Solutions, Acids & Bases
Unit 9
What is a solution?
– a mixture where the
components are uniformly intermingled.
Homogeneous
 Solutions are made up of a solvent and
solute.
• Solute – substance that is dissolved
• Solvent – substance in which materials
are dissolved in. The medium in which
the solute is dissolved in.
 Solution
Solution Composition
 The
solubility of a solute is limited.
• Saturated solution – contains as much
solute as will dissolve at that temperature
• Unsaturated solution – has not reached
the limit of solute that will dissolve
• Oversaturated solution – contains too
much solute in which the rest of solute
remains on the bottom
solution – occurs when a
solution is saturated at an elevated
temperature and then allowed to cool but
all of the solid remains dissolved
 Supersaturated
• Contains more dissolved solid than a saturated
solution at that temperature
• Unstable – adding a crystal causes precipitation
•
Amounts of substances can vary in
different solutions.
– Specify the amounts of solvent and
solutes
– Qualitative measures of concentration
• concentrated – relatively large amount of solute
• dilute – relatively small amount of solute
B. Solution Composition:
An Introduction
Which solution is more concentrated?
B. Solution Composition:
An Introduction
Which solution is more concentrated?
C. Factors Affecting the
Rate of Dissolving
 Surface
area
 Stirring
 Temperature
Solubility Curves
 Graphs
that illustrates how much of
solute will dissolve at any given
temperature.
Solubility Chart
Questions













