Transcript Acids, Bases, and Salts

Acids, Bases, and Salts
You should be able to
Understand the acid-base theories of Arrhenius, Brønsted-Lowry,
and Lewis.
Identify strong acids and bases and calculate their pH’s.
Calculate the pH of a weak acid or base.
Calculate the concentration of a strong or weak acid or base from
its pH.
Calculate the pH and ion concentration in a polyprotic acid.
Predict the pH of a salt from its formula and then calculate the pH
of the salt.
Be familiar with titration curves and selection of an acid-base indicator.
pH scale: measures acidity/basicity
Søren Sorensen
(1868 - 1939)
ACID
BASE
10x10x
100x 10x
0
1
2
3
4
5
6
7
8
9
10
11
12
13
14
NEUTRAL
Each step on pH scale represents a factor of 10.
pH 5 vs. pH 6
pH 3 vs. pH 5
pH 8 vs. pH 13
(10X more acidic)
(100X different)
(100,000X different)
pH = -log [H1+]
Acid
Base
[H+]
Acidic
[H+] = [OH-]
Neutral
pH = 7
Basic
[OH-]
Acid vs. Base
Different
Alike
pH < 7
Affects pH
and
litmus paper
Topic
sour taste
react with
metals
Acid
Different
pH > 7
Topic
Related to
H+ (proton)
concentration
pH + pOH = 14
Base
bitter taste
does not
react with
metals
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
Common Acids and Bases
Strong Acids (strong electrolytes)
HCl
HNO3
HClO4
H2SO4
hydrochloric acid
nitric acid
perchloric acid
sulfuric acid
Weak Acids (weak electrolytes)
CH3COOH
H2CO3
acetic acid
carbonic
Strong Bases (strong electrolytes)
NaOH
KOH
Ca(OH)2
sodium hydroxide
potassium hydroxide
calcium hydroxide
Weak
Weak Base
Base (weak
(weak electrolyte)
electrolyte)
NH
NH43OH
ammonia
NH3 + H2O  NH4OH
Kotz, Purcell, Chemistry & Chemical Reactivity 1991, page 145
Acid + Base  Salt + Water
• Orange juice + milk  bad taste
• Evergreen shrub + concrete  dead bush
• Under a pine tree + fertilizer  white powder
HCl + NaOH  NaCl + HOH
salt
water
Acid-Base Neutralization
1-
1+
+
+
H3O+
OH-
H2O
H2O
Hydronium ion
Hydroxide ion
Water
Water
Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 584
Formation of Sulfuric Acid
+
+
SO2(g) + H2O(l)
H2SO3(aq)
2SO2(g) + O2(g)
2SO3(g)
SO3(g) + H2O(l)
H2SO4(aq)
Sulfuric acid
Catalyzed by atmospheric dust
SO2(g) + H2O2(l)
Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 302
H2SO4(aq)
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
CO2 (g)
H2O (l)
H2CO3 (aq)
Carbon
dioxide
Water
Carbonic
acid
Weak
acid
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Common Acids
Sulfuric Acid
H2SO4
Battery acid
Nitric Acid
HNO3
Used to make fertilizers
and explosives
Phosphoric Acid
H3PO4
Food flavoring
Hydrochloric Acid
HCl
Stomach acid
Acetic Acid
Carbonic Acid
CH3COOH
H2CO3
Vinegar
Carbonated water
Common Acids
Formula
Name of Acid
Name of Negative
Ion of Salt
HF
HBr
HI
HCl
HClO
HClO2
HClO3
HClO4
H2S
H2SO3
H2SO4
HNO2
HNO3
H2CO3
H3PO3
H3PO4
hydrofluoric
hydrobromic
hydroiodic
hydrochloric
hypochlorous
chlorous
chloric
perchloric
hydrosulfuric
sulfurous
sulfuric
nitrous
nitric
carbonic
phosphorous
phosphoric
fluoride
bromide
iodide
chloride
hypochlorite
chlorite
chlorate
perchlorate
sulfide
sulfite
sulfate
nitrite
nitrate
carbonate
phosphite
phosphate
Formation of Hydronium Ions
1+
1+
1+
+
H+
H2O
hydrogen ion
(a proton)
water
H3O+
hydronium ion
Sulfuric Acid, H2SO4
Sulfuric acid is the most commonly produced industrial chemical in the world.
Uses: petroleum refining, metallurgy, manufacture of fertilizer,
many industrial processes: metals, paper, paint, dyes, detergents
Sulfuric acid is used in
automobile batteries.
H2SO4
“oil of vitriol”
Nitric Acid, HNO3
Nitric acid stains proteins yellow (like your skin).
Uses: make explosives, fertilizers, rubber, plastics, dyes, and pharmaceuticals.
HNO3
O
“aqua fortis”
H
N
O
O
Hydrochloric Acid, HCl
The stomach produces HCl to aid in the digestion of food.
Uses: For ‘pickling’ iron and steel.
Pickling is the immersion of metals in acid solution to remove
surface impurities.
A dilute solution of HCl is called muriatic acid (available in many hardware
stores). Muriatic acid is commonly used to adjust pH in swimming pools
and in the cleaning of masonry.
HCl(g) + H2O(l)
hydrogen chloride
water
HCl(aq)
hydrochloric acid
OH1-
Common Bases
hydroxide
ion
Name
Formula
Common Name
Sodium hydroxide
NaOH
lye or caustic soda
Potassium hydroxide
KOH
lye or caustic potash
Magnesium hydroxide
Mg(OH)2
milk of magnesia
Calcium hydroxide
Ca(OH) 2
slaked lime
Ammonia water
NH
H2O
NH43.OH
household ammonia
NH41+ + OH1ammonium hydroxide
Relative Strengths of Acids and Bases
perchloric
hydrogen chloride
nitric
sulfuric
hydronium ion
hydrogen sulfate ion
phosphoric
acetic
carbonic
hydrogen sulfide
ammonium ion
hydrogen carbonate ion
water
ammonia
hydrogen
Metcalfe, Williams, Catska, Modern Chemistry 1966, page 229
Formula
HClO4
HCl
HNO3
H2SO4
H3O+
HSO4H3PO4
HC2H3O2
H2CO3
H2S
NH4+
HCO3H2O
NH3
H2
acid
Conjugate base
Formula
perchlorate ion
chloride ion
nitrate ion
hydrogen sulfate ion
water
sulfate ion
dihydrogen phosphate ion
acetate ion
hydrogen carbonate ion
hydro sulfide ion
ammonia
carbonate ion
hydroxide ion
amide ion
hydride ion
conjugate base + H+
ClO4ClNO3HSO4H2O
SO42H2PO4C2H3O2HCO3HSNH3
CO32OHNH2H-
Decreasing Base Strength
Decreasing Acid Strength
Acid
Binary Hydrogen Compounds
of Nonmetals When Dissolved in Water
(These compounds are commonly called acids.)
The prefix hydro- is used to represent hydrogen, followed by the name
of the nonmetal with its ending replaced by the suffix –ic and the word
acid added.
Examples:
*HCl
Hydrochloric acid
HBr
Hydrobromic acid
*The name of this compound would be hydrogen chloride if it was NOT dissolved in water.
Naming Simple Chemical Compounds
Ionic (metal and nonmetal)
Metal
