Acids, Bases, and Salts

1
Acids, Bases, and Salts

• Properties of Acids, Bases and Salts
• Calculating pH, pOH, H3O, OH-
• Acid-Base Theories, Strengths of Acids and Bases
• Calculating Ka, Kb, and percent dissociated
• Salt Hydrolysis
• Buffers
• Titrations

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Properties of Acids

• Aqueous solutions of acids have a sour taste. For
  example, the sour taste of lemons and other
  citrus fruits is due to citric acid.
• Acids are electrolytes they conduct electricity
  when dissolved in water.
• Acids change the color of acid-base indicators.
  Acids turn blue litmus red.
• Some acids react with active metals to produce
  hydrogen gas.
• Acids can be solids, liquids or gases in their
  pure state.
• Nonmetals and nonmetal oxides tend to form acids
  in water.

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Naming Binary Acids
Binary acids are named by using the prefix hydro
followed by the root and the ic suffix.
Name the following binary acids. 1. HCl 2. HF 3. H
2S 4. HBr

hydrochloric acid
hydrofluoric acid
hydrosulfuric acid
hydrobromic acid
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Writing Formulas for Binary Acids
Remember to balance the charges when writing the
formulas for binary acids.
Write formulas for the following binary
acids. 1. hydroiodic acid 2. hydroselenic acid

H, I-
HI
H, Se2-
H2Se
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Naming and Writing Formulas for Oxyacids
Oxyacids are named based upon the polyatomic ion
making up the acid.
Polyatomic Ion Ending Acid Ending Example Name
per ate per - ic HClO4 perchloric acid
-ate -ic HClO3 chloric acid
-ite -ous HClO2 chlorous acid
hypo ite hypo - ous HClO hypochlorous acid

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Naming Writing Formulas for Oxyacids
Remember to balance the charges when writing the
formulas for oxyacids.

• Write the name or formula for each of the
  following acids.
• 1. HIO3
• HNO2
• sulfuric acid
• 4. phosphorous acid

iodic acid
IO3- iodate
nitrous acid
NO2- nitrite
H, SO42-
H2SO4
H, PO33-
H3PO3
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Properties of Bases

• Aqueous solutions of bases taste bitter. Have
  you ever had to wash out your mouth with soap?
• Dilute aqueous solutions of bases feel slippery.
• Bases change the color of acid-base indicators.
  Bases turn red litmus blue.
• Bases are electrolytes they conduct electricity
  when dissolved in water.
• Most common bases are solids. An exception is
  ammonia (NH3) which is a gas at room temperature.
• Active metals and metal oxides tend to form bases
  in water.

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Naming and Writing Formulas for Bases

• Bases are named by writing the name of the metal
  ion followed by the word hydroxide.
• The charges are balanced when writing the
  formulas for bases.
• Name or write the formula for the following bases
• 1. LiOH
• 2. Magnesium hydroxide

Lithium hydroxide
Mg2, OH-
Mg(OH)2
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Properties of Salts

• A salt is an ionic compound composed of a cation
  (positive ion) and an anion (negative ion).
• Salts are crystalline solids that have high
  melting points.
• Salts are generally soluble in water.
• Salts are electrolytes.

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Naming and Writing Formulas for Salts

• Salts are named by writing the name of the cation
  followed by the name of the anion.
• The charges are balanced when writing the
  formulas for salts.
• Name or write the formula for the following
  salts.
• 1. CaSO4
• 2. Magnesium nitrate
• 3. LiBr
• 4. Copper(II) chloride

Calcium sulfate
Mg2, NO3-
Mg(NO3)2
Lithium bromide
CuCl2
Cu2, Cl-
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Neutralization Reactions

• The reaction of an acid with a base is called a
  neutralization reaction. The products of a
  neutralization reaction are a salt and water.
  Neutralization reactions are double replacement
  reactions.
• Examples
• HCl NaOH ?
• HNO3 KOH ?
• H2SO4 LiOH ?

H2O
NaCl
H2O
KNO3
H2O
2
2
Li2SO4
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Self-Ionization of Water

• The origination of the idea of pH is the
  self-ionization of water. In the self-ionization
  of water, two water molecules collide producing a
  hydronium ion and a hydroxide ion by transfer of
  a proton.
• H2O(l) H2O(l) ? H3O(aq)
  OH-(aq)
• also written as H2O(l) ? H(aq)
  OH-(aq)

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Ion Product Constant for Water (Kw)

• For aqueous solutions, the product of the
  hydronium ion concentration and the hydroxide ion
  concentration equals 1.010-14 M2 (at 25C).
• H3OOH- 1.010-14 M2
• also written as H OH- 1.010-14 M2
• The product of the concentrations of the
  hydronium ions and the hydroxide ions in water is
  called the Ion Product Constant for water (Kw).
• Kw H3OOH- 1 10-14 M2
• also written as Kw HOH- 1
  10-14 M2

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H3O versus OH-

• All aqueous solutions have H3O and OH- present.
  
