Water: The Universal Solvent

1
Water The Universal Solvent

• One of the most valuable properties of water is
  its ability to dissolve.
• An individual water molecule has a bent shape
  with a H-O-H bond angle of approximately 105
  degrees.
• Water is polar thus having positive negative
  partial charges on its ends.

2
Ionic Compounds in Water

• The positive ends of a water molecule are
  attracted to negative cations and the negative
  ends are attracted to positive cations in an
  ionic compound this is called hydration.
• The ions
• become hydrated
• move around
• independently.

3
Covalent Compounds in Water

• Water also dissolves many nonionic substances
  such as ethanol (C2H5OH).
• The reason for this is that ethanol is also
  polar.
• Polar dissolves polar like dissolves like.
• This is the reason water will not dissolve oil.

4
Solutions as Electrolytes

• Strong conductor strong electrolyte
• Weak conductor weak electrolyte
• No conductor - nonelectrolyte

5
Solution Concentration

• Solution can be expressed in a variety of ways
  but the most common method is molarity
• M moles solute
• volume (L) solution
• Example calculate the molarity of a solution
  prepped by dissolving 1.56 g of gaseous
  hydrochloric acid in enough water to make 26.8 mL
  of solution.
• 1.60 M HCl

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Concentration of Ions in Solution

• When an ionic salt dissolves in water ions are
  form in solution. The moles of ions formed must
  be considered in concentration.
• Example what is the concentration each ion in
  (a) 0.50M cobalt II nitrate (b) 2.0M iron III
  perchlorate.
• Co(NO3)2 (s) H2O ? Co2 (aq) 2NO3- (aq)
• Co2 1 x 0.50M 0.50M
• NO3- 2 x 0.50M 1.0M
• Fe(ClO4)3 (s) H2O ? Fe3 (aq) 3ClO4- (aq)
• Fe3 2.0M
• ClO4- 6.0M

7
Concentration of Ions in Solution

• Example Calculate the number of moles of
  chloride ions in 1.75 L of 1.0 x 10-3M zinc
  chloride.
• (step 1) ZnCl2 (s) ? Zn2(aq) 2Cl-(aq)
• (step 2) Cl- 2 x 1.0 x 10-3 2.0 x 10-3M
• (step 3) 1.75 L x 2.0 x 10-3 mole Cl-
• 1 L
• 3.5 x 10-3 moles of chloride ions

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Example Concentration Volume

• Typical blood serum is about 0.14M sodium
  chloride. What volume of blood contains 1.0 mg of
  sodium chloride?
• 0.12 mL of blood
• To analyze the alcohol content of a certain wine,
  a chemist needs 50.00 mL of an aqueous 0.200M
  potassium dichromate solution. How much solid
  potassium dichromate must be weighed out to make
  this solution?
• 2.94 g K2Cr2O7

9
Dilution Formula

• This formula allows a chemist to prepare a
  diluted solution from a concentrated one.
• M1V1 M2V2 or McVc MdVd
• Example what volume of 16 M sulfuric acid must
  be used to prepare 1.5 L of a 0.10 M sulfuric
  acid solution?
• 9.4 mL of H2SO4 must be diluted with 1.5 L of
  water.

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Types of Chemical Reactions
11
Precipitation Reactions

• When 2 aqueous solutions are mixed an insoluble
  precipitate sometimes forms also known as
  double replacement or metathesis reactions.
• It is important to remember that some ions are
  the key players and some are just spectators
• The formula equation gives the overall reaction.
• The complete ionic equation represents all ions
  involved in the reaction.
• The net ionic equation includes only those
  solution components undergoing a change,
  spectator ions are not included.