Forms
only one
positive
ion
Use the
name of
element
Forms
more than
one positive
ion
Covalent (2 nonmetals)
Nonmetal
Single
Negative
Ion
Use element
Use the name
name followed
of the
by a Roman
element, but
numeral to
end with ide
show the charge
First
nonmetal
Second
nonmetal
Before
element name
use a prefix
to match
subscript
Use a prefix
before
element name
and end
with ide
Polyatomic
Ion
Use the
name of
polyatomic
ion (ate or
Ite)
Naming Ternary Compounds
from Oxyacids
The following table lists the most common families of oxy acids.
one more
oxygen atom
HClO4
perchloric acid
most
“common”
HClO3
chloric acid
H2SO4
sulfuric acid
H3PO4
phosphoric acid
HNO3
nitric acid
one less
oxygen
HClO2
chlorous acid
H2SO3
sulfurous acid
H3PO3
phosphorous acid
HNO2
nitrous acid
two less
oxygen
HClO
hypochlorous acid
H3PO2
hypophosphorous acid
(HNO)2
hyponitrous acid
Oxyacids  Oxysalts
If you replace hydrogen with a metal, you have formed an oxysalt.
A salt is a compound consisting of a metal and a non-metal. If the
salt consists of a metal, a nonmetal, and oxygen it is called an
oxysalt. NaClO4, sodium perchlorate, is an oxysalt.
OXYACID
OXYSALT
HClO4
perchloric acid
NaClO4
sodium perchlorate
HClO3
chloric acid
NaClO3
sodium chlorate
HClO2
chlorous acid
NaClO2
sodium chlorite
HClO
hypochlorous acid
NaClO
sodium hypochlorite
ACID
SALT
per stem ic
changes to
per stem ate
stem ic
changes to
stem ate
stem ous
changes to
stem ite
hyper stem ous
changes to
hypo stem ite
HClO3
acid
+
Na1+
cation
NaClO3 + H1+
salt
Acid Definitions
Lewis Acid
Brønsted-Lowry
Arrhenius
acids
Arrhenius Acids and Bases
Acids release hydrogen ions in water.
Bases release hydroxide ions in water.
An acid is a substance that produces hydronium ions, H3O+,
when dissolved in water.
Brønsted-Lowry Definitions
A Brønsted-Lowry acid is a proton donor; it donates a hydrogen ion, H+.
A Brønsted-Lowry base is a proton acceptor; it accepts a hydrogen ion, H+.
Lewis Definitions
A Lewis acid is a substance than can accept (and share) an electron pair.
A Lewis base is a substance than can donate (and share) an electron pair.
Acid – Base Systems
Type
Acid
Base
Arrhenius
H+ or H3O +
producer
OH - producer
BrønstedLowry
Lewis
Proton (H +)
donor
Proton (H +)
acceptor
Electron-pair
acceptor
Electron-pair
donor
Arrhenius Bases and Their Properties
According to the definition of Arrhenius a:
Base - "a substance whose water solution yields...
hydroxide ions (OH-) as the only negative ions."
Are NaOH and NH3 considered to be Arrhenius bases? YES
1) Bases are electrolytes
Dissociation equation for NaOH
NaOH(s)
Na1+(aq) + OH1-(aq)
Dissociation equation for NH3
NH3(g) + H2O(l)
NH41+(aq) + OH1-(aq)
2) Bases cause indicators to turn a characteristic color
3) Bases neutralize acids
NaOH(aq) + HCl(aq)
NaCl(aq) + H2O(l)
4) Water solutions of bases tasted bitter and feel slippery.
Neutralization
Neutralization is a chemical reaction between an acid and a base
to produce a salt (an ionic compound) and water.
NaOH(aq) + HCl(aq)
base
acid
NaCl(aq) + H2O(l)
salt
water
Some neutralization reactions:
H2SO4(aq) + 2 NaOH(aq)
sulfuric acid
2 HC2H3O2(aq) +
acetic acid
sodium hydroxide
Ca(OH)2(aq)
calcium hydroxide
Na2SO4 +
sodium sulfate
2 HOH
water
Ca(C2H3O2)2 + 2 HOH
calcium acetate
water
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.
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
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.
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Salt Formation
strong
base
strong
acid
NaOH
HCl
salt of a strong base and a strong acid
NaCl
NaOH + HCl  NaCl + H2O
strong
base
weak
acid
NaOH
HC2H3O2
NaC2H3O2
salt of a strong base and a weak acid
NaOH + HC2H3O2  NaC2H3O2 + H2O
Note: that in each case H-OH (water) is formed
Salt Formation
weak
base
strong
acid
NH3
H2SO4
salt of a weak base and a strong acid
(NH4) 2 SO4
H2SO4
NH4OH
NH4OH + H2SO4  (NH4)2SO4 + H2O
weak
base
NH43OH
weak
acid
HC2H3O2
NH4 C2H3O2
salt of a weak base and a weak acid
NH4OH + HC2H3O2  NH4C2H3O2 + H2O
Note: that in each case H-OH (water) is also formed
weak
base
strong
acid
NH3
H2SO4
(NH4) 2 SO4
2 NH4OH
ammonium
ion
+
hydroxide
ion
salt of a weak base and a strong acid
H2SO4
(NH4)2SO4
sulfuric
acid
ammonium sulfate
+
2 HOH
water
sulfate ion
1-
1+
H2SO4
NH4OH
21+
1-
1+
2 NH4OH
NH4+
NH4+
1+
OHOH-
+
H2SO4
H2SO4
(NH4)2SO4
(NH4)2SO4
+
2 H 2O
HOH
HOH
Reactions that produce salt
acid
+
base
salt
+
water
H3PO4
+
NH4OH
(NH4)3PO4
+
H2O
phosphoric acid and ammonium hydroxide
HNO3
nitric acid
H2CO3
Mg(OH)2
magnesium hydroxide
KOH
yields ammonium phosphate and water
Mg(NO3)2
magnesium nitrate
K2CO3
carbonic acid
potassium hydroxide
potassium carbonate
HC2H3O2
Al(OH)3
Al(C2H3O2)3
acetic acid
aluminum hydroxide
aluminum acetate
HClO4
perchloric acid
Ba(OH)2
barium hydroxide
H2O
Ba(ClO4)2
barium perchlorate
H2O
H2O
H2O
Brønsted-Lowry Acids and Bases
Acid = any substance that donates a proton.
Base = any substance that accepts a proton.
d+
1-
1+
d-
+
HCl
H2O
H3O+
Cl-
(acid)
(base)
hydronium ion
chloride ion
Brønsted-Lowry Acids and Bases
d-
1-
1+
d+
+
NH3
H2O
(base)
(acid)
NH4+
ammonium ion
OHhydroxide ion
Definitions
 Brønsted-Lowry
• Acids are proton (H+) donors.
• Bases are proton (H+) acceptors.
HCl + H2O 
acid
–
Cl
+
+
H3O
base
conjugate base
conjugate acid
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Definitions
H2O + HNO3  H3O+ + NO3–
B
A
CA
Base
Acid
O
H O
H
N
H O
O
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
CB
Definitions
NH3 + H2O 
B
A
+
NH4
CA
Base
+
OH
CB
Acid
H
H N
H
 Amphoteric
H O
H
- can be an acid or a base.
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Definitions

