Chapter 19 - Acids and Bases
I. Properties of Acids and Bases

ACIDS
1.
2.
3.
4.
5.
Taste sour
Reach with certain metals (Zn, Fe, etc.) to produce hydrogen gas
cause certain organic dyes to change color
react with limestone (CaCO3) to produce carbon dioxide
React with bases to form salts and water

BASES
1.
2.
3.
4.
5.
Taste bitter
feel slippery or soapy
react with oils and grease
cause certain organic dyes to change color
react with acids to form salts and water

Define:
o
o
Acid - a substance that produces protons, H+
Base - a substance that produces hydroxide ions, OH-
II. Reaction of acids and bases with water:



Acids and bases form ions in solution:
HCl(aq)  H+(aq) + Cl-(aq)
H3O+ - hydronium ion H+ and H3O+ are equivalent in aq. solution
When we look at the reactions of acids - can be generalized using hydrogen ion
1. Reaction with zinc yields hydrogen gas
2. Reaction with limestone - produce CO2(g)
3. Acids react with bases to produce a salt

Similarly for bases, produce hydroxide ions
III. Neutralization and Salts



Neutralization - one type of double replacement reaction
Acid + Base  Salt + water
Net ionic equation shows what drives the neutralization reaction
example:
Molecular: HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l)
Total Ionic: H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq)  Na+(aq) + Cl-(aq) + H2O(l)
Net Ionic: H+(aq) + OH-(aq)  H2O(l)

SALT - a salt is formed from the anion of the acid and the cation of the base - usually present
as spectator ions. - not always NaCl
IV. Types of Acids




Monoprotic - a solution that produces one mole of H+ ions per mole of acid HCl , HNO3
Diprotic - a solution that produces two moles of H+ ions per mole of acid H2SO4
Triprotic - a solution that produces three moles of H+ ions per mole of acid H3PO4
Polyprotic - two ore more H+ per mole of acid
V. Polyprotic acids:
 can be Partially neutralized
 acid salt - an ionic compound containing the anion with one or more hydrogens that can be
neutralized with a base
VI. Strengths of Acids and Bases:


STRONG ACIDS
o Acids that are essentially 100% ionized in aqueous solutions
o ex: HCl, HNO3, HClO4
o produce the maximum concentration of H+
o [acid] = [H+]
WEAK ACIDS
o Acids that are partially ionized ( usually less than 5%) in equilibrium.
o
o


HF + H2O(l)
H3O+(aq) + F-(aq)
The forward and the reverse reaction are occurring simultaneously most found as HF.
STRONG BASES
o those compounds that completely ionize in water to produce OH- ions
o NaOH(s)  Na+(aq) + OH-(aq)
o Concentration of base = concentration of hydroxide ions
WEAK BASES
o
o
NH3(aq) + H2O(l)
NH4+(aq) + OH-(aq)
equilibrium lies far to the left (mostly reactants present)
VII. Equilibrium of Water






H2O(l) + H2O(l)
H3O+(aq) + OH-(aq)
Autoionization - produces positive and negative ions from the dissociation of the molecules
of a liquid.
Experimentally, found concentration of ions = 1.0 x 10-7 M at 25 C
[H3O+][OH-] = Kw
at 25 C (1.0 x 10-7)(1.0 x 10-7) = 1.0 x 10-14
Kw = ION PRODUCT - gives us the concentrations of hydronium and hydroxide ions in
pure water and acidic and basic solutions
Neutral
[H3O+] = [OH-] = 1.0 x 10-7 M
Acidic
[H3O+] > 1.0 x 10-7, [OH-] <1.0 x 10-7
Basic
[H3O+] < 1.0 x 10-7, [OH-] >1.0 x 10-7
VIII. pH Scale - another way of writing concentrations.




pH = -log[H3O+] pOH = -log[OH-]
pH = 1.00  [H3O+] = 1.0 x 10-1M
pH = 7.00  [H3O+] = 1.0 x 10-7M
Sig. Figs:
1. The number of sig figs to the right of decimal in pH equals the number of total sig. figs. in
the concentration.
2. The total number of sig. figs. in the concentration equals the number of sig. figs. to the right
of the decimal in the pH.
IX. Similarly for hydroxide



pOH = -log[OH-]
pOH = 1.00  [OH-] = 1.0 x 10-1M
pOH = 7.00  [OH -] = 1.0 x 10-7M

pH + pOH = 14.000
Neutral
pH = 7
pOH = 7
Acidic
pH < 7
pOH > 7
Basic
pH > 7
pOH < 7
X. Brønsted-Lowry Acids and Bases


acid - a proton (H+) donor
base - a proton (H+) acceptor
NH3(aq) + H2O(aq)  NH4+(aq) + OH- (aq)


