The Chemistry of
Acids and Bases
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Acid and Bases
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Acid and Bases
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Some Properties of Acids
Produce H+ (as H3O+) ions in water (the hydronium ion is a hydrogen
ion attached to a water molecule)
Taste sour
Corrode metals
Electrolytes
React with bases to form a salt and water
pH is less than 7
Turns blue litmus paper to red “Blue to Red A-CID”
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Some Properties of Bases
Produce OH- ions in water
Taste bitter, chalky
Are electrolytes
Feel soapy, slippery
React with acids to form salts and water
pH greater than 7
Turns red litmus paper to blue
“Basic Blue”
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Some Common Bases
NaOH
sodium hydroxide
lye
KOH
potassium hydroxide
liquid soap
Ba(OH)2
barium hydroxide
stabilizer for plastics
Mg(OH)2
magnesium hydroxide “MOM” Milk of magnesia
Al(OH)3
aluminum hydroxide
Maalox (antacid)
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Acid/Base definitions
• Definition #1: Arrhenius (traditional)
Acids – produce H+ ions (or hydronium ions
H3O+)
acids can be classified as monoprotic,
diprotic or triprotic
Bases – produce OH- ions
problem: some bases don’t have
hydroxide ions
Nice, simple definitions but has limitations…
restricted to aqueous solutions
couldn’t explain basic properties of NH3
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Acid/Base Definitions
• Definition #2: Brønsted – Lowry
Acids – proton donor
Bases – proton acceptor
A “proton” is really just a hydrogen
atom that has lost it’s electron!
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A Brønsted-Lowry acid is a proton donor
A Brønsted-Lowry base is a proton acceptor
base
acid
conjugate
acid
conjugate
base
Conjugate Pairs
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Bronsted-Lowry A-B
• Because many A/B reactions are
reversible, it is useful to identify
conjugates (opposites) that will react in
the reverse reaction
• conjugate A-B pairs = 2 formulas in an
equation whose formulas differ by a H+
• strong A-B are not equilibrium
expressions, but all other A-B are
reversible
Acids & Base Definitions
Definition #3 – Lewis
Lewis acid - accepts an
electron pair
Lewis base - donates
an electron pair
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Lewis Acid/Base Reaction
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Lewis Acid-Base Interactions
in Biology
Heme group
• The heme group
in hemoglobin
can interact with
O2 and CO.
• The Fe ion in
hemoglobin is a
Lewis acid
• O2 and CO can
act as Lewis
bases
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Brønsted-Lowry vs. Lewis
• All B/L bases are Lewis bases BUT, by
definition, a B/L base cannot donate its
electrons to anything but a proton (H+)
• While B/L is most useful for our purposes,
Lewis allows us to treat a wider variety of
reactions (even if no H+ transfer occurs) as
A/B reactions
Strong Acids/Bases
The strength of an acid (or base) is
determined by the amount of
IONIZATION (Not pH value).
HCl, HBr, HI, HNO3, H2SO4 HClO3 and HClO4
are the strong acids.
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Strong Acids/Bases
The strong bases are:
alkali-metal hydroxides and
alkaline earth metal hydroxides.
Ex. NaOH, KOH, CsOH, Ca(OH)2, etc.
Note: The stronger the acid/base, the
weaker its conjugate pair.
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The pH scale is a way of
expressing the strength
of acids and bases.
Instead of using very
small numbers, we just
use the NEGATIVE
power of 10 on the
Molarity of the H+ (or
OH-) ion.
Under 7 = acid
7 = neutral
Over 7 = base
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pH of Common
Substances
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pH and acidity
1. Acidity or Acid Strength depends on Hydronium Ion
Concentration [H3O+]
2. The pH system is a logarithmic representation of
the Hydrogen Ion concentration (or OH-) as a means
of avoiding using large numbers and powers.
pH
= - log [H3O+]
pOH = - log [OH-]
3. In pure water [H3O+] = 1 x 10-7 mol / L (at 25oC)
 pH = - log(1 x 10-7) = - (0 - 7) = 7
4. pH range of solutions:
0 - 14
pH < 7 (Acidic) [H3O+] > 1 x 10-7 m / L
pH > 7 (Basic) [H3O+] < 1 x 10-7 m / L
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pH calculations – Solving for H+
If the pH of Coke is 3.12, [H+] = ???