Give the conjugate base for each of the following:
-
HF
F
H3PO4
H2PO4
+
H3O
H2O
 Polyprotic
- an acid with more than one H+
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Definitions

Give the conjugate acid for each of the following:
Br
-
HBr
HSO4
H2SO4
2CO3
HCO3
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Definitions
 Lewis
• Acids are electron pair acceptors.
• Bases are electron pair donors.
Lewis
base
Lewis
acid
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Basic
7
Acid
14
Neutral
pH Scale
Acidic
0
Base
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 515
[H+]
pH
10-14
14
10-13
13
10-12
12
10-11
11
10-10
10
10-9
9
10-8
8
10-7
7
10-6
6
10-5
5
10-4
4
10-3
3
10-2
2
10-1
1
100
0
1 M NaOH
Ammonia
(household
cleaner)
Blood
Pure water
Milk
Vinegar
Lemon juice
Stomach acid
1 M HCl
pH of Common Substances
gastric
juice
1.6
vinegar
2.8
carbonated
beverage
3.0
0
1
2
acidic
Timberlake, Chemistry 7th Edition, page 335
urine
6.0
4
5
bile
8.0
6
7
neutral
[H+] = [OH-]
8
ammonia
11.0
bleach
12.0
seawater
8.5
9
1.0 M
NaOH
(lye)
14.0
milk of
magnesia
10.5
detergents
8.0 - 9.0
milk
6.4
tomato
4.2
coffee
5.0
3
blood
7.4
potato
5.8
apple juice
3.8
lemon
juice
2.2
drinking water
7.2
bread
5.5
orange
3.5
1.0 M
HCl
0
water (pure)
7.0
soil
5.5
10
11
basic
12
13
14
pH of Common Substance
More acidic
More basic
pH
NaOH, 0.1 M
Household bleach
Household ammonia
Lime water
Milk of magnesia
Borax
Baking soda
Egg white, seawater
Human blood, tears
Milk
Saliva
Rain
Black coffee
Banana
Tomatoes
Wine
Cola, vinegar
Lemon juice
Gastric juice
14
13
12
11
10
9
8
7
76
5
4
3
2
1
0
[H1+]
[OH1-]
pOH
1 x 10-14
1 x 10-13
1 x 10-12
1 x 10-11
1 x 10-10
1 x 10-9
1 x 10-8
1 x 10-7
1 x 10-6
1 x 10-5
1 x 10-4
1 x 10-3
1 x 10-2
1 x 10-1
1 x 100
1 x 10-0
1 x 10-1
1 x 10-2
1 x 10-3
1 x 10-4
1 x 10-5
1 x 10-6
1 x 10-7
1 x 10-8
1 x 10-9
1 x 10-10
1 x 10-11
1 x 10-12
1 x 10-13
1 x 10-14
0
1
2
3
4
5
6
8
9
10
11
12
13
14
Acid – Base Concentrations
concentration (moles/L)
10-1
pH = 3
pH = 11
OH-
H3O+
pH = 7
10-7
H3O+
OH-
OH-
H3O+
10-14
Timberlake, Chemistry 7th Edition, page 332
[H3O+] > [OH-]
[H3O+] = [OH-]
acidic
solution
neutral
solution
[H3O+] < [OH-]
basic
solution
pH
pH = -log [H1+]
Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 285
pH Calculations
pH
pH = -log[H3O+]
[H3O+]
[H3O+] = 10-pH
[H3O+] [OH-] = 1 x10-14
pH + pOH = 14
pOH
pOH = -log[OH-]
[OH-]
[OH-] = 10-pOH
pH = - log [H+]
Given: pH = 4.6
pH = - log [H+]
choose proper equation
4.6 = - log [H+]
substitute pH value in equation
- 4.6 =
2nd
log
determine the [hydronium ion]
- 4.6 =
log [H+]
log [H+]
[H+] = 2.51x10-5 M
multiply both sides by -1
take antilog of both sides
Recall, [H+] = [H3O+]
10x
antilog
You can check your answer by working backwards.
pH = - log [H+]
pH = - log [2.51x10-5 M]
pH = 4.6
Acid Dissociation
monoprotic
e.g. HCl, HNO3
HA(aq)
0.03 M
H1+(aq) + A1-(aq)
0.03 M
0.03 M
pH = ?
pH = - log [H+]
pH = - log [0.03M]
pH = 1.52
diprotic
e.g. H2SO4
H2A(aq)
0.3 M
2 H1+(aq) + A2-(aq)
0.6 M
0.3 M
pH = - log [H+]
pH = - log [0.6M]
pH = 0.22
polyprotic
e.g. H3PO4
H3PO4(aq)
?M
3 H1+(aq) + PO43-(aq)
xM
Given: pH = 2.1
find [H3PO4]
assume 100%
dissociation
Given: pH = 2.1
3 H1+(aq) + PO43-(aq)
0.00794 M
H3PO4(aq)
XM
find [H3PO4]
assume 100%
dissociation
Step 1) Write the dissociation of phosphoric acid
Step 2) Calculate the [H+] concentration
pH = - log [H+]
[H+] = 10-pH
2.1 = - log [H+]
[H+] = 10-2.1
- 2.1 = log [H+]
[H+] = 0.00794 M
2nd
7.94 x10-3 M
log
- 2.1 =
2nd
log log [H+]
[H+] = 7.94 x10-3 M
Step 3) Calculate [H3PO4] concentration
Note: coefficients (1:3) for (H3PO4 : H+)
7.94 x10-3 M = 0.00265 M H PO
3
4
3
How many grams of magnesium hydroxide are needed to add to 500 mL of H2O
to yield a pH of 10.0?
Step 1) Write out the dissociation of magnesium hydroxide
Mg2+(aq) + 2 OH1-(aq)
1 x10-4 M
Mg(OH)2(aq)
-4 M
0.5
5 x10-5
Step 2) Calculate the pOH
Mg2+ OH1-
Mg(OH)2
pH + pOH = 14
10.0 + pOH = 14
pOH = 4.0
pOH = - log [OH1-]
Step 3) Calculate the [OH1-]
[OH1-] = 10-OH
[OH1-] = 1 x10-4 M
Step 4) Solve for moles of Mg(OH)2
M
mol
5 x10
-5
M
L
x mol
0.5 L
x = 2.5 x 10-5 mol Mg(OH)2
Step 5) Solve for grams of Mg(OH)2
x g Mg(OH)2 = 2.5 x 10-5 mol Mg(OH)2
58 g Mg(OH)2
= 0.00145 g Mg(OH)2
1 mol Mg(OH)2
I wish I had
sweat glands.
In a chicken… CaO
+
CO2
In summer, [ CO2 ] in a chicken’s blood
-- shift
CaCO3
(eggshells)
due to panting.
; eggshells are thinner
How could we increase eggshell thickness in summer?
[ CO2 ]
, shift
-- give chickens carbonated water
[ CaO ]
, shift
-- put CaO additives in chicken feed
-- air condition the chicken house
TOO much $$$
-- pump CO2 gas into the chicken house
would kill all the chickens!
LeChatelier’s Principle
N2 + 3 H2
2 NH3 + heat
…favors the endothermic reaction (the reverse
reaction) in which the rise in temperature is
counteracted by the absorption of heat.
Raising the temperature…
Increasing the pressure…
…favors the forward reaction in which 4 mol
of gas molecules is converted to 2 mol.
Decreasing the concentration
of NH3…
Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 532
…favors the forward reaction in order to
replace the NH3 that has been removed.
Animation by Raymond Chang
All rights reserved
Equilibrium Expression
Haber Process
N2 + 3 H2
K eq 
products 
reactants 
Dorin, Demmin, Gabel, Chemistry The Study of Matter 3rd Edition, page 532
2 NH3 + heat
NH 3 
3
N 2 H 2 
2
K eq 
reversible reaction:
Reactant
 Product
Reactant
and
P  R
Product
Acid dissociation is a reversible reaction.
H2SO4
equilibrium:
2 H1+ + SO42–
Rate at which
Rate at which
R  P
P  R
=
looks like nothing is happening, however…
system is dynamic, NOT static
Le Chatelier’s principle
Le Chatelier’s principle:
When a system at equilibrium is disturbed, it shifts to a
new equilibrium that counteracts the disturbance.
N2(g) + 3 H2(g)
Disturbance
2 NH3(g)
Equilibrium Shift
Add more N2…………………..
“
“
H2…………………..
“
“
NH3…………………
Remove NH3…………………..
Add a catalyst…………………
Increase pressure…………….
Fritz Haber
no shift
Light-Darkening Eyeglasses
AgCl + energy
(clear)
Go outside…
Ago + Clo
(dark)
Sunlight more intense than inside light;
“energy”
shift
to a new equilibrium: GLASSES DARKEN
Then go inside…
“energy”
shift
to a new equilibrium: GLASSES LIGHTEN
Maintaining Blood pH
Carbon dioxide is exhaled