How many grams of KClO3 is diluted in water at 90 degrees?
How many grams of K2Cr2O7 is diluted in water at 60 degrees?
How many grams of Pb(NO3)2 is diluted in water at 10 degrees?
70g of CaCl2 will dissolve at what temperature to form a saturated solution?
40g of NaCl will dissolve at what temperature to form a saturated solution?
90g of KNO3 will dissolve at what temperature to form a saturated solution?
At 30 degrees, you dissolve 25 grams of KNO3 in water. How much more KNO3 can
you add to make it a saturated solution?
At 10 degrees, you dissolve 25 grams of NaNO3 in water. How much more NaNO3 can
you add to make it a saturated solution?
At 90 degrees, you dissolve 10 grams of KClO3 in water. How much more KClO3 can
you add to make it a saturated solution?
Which salt has solubility values that are least affected by temperature?
At 75 degrees, I attempted to dissolve 50g of KCl. Is this a saturated, unsaturated, or
supersaturated solution?
At 40 degrees, I attempted to dissolve 50g of NaCl. Is this a saturated, unsaturated, or
supersaturated solution?
At 55 degrees, I attempted to dissolve 50g of Pb(NO3)2. Is this a saturated,
unsaturated, or supersaturated solution?
B. Solution Composition:
Molarity
 Concentration
of a solution is the
amount of solute in a given volume of
solution.
B. Solution Composition:
Molarity
 Consider
both the amount of solute and
the volume to find concentration.
Example Problem
What is the concentration (or Molarity) of a
solution that contains 432 mole NaCl in 3.0
L of water?
𝑚𝑜𝑙𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒
𝐿𝑖𝑡𝑒𝑟𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
=
432 𝑚𝑜𝑙 𝑁𝑎𝐶𝑙
3.0 𝐿 𝐻2 𝑂
= 144M
Example Problem
 What
is the concentration of a solution
that contains 11.2g of LiCl in a 0.50L
solution?
• Make sure that you convert the grams
to moles before you find the molarity.
First, convert the grams to moles.
11.2g LiCl
1 mol LiCl
= 0.264 mol LiCl
42.39g LiCl
Use the moles to find the molarity.
0.264𝑚𝑜𝑙 𝐿𝑖𝐶𝑙
0.50𝐿
= 0.53M
B. Solution Composition:
Molarity
 To
find the moles of solute in a given
volume of solution of known molarity
use the definition of molarity.
 How
many grams of HCl was mixed into
a 3L of a 9M solution?
Remember that:
𝑚𝑜𝑙𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒
= 𝐶𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 (𝑀)
𝐿 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
 How
many grams of HCl was mixed into
a 3L of a 9M solution?
𝑚𝑜𝑙𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒
= 𝐶𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 (𝑀)
𝐿 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
 What
do we know?
• Concentration = 9M
• L of solution = 3
• Mols of solute = ?
 How
many grams of HCl was mixed into
a 3L of a 9M solution?
𝑚𝑜𝑙𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒
= 𝐶𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 (𝑀)
𝐿 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
𝑚𝑜𝑙𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒
= 9𝑀
3𝐿
Solve for the moles of solute
 How
many grams of HCl was mixed into
a 3L of a 9M solution?
𝑚𝑜𝑙𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒
= 9𝑀
3𝐿
Mols of solute = (9M)(3L) = 27 mols of HCl
You need to convert the mols of HCl to grams of HCl.
27 mols HCl
34.46g HCl
1 mol HCl
= 930.42g HCl
Practice Problems
Calculate the molarity of a solution
prepared by dissolving 11.5g of solid
NaOH in enough water to make 1.50L
of solution.
2. 2. Calculate the molarity of a solution
prepared by dissolving 1.56g of
gaseous HCl into enough water to
make 0.0268L of solution.
1.
Practice Problems
3. What is the volume of a 3.2M solution of
H2SO4 in a 0.30 L beaker?
4. What is the mass (g) of a 1.2M solution
of CaCO3 in a 0.50 L flask?
Dilution
 Water
can be added to an aqueous
solution to dilute the solution to a lower
concentration.
 Only water is added in the dilution – the
amount of solute is the same in both the
original and final solution.
D. Dilution
 Diluting
a solution
• Transfer a measured amount of
original solution to a flask containing
some water.
• Add water to the flask to the mark
(with swirling) and mix by inverting the
flask.
D. Dilution
 We
can use a simple equation can be
used in order to correctly dilute a
solution to any concentration
M1V1 = M2V2
M1 = the molarity of the initial solution
V1 = volume of the initial solution
M2 = the molarity of the diluted solution
V2 = volume of the diluted solution
Example Problem
1. Suppose we want to prepare 500. mL
of 1.00 M acetic acid, HC2H3O2, from a
17.5M stock solution of acetic acid. What
volume of the stock solution is required?
M1 = 17.5M
V1 = ?
M2 = 1.00M
V2 = 500 mL
M1V1 = M2V2
(17.5)V1 = (1.00)(500)
V1 =
(1.00)(500)
17.5
= 28.6mL
Practice Problems
What volume of 16M sulfuric acid must
be used to prepare 1.5L of a 0.10M
H2SO4 solution?
2. What is the concentration when you
dilute 0.3L of 10M of HCl in to a .500L
solution?
6M
3. What is the concentration of a 250mL
solution that was used to make a 6M
600mL solution of 2.5M of NaCl?
1.
Acids & Bases
I. Introduction to
Acids & Bases
A. Properties

electrolytes
 electrolytes

sour taste

bitter taste

turn litmus red

turn litmus blue

react with metals
to form H2 gas

slippery feel

vinegar, milk, soda,
apples, citrus fruits

ammonia, lye,
antacid, baking soda
ChemASAP
Acids
(Svante Arrhenius’
Definition)
Acids form
hydrogen ions or
hydronium ions
when dissolved in
water H+ H O+
3
B. Definitions
 Arrhenius
- In aqueous solution…
• Acids form hydronium ions (H3O+)
HCl + H2O 
+
H3O
H
H
Cl
acid
O
H
H
–
+
O
H
+
Cl
H
–
Cl
Common Acids

