Background
A(g) + 2 B(g)
3 C(g) + D(g)
Equilibrium constant (Keq)
Keq =
=
[Products]
[Reactants]
[C] 3[D]
[A][B] 2
A florence flask was getting dressed
for the opera. All of a sudden she
screamed: "Erlenmeyer, my joules!
Somebody has stolen my joules!".
The husband replied: "Take it easy
honey, do not overreact. We'll find a
solution".
LeChatelier’s Principle
(lu-SHAT-el-YAY’s)
Acids, Bases, and Salts
Ch. 19
Acids
• Properties
– Taste sour or tart
– Change the color of an acid-base indicator
– Can be strong or weak electrolytes in
aqueous solution
Bases
• Properties
Q: Why do chemistry professors like to
teach about ammonia?
A: Because it's basic stuff.
– Taste bitter
– Feel slippery
– Will change the color of an acid-base indicator
– Can be strong or weak electrolytes in aqueous
solution
During class, the chemistry professor was
demonstrating the properties of various
acids. “Now I’m going to drop this silver coin
into this glass of acid. Will it dissolve?”
“No sir,” one student called out.
“No?” queried the professor. “Perhaps you
can explain why the silver won’t dissolve in
this particular acid.”
“Because if it would, you wouldn’t have
dropped it in!”
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 Base (weak electrolyte)
NH3
ammonia
NH3 + H2O  NH4OH
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
Common Bases
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
NH3 H2O
household ammonia
Have you heard the one about a chemist
who was reading a book about helium
and just couldn't put it down?
Arrhenius Acids and Bases
• Arrhenius said that acids are hydrogen-containing compounds
that ionize to yield hydrogen ions (H+) in aqueous solution
• Bases are compounds that ionize to yield hydroxide ions (OH-) in
aqueous solution
• Monoprotic acids = acids that contain 1 ionizable hydrogen like
nitric acid (HNO3)
• Diprotic acids = acids that contain 2 ionizable hydrogens like
sulfuric acid (H2SO4)
• Triprotic acids = acids that contain 3 ionizable hydrogens like
phosphoric acid (H3PO4)
• Hydroxides of group I metals are very soluble in water and caustic
to skin, hydroxide of group II metals are not very soluble in water
and very dilute even when saturated (can be taken internally)
Q: if both a bear in Yosemite and one in Alaska
fall into the water
which one disolves faster?
A: The one in Alaska because it is Polar.
Bronsted-Lowry Acids and Bases
• But what about bases like sodium carbonate
(Na2CO3) and ammonia (NH3)???
– The Bronsted-Lowry theory defines an acid as a
hydrogen-ion donor, and a base as a hydrogenion acceptor
– More complete definition
Conjugate Acids and Bases
• Conjugate acid = the particle formed when a base
gains a hydrogen ion
• Conjugate base = the particle that remains when
an acid has donated a hydrogen ion
– Conjugate acids and bases are always paired w/ a
base or an acid, respectively
• Conjugate acid-base pair = consists of 2
substances related by the loss or gain of a single
hydrogen ion.
Cont…
• A water molecule that gains a
hydrogen ion becomes a
positively charged
Hydronium ion (H3O+)
• Amphoteric = a substance
that can act as both an acid
and a base – EX: Water
Lewis Acids and Bases
• Lewis proposed that an acid accepts a pair of
electrons during a reaction while a base
donates a pair of electrons
– More general than either of the other 2 theories
• Lewis acid = a substance that can accept a pair
of electrons to form a covalent bond
• Lewis base = a substance that can donate a
pair of electrons to form a covalent bond
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
• Many Lewis acids are also Bronsted-Lowry
acids and vice versa but not all
• *PP 1-2, 19.1 sect. assessment #8 pg. 593
Copper leaves Copper Sulfate and says see
you: he answers CuSO4!!!!!
Ion Product Constant for Water
• Self-ionization of water = the reaction in which water molecules
produce ions
• For aqueous solutions, the product of the hydrogen-ion
concentration and the hydroxide-ion concentration equals
1.0 x 10-14
– [H+] x [OH-] = 1.0 x 10-14
• ion-product constant for water (KW) = the product of the
concentrations of the hydrogen ions and hydroxide ions in water
• Acidic solution = one in which [H+] is greater than [OH-]
– *The [H+] is greater than 1 x 10-7 M*
• Basic solution = one in which [H+] is less than [OH-]
– *The [H+] is less than 1 x 10-7 M*
• *SP 19.1, PP 9-10 pg. 596
Basic
7
Acid
14
Neutral
pH Scale
Acidic
0
Base
[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
urine
6.0
acidic
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
Søren Sorensen
(1868 - 1939)
pH Concept
• pH = the negative logarithm of the hydrogenion concentration of a solution
• A solution in which [H+] if greater than
1 x 10-7 M has a pH less than 7.0 and is acidic.
The pH of pure water or a neutral aqueous
solution is 7.0. A solution with a pH greater
than 7 is basic and has a [H+] of less than 1 x
10-7 M.
Cont…
• The pOH of a solution equals
the negative logarithm of the
hydroxide-ion concentration
– A solution w/ a pOH less than 7
is basic, greater than 7 is acidic
• For pH calculations, you should
express the hydrogen-ion
concentration in scientific
notation
• *SP 19.2, PP 11 pg. 599
• *Given pH = 4.6 determine the
hydronium ion
• *SP 19.3-19.4, PP 13-16 pg.