Acid entering the blood stream
HCO31-
+
H+
H2CO3
H2O + CO2
Bicarbonate ion circulates in the blood stream where it is in equilibrium with H+ and OH-.
In the lungs, bicarbonate ions combine with a hydrogen ion and lose a water molecule
to form carbon dioxide, which is exhaled.
Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 291
Alkalosis
If our breathing becomes too fast (hyperventilation)…
Carbon dioxide is removed from the blood too quickly.
This accelerates the rate of degradation of carbonic acid into carbon dioxide and water.
The lower level of carbonic acid encourages the combination of hydrogen ions and
bicarbonate ions to make more carbonic acid. The final result is a fall in blood H1+
levels that raises blood pH which can result in over-excitability or death.
Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 291
Acidosis
If breathing becomes too slow (hypoventilation)…
…free up acid, pH of blood drops, with associated health risks such as depression
of the central nervous system or death.
The normal pH of blood is between 7.2 – 7.4.
This pH is maintained by the bicarbonate ion and other buffers.
Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 291
Acids: Concentration vs. Strength
WEAK
STRONG
CONCENTRATED
H+ A- H+ A- H+ A- H+ A- HA
A- H+ A- H+ A- H+ A- H+ A H+ A- HA H+ A- H+ A- H+ AA- H+ A- H+ A- H+ A- H+ A- H+
H+ A - H + A - H + A - HA H + A A- H+ A- H+ A- H+ A- H+ A–
H+ A- H+ A- H+ A- H+ A- H+
A- H+ A- H+ A- H+ A- H+ AHA A- H+ A- H+ A- H+ A- H+
HA HA H+ A- HA HA
HA HA HA HA HA
H+ A- HA HA HA HA
HA HA H+ A- HA HA
HA HA HA H+ A- HA
H+ A- HA HA HA HA
HA HA HA H+ A- HA
H+ A- HA HA HA HA
HA HA H+ A- HA HA
HA
H+ AHA
HA
HA
HA
HA
HA
HA
DILUTE
H+
A-
H+
A-
HA
A-
H+
A-
H+
A–
H+
A-
H+
A-
H+
A-
H+
HA
H+
A-
HA
H+
HA
A-
HA
HA
HA
H+ A -
HA
HA
HA
HA
H+A–
A-
Dissociate nearly 100%
HA
H1+
+
A-
+
A-
H+
HA
HA
HA
H+A–
HA
STRONG ACIDS
HA
WEAK ACIDS
Dissociate very little
HA
H1+
Comparison of Strong and Weak Acids
Type of acid, HA
Reversibility
of reaction
Ka value
Ions existing when acid,
HA, dissociates in H2O
Strong
Not
reversible
Ka value very large
H+ and A-, only.
No HA present.
Weak
reversible
Ka is small
H+, A-, and HA
HA(aq) + H2O(l)
H3O+(aq) + A-(aq)
The equilibrium expression for the reaction is
Ka =
[H3O+] [A-]
[HA]
Note: H3O+ = H+
Strong vs. Weak Acid
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 508
Concentrated vs. Dilute
0.3 M HCl
10.0 M CH3COOH
Dilute, strong acid
Concentrated, weak acid
2.0 M HCl
Concentrated, strong acid
OR Dilute, strong, acid
12.0 M HCl
Concentrated, strong acid
Naming Acids
Anion
Acid
_________ ide
(chloride, Cl1-)
add H+
_________ ate
(chlorate, ClO3-)
(perchlorate, ClO4-)
add H+
_________ite
(chlorite, ClO2-)
(hypochlorite, ClO-)
add H+
ions
ions
ions
Hydro____ ic acid
(hydrochloric acid, HCl)
_________ic acid
(chloric acid, HClO3)
(perchloric acid, HClO4)
______ous acid
(chlorous acid, HClO2)
(hypochlorous acid, HClO)
Equilibrium and pH Calculations
Weak acid
Strong acid
H3O+ + A-
HA + H2O
HA + HHA
2O
H3+O++ A
+ - A-
acid-dissociation
constant calculations
Ka =
[HA] = [H3O+]
[A-] [H3O+]
[H3O+]
[HA]
+
antilog(-pH)
7
[OH-]
-log [H3O+]
pH
0
1 x 10-14
=
[OH-]
14
1 x 10-14
+]
[H
O
-14
+
3
1 x 10 = [H3O ][OH ]
Kw = [H3O+][OH-]
Tocci, Viehland, Holt Chemistry Visualizing Matter 1996, page 525
=
-
Strengths of Conjugate Acid-Base Pairs
Acid strength increases
strong
HCl H2SO4
medium
weak
very weak
HNO3 H3O+ HSO4- H3PO4 HC2H3O2 H2CO3 H2S H2PO4- NH4+ HCO3- HPO42- H2O
Cl- HSO4-
negligible
NO3
H2O
SO42-
H2PO4-
very weak
C2H3O2- HCO3- HS- HPO42-
weak
Base strength increases
NH3
medium
CO32- PO43- OH-
strong
Kw = [H3O+][OH-]
1 x 10-14 = [H3O+][OH-]
Keq
equilibrium constant
Kw
Ka
Kb
water dissociation
constant
acid dissociation
constant
base dissociation
constant
CA
base
+ ++
H
NH
4
NH3H+
+
acid
NH34+
NH
CB
HA
H+
0.1 M
0.1 M
HA
H+
0.1 M
?M
+
A-
strong acid
0.1 M
+
A-
weak acid
Conjugate
Acid Strength
Relative
acid
strength
Relative
conjugate
base
strength
Very
strong
Very
weak
Strong
HA
H+ + A-
Weak
[H+] [A-]
pKa =
[HA]
Weak
Strong
Very
weak
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 508
Very
strong
Weak Acids (pKa)
Weak Acids – dissociate incompletely (~20%)
Strong Acids – dissociate completely (~100%)
A(g) + 2 B(g)
3 C(g) + D(g)
Equilibrium constant (Keq) =
Keq =
[Products]
[Reactants]
[C] 3[D]
[A][B] 2
LeChatelier’s Principle
(lu-SHAT-el-YAY’s)
H+(aq) + C2H3O21-(aq)
HC
CH
3COOH
2H3O2(aq)
Equilibrium constant
Keq =
[H+[Product]
][C2H3O21-]
[HC
[Reactant]
2H3O2]
= Ka = Acid dissociation constant
Ka = 1.8 x 10-5 @ 25 oC for acetic acid
Ka
1.8 x
10-5
[H+][C2H3O21-]
=
[HC2H3O2]
[X2][X
] 21-]
[H+][C
H3O
=
[HC
[0.1
] 2]
2HM
3O
Assume we begin with 0.1 M acetic acid.
pH = -log[H+]
X2 = 1.8 x 10-6 M
[H
X+] = 1.34 x 10-3 M
pH = -log[1.34 x10-3]
pH = 2.87
Ionization Constants for Acids
Ka
HCl
H+ + Cl1-
very large
HNO3
H+ + NO31-
very large
H2SO4
H+ + HSO41-
HC2H3O2
H2S
large
H+ + C2H3O21-
1.8 x 10-5
H+
9.5 x 10-8
+
HS1-
Sample 1)
One gram of concentrated sulfuric acid (H2SO4) is diluted to a 1.0 dm3 volume
with water. What is the molar concentration of the hydrogen ion in this solution?
What is the pH?
Solution)
First determine the number of moles of H2SO4
x mol H2SO4 = 1 g H2SO4
H2SO4
H+ + HSO41-
1 mol H2SO4
98 g H2SO4
&
= 0.010 mol H2SO4
HSO41-
H+ + SO42-
OVERALL:
H2SO4
0.010 M
2 H+ + SO42-
in dilute solutions...occurs ~100%
0.020 M
pH = - log [H+]
substitute into equation
pH = - log [0.020 M]
pH = 1.69
A volume of 5.71 cm3 of pure acetic acid, HC2H3O2, is diluted with water at
25 oC to form a solution with a volume of 1.0 dm3.
What is the molar concentration of the hydrogen ion, H+, in this solution?
(The density of pure acetic acid is 1.05 g/cm3.)
Step 1) Find the mass of the acid
Mass of acid = density of acid x volume of acid
= 1.05 g/cm3 x 5.71 cm3
= 6.00 g
Step 2) Find the number of moles of acid. (From the formula of acetic acid,
you can calculate that the molar mass of acetic acid is 60 g / mol).
x mol acetic acid = 6.00 g HC2H3O2
Molarity: M = mol / L
Substitute into equation
1 mol HC2H3O2
= 0.10 mol acetic acid (in 1 L)
60 g HC2H3O2
M = 0.10 mol / 1 L
M = 0.1 molar HC2H3O2
Step 3) Find the [H+]
Ka =
HC2H3O2
Step 3) Find the
0.1 M
[H+]
weak acid
H+ + C2H3O21-
0.1? M
Ka = 1.8 x 10-5 @ 25 oC for acetic acid
1

