Chapter 17 Principles of Reactivity: Chemistry of Acids and Bases

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Chapter 17 Principles of Reactivity Chemistry
of Acids and Bases
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Acids Bases

• Acids are some of peoples' favorite chemicals.
  Everyone's favorite soft drink is a dilute acid
  solution. Your own stomach contains the strong
  acid HCl. Citrus fruits contain citric acid. If
  wine is too aged - exposed to oxygen, it turns
  sour - it forms acetic acid. Sulfuric Acid is the
  top commercially produced chemical in the United
  State. Although much of it is used in the steel
  and petroleum refining industries, several
  million tons of sulfuric acid are used to make
  Jello.

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PROPERTIES OF ACIDS AND BASES

• ACIDS

• BASES

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Arrhenius definition

• acid
• produces hydronium ion (H3O) in aqueous solution
• base
• produces hydroxide ion (OH) in aqueous solution

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Brønsted definition
donates a proton (hydrogen ion, H) accepts a
proton
A conjugate acid is formed by adding a proton to
something. A conjugate base is formed by
removing a proton from something.
acid
conjugate acid
base
conjugate base
base
conjugate acid
acid
conjugate base
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Relative strengths of conjugate acid-base pairs
???? If HA is a stronger acid then A is a
weaker base. If HA is a weaker acid then A is a
stronger base.
???? If B is a stronger base then BH is a
weaker acid. If B is a weaker base then BH is a
stronger acid.
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• You should memorize the names and formulas of the
  6 STRONG ACIDs , i.e., HCl, HBr, HI, HClO4,
  HNO3, and H2SO4.
• The organic acid present in vinegar, acetic acid,
  is a common WEAK ACID.
• The common STRONG BASES contain the hydroxide ion
  (OH-).
• Ammonia (NH3), a common WEAK BASE, is that smelly
  stuff your Grandma used in a dilute solution to
  clean windows

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Ion Product Constant of Water

• Water is an important solvent.
• Universal solvent
• Biological solvent
• Small size
• Density of water is greater than ice
• Very polar
• Hydrogen Bonding

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Self-ionization of Water

• Water is an amphiprotic substance that can act
  either as an acid or a base.
• HC2H3O2(aq) H2O(l) H3O
  C2H3O2-(aq)
• acid base acid
  base
• H2O(l) NH3(aq) NH4(aq)
  OH-(aq)
• acid base acid
  base

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Self-ionization of Water

• When water molecules react with one another to
  form ions.
• H2O(l) H2O(l) H3O(aq)
  OH-(aq)
• 
  (10-7M) (10-7M)
• Kw H3O OH-
• 1.0 x 10-14 at 25oC
• Note H2O is constant and is already
• included in Kw.

ion product of water
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pH and pOH

• We need to measure and use acids and bases over a
  very large concentration range.
• pH and pOH are systems to keep track of these
  very large ranges.
• pH - log H3O
• pOH - log OH-
• pH pOH 14

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pH Calculations

• Determine the following. pH - log
  H
• pH of 6.7x10-3 M H
• 
  2.2
• pH of 5.2x10-12 M H
• 
  11.3
• H, if the pH is 4.5
• 
  3.2 x 10-5 M H

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pOH Examples

• Determine the following.
• pOH - log OH-
  14 - pH
• pOH of 1.7 x 10-4 M NaOH
• pOH 3.8
  pH 10.2
• pOH of 5.2 x 10-12 M H
• pOH 2.7
  pH 11.3
• OH- , if the pH is 4.5
• pOH 9.5
  
• 
  OH- 3.2 x 10-10 M

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pH Scale

• A log based scale used to keep track of the large
  change important to acids and bases.

14 7
0
10-14 M 10-7 M
1 M Very Neutral
Very Basic Acidic
When you add an acid, the pH gets smaller. When
you add a base, the pH gets larger.
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pH of SomeCommon Materials

• Substance pH
• 1 M HCl 0.0
• Lemon juice 2.3
• Coffee 5.0
• Pure Water 7.0
• Blood 7.35 - 7.45
• Milk of Magnesia 10.5
• 1M NaOH 14.0

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Definitions Ka, pKa
HA H2O
H3O A
Kc
pKa -log Ka if pKa 5 then Ka 105 if
pKa 8 then Ka 108
stronger acid weaker acid
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Acid Ionization Constant, Ka

• Acid ionization constants let us define weak,
  moderate and strong acids.
• Ka lt 10-3 it is a weak acid.
• Ka 10-3 to 1 it is a moderate acid.
• Ka gt 1 it is a strong acid.

