Chapter 4 Aqueous Reactions and Solution Stoichiometry

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Chapter 4Aqueous Reactions and Solution
Stoichiometry
Chemistry, The Central Science, 10th
edition Theodore L. Brown H. Eugene LeMay, Jr.
and Bruce E. Bursten

• John D. Bookstaver
• St. Charles Community College
• St. Peters, MO
• ? 2006, Prentice Hall, Inc.

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Chapter 4 HW

• Visualizing 1,3,5,7,9,10
• Electrolytes 11,13,15,17
• Pp rx and net ionic eq 19,23,25
• Acid-base rx 29,35,37,39
• Redox 49,51,53
• Molarity 61,65,67,71
• Titrations 77,81,85

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Aqueous Solutions

• Water is the dissolving medium, or solvent.

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Figure 4.1 The Water Molecule is Polar
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Some Properties of Water

• Water is bent or V-shaped.
• The O-H bonds are covalent.
• Water is a polar molecule.
• Hydration occurs when salts dissolve in water.

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Solutions

• Homogeneous mixtures of two or more pure
  substances.
• The solvent is present in greatest abundance.
• All other substances are solutes.

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A Solute

• dissolves in water (or other solvent)
• changes phase (if different from the solvent)
• is present in lesser amount (if the same phase as
  the solvent)

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A Solvent

• retains its phase (if different from the solute)
• is present in greater amount (if the same phase
  as the solute)

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Polar Water Molecules Interact with the Positive
and Negative Ions of a Salt
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BaCI2 Dissolving
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Dissociation

• When an ionic substance dissolves in water, the
  solvent pulls the individual ions from the
  crystal and solvates them.
• This process is called dissociation.

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Electrolytes

• Substances that dissociate into ions when
  dissolved in water.
• A nonelectrolyte may dissolve in water, but it
  does not dissociate into ions when it does so.

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Electrolytes and Nonelectrolytes

• Soluble ionic compounds tend to be electrolytes.

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Electrolytes and Nonelectrolytes

• Molecular compounds tend to be nonelectrolytes,
  except for acids and bases.

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Electrolytes

• A strong electrolyte dissociates completely when
  dissolved in water.
• A weak electrolyte only dissociates partially
  when dissolved in water.

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Strong Electrolytes Are

• Strong acids

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Strong Electrolytes Are

• Strong acids
• Strong bases

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Strong Electrolytes Are

• Strong acids
• Strong bases
• Soluble ionic salts

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Electrolytes

• Strong - conduct current efficiently
• NaCl, HNO3
• Weak - conduct only a small current
• vinegar, tap water
• Non - no current flows
• pure water, sugar solution

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Strong Electrolytes

1. Soluble salts- eg. NaCl, Pb(ClO3)2
2. Strong acids- eg. HCl, H2SO4
3. Strong bases- eg. NaOH, KOH

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Weak or Nonelectrolytes

1. Insoluble or only slightly soluble salts. Eg.
   AgCl, CaSO4
2. Weak Acids- eg. CH3COOH, HF
3. Weak Bases- eg. NH3, amines
4. Water

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September 23-Section 4.2Precipitation Reactions

• SOLUBILITY
• PRECIPITATION REACTIONS
• DOUBLE REPLACEMENT (METATHESIS EXCHANGE )
  REACTION
• NET IONIC EQUATIONS
• SPECTATOR IONS

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Precipitation Reactions

• When one mixes ions that form compounds that are
  insoluble (as could be predicted by the
  solubility guidelines), a precipitate is formed.

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SOLUBILITY

• The amount of a substance that can be dissolved
  in a given quantity of solvent at a given
  temperature.
• Example
• Solubility of KNO3
• 65 g/100 ml H2O of KNO3 at 40 C
• It is determined experimentally.

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Solubility Rules Used to determine what
reaction occurs, if any 1. Separate all
ions. 2. Determine all possible compounds
formed. 3. Determine which, if any, of the
compounds will precipitate. 4. Write the
appropriate chemical equation.
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• Rules for Solubility of Ionic Compounds in Water
• All common Group 1 and ammonium salts are
  soluble.
• All nitrates, chlorates, and acetates are soluble
  except silver acetate.
• All halide salts (except fluorides) are soluble
  except those of silver, mercury (I), and lead.
• All sulfates are soluble except those of silver,
  mercury (I or II), lead, calcium, strontium, and
  barium.
• Calcium, strontium, and barium hydroxides (as
  well as group 1 hydroxides) are soluble other
  hydroxides generally are not.
• Most other ionic compounds are generally
  insoluble.