Ka =
[ H ][C 2 H 3O 2 ]
[ H C 2 H 3O 2 ]
1

1.8 x
10-5
1.8 x 10
1.8 x 10
-5
-5
=


[ H ][C 2 H 3O 2 ]
[ H C 2 H 3O 2 ]
[x][x]
How do the concentrations of
H+ and C2H3O21- compare?
Substitute into equation:
[HC 2 H 3 O 2 ]
x
2
[0.10 M ]
pH = - log[H+]
x2 = 1.8 x 10-6 M
x = 1.3 x 10-3 molar
pH = - log [1.3 x10-3 M]
= [H+]
pH = 2.9
H+ Concentrations
…Strong vs. Weak Acid
Moles of Acid used to make
1 L of solution
H+
pH
0.010 mol H2SO4
0.0200 M
1.7
Strong acid
0.100 mol HC2H3O2
0.0013 M
2.9
Weak acid
Note: although the sulfuric acid is 10x less
concentrated than the acetic acid...
…it produces > 10x more H+
pH = - log[H+]
Practice Problems:
1a) What is the molar hydrogen ion concentration in a 2.00 dm3 solution
of hydrogen chloride in which 3.65 g of HCl is dissolved?
1b) pH
2a) What is the molar concentration of hydrogen ions in a solution
containing 3.20 g of HNO3 in 250 cm3 of solution?
2b) pH
3a) An acetic acid solution is 0.25 M. What is its molar concentration of
hydrogen ions?
3b) pH
4) A solution of acetic acid contains 12.0 g of HC2H3O2 in 500 cm3
of solution. What is the molar concentration of hydrogen ions?
1a) 0.0500 M
1b) pH = 1.3
2a) 0.203 M
2b) pH = 0.7
3a) 2.1 x 10-3 M
3b) pH = 2.7
4) 2.7 x 10-3 M
Weak Acids
Cyanic acid is a weak monoprotic acid. If the pH of 0.150 M cyanic acid is 2.32.
calculate Ka for cyanic acid.
H3O+(aq)
4.8 x 10-3 M
0.150 M
Ka =
Ka =
+ CN1-(aq)
H+(aq)
HCN(aq)
[Products]
[Reactants]
Ka =
4.8 x 10-3 M
[H3O+] [CN1-]
pH = -log[H3O+]
[HCN]
10-pH = [H3O+]
[4.8 x 10-3 M][4.8
[CNx1-10
] -3 M]
10-2.32 = [H3O+]
[0.150 M]
4.8 x10-3 M = [H3O+]
Ka = 1.54 x 10-4
Acid Dissociation
+
H
1-
HCl
Cl
Acid
Conjugate base
Conjugate pair
Kelter, Carr, Scott, Chemistry A World of Choices 1999, page 280
Conjugate Acid-Base Pairs
conjugates
HCl
+
base
acid
H2O
H3O+
+
acid
Clbase
conjugates
HCl
acid
+
H 2O
H3O+
base
CA
+
ClCB
Conjugate Acid-Base Pairs
conjugates
acid
NH3
+
H2O
base
base
NH41+
+
OH-
acid
conjugates
NH3
base
+
H2O
acid
NH41+
CA
+ OHCB
Water is Amphoteric
Amphoteric or Amphiprotic substances:
Substances which can act as either proton donors (acids) or
proton acceptors (bases) depending on what substances are present.
HCl
+
acid
NH3
base
+
H 2O
H3O+
base
CA
H2O
acid
NH41+
CA
+
ClCB
+ OHCB
Amphoteric
A substance that can act as either an acid or a base.
11+
+
H3O+
hydronium ion
+
HSO4hydrogen sulfate
ion
1-
H2O
water
H2SO4
sulfuric acid
2-
1-
+
HSO4hydrogen sulfate
ion
OHhydroxide ion
+
SO42-
H2O
sulfate ion
water
Amphoteric
A substance that can act as either an acid or a base.
11+
+
H3O+
hydronium ion
+
HSO4hydrogen sulfate
ion
(HSO4- as a base)
H2SO4
sulfuric acid
H2O
water
Amphoteric
A substance that can act as either an acid or a base.
1-
21-
+
+
HSO4-
OH-
SO42-
H2O
hydrogen sulfate
ion
hydroxide ion
sulfate ion
water
(HSO4- as an acid)
Conjugate Acid-Base Pairs
conjugates
base2
HC2H3O2 + H2O
acid2
H3O1+
+
acid1
C2H3O2base1
conjugates
HC2H3O2 + H2O
acid
base
H3O1+
CA
+
C2H3O2CB
The reaction proceeds in the direction such that the stronger acid
donates its proton to the stronger base.
Litmus Paper
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
pH Paper
pH 0
1
2
3
4
5
6
pH 7
8
9
10
11
12
13
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Desired Features of Sensors
pH paper
1904
pH 0
1
2
3
4
5
pH 7
8
9
10
11
12
6
13
Detection limit
Low deflection
High sensitivity
High selectivity
Wide dynamic
range
Simple to use
Cost-effective
Range and Color Changes of Some
Common Acid-Base Indicators
pH Scale
1
2
3
4
5
6
7
8
9
10
11
12
13
14
Indicators
Methyl orange
Methyl red
Bromthymol blue
Neutral red
Phenolphthalein
3.1 – 4.4
red
red
4.4
yellow
yellow
6.2
6.2
red
colorless
6.8
yellow
7.6
8.0
8.0
blue
yellow
10.0
red
Bromthymol blue indicator would be used in titrating a strong acid with a strong base.
Phenolpthalein indicator would be used in titrating a weak acid with a strong base.
Methyl orange indicator would be used in titrating a strong acid with a weak base.
colorless beyond 13.0
Indicator
1
2
Orange IV
phenolphthalein
4
5
Colorless
Phenolphthalein
Methyl Red
3
pH
6 7 8
Red
Orange
Peach
methyl red
Orange
9 10 11 12
Pink
Red
Yellow
Yellow
methyl orange
Red Cabbage Indicator
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
H+
Phenolphthalein
Indicator
Colorless = Acidic pH
Pink = Basic pH
How to read a buret volume
23
24.55 mL?
23.45 mL
(not 24.55 mL)
24
Titration
standard solution
• Titration
– Analytical method in which
a standard solution is
used to determine the
concentration of an
unknown solution.
unknown solution
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Titration