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Definitions Kb, pKb
B H2O
BH OH
Kc
pKb -log Kb if pKb 4 then Kb 104 if
pKb 9 then Kb 109
stronger base weaker base
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Ka and Kb Values

• For weak acids and bases
• Ka and Kb always have values that are smaller
  than one.
• Acids with a larger Ka are stronger than ones
  with a smaller Ka.
• Bases with a larger Kb are stronger than ones
  with a smaller Kb.
• Ka x Kb Kw
• Most acids and bases are considered weak.

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pKa and pKb Concepts

• The negative logarithms of Ka and Kb are useful
  in the same way as pH.
• pKa - log Ka
• pKb - log Kb
• pKa pKb 14.00
• The larger that the value of pKa is, the weaker
  the acid.
• The larger that the value of pKb is, the weaker
  the base.

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Kw autodissociation of water
H2O H2O
H3O OH
pH 14 - pOH
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Ka and Kb for conjugate acid-base pairs
HA
H3O A
Ka
H3OOH Kw
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III. pH Calculations
A. Strong acids and bases
100 dissociated ? for strong acid H3Oeq
HAI base OHeq Bi
e.g., 1.0 x 103 M HCl H3O pH OH pOH

e.g., 2.5 x 102 M NaOH OH pOH H3O
pH
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III. pH Calculations
B. Weak acids and bases
H3OA HA
HA
H3O A
Ka
Solve equilibrium expressions
BHOH B
B
BH OH
Kb
e.g., What is the pH of 0.10 M HC2H3O2? (Ka
1.8 x 105)
x H3O 1.3 x 103 M (assumption valid) pH
2.87
assume x ltlt 0.10
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Buffers

• Solutions that resist change in pH when small
  amounts of acid or base are added.
• Two types
• weak acid and its salt.
• weak base and its salt.
• HA(aq) H2O(l) H3O(aq)
  A-(aq)
• Add OH- Add
  H
• shift to right shift to
  left
• Based on Le Châteliers Principle.

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III. pH Calculations
C. Polyprotic acids
H2SO4, H2SO3, H2CO3, etc.
Lose their protons in separate steps
(Usually, Ka1 gtgt Ka2)
Assume 1) H2A, H3O, and HA can be
determined from the 1st step. (i.e., HA
dissociates only very little.) 2) A2 can be
determined from the 2nd step.
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Buffers and Blood

• Control of blood pH.
• Oxygen is transported primarily by hemoglobin in
  the red blood cells.
• CO2 transported both in plasma and the red blood
  cells.
• CO2 (aq) 2 H2O
• H2CO3 (aq)
• H3O(aq) HCO3-(aq)

Carbonate Buffer
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Buffers and Blood

• The amount of CO2 helps control blood pH.
• Too much CO2 - Respiratory arrest,
• pH goes down, acid level goes up.
• acidosis
• Solution - ventilate and give bicarbonate
  via IV.
• Too little CO2 - Hyperventilation, anxiety,
• pH goes up, acid level goes down.
• alkalosis
• Solution - re-breathe CO2 in paper bag to raise
  level.

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Quantitative Aspects of Buffers

• Ka for a weak acid
• HA H A-
• Ka H A-
• HA
• Henderson-Hasselbalch Equation
• pH pKa log anion
• acid

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Neutralization

• The reaction of an acid with a base to produce a
  salt and water.
• HCl NaOH NaCl H2O
• We do this when we use antacids.
• Neutralization can be used to determine the
  amount of acid or base in a sample. -
  titrations

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Titrations

• Analytical methods based on measurement of
  volume.
• If the concentration of an acid is known, the
  concentration of the base can be found.
• If we know the concentration of the base, then we
  can determine the amount of acid.
• All that is needed is some calibrated glassware
  and either an indicator or pH meter.

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TItrations
Buret - volumetric glassware used for
titrations. It allows you to add a known
amount of your titrant to the solution you are
testing. If a pH meter is used, the
equivalence point can be measured. An
indicator will give you the endpoint.
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Indicator Examples

• Acid-base indicators are weak acids that undergo
  a color change at a known pH.

phenolphthalein
methyl red
bromothymol blue