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• Double Replacement (Metathesis) Reactions
• Metathesis or double displacement/replacement
  reactions involve swapping ions in solution
• AX BY ? AY BX.
• Metathesis reactions will lead to a change in
  solution if one of three things occurs
• an insoluble solid is formed (precipitate),
• weak or nonelectrolytes are formed, or
• an insoluble gas is formed.

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Molecular Equation

• The molecular equation lists the reactants and
  products in their molecular form.

• AgNO3 (aq) KCl (aq) ?? AgCl (s) KNO3 (aq)

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Ionic Equation

• In the ionic equation all strong electrolytes
  (strong acids, strong bases, and soluble ionic
  salts) are dissociated into their ions.
• This more accurately reflects the species that
  are found in the reaction mixture.

• Ag (aq) NO3- (aq) K (aq) Cl- (aq) ??
• AgCl (s) K (aq) NO3- (aq)

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Net Ionic Equation

• To form the net ionic equation, cross out
  anything that does not change from the left side
  of the equation to the right.

• Ag(aq) NO3-(aq) K(aq) Cl-(aq) ??
• AgCl (s) K(aq) NO3-(aq)

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Net Ionic Equation

• To form the net ionic equation, cross out
  anything that does not change from the left side
  of the equation to the right.
• The only things left in the equation are those
  things that change (i.e., react) during the
  course of the reaction.

• Ag(aq) Cl-(aq) ?? AgCl (s)

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Net Ionic Equation

• To form the net ionic equation, cross out
  anything that does not change from the left side
  of the equation to the right.
• The only things left in the equation are those
  things that change (i.e., react) during the
  course of the reaction.
• Those things that didnt change (and were deleted
  from the net ionic equation) are called spectator
  ions.

• Ag(aq) NO3-(aq) K(aq) Cl-(aq) ??
• AgCl (s) K(aq) NO3-(aq)

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Writing Net Ionic Equations

1. Write a balanced molecular equation.
2. Dissociate all strong electrolytes.
3. Cross out anything that remains unchanged from
   the left side to the right side of the equation.
4. Write the net ionic equation with the species
   that remain.

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• Examples Write the molecular, complete ionic
  and net ionic equations for each of the following
  reactions.
• Aqueous solutions of sodium sulfide and calcium
  nitrate are mixed.

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• 2. Aqueous solutions of barium chloride and
  potassium sulfate are mixed.

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• 3. Aqueous solutions of silver nitrate and
  ammonium chloride are mixed.

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Homework

• Page 157 answer 4.7 to 4.9
• Page 158 4.19 to 4.28 odd
• Read Section 4.3

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September 24- Section 4.3 Acid Bases Reactions

• Strong and weak Acid and Bases
• Neutralization Reactions

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Classify as strong, weak or nonelectrolyte

• CaCl2
• HNO3
• C2H5OH
• HC2H3O2

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Rank the following solutions in order of
increasing electrical conductivity

• Ca(NO3)2
• C6 H12 O6
• NaC2 H3O2
• HC2 H3O2

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Acids

• Substances that increase the concentration of H
  when dissolved in water (Arrhenius).
• Proton donors (BrønstedLowry).

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Acids

• There are only seven strong acids
• Hydrochloric (HCl)
• Hydrobromic (HBr)
• Hydroiodic (HI)
• Nitric (HNO3)
• Sulfuric (H2SO4)
• Chloric (HClO3)
• Perchloric (HClO4)

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• Acids
• Dissociation pre-formed ions in solid move
  apart in solution.
• Ionization neutral substance forms ions in
  solution.
• Acids substances that ionize to form H in
  solution (e.g. HCl, HNO3, CH3CO2H, lemon, lime,
  vitamin C). Proton donors.
• Acids with one acidic proton are called
  monoprotic (e.g., HCl).
• Acids with two acidic protons are called diprotic
  (e.g., H2SO4).
• Acids with many acidic protons are called
  polyprotic.