• Equivalence point (endpoint)
– Point at which equal amounts of
H3O+ and OH- have been added.
– Determined by…
• indicator color change
• dramatic change in pH
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Titration
sunnyside
muriatic acid
30.0 mL of 2.0 M of NaOH
? M of HCl
10.5 mL
HCl must be ~ __x
6
more concentrated
than the NaOH.
If it requires 10.5 mL of ? M HCl to titrate 30.0 mL of 2.0 M NaOH to its endpoint:
what is the concentration of the HCl?
(x M)(10.5 mL) = (2.0 M)(30.0 mL)
M1V1 = M2V2
X = 5.7 M
10.5 mL of HCl
HCl(aq)  H+(aq) + Cl-(aq)
MH VH = MOH VOH
+
+
-
-
0.1 M
MH VH n = MOH VOH n
+
+
-
-
30.0 mL of NaOH with bromthymol blue indicator
Endpoint of titration is reached…color change.
Al(OH)3(aq)  Al3+(aq) + 3 OH-(aq)
0.1 M
0.1 M
H2SO4(aq)  2 H+(aq) + SO42-(aq)
0.1 M
“0.2 M”
0.1 M
proper term is Normality (N)
0.1 molar H2SO4 is 0.2 normal
Titration
+
O
moles H3 = moles
MVn = MVn
M: Molarity
V: volume
n: # of H+ ions in the acid
or OH- ions in the base
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
OH
Titration
Solution
Solution
of
of KOH
NaOH
42.5 mL of 1.3M KOH are required
to neutralize 50.0 mL of H2SO4
Find the molarity of H2SO4.
H3O+
OH-
M=?
M = 1.3M
V = 50.0 mL
n=2
V = 42.5 mL
n=1
Solution
of H2SO4
50.0 mL
MV# = MV#
M(50.0mL)(2)
=(1.3M)(42.5mL)(1)
M = 0.55M H2SO4
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Acid-Base Titration
Data Table
0.10 M HCl
Base (mL)
Calibration Curve
0.00 mL
1.00 mL
2.00 mL
4.00 mL
9.00 mL
17.00 mL
27.00 mL
42.00 mL
? M NaOH
1.00 mL
1.00 mL
2.00 mL
5.00 mL
8.00 mL
10.0 mL
15.0 mL
Acid (mL)
1)
2)
3)
Solution
Solution
of
of NaOH
NaOH
Create calibration curve of six data points
Using [HCl], determine concentration of NH3
Determine vinegar concentration using [NaOH]
determined earlier in lab
Solution
of HCl
5 mL
Titration
Curve
Calibration Curve
endpoint
Base (mL)
pink
pH
7
equivalence
point
Acid (mL)
Pirate…”Walk the plank”
once in water, shark eats and
water changes to pink color
base
indicator - changes color to indicate pH change
e.g. phenolphthalein is colorless in acid
and pink in basic solution
Titration Curve
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 527
Acid-Base Titrations
Titration of a Strong Acid With a Strong Base
14.0
12.0
Solution
of NaOH
10.0
OHNa+
Na+
pH
-
OH- OH
Na+ Na+
OH-
8.0
equivalence point
6.0
4.0
Solution
of HCl
H+
Cl-
2.0
Cl
H+
H+
Cl-
H+
Cl-
0.0
0.0
10.0
20.0
30.0
40.0
Volume of 0.100 M NaOH added
(mL)
Additional
Adding
additional
NaOH
NaOH
from
isNaOH
added.
the buret
is added.
pH
to increases
hydrochloric
pH rises
and
as
acid
theninlevels
the flask,
off as
the
NaOH
a strong
equivalence
is acid.
addedInbeyond
point
the beginning
is the
approached.
equivalence
the pH increases
point. very slowly.
Titration Data
pH
0.00
10.00
20.00
22.00
24.00
25.00
26.00
28.00
30.00
40.00
50.00
1.00
1.37
1.95
2.19
2.70
7.00
11.30
11.75
11.96
12.36
12.52
Solution
of NaOH
Na+
OH-
OHNa+ Na+
OH-
Solution
of HCl
H+
Cl-
25 mL
14.0
12.0
10.0
8.0
equivalence point
6.0
4.0
2.0
OHNa+
Titration of a Strong Acid With a Strong Base
pH
NaOH added
(mL)
0.0
0.0
10.0
20.0
30.0
40.0
Volume of 0.100 M NaOH added
(mL)
ClH+
H+
Yellow
Blue
Cl-
H+
Cl-
Bromthymol blue is best indicator: pH change 6.0 - 7.6
Titration of a Strong Acid With a Strong Base
(20.00 mL of 0.500 M HCl by 0.500 M NaOH)
14.0
12.0
Color change
alizarin yellow R
10.0
Color change
phenolpthalein
pH
8.0
Color change
bromthymol blue
equivalence
point
6.0
Color change
bromphenol blue
4.0
Color change
methyl violet
2.0
0.0
0.0
10.0
20.0
30.0
Volume of 0.500 M NaOH added
(mL)
Hill, Petrucci, General Chemistry An Integrated Approach 2nd Edition, page 680
Titration of a Weak Acid With a Strong Base
Titration of a Weak Acid With
a Strong Base
Titration Data
14.0
NaOH added
(mL)
12.0
10.0
pH
equivalence point
8.0
6.0
4.0
2.0
0.0
0.0
10.0
20.0
30.0
Volume of 0.100 M NaOH added
(mL)
40.0
0.00
5.00
10.00
12.50
15.00
20.00
24.00
25.00
26.00
30.00
40.00
pH
2.89
4.14
4.57
4.74
4.92
5.35
6.12
8.72
11.30
11.96
12.36
Phenolphthalein is best indicator: pH change 8.0 - 9.6
Titration of a Weak Base With a Strong Acid
Titration of a Weak Base With a Strong Acid
Titration Data
14.0
HCl added
(mL)
pH
0.00
10.00
20.00
30.00
40.00
45.00
47.00
48.00
49.00
50.00
51.00
11.24
9.91
9.47
8.93
8.61
8.30
7.92
7.70
7.47
5.85
3.34
12.0
pH
10.0
8.0
6.0
equivalence point
4.0
2.0
0.0
0.0
10.0
20.0
30.0
40.0
Volume of 0.100 M HCl added
(mL)
50.0
7. What is the pH of a solution made by dissolving 2.5 g NaOH in 400 mL water?
Determine number of moles of NaOH
 1 mol NaOH 