• A solution can be classified as acidic, basic or
  neutral by comparing the number of hydronium ions
  in solution to the number of hydroxide ions.
• Neutral H3O OH- 110-7 M
• Acidic H3O gt OH- H3O gt 110-7 M
• Basic (alkaline) H3O lt OH- H3O lt
  110-7 M

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Example Problems.

• Calculate the OH- if the H3O is 1.0 x 10-5
  M. Is the solution acidic, basic or neutral?

H3O gt OH- acidic  
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Example Problems.

• Calculate the OH- if the H3O is 1.8 10-8
  M. Is the solution acidic, basic or neutral?

H3O lt OH- basic
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pH

• An easier way to determine whether a solution is
  acidic, basic or neutral is by determining the pH
  of the solution.
• The pH scale was developed by Sorenson in 1909.
  The pH scale ranges from 0 to 14 at 25C.
• acidic solution
• neutral solution
• basic solution
• pH -log H3O H3O 10x (-pH)
• pOH -log OH- OH- 10x (-pOH)
• pH pOH 14

pH lt 7
pH 7
pH gt 7
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Example Problems

• Calculate the pH of a solution that has a H3O
  of 3.4 10-5 M.
• pH -log H3O
• pH 4.47
• Calculate the pOH of a solution that has a OH-
  of 2.5 10-2 M.
• pOH -log OH-
• pOH 1.60

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Example Problems

• Calculate the H3O of a solution that has a pH
  of 3.5.
• H3O 10x (-pH)
• H3O 3.16 10-4
• Calculate the OH- of a solution that has a pOH
  of 2.3.
• OH- 10x (-pOH)
• OH- 5.01 10-3
• Calculate the pH of a solution that has a pOH of
  12.
• pH pOH 14
• pH 14-12 2

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Complete the following table.
pH -log H3O H3O 10x (-pH) pOH -log
OH- OH- 10x (-pOH) pH pOH 14
pH pOH H3O OH- Acidic, Basic or Neutral
4.00 1.0 10-10 Acidic
8.00 1.0 10-6 Acidic
7.00 1.0 10-7 Neutral
5.00 1.0 10-5 1.0 10-9
3.50 3.2 10-4 Acidic
13.90 0.79 1.26 10-14
4.5 10-11 2.51 10-4 Basic
9.15 1.4 10-5 Basic
10.00
1.0 10-4
6.00
1.0 10-8
7.00
1.0 10-7
9.00
Acidic
10.50
3.2 10-11
0.10
Acidic
3.60
10.40
4.85
7.08 10-10
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Acid-Base Theories- Arrhenius Acids and Bases

• Arrhenius theorized that acids and bases must
  produce ions in solution.
• acid produces hydrogen ions in water solutions.
• Examples HCl, H2SO4
• base produces hydroxide ions in water
  solutions.
• Examples NaOH, Ba(OH)2

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Acid-Base Theories-Brønsted-Lowry Acids and Bases

• Brønsted-Lowry defined an acid as a proton
  donor
• Brønsted-Lowry defined a base as a proton
  acceptor

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Acid Base ? conjugate base conjugate acid

• A conjugate base is the remaining part of an acid
  after it has released a proton.
• A conjugate acid is the acid formed when a base
  accepts a proton.
• NH3 H2O ? NH4 OH-

base
acid
CB
CA
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Label the acid, base, CA, and CB

• HNO2 H2O ? H3O NO2-
• H2O C2H3O2- ? HC2H3O2 OH-

acid
CA
CB
base
acid
base
CA
CB
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Strengths of Acids and Bases

• When you refer to the strength of an acid or a
  base, you are talking about the degree to which
  it is ionized in aqueous solutions.
• Strong acids are completely ionized into aqueous
  solution. This makes them strong electrolytes.
  Examples HCl HBr, HI, HClO4, HClO3, HNO3, and
  H2SO4
• Example HCl H2O ? H3O
  Cl-
• The concentration of H3O present after the HCl
  ionizes is equal to the original concentration of
  HCl.

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Weak Acids

• Weak acids ionize only slightly in aqueous
  solution. Weak acids are weak electrolytes.
• Ex. HCN, HF, HC2H3O2
• HC2H3O2 H2O ? H3O C2H3O2-
• The initial concentration of HC2H3O2 is much
  greater than the concentration of H3O at
  equilibrium.
• The acid dissociation (ionization) constant, Ka,
  can be written for a weak acid. The acid
  dissociation constant is a ratio of the
  dissociated form of an acid to the undissociated
  form.