NH3 and NH4+ are conjugate acid-base pairs
H2O and OH- are conjugate acid-base pairs

Amphiprotic - a compound or ion that can either donate or accept H+ ions.
H2O, HSO4- , HPO42-, HSO3- etc.
XI. Predicating acid base reactions in water:





Acid-Base reactions always yield conj. acid-base
Strong Acid  weak conj. base
Strong Base  weak conj. acid
Weak Acid  strong conj. base
Weak Base  strong conj. acid


The strength of the reactant compared to the strength in the product determines which
direction the equilibrium lies.
Three predictions can be made:
o The reactant may Not react at all, leaving essentially all reactants (negligible)
o The reactants may Slightly react, leaving mostly reactants (limited)
o The reactants may react (essentially) completely, leaving little or no reactants
(favorable)
XI. Acidic and Basic Salt solutions:





Hydrolysis - the reaction of an anion with water to produce OH- or the reaction of a cation to
produce H3O+ .
Neutral solutions of salts: Cation does not undergo hydrolysis
Anion does not undergo hydrolysis
Basic solutions of salts: Cation same as above
Anion undergoes some hydrolysis
Acidic solutions of salts: Cation undergoes some hydrolysis
Anion does not
Complex solutions: Cation and anion undergoes hydrolysis
Then you need to know the relative strength of each.
XII. Buffer solutions