Because pH = - log [H+] then
- pH = log [H+]
Take antilog (10x) of both
sides and get
10-pH = [H+]
[H+] = 10-3.12 = 7.6 x 10-4 M
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pOH
• Since acids and bases are
opposites, pH and pOH are
opposites!
• pOH does not really exist, but it is
useful for changing bases to pH.
• pOH looks at the perspective of a
base
pOH = - log [OH-]
Since pH and pOH are on opposite
ends,
pH + pOH = 14
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pH
[H+]
[OH-]
pOH
pH indicators
• Indicators are dyes that can be
added that will change color in
the presence of an acid or base.
• Some indicators only work in a
specific range of pH
• Once the drops are added, the
sample is ruined
• Some dyes are natural, like radish
skin or red cabbage
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Indicators
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Wait, water can go both ways?
• amphoteric substances can behave as either
an acid or base depending on what they react
with.
• water and anions with protons (H+) attached
are most common amphoterics
Autoionization of Water
H2O + H2O
OH- + H3O+
OH -
H3O+
@ equilibrium Keq = [H3O+] [OH-] = so Keq
becomes Kw and equals 1.00 x 10-14 at 25 oC
In a neutral solution [H3O+] = [OH-]
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Weak Acids/Bases
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• Only partially ionize (less than 100%) in water.
• Acids (symbol=HA)
– @ equilibrium Keq = [H3O+] [OH-]/[HA] = so Keq
becomes Ka
– The bigger the Ka the stronger the acid
– Polyprotic acids have more than 1 H+ available to
ionize
» Have multiple Ka values (Ka1 Ka2 Ka3)
» Always easier to remove 1st H+ than 2nd
» If Ka2 is different than Ka1 by 103 (or more) consider only
Ka1 to get pH
Weak Acids/Bases
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• Bases (symbol=B)
– 2 general categories
» Contain neutral substances that have an atom with
non-bonding electrons Ex. NH3  NH4+
» Anions of weak acids Ex. ClO-  HClO
– @ equilibrium Keq = [HB+] [OH-]/[B] = so Keq becomes
Kb
• Relationship between Ka and Kb
– Ka x Kb = Kw
– pKa + pKb = pKw = 14.00
Equilibrium Constants
for Weak Acids
Weak acid has Ka < 1
Leads to small [H3O+] and a pH of 2 - 7
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Equilibrium Constants
for Weak Bases
Weak base has Kb < 1
Leads to small [OH-] and a pH of 7-12
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Relation
of Ka, Kb,
[H3O+]
and pH
Equilibria Involving A Weak Acid
You have 1.00 M HA. Calc. the equilibrium
concs. of HA, H3O+, A-, and the pH.
H3O+ + A-
HA + H2O
Step 1. Define equilibrium concs. in ICE
table
[HA]
[H3O+]
[A-]
initial
1.00
0
0
change
-x
+x
+x
equilib
1.00-x
x
x
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Equilibria Involving A Weak Acid
Step 2. Solve equilibrium (Ka) expression
+
-
2
[H3O ][A ] x
Ka 1.8 x 10 =

[HA] 1.00 - x
-5
This is a quadratic. Solve using quadratic formula.
because Ka for most weak acids is less than 10-3, 1x is about equal to 1, so Ka = x2/1.00

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Equilibria Involving A Weak Acid
So the Ka expression
+
-
2
[H3O ][A ] x
Ka 1.8 x 10 =

[HA] 1.00 - x
-5
becomes
Ka  1.8 x 10-5 =

x2
1.00
Now we can more easily solve this approximate
expression.
