = _____ M H+
 0.1 M NaOH =>
Mb = _____ M OH-
 0.5 M H3PO4 =>
Ma = _____ M H+
 0.15 M Sr(OH)2 =>
Mb = _____ M OH-
How much of 0.25M HCl is required to completely neutralize
1. Write reaction.
42 mL of a 0.10 M NaOH solution?
2. Write down “knowns”.
3. Write formula and substitute.
How much 0.25 M H2SO4 is required to completely neutralize
1. Write reaction.
42 mL of a 0.10 M NaOH solution?
2. Write down “knowns”.
3. Write formula and substitute.
A reaction was done to neutralize 25mL of a 0.10M HBr solution. If it took 31mL of LiOH, determine
the concentration of the LiOH solution.
A reaction was done to neutralize 25mL of a 0.10M H3PO4 solution. If it took 31mL of LiOH, determine
the concentration of the LiOH solution.
**HydrolysisReverse of __________Salts can be acidic or basic or neutral so pH after neutralization may not have a pH of 7 (neutral) because
of the salt produced. *Not on Regents but VERY important concept in the real world!
IV. pH
 pH is a measure of acidity (or alkalinity)
_______________
([H+] = 10-pH)
 Every decrease in 1 pH unit represents a
_________ increase in the [H+]
 A sol’n with a pH=_______ is 100x more acidic
then a sol’n with pH=7.00.
 Find the [H+]- remember it is not always the same
as the [acid]. Why?
 Sometimes writing [H+] in scientific notation makes the math easier.
 What is the pH of 0.1 M HCl(aq)?
 [H+] = 0.1 M
 pH = -log[H+]
(0.1 = 1x10-1)
 pH =
 Calculate the [H+] and pH
a) 0.01 M HNO3
 What is the pH of 0.01 M HCl(aq)?
 .
 .
 .
b) 0.05 M H2SO4
*c) 0.25 M H3PO4
[H+] v. [OH-]
 There will always be H+ and OH- in aqueous solutions.
Acidic(lower pH)
Basic *pOH = -log[OH-] and [H+][OH-] = 1x10-14
 pH + pOH = _____ for any (aq) sol’n
 What is the pH of a 0.1M NaOH solution?
 [OH-] = 0.1 M
 pOH = - log [0.1]
 pOH =
 pH =
Complete the following table.
Acid/Base
pH
[H+]
pOH
6
5
1x10-3
[OH-]
(higher pH)
Neutral-
V. Bronsted-Lowery Definition
Acid- donates _____ ________
Base- accepts _____
Ex:
HCl
+
H2O 
H3O+ +
Cl-
NH3
+
H2O 
NH4+ +
OH-
Amphoteric-


Acid, Base or Salt
#1
Classify each of the following compounds as an acid, base or salt and
write the ionization equation.
HNO3 
H+ + NO3-
Electrolytes
#2
Electrolytes are substances that break up (dissociate or ionize) in water to produce
ions. These ions are capable of conducting an electric current. Generally,
electrolytes consist of acids, bases and salts (ionic compounds). Nonelectrolytes
are usually covalent compounds, with the exception of acids. Classify the
1.
HNO3 Acid
2.
NaOH_____
______________________
3.
NaNO3
_____
______________________
4.
HCI
_____
______________________
5.
KCI
_____
______________________
6.
Ba(OH)2
_____
______________________
7.
KOH
_____
______________________
8.
H2S
_____
______________________
9.
AI(NO3)3
_____
______________________
10.
H2SO4
_____
______________________
6. NaOH
11.
CaCl2
_____
______________________
7. C2H5OH
12.
H3PO4
_____
______________________
8. CH3COOH
13.
Na2SO4
_____
______________________
14.
Mg(OH)2
Base
Mg(OH)2  Mg2+ + 2(OH)-
15.
H2CO3
_____
______________________
16.
NH4OH
_____
______________________
17.
NH4Cl
_____
______________________
12. C3H5(OH)3
18.
CH3COOH _____
______________________
13. Ca(OH)2
19.
FeBr3
_____
______________________
14. H3PO4
20.
HF
_____
______________________
following compounds as either an electrolyte or a nonelectrolyte and give a
reason for your selection. The first two are done for you as examples!