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Bases

• Substances that increase the concentration of OH-
  when dissolved in water (Arrhenius).
• Proton acceptors (BrønstedLowry).

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• Bases
• Bases substances that react with the H ions
  formed by acids (e.g. NH3, Drano, Milk of
  Magnesia). Proton acceptors (Bronsted-Lowry).

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• Strong and Weak Acids and Bases
• Strong acids and bases are strong electrolytes.
• They are completely ionized in solution.
• Strong acids are HCl, HBr, HI, HNO3, H2SO4, HClO3
  and HClO4.
• Strong bases are group 1 hydroxides and soluble
  group 2 hydroxides.
• Weak acids and bases are weak electrolytes.
• They are partially ionized in solution.

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Acid-Base Reactions

• In an acid-base reaction, the acid donates a
  proton (H) to the base.

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• Neutralization Reactions and Salts
• Neutralization occurs when a solution of an acid
  and a base are mixed
• HCl(aq) NaOH(aq) ? H2O(l) NaCl(aq)
• We form a salt (NaCl) and water.
• Salt ionic compound whose cation comes from a
  base and anion from an acid.
• Neutralization between acid and metal hydroxide
  produces water and a salt.
• In net ionic equations, strong acids and bases
  are written dissociated, while weak acids and
  bases are written associated (non-dissociated)

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Neutralization Reactions

• Observe the reaction between Milk of Magnesia,
  Mg(OH)2, and HCl.

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September 27 Section 4.4Oxidation-Reduction
Reactions

• Acid-Base reactions with gas formation
• Redox concept LEO GER
• Oxidation number
• Oxidizing-Reducing Agents
• Single replacement reactions (oxidation of metals
  by acids and salts)- Displacement reactions
• Activity Series
• Balancing redox in acid and alkaline media.

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Metathasis Reactions that produce gas
GAS REACTANTS
H2S (g) Any sulfide plus any acid
CO2 (g) Any carbonate plus acid
SO2 (g) Any sulfite plus acid
NH3 (g) Any ammonium salt plus strong hydroxide and heat
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• Acid-Base Reactions with Gas Formation
• Sulfide ions can react with H producing H2S
  gas.
• 2HCl(aq) Na2S(aq) ? H2S(g) 2NaCl(aq)
• 2H(aq) S2-(aq) ? H2S(g)
• Na2S (aq) H2SO4 (aq) ?? Na2SO4 (aq) H2S (g)

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Carbonate ions produce CO2(g)and H2O

• HCl(aq) NaHCO3(aq) ? NaCl(aq) H2O(l) CO2(g)
• CaCO3 (s) HCl (aq) ??CaCl2 (aq) CO2 (g)
  H2O (l)
• NaHCO3 (aq) HBr (aq) ??NaBr (aq) CO2 (g)
  H2O(l)

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Sulfite ions produce SO2 and H2O

• SrSO3 (s) 2 HI (aq) ??SrI2 (aq) SO2 (g) H2O
  (l)

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Ammonium salts and soluble bases produce NH3 when
solution is warmed

• NH4Cl(aq) NaOH (aq) --gt NH3 (g) H2O (l)
  NaCl (aq)
• Theorically NH4OH is produced but is unstable and
  decomposes into ammonia and water.
• NH4OH (aq) ? NH3 (g) H2O (l)

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LEO GER

• LOSING
• ELECTRONS
• OXIDATION
• O.N. increases in oxidation

• GAINING
• ELECTRONS
• REDUCTION
• O.N. decreases in reduction

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Oxidation-Reduction Reactions

• An oxidation occurs when an atom or ion loses
  electrons.
• A reduction occurs when an atom or ion gains
  electrons.

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Oxidation-Reduction Reactions

• One cannot occur without the other.

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Oxidizing and Reducing Agents

• Oxidizing agents cause oxidation to occur. How?
• Reducing agents cause reduction to occur. How?