HCl Hydrochloric

H SO



2
Sulfuric
4
HNO
Nitric
3
H PO
Phosphoric
3 4
H CO
Carbonic
2 3
Acid Nomenclature
Compound – compound
consisting of 2 elements.
 An acid binary compound consists of 2
elements in which one of them is
ALWAYS hydrogen.
• HF
• HCl
• HBr
 Binary
Acid Nomenclature
 To
name acid binary compounds
• Add hydro- as a prefix
• Use the root of the element and add –
ic acid as a suffix
Acid Nomenclature
 HF
• Hydrofluoric acid
 HCl
• Hydrochloric acid
 H2S
• Hydrosulfuric acid
 H3P
• Hydrophosphuric acid
Acid Nomenclature
 Writing
the name from chemical formulas
• Figure out the charge of the anion
• Hydrobromic acid – H+ and Br• Br has a negative 1 charge
• Add enough H+ until the compound is
neutral
Acid Nomenclature
 Hydronitric
acid
• H3N
 Hydrofluoric acid
• HF
 Hydroselenic acid
• H2Se
Acid Nomenclature
 Acid
polyatomic compounds
Remember polyatomic compounds are compounds
that have a charge.
 When
naming polyatomic acids…
• DO NOT USE HYDRO- as preffix
• Use either –ic acid or –ous acid
Acid Nomenclature
 ate—ic
/ ite—ous Rule
• If the polyatomic ion ends with an –ate
replace it with an –ic acid
• H2SO4 – SO4 is sulfate Add –ic
H2SO4 – Sulfuric acid
Acid Nomenclature
 ate—ic
/ ite—ous Rule
• If the polyatomic ion ends with an –ite
then replace it with an –ous acid
• HNO2 – NO2 is nitrite Add -ous
HNO2 – Nitrous acid
Example Problems
1.
2.
3.
4.
5.
6.
7.
H2SO3
H3PO4
HClO2
HC2H3O2
H2Cr2O7
H2CO3
HNO2
Sulfurous acid
Phosphoric acid
Chlorous acid
Acetic acid
Dichromic acid
Carbonic acid
Nitrous acid
Acid Nomenclature
 Writing
the name from chemical formulas
• Figure out the charge of the
polyatomic ion
• Sulfuric acid – H+ and SO4-2
• SO4 has a negative 2 charge
• Add enough H+ until the compound is
neutral
• H2SO4
Example Problems
1.
2.
3.
4.
5.
Hypochlorous acid
Perchloric acid
Nitric acid
Permanganic acid
Sulfurous acid
HClO
HClO4
HNO3
HMnO4
H2SO3
More Example Problems
1.
2.
3.
4.
5.
6.
7.
8.
HF
Hydrochloric acid
H2Se
Chromic acid
HClO3
Chlorous acid
HNO3
H3N
Hydrofluoric acid
HCl
Hydroselenic acid
H2CrO4
Chloric acid
HClO2
Nitric acid
Hydronitric acid
Bases
(Arrhenius’ Definition)
Bases form
hydroxide ions
when dissolved in
water
OH
B. Definitions
 Arrhenius
- In aqueous solution…
• Bases form hydroxide ions (OH-)
NH3 + H2O 
+
NH4
H
H
H
N
H
base
O
H
H
–
+
O
N
H
+
OH
H
H
H
Common Bases
 NaOH
 NH
Sodium Hydroxide
OH Ammonium Hydroxide
4
 Ca(OH)
2
Calcium Hydroxide
B. Definitions
 Brønsted-Lowry
• Acids are proton (H+) donors.
• Bases are proton (H+) acceptors.
HCl + H2O 
acid
–
Cl
+
+
H3O
base
conjugate base
conjugate acid
B. Definitions
H2O + HNO3  H3O+ + NO3–
B
A
CA
CB
B. Definitions
NH3 + H2O 
B
A
 Amphoteric
+
NH4
CA
+
OH
CB
- can be an acid or a base.
B. Definitions