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Weak Acids

• Write the Ka expression for acetic acid.
• HC2H3O2 H2O ? H3O C2H3O2-

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Comparing Ka values

• The smaller the value of Ka , the weaker the
  acid.
• Which of the following acids is the weakest?
• carbonic acid Ka 4.2 10-7
• formic acid Ka 1.8 10-4
• benzoic acid Ka 6.3 10-5
• Carbonic Acid

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Bases

• Strong bases ionize completely in aqueous
  solution. Strong bases are strong electrolytes.
• Examples Ca(OH)2, NaOH, KOH, LiOH, Ba(OH)2,
  Sr(OH)2
• Weak bases partially ionize in aqueous solution.
  They are weak electrolytes.
• Examples NH3, Al(OH)3

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Kb values

• The base dissociation constant, Kb, can be
  written for a weak base.
• The base dissociation, Kb, is the ratio of the
  dissociated form of a base to the undissociated
  form.

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Kb Expressions

• Write an ionization equation and a Kb expression
  for hydrazine, N2H4.
• N2H4 H2O ? N2H5
  OH-

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Calculating the pH of Solutions of Strong Acids
and Strong Bases

1. Calculate the pH of a 1.00 M HI solution.
2. Calculate the pH of a 1.0 M KOH solution.

Since it is a strong acid it completely
dissociates in water. HI H3O pH -log
H3O pH -log1.00 0
Since it is a strong base it completely
dissociates in water. KOH OH- pOH -log
OH- pOH -log1.00 0 pH 14 0 14
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Calculations Involving Weak Acids and Bases

• A 0.100 M solution of acetic acid (HC2H3O2) is
  only partially ionized. The Ka of acetic acid is
  1.810-5.
• Write a dissociation reaction for acetic acid.
• HC2H3O2 H2O ? C2H3O2-
  H3O

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Calculations Involving Weak Acids and Bases

• A 0.100 M solution of acetic acid (HC2H3O2) is
  only partially ionized. The Ka of acetic acid is
  1.810-5.
• b. Write a Ka expression for acetic acid.

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Calculations Involving Weak Acids and Bases

• c. Calculate the pH of a 0.15 M solution of
  HC2H3O2.

HC2H3O2 H3O C2H3O2-
Initial
Change
Equilibrium
0.15 M
0.0 M
0.0 M
x
x
-x
x
x
0.15 M-x 0.15 M
The Ka value is so small that it is assumed that
the amount the HC2H3O2 changes is negligible.
x22.7310-6, x 0.00165 M
pH -logH3O -log(0.00165) 2.78
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Calculations Involving Weak Acids and Bases

• Another way of expression how much of a weak acid
  (or base) is in ionic form is to give the percent
  dissociated (also called percent ionized).
• HA H2O ? H3O A-

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Calculations Involving Weak Acids and Bases

• e. Find the percent of dissociation of the 0.15 M
  HC2H3O2 solution.
• HC2H3O2 H2O ? H3O C2H3O2-

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• 2. Nitrous acid (HNO2) is a weak acid with a Ka
  of 4.610-4 at 25C.
• Write a dissociation reaction for nitrous acid.
• HNO2 H2O ? NO2- H3O

b. Write a Ka expression for nitrous acid.
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• c. Calculate the pH of a 0.0450 M nitrous acid
  solution.

HNO2 H3O NO2-
Initial
Change
Equilibrium
0.0450 M
0.0 M
0.0 M
-x
x
x
x
x
0.0450 M-x 0.0450 M
x2 2.0710-5 x 0.00455
pH -log(0.00455) 2.34
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• Lactic acid (HC3H5O3) is a waste product that
  accumulates in muscle tissue during exertion,
  leading to pain and a feeling of fatigue.
• a. Write a dissociation reaction for lactic
  acid.
• HC3H5O3 H2O ? C3H5O3- H3O
• b. Write a Ka expression for lactic acid.
• c. In a 0.100 M aquous solution, lactic acid is
  3.7 dissociated. Calculate H3O.

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• d. Calculate the value of Ka for lactic acid.

HC3H5O3 H3O C3H5O3-
Initial
Change
Equilibrium
0.100 M
0.0 M
0.0 M
-3.710-3
3.710-3
3.710-3
3.710-3
0.0963 M
3.710-3
In this case, the H3O (we calculated in part
C) represents the amount the initial
concentrations change.
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Salt Hydrolysis

• A salt is made by neutralizing an acid with a
  base.
• When a salt dissolves in water, it releases ions
  having an equal number of positive and negative
  charges.
• Thus a solution of a salt should be neither
  acidic nor basic.
• Some salts do form neutral solutions, but other
  react with water (hydrolyze) to form acidic or
  basic solutions.