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A Summary of an Oxidation-Reduction Process
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September 28

• How to determine the Oxidation Number
• of elements in compounds.
• How to predict if a reaction will ocurr using the
  reactivity series
• Single replacement reactions
• Balancing redox reactions

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Oxidation Numbers

• To determine if an oxidation-reduction reaction
  has occurred, we assign an oxidation number to
  each element in a neutral compound or charged
  entity.

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• Oxidation Numbers
• Oxidation numbers are assigned by a series of
  rules
• If the atom is in its elemental form, the
  oxidation number is zero. e.g., Cl2, H2, P4.
• For a monoatomic ion, the charge on the ion is
  the oxidation state.
• Elements of Group I have always O.N. 1
• Elements of Group II have alwas O.N. 2

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• 5. Oxidation number of O is usually 2. The
  peroxide ion, O22-, has oxygen with an oxidation
  number of 1 (H2O2, Na2O2).
• Oxygen with F- has O.N. 2
• 6. Oxidation number of H is 1 when bonded to
  nonmetals and 1 when bonded to metals in metal
  Hydrides
• 7. The oxidation number of F is 1

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Oxidation Numbers

• The sum of the oxidation numbers in a neutral
  compound is 0.
• The sum of the oxidation numbers in a polyatomic
  ion is the charge on the ion.

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• Examples Determine the oxidation numbers of
  each element in each of the following
• H2SO4
• KMnO4
• NO3-
• C2H6
• CH3OH

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Displacement ReactionsSingle Replacement(always
redox reactions!)

• In displacement reactions, ions oxidize an
  element.
• The ions, then, are reduced.

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Displacement Reactions

• In this reaction,
• silver ions oxidize
• copper metal.
• Cu (s) 2 Ag (aq) ?? Cu2 (aq) 2 Ag (s)

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Displacement Reactions

• The reverse reaction,
• however, does not
• occur.
• Cu2 (aq) 2 Ag (s) ?? Cu (s) 2 Ag (aq)

x
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• Oxidation of Metals by Acids and Salts
• Metals are oxidized by acids to form salts
• Mg(s) 2HCl(aq) ? MgCl2(aq) H2(g)
• During the reaction, 2H(aq) is reduced to H2(g).
• Metals can also be oxidized by other salts
• Fe(s) Ni2(aq) ? Fe2(aq) Ni(s)
• Notice that the Fe is oxidized to Fe2 and the
  Ni2 is reduced to Ni.

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• Activity Series
• Some metals are easily oxidized whereas others
  are not.
• Activity series a list of metals arranged in
  decreasing ease of oxidation.
• The higher the metal on the activity series, the
  more active that metal.
• Any metal can be oxidized by the ions of elements
  below it.

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• Examples Write complete ionic and net ionic
  equations for the following
• Aluminum metal is added to an aqueous solution of
  copper chloride.
• Zinc metal is added to a solution of hydrobromic
  acid.
• Chromium metal is placed in a solution of
  potassium nitrate.

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Balancing by Half-Reaction Method

• 1. Write separate reduction, oxidation
  reactions.
• 2. For each half-reaction
• ? Balance elements (except H, O)
• ? Balance O using H2O
• ? Balance H using H
• ? Balance charge using electrons

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Balancing by Half-Reaction Method (continued)

• 3. If necessary, multiply by integer to
  equalize electron count.
• 4. Add half-reactions.
• 5. Check that elements and charges are balanced.

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Example

• Balance the following oxidation reduction using
  the half-reaction method
• MnO2 Cl- ? Mn2 Cl2

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Half-Reaction Method - Balancing in Base

• 1. Balance as in acid.
• 2. Add OH? that equals H ions (both sides!)
• 3. Form water by combining H, OH?.
• 4. Check elements and charges for balance.

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Example

• Balance the following reaction which occurs in
  alkaline solution.
• Ag (s) CN-(aq) O2(g) ? Ag(CN)2-(aq)

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September 29 Section 4.5Concentration of
solutions

• Molarity Concept and problems
• Find moles in a volume of solution
• How to prepare a solution of a given molarity
• Conversions volume to moles, to mass etc.
• Expressing concentration of electrolytes
• DILUTIONS

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NEXT WEEK LAB 1

• BOOKS ARE IN NEED TO PURCHASE LAB BOOK BY
  WEDNESDAY
• EQUATIONS BOOK ARE IN
• STILL WAITING FOR TEXTBOOKS!
• HW 61, 65, 67, 71
• TEST ON CH 1TO 4 ON MONDAY

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Molarity

• Two solutions can contain the same compounds but
  be quite different because the proportions of
  those compounds are different.
• Molarity is one way to measure the concentration
  of a solution.