Buffer solution - resists changes in pH caused by the addition of limited amounts of a strong
acid or a strong base.
A buffer solution must contain:
A weak acid + its conjugate base
or A weak base + its conjugate acid
ACIDS AND BASES
1.
Define the following terms in your own words:
acid
base
neutral
2.
pH
neutralise
indicator
hydrogen ion
[H+ ]
Indicate which of the following are characteristics of acids or bases:
has sour taste
turns blue litmus red
neutralises acids
turns red litmus blue
has a soapy feel
releases hydrogen ions in water
3.
Worksheet
proton donor
proton acceptor
Write an equation for the dissociation of the following strong acids in
water:
HCl
H2SO4
4.
Write an equation for the dissociation of the weak acid, H2CO3
5.
Briefly outline the difference between strong and weak acids.
6.
Write an equation to show the dissociation of the strong base NaOH in water.
7.
Write an equation for the dissociation of the weak base NH3 ,ammonia
in water.
8.
Briefly outline the difference between strong and weak bases.
9.
State whether the following solutions are acidic, basic or neutral:
blood, pH = 7.4 milk, pH = 7
pancreatic juice, pH = 8.4
gastric juice, pH = 1.6
10.
in water.
caustic soda(NaOH), pH= 14
coffee, pH = 5.5detergent, pH= 8 to 9
vinegar, pH =2.8
Arrange the above substances in order from the most acidic to the most
basic.
Explain briefly why gastric juice is acidic but pancreatic juice is basic.
Extension
11.
12.
a.
Calculate the pH of the following solutions:
b.
State whether they are acids, bases or neutral.
[H+ ]= 7
[H+ ]= 2
[H+ ]= 5
[H+ ]= 14
I M HCl
0.1 M HCl
1 M NaOH
0.1 M NaOH 0.01 NaOH
Write down the [H+ ] and [OH- ] in pure water.
[H+ ]= 9
Acids and bases 2
Neutralisation and buffers
1. Explain how antacids neutralize excess stomach acids. Use an equation in your
explanation.
2. Complete and balance the following acid and base neutralization reactions:
KOH + HCl 
NaOH + H2SO4 
Mg(OH) 2 + 2HCl 
Na HCO3 + HCl 
3. CO2 is transported in the blood as carbonic acid (H2CO3). It then further breaks
down to form bicarbonate ions.
Write an equation to show this reaction.
Why is it important that this is a reversible reaction?
4. What conditions can cause a build up of CO2 ?
What effect would this have on blood pH?
5. What effect does hyperventilation have on blood pH?
6. What is the pH of the blood?
7. What is meant by the term acid-base buffer system and why does the blood need
an acid-base buffer system?
8. Name the major buffer system present in the blood.
9. Show how the blood acid-base buffer system reacts when
a. an acid (H3O +)enters the blood
b. a base (OH-) enters the blood
Solubility Worksheet
Definition:
Solubility of a solution is the exact amount of solute required to form a saturated
solution in a particular solvent at a certain temperature.
Solubility:
1 ppm = 1 mg/L
one parts per million = one milligram per liter
Solubility in water:
Substance
Salt
Water
20oC
36.0g/100mL
170g/100mL
50oC
36.7g/100mL
260g/100mL
Effect of temperature on Solubility:
The relationship between temperature and solubility is NOT simple.
For example:
Solute
Calcium acetate
Sodium chloride
Sugar
Temperature
↑
↑
↑
Solubility
↓
same
↑
Interpreting Solubility Graphs:
A—the amount of solute needed to form a saturated solution increases with increasing
temperature.
B—the amount of solute needed to form a saturated solution decreases with increasing
temperature.
C—omit
D—the amount of solute needed to form a saturated solution remains the same with
increasing temperature.
Pressure:
-Is another factor affecting the solubility of a substance.
-pressure has an effect on the solubility of carbon dioxide in water.
-when you open a bottle or can of pop, the gas solute comes out of solution due to a
decrease in pressure.
Solubility Practice Question:
The following graph illustrates the solubility of salt (sodium chloride) water (H2O) at varying
temperatures. Answer the questions using the graph.
Solubility [g/100mL]
Solubility of Salt in Water
14
13
12
11
10
9
8
7
6
5
4
3
2
1
0
0
5
10 15 20 25 30 35 40 45 50 55 60 65 70 75
o
Temperature [ C]
i. How many grams of salt will dissolve in 100g of water at:
a. 10C? = ________
b. 25C? = ________ (interpolate-estimate between to known values)
c.
65C? = ________ (extrapolate-estimate outside of known values)
ii. At what temperature will 8g of salt dissolve in 100g of water to form a saturated
solution? = __________
iii. What relationship can you see between temperature and solubility from this graph?
____________________________________________________________
____________________________________________________________
SOLUTIONS AND SOLUBILITY Week 11
1.
Define the following terms in your own words:
solvent
solute
solution
saturated
unsaturated
solubility
suspension
concentration
ions
electrolytes
colloid
suspension
concentrated
dilute
2.
What is the average percentage of water found in
adult females
adult males
babies
Why does the percentage of water vary?
.
Distinguish between extracellular fluid and intracellular fluid ?
What are the percentages of extracellular fluid and intracellular fluid of total fluids in the
average adult body?
3.
Distinguish between solutions, colloids and suspensions. Give an example of each.
4.
What is an electrolyte?
Name the 5 major electrolytes found in the human body.
Which are anions and which are cations?
5.
Identify the solute and the solvent in the following:
10 g of NaCl in 100 ml of H2O
10 ml of ethanol in 50 ml of H2O
O2 in 50 ml of H2O
2.0 L O 2 and 8.0 L of N2
6.
What is meant by the terms polar and non-polar molecules? Give an example
each.
of
Fill in the gaps in this statement about water dissolving the solute KCl :
7.
The salt
KCl consists of positive _____ and negative______. The
_______ end of
the water
molecule is attracted to the K+ ions, pulling the ions away from the salt
crystal and
into the solution. The __________end of the water molecule will be
attracted to the
Cl-
ions in the salt crystal. The dissolved potassium and