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Types of Salt Solutions

• 1. A neutral solution results when the salt
  produced from a strong acid and strong base is
  dissolved in water. 
• ex. HCl NaOH ? NaCl H2O
• An acidic solution results when the salt produced
  from a strong acid and a weak base is dissolved
  in water.
• ex. 3HCl Al(OH)3 ? AlCl3 3H2O
• A basic solution results when the salt produced
  from a weak acid and a strong base is dissolved
  in water.
• ex. H2CO3 2NaOH ? Na2CO3
  2H2O
• 4. The salt produced from a weak acid and a weak
  base may form an acidic, basic, or neutral
  solution. (You would have to compare the Ka and
  Kb values to determine whether or not the salt
  formed from a weak acid and base was acidic,
  basic, or neutral. We will not be doing that in
  this class.)

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• Identify the acid and base from which each of the
  following salts was formed and then classify the
  solution as acidic, basic, or neutral.
• a. 0.5 M NH4ClO4
• NH3 - weak base
• HClO4 strong acid 
• acidic
• b. 1.0 M BaSO4
•   Ba(OH)2 - strong base
• H2SO4 strong acid 
• neutral
• c. 0.4 M K2CO3
• KOH - strong base
• H2CO3 weak acid 
• basic

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Buffers

• A buffer system is a solution that can absorb
  moderate amounts of acid or base without a
  significant change in its pH.
• A buffer provides ions that will react with H3O
  or OH- ions if they are introduced into the
  solution. Because the added H3O or OH- ions are
  thereby neutralized, the pH of the system remains
  nearly constant.
• Buffer solutions are prepared by using a weak
  acid with one of its salts or a weak base with
  one of its salts.

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Example of a Buffer System

• Many of the fluids in your body must be
  maintained within a very narrow pH range if you
  are to remain healthy. Lets look at the buffer
  system that is present in your blood.
• This buffer contains
• HCO3- (from the salt) and H2CO3 (carbonic acid)
• When excess hydronium ions enter the blood, the
  hydrogen carbonate ion undergoes the following
  reaction to reduce the H3O.
• HCO3-(aq) H3O(aq) ? H2CO3(aq) H2O(l)
• When excess hydroxide ions form in the blood the
  following reaction occurs.
• H2CO3(aq) OH-(l) ? HCO3-(aq) H2O(l)

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Titrations

• Sometimes we want to know the concentration of an
  unknown solution of an acid or base. The
  concentration of an acid (or base) in a solution
  is determined by carrying out a neutralization
  reaction.

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Titration Steps

1. A measured amount of an acid of unknown
   concentration is added to an Erlenmeyer flask.
2. An appropriate indicator (such as
   phenolphthalein) is added to the solution.
3. Measured amounts of a base of known concentration
   are mixed into the acid. The solution of known
   concentration is called the standard solution.
   The addition of the base is carried out using a
   buret. This process is continued until the
   indicator indicates that the end point has been
   reached.
4. The point at which the two solutions used in a
   titration are present in chemically equivalent
   amounts is the equivalence point.

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Titration Graphs

• A graph can be made of pH versus volume of
  standard solution.
• The equivalence point of the titration
  corresponds to the middle of that portion of the
  graph showing a very large change in pH with the
  addition of a small amount of the standard
  solutions.

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Titration Graphs
51
Titration Graphs

• The equivalence point is not always at a pH of 7.
  Why not?
• It will not be 7 if a weak acid and a strong
  base are combined or if a strong acid and a weak
  base are combined.

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Titration Calculations

• When the acid and base used in a titration react
  in a 11 mole ratio, the following relationship
  can be used to determine the concentration of the
  unknown solution or the volume of known needed to
  neutralize the unknown solution.
• MaVa MbVb

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Titration Calculations

• When the acid and base do not react in a 11 mole
  ratio, a mole factor must be used.
• 1. H2SO4 2NaOH ? Na2SO4 2H2O
• 2(MaVa) MbVb
• 2. 2HCl Ba(OH)2 ? BaCl2 2H2O
• MaVa 2(MbVb)

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Example problems

• How many mL of 0.50 M HCl must be added to 25.0
  mL of 2.0 M KOH to make a neutral solution?
• HCl KOH ? KCl H2O
• Since it is a 11 ratio you can use the
  relationship MaVaMbVb

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Example problems

• What is the molarity of a solution of HNO3 if
  30.0 mL of 1.5 M Ba(OH)2 are required to
  neutralize 10 mL of the acid?
• 2HNO3 Ba(OH)2 ? Ba(NO3)2 2H2O
• Since it is not a 11 ratio you must use a mole
  factor. MaVa2MbVb