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• Examples
• Calculate the molarity of a solution prepared by
  dissolving 3.89 g of sodium sulfate in enough
  water to prepare 250. mL of solution.
• Step 1 Convert 3.89 g Na2SO4 to mol
• FM 142 g/mol 3.89g/142 g/mol 0.027 mol
• Step 2 Set up ratio

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• How many grams of magnesium chlorate are required
  to prepare 100.0 mL of a 0.500 M solution?
• Step 1 Find the mol in V of solution
• Step 2 Convert mass to mol

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• How many g of KMnO4 are needed to prepare 100 ml
  of a 2M solution?
• Step 1 Find the of mol
• Step 2 Convert mol to mass

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Mixing a Solution
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Dilution
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• Dilution
• We recognize that the number of moles are the
  same in dilute and concentrated solutions.
• So
• MdiluteVdilute moles MconcentratedVconcentrate
  d
• M1V1 M2V2

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• Find the molarity of the solution if 1 ml of the
  solution are diluted to 500 ml
• Step 1 M V M V

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Example What volume of 18.0 M H2SO4 solution is
required to prepare 1000.0 mL of 1.00 M H2SO4?
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October 1 -Section 4.6Solution Stoichiometry and
Chemical analysis

• Titration
• Problems

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• Solution Stoichiometry and Chemical Analysis
• There are two different types of units
• laboratory units (macroscopic units measure in
  lab)
• chemical units (microscopic units relate to
  moles).
• Always convert the laboratory units into chemical
  units first.
• Grams are converted to moles using molar mass.
• Volume or molarity are converted into moles using
  M mol/L.

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• Use the stoichiometric coefficients to move
  between reactants and product.

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• Titrations
• Suppose we know the molarity of a NaOH solution
  and we want to find the molarity of an HCl
  solution.
• We know
• molarity of NaOH, volume of HCl.
• What do we want?
• Molarity of HCl.
• What do we do?
• Take a known volume of the HCl solution, measure
  the mL of NaOH required to react completely with
  the HCl.

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• Titrations
• What do we get?
• Volume of NaOH. We know molarity of the NaOH,
  we can calculate moles of NaOH.
• Next step?
• We also know HCl NaOH ? NaCl H2O. Therefore,
  we know moles of HCl.
• Can we finish?
• Knowing mol(HCl) and volume of HCl (20.0 mL
  above), we can calculate the molarity.

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• Examples
• Standardization of NaOH A primary standard
  called potassium hydrogen phthalate is used to
  determine the exact molarity of a solution of
  base. KHP is a monoprotic weak acid (MW
  204.22). In one example, 0.4977 g of KHP is
  dissolved in 100.0 mL of water. This solution
  requires 14.86 mL of a solution of NaOH to
  neutralize it. What is the molarity of the NaOH
  solution?
• Solid Acid An unknown solid acid is analyzed by
  titration. 1.75 g of the acid is weighed and
  dissolved in 150.0 mL of distilled water. This
  solution is titrated with 0.250 M KOH, and
  requires 19.07 mL of the KOH solution to reach
  the endpoint. What is the molar mass of the
  solid, assuming it is monoprotic?

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3. Unknown Solution A solution of the weak
base, Na2SO4, is analyzed by titration with
0.1000 M HCl. A 10.00 mL sample of the base
requires 38.04 mL of HCl to reach the endpoint.
What is the molarity of the sodium sulfate
solution?
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Titration

• The analytical technique in which one can
  calculate the concentration of a solute in a
  solution.

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Key Titration Terms

• Titrant - solution of known concentration used
  in titration
• Analyte - substance being analyzed
• Equivalence point - enough titrant added to
  react exactly with the analyte
• Endpoint - the indicator changes color so you
  can tell the equivalence point has been reached.

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Titration