chloride ions are
surrounded by ____ molecules. This is called ______ation.
Why are most mineral (ionic) salts soluble in water?
Why is water often called the universal solvent?
State whether each of the following solutes will be more soluble in
water (polar solvent) or hexane(non-polar solvent):
NaCl
KCl
8.
sucrose(polar)
vegetable oil
State whether the following refer to saturated or unsaturated solutions:
A sugar cube dissolves when added to a cup of coffee.
A layer of sugar forms on the bottom of a cup of cold tea.
Crystals of uric acid build up in the kidneys
9.
State whether the solubility of the solute will increase or decrease in each of the
following situations:
Increasing the temperature of water when sugar is dissolved.
Increasing the temperature of river water, where O2 is dissolved.
Removing the cork from a bottle of champagne ( CO2 is dissolved).
A deep sea diver moves rapidly to the surface. (N2 is dissolved in the blood and joints)
10.
What is meant by the term percentage concentration (w/v) ?
Name two examples of other measures of concentration.
What is the weight of 1 ml of water?
11.
Calculate the percentage (w/v)concentration of the following solutions:
20g of sucrose in 100 ml of solution
2 g of sucrose in 100 ml of solution
0.2 g of sucrose in 100 ml of solution
20 g of sucrose in 200 ml of solution
20 g of sucrose in 50 ml of solution
20 g of sucrose in 0.5 L of solution
75 g of NaCl in 0.5 L of solution
0.3 kg of glucose in 5 L of solution.
12.
A person receives 100 mL 0f 20% (w/v) mannitol solution every hour.
How many grams of mannitol are given in one hour?
How many grams are given in one day?
Reading a Solubility Chart
1) The curve shows the # of grams of solute in a saturated solution containing 100 mL or 100 g of water
at a certain temperature.
Solubility Curves of Pure Substances
2) Any amount of solute below the
line indicates the solution is
unsaturated at a certain
temperature
150
140
KI
130
3) Any amount of solute above the line in
which all of the solute has dissolved
shows the solution is supersaturated.
120
110
NaNO3
100
grams solute per 100 grams H2O
4) If the amount of solute is above
the line but has not all
dissolved, the solution is
saturated and the # grams of
solute settled on the bottom of
the container = total # g in
solution – # g of a saturated
solution at that temperature.
(according to the curve)
5) Solutes whose curves move
upward w/ increased
temperature are typically solids
b/c the solubility of solids
increases w/ increased
temperature.
90
KNO3
80
70
NH4Cl
NH3
60
50
KCl
40
NaCl
30
20
6) Solutes whose curves move
downward w/ increased
temperature are typically gases
b/c the solubility of gases
decreases with increased
temperature.
KClO3
10
Ce2(SO4)3
0
0
10
20
30
40
50
60
70
80
90
100
Temperature/Celsuis
Solubility Problems to solve
1. At 10oC, 80 g of NaNO3 will dissolve in
100 mL (a saturated solution)
2. To find the # grams needed to saturate a solution when the volume is NOT 100 mL use the
following strategy to find answer:
Start w/ known vol. x Solubility/100mL at set temp. = amount of Solute needed to saturate
Ex. 60 mL H2O x 80 g NaNO3 = 48 g NaNO3 needed to saturate solution
100 mL H2O
or if the chart is in units of 100 g of H2O use the density of water conversion 1mL H2O= 1 g H2O
Ex.
60 mL H2O
x
1 g H2O x 80 g NaNO3 = 48 g NaNO3
1 mL H2O
100 g H2O
WS - Reading the Solubility Chart Problems
grams solute per 100 grams H2O
1. Which of the salts shown on the graph is the least soluble in water at 10oC?
2. Which of the salts shown on the graph has the greatest increase in solubility as the temperature increases
from 30 degrees to 60 degrees?
Solubility Curves of Pure Substances
3. Which of the salts has its solubility
affected the least by a change in
150
temperature?
4. At 20oC, a saturated solution of
140
sodium nitrate contains 100 grams
of solute in 100 ml of water. How
KI
130
many grams of sodium chlorate
must be added to saturate the
120
solution at 50oC?
5. At what temperature do saturated
110
solutions of potassium nitrate and
NaNO3
100
sodium nitrate contain the same
weight of solute per 100 mL of
90
water?
KNO3
6. What two salts have the same degree
80
o
of solubility at approximately 19 C?
7. How many grams of potassium
70
chlorate must be added to 1 liter of
water to produce a saturated solution
NH4Cl
NH3
60
at 50oC?
8. A saturated solution of potassium
50
KCl
nitrate is prepared at 60oC using
100.mL of water. How many grams
40
NaCl
of solute will precipitate out of
solution if the temperature is
30
suddenly cooled to 30oC?
20
9. What is the average rate of increase
KClO3
for the solubility of KNO3 in grams
10
per 100 mL per degree Celsius in the
Ce2(SO4)3
temperature range of 60oC to 70oC?
0
10. If 50. mL of water that is saturated
0
10
20
30
40
50
60
70
80
90
100
with KClO3 at 25oC is slowly
Temperature/Celsuis
evaporated to dryness, how many
grams of the dry salt would be recovered?
11. Thirty grams of KCl are dissolved in 100 mL of water at 45oC. How many additional grams of KCl are
needed to make the solution saturated at 80oC?
12. What is the smallest volume of water, in mL, required to completely dissolve 39 grams of KNO3 at 10oC?
13. What is the lowest temperature at which 30. grams of KCl can be dissolved in 100 mL of water?
14. Are the following solutions saturated, unsaturated or supersaturated (assume that all three could form
supersaturated solutions)
a. 40. g of KCl in 100 mL of water at 80oC
b. 120. g of KNO3 in 100 mL of water at 60oC
c. 80. g of NaNO3 in 100 mL of water at 10oC
15. Assume that a solubility curve for a gas such as ammonia, at one atmosphere of pressure, was plotted on
the solubility curve graph. Reading from left to right, would this curve would _____
a. slope upward b. slope downward
c. go straight across
Answers: 1) KClO3, 2) KNO3, 3) NaCl 4) 14 g, 5) 72  2oC, 6) KNO3 and KCl, 7) 210 g, 8) 55  2g,
9) 2.5 g/oC, 10) 5 g, 11) 20 g, 12) 170 mL, 13) 10oC, 14) a) unsat., b) supersat., c) sat. 15) b