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Equilibria Involving A Weak Acid
Step 3. Calculate the pH
Ka  1.8 x 10-5 =
x2
1.00
x = [H3O+] = [A-] = 4.2 x 10-3 M
pH = - log [H3O+] = -log (4.2 x 10-3) = 2.37
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For each problem, identify if the
short-cut is a valid method
• What is the pH of a 0.12 M solution of
hypochlorous acid?
• What is the pH of a 0.12 M solution of lactic
acid?
• What is the pH of 0.12 M solution of chlorous
acid?
• What is the pH of a 0.12 M solution
hydrochloric acid?
Equilibria Involving A Weak Base
You have 0.010 M of B. Calc. the pH.
B + H2 O
OH- + HB+
Kb = 1.8 x 10-5
Step 1. Define equilibrium concs. in ICE table
[B]
[HB+]
[OH-]
initial
0.010
0
0
change
-x
+x
+x
equilib
0.010 - x
x
x
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
Equilibria Involving A Weak Base
Step 2. Solve the equilibrium expression
+
-
2
[HB ][OH ]
x
Kb  1.8 x 10 =
=
[B]
0.010 - x
-5
b/c Kb for most weak bases is less than 10-3, 0.010x is about equal to 0.010, so Kb = x2/0.010
x = [OH-] = [HB+] = 4.2 x 10-4 M and [B] = 0.010 4.2 x 10-4 ≈ 0.010 M
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Equilibria Involving A Weak Base
Step 3. Calculate pH
[OH-] = 4.2 x 10-4 M
so pOH = - log [OH-] = 3.37
Because pH + pOH = 14,
pH = 10.63
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A-B properties of salt solutions
• for the most part anions are slightly basic
(because they attract protons) and cations are
slightly acidic (because they can donate protons)
• Anions
–
–
–
–
Look at strength of the acid to which it is a conjugate
If acid is strong, anion won’t really take H+ from water
If acid is weak, anion will react to a small extent ( pH)
If amphoteric need to look at magnitudes of Ka and Kb
• Cations
– Polyatomic cations can be the conjugate acid of a weak base
– Usually react to make H3O+ (  pH)
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Predict whether each of the
following salts will result in an
acidic, basic, or neutral solution
when dissolved in water
•
•
•
•
Fe(NO3)3
BaCl2
NaHSO4
LiF
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A/B Behavior & Chemical
Structure
1. Polarity
-
if H is the positive end = acid
if H is negative end (NaH) = base
if non-polar, neither acid or base
2. Strength of bond
-
Look at what H is bonded to
How easy is it to dissociate into ions?
3. Stability of conjugate: the greater the
stability of the conjugate, the stronger the
acid
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A/B Behavior & Chemical Structure
Binary Acids
1. HX bond strength is most important
2. acidity increases down a group
- larger elements
- longer and weaker bonds
3. in same row, look at bond polarity
- across from L to Right increases EN
- Across from L to R increases acidity
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A/B Behavior & Chemical Structure
Oxyacids (have O in compound)
1. Have same # of OH groups and the same #
of O atoms, acid strength  with  in EN of
center atom
2. Have same center atom (Y) acid strength 
as the # of O atoms attached to Y 
3. in a series of oxyacids, acidity  with as
the oxidation # of the center atom 
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A/B Behavior & Chemical
Structure
Carboxylic Acids (have –COOH in compound)
1. H on OH group is ionized
-
the other O atom draws elctron density away from OH bond
increasing polarity
the conj. base (COO-) can exhibit resonance making it
more stable
2. Acid strength of COOH group increases as the # of
electronegative atoms in the acid increase.
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Arrange the following series in
order of increasing acid
strength
• AsH3, HI, NaH, H2O
• H2SeO3, H2SeO4, H2O
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In each pair, choose the
compound that leads to the
more acidic solution
• HBr, HF
• PH3, H2S
• HNO2, HNO3
• H2SO3, H2SeO3