Compound
1. NaCl
2. CH3OH
3. CaCl2
4. HCl
5. C6H12O6
9. NH4OH
(NH3+H2O)
10. H2SO4
11. FeBr3
15. C12H22O11
16. LiNO3
Electrolyte
*(strong or weak)
Nonelectrolyte
Strong- soluble salt
Covalent only, alcohol
pH
#3
pH is a scale that measures the hydronium ion concentration of
a solution. Therefore, the pH scale can be used to determine
the acidity of a solution. A pH of less than 7 indicates an acidic
solution, a pH of 7 is neutral, and a pH of greater than 7 up to
14 is basic. The lower the pH, the higher the acidity. The higher
the pH, the lower the acidity. Indicators are substances that
change color at a different pH levels.
Bronsted-Lowry Acids and Bases
#4H
According to Bronsted-Lowry theory, an acid is a proton (H+) donor,
and a base is a proton acceptor.
Label the Bronsted-Lowry acids and bases in the following
reactions and show the direction of proton transfer.
1.
H2O + H2O ↔ H3O+ + OH –
acid
base
2.
H2SO4 + OH - ↔ HSO4- + H2O
3.
HSO4- + H2O ↔ SO4-2 + H3O+
4.
OH - + H3O+ ↔ H2O + H2O
acid
base
5.
NH 4+ + OH – ↔ NH3 + H2O
6.
HCl is considered a Arrhenius acid and a Bronsted Lowry acid.
Explain this in terms of the ions and definitions.
7.
NaOH is considered an Arrhenius base. The OH- that is produced is
considered a Bronsted Lowry base. Explain this in terms of the ions
and definitions.
8.
What is the conjugate base for HCO3-?
9.
What is the conjugate acid for HSO3-?
Using RT M, complete the following chart.
pH
Acid,
Base,
Neutral
Phenolphthalein
Litmus
Methyl
Orange
2
8
4
7
13
11
Colorless
Purple
Yellow
Light pink
Purple
Yellow
NH3?
NH3?
Ma and Mb
Acid-Base Titration #5
To determine the concentration of an acid (or base), we can react it with a base (or
acid) of known concentration until it is completely neutralized. This point of exact
neutralization, known as the endpoint, is noted by the change in color of the indicator.
We use the following equation:
Ma = concentration of the hydrogen ions
= concentration of the acid x the number of H+.
= [acid] x (# H+)
Mb = concentration of the hydroxide ions
= concentration of the base x the number of OH-.
(# OH-)
= [base] x
Solve the problems below.
1.
A 25.0 mL sample of HCI was titrated to the endpoint with 15.0 mL of
2.0 M NaOH. What was the molarity of the HCI?
(1.2 M)
2.
What is the molarity of a NaOH solution if 15.0 mL is exactly
neutralized by 7.5 mL of a 0.020 M HC2H3O2 solution? (0.01M)
3.
How much 1.5 M NaOH is necessary to exactly neutralize 20.0
mL of 2.5 M H3PO4?
(100 mL)
4.
How much of 0.5 M HNO3 is necessary to titrate 25.0 mL of 0.05
M Ca(OH)2 solution to the endpoint?
(5 mL)
5.
A 10.0 mL sample of H2SO4 was exactly neutralized by 13.5 mL of
1.0 M KOH. What is the Ma? What is the molarity of the H2SO4?
Solve the problems below.
1. What is the Mb of a 2.0 M NaOH solution?
2. What is the M a of a 0.001 M HC2H3O2 solution?
3. What is the Ma of a 2.0 M H3PO4 solution?
4. What is the M b of a 0.01 M Ca(OH) 2 solution?
5. What is the Ma of a 0.23 M sulfuric acid solution?
6. A solution of H 2SO4 has an Ma of 3.0 M. What is its
concentration of the H 2SO4?
(0.675 M)
pH and pOH
#6
The pH of a solution indicates how acidic or basic that solution is.
pH range of 0 - 7 acidic
7 neutral
7-14 basic
Since [H+] [OH-] = 10-14 at 25° C, if [H+] is known, the [OH -] can be calculated and
vice versa.
pH = -log [H+] So if [H+] = 10-6 M, pH = 6.
pOH = - log [OH -]
So if [OH -] = 10-8 M, pOH = 8.
Together, pH + POH = 14.
Complete the following chart.
1.