Give the conjugate base for each of the following:
-
HF
F
H3PO4
H2PO4
+
H3O
H2O
 Polyprotic
- an acid with more than one H+
B. Definitions

Give the conjugate acid for each of the following:
Br -
HBr
HSO4-
H2SO4
CO32-
HCO3
C. Strength
 Strong
Acid/Base
• 100% ionized in water
• strong electrolyte
HCl
HNO3
H2SO4
HBr
HI
HClO4
-
+
NaOH
KOH
Ca(OH)2
Ba(OH)2
C. Strength
 Weak
Acid/Base
• does not ionize completely
• weak electrolyte
HF
CH3COOH
H3PO4
H2CO3
HCN
-
+
NH3
pH
A. Ionization of Water
At any one time water
solution contains
hydronium and
hydroxide ions.
H 2O + H 2 O
Kw =
+
[H3O ][OH ]
H3
+
O
+
= 1.0 
OH
-14
10
A. Ionization of Water
 Find
the hydroxide ion concentration of
3.0  10-2 M HCl.
[H3O+][OH-] = 1.0  10-14
[3.0  10-2][OH-] = 1.0  10-14
[OH-] = 3.3  10-13 M
Acidic or basic? Acidic
B. pH Scale
14
0
7
INCREASING
ACIDITY
pH =
NEUTRAL
+
-log[H3O ]
pouvoir hydrogène (Fr.)
“hydrogen power”
INCREASING
BASICITY
B. pH Scale
pH of Common Substances
B. pH Scale
Important Equations
pH =
+
-log[H3O ]
pOH =
-log[OH ]
pH + pOH = 14
pH = -log[H3O+]
pOH = -log[OH-]
pH + pOH = 14
[H3O+] =
concentration of
the acid
[OH-] =
concentratio
n of the base
The pH of sea water is about 7.8. What is the
pOH?
What do we know?
pH = 7.8
What are we trying to figure out?
pOH = ?
What equation should we use?
pH + pOH = 14
The pH of sea water is about 7.8. What is the pOH?
pH + pOH = 14
7.8 + pOH = 14
Solve for pOH
pOH = 14 – 7.8
Calculate
6.2
B. pH Scale
 What
is the pH of 0.050 M HNO3?
pH = -log[H3O+]
pH = -log[0.050]
pH = 1.3
Acidic or basic? Acidic
B. pH Scale
 What
is the molarity of HBr in a solution
that has a pOH of 9.6?
pH + pOH = 14
pH = -log[H3O+]
pH + 9.6 = 14
4.4 = -log[H3O+]
pH = 4.4
-4.4 = log[H3O+]
Acidic
[H3O+] = 4.0  10-5 M HBr
Example Problems
1.
2.
3.
4.
5.
What is the pH of a substance that has a
pOH of 10.3? 3.7
What is the concentration of an acid that has
the pH of 4.3? 5.0 x 10-5 M
What is the concentration of a base that has
the pH of 8.9?
7.9 x 10-6 M
What is the pH of an substance with the
[H3O+] concentration of 5.0 x 10-5 M? 4.3
What is the pOH of a substance with the
[H O+] concentration of 4.2 x 10-13? 1.6
Titrations
A. Neutralization
 Chemical
reaction between an acid and
a base.
 Products are a salt (ionic compound)
and water.
A. Neutralization
ACID + BASE  SALT + WATER
HCl + NaOH  NaCl + H2O
strong
strong
neutral
HC2H3O2 + NaOH  NaC2H3O2 + H2O
weak
strong
basic
• Salts can be neutral, acidic, or basic.
• Neutralization does not mean pH = 7.
B. Titration
 Titration
standard solution
• Analytical method
in which a standard
solution is used to
determine the
concentration of an
unknown solution.
unknown solution
B. Titration
 Equivalence
point (endpoint)
• Point at which equal
amounts of H3O+ and OHhave been added.
• Determined by…
• indicator color change
• dramatic change in pH