 40 g NaOH 
x mol NaOH = 2.5 g NaOH 
0.0625 mol NaOH
Calculate the molarity of the solution
M
mol
L

0.0625 mol NaOH
[Recall 1000 mL = 1 L]
0.4 L
MNaOH = 0.15625 molar
NaOH
0.15625 molar
Na1+ +
0.15625 molar
pOH = -log [OH-]
OH10.15625 molar
or
kW = [H+] [OH-]
pOH = -log [0.15625 M]
1 x 10-14 = [H+] [0.15625 M]
pOH = 0.8
[H+] = 6.4 x 10-14 M
pOH + pH = 14
pH = -log [H+]
0.8 + pH = 14
pH = 13.2
pH = -log [6.4 x 10-14 M]
What volume of 0.5 M HCl is required to titrate 100 mL of 3.0 M Ca(OH)2?
2 HCl
x mL
0.5 M
+
"6.0 M"
Ca(OH)2
100 mL
3.0 M
CaCl2
M1V1 = M2V2
(0.5 M) (x mL) = (3.0 M) (100 mL)
+ 2 HOH
M1V1 = M2V2
(0.5 M) (x mL) = (6.0 M) (100 mL)
x = 600 mL of 0.5 M HCl
mol
M
L
HCl
0.3 mol
x = 1200 mL of 0.5 M HCl
HCl
molHCl = M x L
Ca(OH)2
mol Ca(OH)2 = M x L
mol = (0.5 M)(0.6 L)
mol = (3.0 M)(0.1 L)
mol = 0.3 mol HCl
mol = 0.3 mol Ca(OH)2
H1+ +
0.3 mol
Cl1-
Ca(OH)2
0.3 mol
0.3 mol
[H+] = [OH-]
Ca2+ +
2OH1-
0.3 mol
0.6 mol
6.
10.0 grams vinegar titrated with 65.40 mL of 0.150 M NaOH
(acetic acid + water)
moles HC2H3O2 =
A)
moles NaOH
mol
M
L
NaOH
molNaOH = M x L
therefore, you have ...
0.00981 mol HC2H3O2
mol = (0.150 M)(0.0654 L)
mol = 0.00981 mol NaOH
 60 g HC 2 H 3 O 2 
 
1
mol
HC
H
O

2 3
2 
B) x g HC2H3O2 = 0.00981 mol HC2H3O2 
C)
% = 
0.59 g HC2H3O2

 x 100%
 whole 
part
 0.59 g acetic acid 
 x 100%
 10.0 g vinegar

% = 
% =
5.9 % acetic acid
Commercial vinegar is sold as 3 - 5 % acetic acid
1) What is the pH of a solution made by combining 49 mL of 0.2 M HCl
with 50 mL of 0.2 M NaOH?
H2O + NaCl
HCl + NaOH
49 mL 0.2 M HCl
+
50 mL 0.2 M NaOH
B) molNaOH = M . L
A) molHCl = M . L
molHCl = (0.2 M) . (0.049 L)
molNaOH = (0.2 M) . (0.05 L)
molHCl = 0.0098 mol
molNaOH = 0.010 mol
49 mL
0.2 M HCl
50 mL
0.2 M NaOH
99 mL H2O
1 mL of 0.2 M NaOH
0.010 mol OH1- 0.0098 mol H1+
“net”
0.0002 mol OH1-
1) What is the pH of a solution made by adding 1mL
of 0.2 M NaOH with 99 mL H2O?
1) What is the pH of a solution made by adding
1mL of 0.2 M NaOH with 99 mL H2O?
NaOH  Na1+ + OH1-
49 mL
0.2 M HCl
50 mL
0.2 M NaOH
99 mL H2O
1 mL of 0.2 M NaOH
Calculate the molarity of the solution
M=
mol
M
mol
L
M=
0.0002 mol NaOH
0.0099 L
[Recall 1000 mL = 1 L]
0.010 mol OH1- 0.0098 mol H1+
MNaOH = 0.002020 molar
L
NaOH
0.002020 molar
Na1+ +
OH1-
“net”
0.0002 mol OH1-
0.002020 molar 0.002020 molar
pOH = -log [OH-]
or
kW = [H+] [OH-]
pOH = -log [0.002020 M]
1 x 10-14 = [H+] [0.002020 M]
pOH = 2.7
[H+] = 4.95 x 10-12 M
pOH + pH = 14
pH = -log [H+]
2.7 + pH = 14
pH = 11.3
pH = -log [4.95 x 10-12 M]
Carboxylic Acid
HC2H3O2
= acetic acid
H
O
H C C
1O
H
H
H+
CH3COOH
R - COOH
carboxylic acid
C2H4O2
Lactic Acid
OH
H3C
C
CO2H
H
Lactic acid
C3H6O3
Titration
?
1.0 M HCl
titrate with
? M NaOH
2.00 mL
1.00 mL
M1 V1 = M2 V2
(1.0 M)(1.00 mL) = (x M)(2.00 mL)
2.0 M
H1+
1.0 M H2SO4
X = 0.5 M NaOH
titrate with
? M NaOH
2.00 mL
1.00 mL
M1 V1 = M2 V2
(1.0 M)(1.00 mL) = (x M)(2.00 mL)
X = 0.5 M NaOH
Base
vinegar
ammonia
Calibration Curve
Vinegar
1 mL
3 mL
5 mL
10 mL
15 mL
Ammonia
ammonia
Acid
vinegar
Using 3 mL vinegar… titrate with 0.130 M NaOH solution.
% acetic(M)
acidofinacetic
vinegar.
=1part
Calculate molarity
acid.% M
V1/=
M2x100
V2
whole
A)
moles HC2H3O2 =
moles NaOH
NaOH
molNaOH = M x L
mol
M
mol = (0.130 M)(0.0196 L)
L
mol = 0.002548 mol NaOH
therefore, you have ...
0.002548 mol HC2H3O2
 60 g HC 2 H 3 O 2 