[H+]
pH
[OH-]
pOH
Acidic or
Basic
1x10-5 M
5
1x10-9 M
9
Acidic
.
7
4.
#7
In an aqueous solution at 25oC, the product of the [H+] and [OH-] = 1 x 10-14. This fact
allows the calculation of either ion in solution, given the other concentration. The
negative logarithm of [H+] is called pH.
Solve the problems below.
1. Given pure water:
a. What is the hydrogen ion concentration?
b. What is the hydroxide ion concentration?
c. What is the pH?
2. Given a 0.10 M NaOH solution:
a. What is the hydroxide ion concentration?
b. What is the hydrogen ion concentration?
c. What is the pH?
3. Given a 0.001 M HCl solution:
a. What is the hydrogen ion concentration?
b. What is the hydroxide ion concentration?
c. What is the pH?
10-4 M
3.
Hydrogen and Hydroxide Ion
Concentrations (M a and Mb )
4. What is the pH of a 0.010 M HNO3 solution?
10-2 M
5. What is the pH of a 0.010 M KOH solution?
5.
11
6.
10-5 M
7.
8.
10-11 M
7. A container of water has a pH of 4.0. The hydrogen ion concentration
increased by 100x. What is the new pH?
9.
10.
6. Given a solution whose pH = 5.
a. What is the concentration of hydrogen ions?
b. What is the concentration of hydroxide ions?
c. Is this solution acidic, basic or neutral
i. Explain in terms of pH.
ii. Explain in terms of ion concentrations.
12
13
6
8.A solution was changed from a pH of 6.5 to 9.5. Describe the change in
the hydronium concentration between the 2 solutions. What type of
substance was added to cause this change?
Bronsted-Lowry Acids and Bases
#8
According to Bronsted-Lowry theory, an acid is a proton (H+) donor,
and a base is a proton acceptor.
Label the Bronsted-Lowry acids and bases in the following
reactions and show the direction of proton transfer.
10.
H2O + H2O ↔ H3O+ + OH –
acid
base
11. H2SO4 + OH - ↔ HSO4- + H2O
12. HSO4- + H2O ↔ SO4-2 + H3O+
13. OH - + H3O+ ↔ H2O + H2O
acid
base
1.
What is happening to an ionic solute when it is dissolving?
2.
Ionic compounds are also known as _____________________,
covalent substances are known as ______________________.
3.
In terms of particles, why do metals conduct electricity?
4.
Why do ionic solutions conduct electricity?
5.
Why are covalent acid solutions able to conduct electricity?
6.
Why are ionic solids NOT able to conduct electricity?
7.
What do the following abbreviations/units mean?
mol
14. NH
+
4
#8
Review Concepts
M
–
+ OH ↔ NH3 + H2O
8.
What 2 formulas on Table T use molarity? When do you use each one?
What will the key “gove-aways” be?
9.
In a system that is at equilibrium, which way does the reaction shift as
you increase the concentration of a reactant? What happens to the
concentration of the products as a result?
15. HCl is considered a Arrhenius acid and a Bronsted Lowry acid.
Explain this in terms of the ions and definitions.
7. NaOH is considered an Arrhenius base. The OH- that is produced is
considered a Bronsted Lowry base. Explain this in terms of the ions
and definitions.
Antacid Effectiveness: Which is more effective; Milk of Magnesia or Tums?
Active Ingredient
Active Ingredient per unit
Recommended dose
Milk of Magnesia
Mg(OH)2
400 mg / teaspoon
2 teaspoons
Tums
CaCO3
1000 mg / tablet
2 tablets
We can compare the effectiveness of these antacids by figuring out which one can neutralize more grams of
stomach acid (HCl) per dose. Use the information above to follow the steps below in order to obtain your
final conclusion.
Milk of Magnesia
1) Write out the complete equation for the neutralization of HCl by this antacid:
2) Calculate how many grams of the active ingredient are in each dose:
3) Use stoichiometry to relate the grams of antacid to grams of HCl
Tums
1) Write out the complete equation for the reaction between of HCl and this antacid:
2) Calculate how many grams of the active ingredient are in each dose.
3) Use stoichiometry to relate the grams of antacid to grams of HCl
Which antacid can neutralize more HCl per dose? __________________________
What other factors might you look at to choose an antacid?
Which antacid might work the fastest, based on your knowledge of kinetics? __________________