600-601
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
Strength of Acids and Bases
• Strong Acids = completely ionized in aqueous
solution
– HCl and H2SO4
• Weak Acids = ionize only slightly in aqueous
solution
– Ethanoic acid (acetic acid)
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+
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
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
Acid Dissociation Constant
• For dilute solutions, the conc. of water is a constant. It can
be combined w/ Keq to give the acid dissocation constant.
• Acid dissociation constant = the ratio of the concentration
of the dissociated (or ionized) form of an acid to the
concentration of the undissociated (nonionized) form.
• Weak acids have small Ka values. The stronger an acid is,
the larger is its Ka value.
– Nitrous acid (HNO2) has a Ka of 4.4 x 10-4, acetic acid has a
Ka of 1.8 x 10-5 so nitrous acid is more ionized and has a
higher [H3O+] or [H+] thus is a stronger acid
• Di and triprotic acids lose each H separately so they have
multiple dissociation constants
Equilibrium and pH Calculations
Weak acid
Strong acid
HA
H + + AHA + H2O
H3O+ + A-
H3O+ + A-
HA + H2O
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-]
=
-
Base Dissociation Constant
• Strong bases = dissociate completely into metal
ions and hydroxide ions in aqueous solution
– Ex: Ca(OH)2
• Weak bases = react w/ water to form the
hydroxide ion and the conjugate acid of the base
– Ex: ammonia NH3
• Base dissociation constant (Kb) = the ratio of the
concentration of the conjugate acid times the
conc. of the hydroxide ion to the conc. of the
base
Calculating Dissociation Constants
• To find the Ka of a weak acid or the Kb of a
weak base, substitute the measured
concentration of all the substances present at
equilibrium into the expression for Ka or Kb.
• *SP 19.5, PP 22-23 pg. 610
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)
Formula
Name
Value of Ka*
HSO4HClO2
HC2H2ClO2
HF
HNO2
HC2H3O2
HOCl
HCN
NH4+
HOC6H5
hydrogen sulfate ion
chlorous acid
monochloracetic acid
hydrofluoric acid
nitrous acid
acetic acid
hypochlorous acid
hydrocyanic acid
ammonium ion
phenol
1.2 x 10-2
1.2 x 10-2
1.35 x 10-3
7.2 x 10-4
4.0 x 10-4
1.8 x 10-5
3.5 x 10-8
6.2 x 10-10
5.6 x 10-10
1.6 x 10-10
*The units of Ka are mol/L but are customarily omitted.
Increasing acid strength
Values of Ka for Some Common Monoprotic Acids
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
[H  ][C2H3O2 ]
Ka =
[HC2H3O2 ]
1
1.8 x
10-5
[H  ][C2H3O2 ]
=
[HC2H3O2 ]
How do the concentrations of
H+ and C2H3O21- compare?
[x][x]
[HC2H3O2 ]
Substitute into equation:
1.8 x 10-5 
1.8 x 10
-5
x2

[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
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 =
[Products]
Ka =
[Reactants]
[4.8 x 10-3 M]
[4.8 x 10-3 M] [CN1-]
Ka =
+ CN1-(aq)
H+(aq)
HCN(aq)
[0.150 M]
Ka = 1.54 x 10-4
4.8 x 10-3 M
[H3O+] [CN1-]
pH = -log[H3O+]
[HCN]
10-pH = [H3O+]
10-2.32 = [H3O+]
4.8 x10-3 M = [H3O+]
Titration
Q: How did the chemist
survive the famine?
A: By subsisting on titrations.
• Neutralization Reaction = a reaction b/w an acid and a
base in aqueous solution to produce salt and water
• Equivalence point = when the # of moles of hydrogen
ions equals the number of moles of hydroxide ions
• Titration = the process of adding a known amt. of
solution of known conc. to determine the conc. of
another solution
– Standard solution = the solution of known conc.
– End point = the point at which the indicator changes color
• The point of neutralization is the end pt. of the
titration
• *SP 19.7, PP 32-33 pg. 616 Q: How do you get lean molecules?
A: Feed them titrations.
Q: What did one titration say to the other?
A: Let's meet at the endpoint!
Buffers
• Buffer = a solution in which the pH remains
relatively constant when small amts. of acid or
base are added
• A buffer is a solution of weak acids and one of
its salts or a solution of a weak base and one of
its salts
• A buffer solution is better able to resist drastic
changes in pH than is pure water
• Buffer capacity = the amt. of acid or base that
can be added to a buffer solution before a
significant change in pH occurs
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)
Review
A physicist, biologist and a chemist were going
to the ocean for the first time.
What does one do with a dead body? Barium
in a krypt-on
Maybe he was killed oxydentally.
They should have seen the doctor first, he'd
Curium.
Ah, barium anyway, just to see how he reacts.
better though to have helium.
Perhaps with a houseplant, a Germanium.
And if they stole it, the police would Cesium.
Locked up for life, in Irons.
They would go crazy in jail, a Silicon.
The physicist saw the ocean and was
fascinated by the waves. He said he wanted to
do some research on the fluid dynamics of the
waves and walked into the ocean. Obviously he
was drowned and never returned.
The biologist said he wanted to do research on
the flora and fauna inside the ocean and
walked inside the ocean. He too, never
returned.
The chemist waited for a long time and
afterwards, wrote the observation, "The
physicist and the biologist are soluble in ocean
water".
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
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+ + PO43H + + F-
Acid
+
Base
Salt
+
Water
How would you make calcium sulfate in the lab?
H2SO4
+
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