 1 mol HC 2 H 3 O 2 
B) x g HC2H3O2 = 0.002548 mol HC2H3O2 
C)
0.153 g HC2H3O2
 part

 x 100%
 whole 
% = 
 0.1529 g acetic acid 
 x 100%
3.0
g
vinegar


% =
% =
5.1 % acetic acid
Commercial vinegar is sold
as 3 - 5 % acetic acid
Base
vinegar
ammonia
Calibration Curve
Vinegar
1 mL
3 mL
5 mL
10 mL
15 mL
Ammonia
ammonia
Acid
vinegar
Using 3 mL vinegar… titrate with 0.130 M NaOH solution.
Calculate molarity (M) of acetic acid. M1V1 = M2V2
It required 19.6 mL of NaOH to reach the endpoint.
M1 V1 = M2 V2
(Macetic acid)(3.0 mL) = (0.130 MNaOH )(19.6 mL)
Macetic acid = 0.8493 molar
HCl
Hydrochloric acid
stomach acid, pickling metal
H2SO4
Sulfuric acid
battery acid, # 1 selling chemical
H3PO4
Phosphoric acid
food flavoring
HNO3
Nitric acid
fertilizer, explosives
CH3COOH
Acetic acid
vinegar
HF
Hydrofluoric acid
etch glass
NaOH
Ca(OH)2
NH4OH
sodium hydroxide
calcium hydroxide
ammonium hydroxide
Soren Sorenson developed pH scale
7
neutral
pH scale
0
[H+] = [OH-]
acid
14
base
(alkalinity)
Arnold Beckman invented the pH meter
pH = -log [H+]
pOH = -log [OH-]
pH + pOH = 14
kW =
[H+]
[OH-]
kw = 1 x 10-14
H+ + H2O
proton
H3O+
hydronium ion
Concentrated vs. Dilute
Concentration:
Molarity
molality
mol
M= L
mol
m = kg
H2SO4
2 H1+
3M
“6 M”
+
SO42-
Normality
Strong / Weak Acid
Strong
HA
H+
+
A-
(~100% dissociation)
Weak
HA
H+
+
A-
(~20% dissociation)
H2A
2 H+
[Product]
Ka = [Reactant]
+
Ka =
A-
[H+]2 [A-]
[H2A]
acid dissociation constant
Ka
0.8
0.0021
H3PO4
HF
3H+ + PO43-
H+ + F -
Acid
+
Base
Salt
+
Water
How would you make calcium sulfate in the lab?
H2?SO4
+
Ca(OH)
?
2
BASE
ACID
Sour taste, litmus
CaSO4 + 2 H2O
red
bitter taste, litmus
blue
Arrhenius – H+ as only ion in water
Arrhenius – OH- as only ion in water
Brønsted-Lowry – proton donor
Brønsted-Lowry – proton acceptor
Indicators
colorless
weak acid
phenolphthalein
yellow
strong acid
bromthymol blue
universal indicator
&
blue
strong base
R O Y G B I V
pH 4
litmus paper
pink
strong base
pH paper
7
12
Buffers - salts of weak acids and weak bases that maintain a pH
e.g. Aspirin (acetyl salicylic acid) vs. Bufferin
low pH upsets stomach
LeChatelier’s Principle
- acidosis & alkalosis (bicarbonate ion acts as buffer)
- darkening glasses
- egg shells thinner in summer (warm)
Amino Acids – Functional Groups
Amine
Base Pair
Carboxylic Acid
R- COOH
NH21lose H+
NH21-
H+
NH3
amine
ammonia
NH41+
ammonium ion
1+
:
1-
N
N
N
H
H
:
:
H
H
H
H
H
H
H
Water – Amphiprotic
lose H+
OH1-
H+
hydroxide
H2O
H3O1+
water
hydronium ion
:
N
H
H
:
:
N
H
1+
H
1-
N
H
H
H
H
H
Water – Also Amphiprotic
Amphiprotic – Act as an acid (proton donor) or base (proton acceptor)
OH1-
hydroxide
lose H+
H+
H2O
H3O1+
water
hydronium ion
+
d+ H
H+ O2- d
Range and Color Changes of Some
Common Acid-Base Indicators
pH Scale
1
2
3
4
5
6
7
8
9
10
11
12
13
14
Indicators
Methyl orange
Methyl red
Bromthymol blue
3.1 – 4.4
red
red
4.4
yellow
Neutral red
red
Phenolphthalein
colorless
yellow
6.2
6.2
6.8
yellow
7.6
blue
8.0
8.0
yellow
10.0
red
colorless beyond 13.0
Neutralization of Bug Bites
Wasp - stings with base
Red Ant - bites with acid
(neutralize with lemon juice or vinegar)
(neutralize with baking soda)
Strength
 Strong
Acid/Base
• 100% ionized in water
• strong electrolyte
HCl
HNO3
H2SO4
HBr
HI
HClO4
-
+
NaOH
KOH
Ca(OH)2
Ba(OH)2
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Strength
 Weak
Acid/Base
• does not ionize completely
• weak electrolyte
HF
CH3COOH
H3PO4
H2CO3
HCN
-
+
NH3
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Ionization of Water
H 2O + H 2O
Kw =
+
[H3O ][OH ]
H3
+
O
+
= 1.0 
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
OH
-14
10
Why is pure water pH = 7?
1
in 500,000,000 water molecules will
autoionize.
 H2O + H2O  H3O+ + OH1 This yields a hydronium ion concentration
of 1 x 10-7 M H3O+ per liter of solution
 pH = -log[H3O+]
 pH = -log[1 x 10-7] or pH = 7
H
H
O
1-
H
O
H
H
H
H
H
1+
O
H
O
O
O
H
O
H
H
H
H
H
O
H
H
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
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
pH Scale
14
0
7
INCREASING
ACIDITY
pH =
NEUTRAL
INCREASING
BASICITY
+
-log[H3O ]
pouvoir hydrogène (Fr.)
“hydrogen power”
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
pH Scale
pH =
+
-log[H3O ]
pOH =
-log[OH ]
pH + pOH = 14
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
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
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
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
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem