Reactions in Aqueous Solutions II: Calculations

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CHAPTER 11

• Reactions in Aqueous Solutions II Calculations

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Chapter Goals

• Aqueous Acid-Base Reactions
• Calculations Involving Molarity
• Titrations
• The Mole Method and Molarity
• Equivalent Weights and Normality
• Oxidation-Reduction Reactions
• The Half-Reaction Method
• Adding in H, OH- , or H2O to Balance Oxygen or
  Hydrogen
• Stoichiometry of Redox Reactions

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Calculations Involving Molarity

• Example 11-1 If 100.0 mL of 1.00 M NaOH and
  100.0 mL of 0.500 M H2SO4 solutions are mixed,
  what will the concentration of the resulting
  solution be?
• What is the balanced reaction?
• It is very important that we always use a
  balanced chemical reaction when doing
  stoichiometric calculations.

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Calculations Involving Molarity
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Calculations Involving Molarity

• What is the total volume of solution?
• 100.0 mL 100.0 mL 200.0 mL
• What is the sodium sulfate amount, in mmol?
• 50.0 mmol
• What is the molarity of the solution?
• M 50 mmol/200 mL 0.250 M Na2SO4

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Calculations Involving Molarity

• Example 11-2 If 130.0 mL of 1.00 M KOH and 100.0
  mL of 0.500 M H2SO4 solutions are mixed, what
  will be the concentration of KOH and K2SO4 in the
  resulting solution?
• What is the balanced reaction?

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Calculations Involving Molarity
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Calculations Involving Molarity

• What is the total volume of solution?
• 130.0 mL 100.0 mL 230.0 mL
• What are the potassium hydroxide and potassium
  sulfate amounts?
• 30.0 mmol 50.0 mmol
• What is the molarity of the solution?
• M 30.0 mmol/230.0 mL 0.130 M KOH
• M 50.0 mmol/230.0 mL 0.217 M K2SO4

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Calculations Involving Molarity

• Example 11-3 What volume of 0.750 M NaOH
  solution would be required to completely
  neutralize 100 mL of 0.250 M H3PO4?
• You do it!

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Calculations Involving Molarity
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Titrations

• Acid-base Titration Terminology
• Titration A method of determining the
  concentration of one solution by reacting it with
  a solution of known concentration.
• Primary standard A chemical compound which can
  be used to accurately determine the concentration
  of another solution. Examples include KHP and
  sodium carbonate.
• Standard solution A solution whose
  concentration has been determined using a primary
  standard.
• Standardization The process in which the
  concentration of a solution is determined by
  accurately measuring the volume of the solution
  required to react with a known amount of a
  primary standard.

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Titrations

• Acid-base Titration Terminology
• Indicator A substance that exists in different
  forms with different colors depending on the
  concentration of the H in solution. Examples
  are phenolphthalein and bromothymol blue.
• Equivalence point The point at which
  stoichiometrically equivalent amounts of the acid
  and base have reacted.
• End point The point at which the indicator
  changes color and the titration is stopped.

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Titrations

• Acid-base Titration Terminology

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The Mole Method and Molarity

• Potassium hydrogen phthalate is a very good
  primary standard.
• It is often given the acronym, KHP.
• KHP has a molar mass of 204.2 g/mol.
• A very common mistake is for students to see the
  acronym KHP and think that this compound is made
  of potassium, hydrogen, and phosphorous.

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The Mole Method and Molarity

• Example 11-4 Calculate the molarity of a NaOH
  solution if 27.3 mL of it reacts with 0.4084 g of
  KHP.

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The Mole Method and Molarity

• Example 11-5 Calculate the molarity of a
  sulfuric acid solution if 23.2 mL of it reacts
  with 0.212 g of Na2CO3.

You do it!
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The Mole Method and Molarity

• Example 11-6 An impure sample of potassium
  hydrogen phthalate, KHP, had a mass of 0.884 g.
  It was dissolved in water and titrated with 31.5
  mL of 0.100 M NaOH solution. Calculate the
  percent purity of the KHP sample. Molar mass of
  KHP is 204.2 g/mol.
• NaOH KHP ? NaKP H2O
• You do it!

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The Mole Method and Molarity
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Equivalent Weights and Normality

• Normality is another method of expressing
  concentration.
• Normality is defined as the number of equivalent
  weights of solute per liter of solution.

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Equivalent Weights and Normality

• The equivalent weight of an acid is the mass in
  grams of the acid necessary to furnish Avogadros
  number of H ions.
• For monoprotic acids like HCl 1 mol 1 eq
• For diprotic acids like H2SO4 1 mol 2 eq
• For triprotic acids like H3PO4 1 mol ? eq
• You do it!
• For triprotic acids like H3PO4 1 mol 3 eq

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Equivalent Weights and Normality

• Similarly for bases
• 1 mol 1 eq for NaOH
• 1 mol 2 eq for Ba(OH)2
• 1 mol 3 eq for Fe(OH)3
• The equivalent weights of some acids and bases
  are given in Table 11-1 in your text.
• Any problem that can be done in normality can
  also be done in molarity.
• The difference between the two is that in
  molarity you must remember to use the reaction
  stoichiometry but in normality the stoichiometry
  is included in the solution concentrations.

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Equivalent Weights and Normality

• Example 11-7 Calculate the normality of a
  solution that contains 196 g of sulfuric acid in
  1.500 x 103 mL of solution.

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Equivalent Weights and Normality

• Example 11-8 Calculate the molarity and
  normality of a solution that contains 34.2 g of
  barium hydroxide in 8.00 liters of solution.
• You do it!

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Equivalent Weights and Normality

• Since M x L moles then
• N x L number of equivalents or
• N x mL number of milliequivalents
• Example 11-9 What volume of 6.00 M phosphoric
  acid solution is required to prepare 9.00 x 102
  mL of 0.200 N phosphoric acid solution?

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Equivalent Weights and Normality
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Equivalent Weights and Normality

• In reactions, including acid-base reactions, 1
  equivalent of an acid will react with 1
  equivalent of a base.
• Example 11-10 What is the normality of a
  sulfuric acid solution if 31.3 mL of it reacts
  with 0.318 g of sodium carbonate?

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Equivalent Weights and Normality
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Equivalent Weights and Normality

• Example 11-11 30.0 mL of 0.0750 N nitric acid
  solution required 22.5 mL of calcium hydroxide
  solution for neutralization. Calculate the
  normality and the molarity of the calcium
  hydroxide solution.
• You do it!

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Equivalent Weights and Normality
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Oxidation-Reduction Reactions

• We have previously gone over the basic concepts
  of oxidation reduction in Chapter 4.
• Rules for assigning oxidation numbers were also
  introduced in Chapter 4.
• Refresh your memory as necessary.
• We shall learn to balance redox reactions using
  the half-reaction method.

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The Half-Reaction Method

• Half reaction method rules
• Write the unbalanced reaction.
• Break the reaction into 2 half reactions
• One oxidation half-reaction and
• One reduction half-reaction
• Each reaction must have complete formulas for
  molecules and ions.
• Mass balance each half reaction by adding
  appropriate stoichiometric coefficients. To
  balance H and O we can add
• H or H2O in acidic solutions.
• OH- or H2O in basic solutions.

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The Half-Reaction Method

• Charge balance the half reactions by adding
  appropriate numbers of electrons.
• Electrons will be products in the oxidation
  half-reaction.
• Electrons will be reactants in the reduction
  half-reaction.
• Multiply each half reaction by a number to make
  the number of electrons in the oxidation
  half-reaction equal to the number of electrons
  reduction half-reaction.
• Add the two half reactions.
• Eliminate any common terms and reduce
  coefficients to smallest whole numbers.

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The Half-Reaction Method

• Example 11-12 Tin (II) ions are oxidized to tin
  (IV) by bromine. Use the half reaction method to
  write and balance the net ionic equation.

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The Half-Reaction Method
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The Half-Reaction Method
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The Half-Reaction Method

• Example 11-19 Dichromate ions oxidize iron (II)
  ions to iron (III) ions and are reduced to
  chromium (III) ions in acidic solution. Write
  and balance the net ionic equation for the
  reaction.

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The Half-Reaction Method
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The Half-Reaction Method
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The Half-Reaction Method

• Example 11-20 In basic solution hydrogen
  peroxide oxidizes chromite ions, Cr(OH)4-, to
  chromate ions, CrO42-. The hydrogen peroxide is
  reduced to hydroxide ions. Write and balance the
  net ionic equation for this reaction.
• You do it!

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The Half-Reaction Method
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The Half-Reaction Method

• Example 11-21 When chlorine is bubbled into
  basic solution, it forms hypochlorite ions and
  chloride ions. Write and balance the net ionic
  equation.
• You do it!
• This is a disproportionation redox reaction. The
  same species, in this case Cl2, is both reduced
  and oxidized.

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The Half-Reaction Method
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Stoichiometry of Redox Reactions

• Just as we have done stoichiometry with acid-base
  reactions, it can also be done with redox
  reactions.
• Example 11-22 What volume of 0.200 M KMnO4 is
  required to oxidize 35.0 mL of 0.150 M HCl? The
  balanced reaction is

You do it!
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Stoichiometry of Redox Reactions
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Stoichiometry of Redox Reactions

• Example 11-23 A volume of 40.0 mL of iron (II)
  sulfate is oxidized to iron (III) by 20.0 mL of
  0.100 M potassium dichromate solution. What is
  the concentration of the iron (II) sulfate
  solution? From Example 11-19 the balanced
  equation is

You do it!
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Stoichiometry of Redox Reactions
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End of Chapter 11

• Redox reactions are very important commercially.