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The reaction of Pb(NO3)2 NaI.Write the formula
equation, ionic net ionic equations
13
Stoichiometry of Precipitation Reactions

• Calculate the mass of solid sodium chloride that
  must be added to 1.50 L of a 0.100 M silver
  nitrate solution to precipitate all the silver
  ions in the form of silver chloride.
• First find the of moles Na necessary for the
  of moles of silver ions already present, then
  convert to grams.
• 8.77 g NaCl

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Example Determine Mass of Product Formed

• When aqueous solutions of sodium sulfate and lead
  II nitrate are mixed a precipitate forms. Give
  the net ionic equation for the reactions and
  calculate the mass of this precipitate when 1.25
  L of 0.0500 M lead II nitrate and 2.00 L of
  0.0250 M sodium sulfate.
• 15.2 g PbSO4

15
Acid-Base Reactions

• An acid is a substance that produces H ions when
  dissolved in water.
• A base is a substance that produces OH- ions when
  dissolved in water.

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Acids and bases as Electrolytes

• A, Strong acids and bases are strong
  electrolytes, as indicated by the brightly lit
  bulb. B, Weak acids and bases are weak
  electrolytes.

18
Neutralization Reactions
19
Strong AcidStrong Base

• Because both ionize completely, the H ions and
  OH- ions react with each other to form water
  molecules.
• Basic HNO3 (aq) NaOH (aq) ?
• Ionic H (aq) NO3- (aq) Na (aq)
  OH- (aq) ?
• Net Ionic H (aq) OH- (aq) ?
  H2O (l)

20
Weak Acid Strong Base

• A two-step process occurs
• 1. The ionization of the weak acid
• HB (aq) ?? H (aq) B- (aq)
• 2. The neutralization of the H ion by the OH-
  from the strong base
• H (aq) OH- (aq) ? H2O (l)
• Net Ionic HX OH- ? X-
  H2O

21
Strong Acid Weak Base

• This is also a two-step process
• 1.The first step occurs when the weak base
  reacts with water to produce OH- ions.
• NH3 H2O ?? NH4 OH-
• 2. In the second step the H ions from the
  strong acid neutralize the OH- ions to form
  water.
• H OH- ? H2O
• Net Ionic H X ? XH

22
Example Neutralization Reactions

• 1)What volume of a 0.100 M HCl is needed to
  neutralize 25.0 mL of 0.350 M NaOH?
• 8.75 x 10-2 L
• 2)In a certain experiment, 28.0 mL of 0.250 M
  nitric acid and 53.0 mL of 0.320 M potassium
  hydroxide are mixed. Calculate the moles of water
  formed in the resulting reaction. What is the
  H and OH- after the reaction goes to
  completion?
• 0.024 moles H2O, H 0, OH- 0.123 M

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Acid-Base Titration

• This is a volumetric analysis technique for
  determining the amount (usually concentration) of
  a substance.
• This involves the delivery (from a buret) of a
  measured volume of a solution of known
  concentration (the titrant) into a solution
  containing the substance being analyzed (the
  analyte).
• The neutralization point is known as the
  equivalence point. This point is marked by an
  indicator.
• The point when the indicator changes color is
  known as the end point.
• The goal is to choose an indicator which has a
  similar endpoint as the equivalence point your
  reaction.

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Titration Example 1

• You perform an acid-base titration to
  standardize an HCl solution by placing 50.00 mL
  of HCl in a flask with a few drops of indicator
  solution. You put 0.1524 M NaOH into the buret,
  and the initial reading is 0.55 mL. At the end
  point, the buret reading is 33.87 mL. What is the
  concentration of the HCl solution?

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Titration Example 2

• In a titration, it is found that 25.0 mL of
  0.500 M NaOH is required to react with
• (a) a 15.0-mL sample of HCl. What is the molarity
  of HCl?
• 0.833 M
• (b) a 15.0-mL sample of a weak acid, H2A. What is
  the molarity of H2A, assuming the reaction to be
• H2A(aq) 2OH-(aq) ? 2H2O A2-(aq)?
• 0.417 M
• (c) an aspirin tablet weighing 2.50 g. What is
  the percentage of acetylsalicylic acid, HC9H7O4,
  in the aspirin tablet? The reaction is
• HC9H7O4 (s) OH- (aq) ? H2O C9H7O4 - (aq)
• 90.0

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Oxidation-Reduction (Redox)

• Oxidation means the losing of electrons (an
  increase in the oxidation ) and reduction means
  the gaining of electrons (a decrease in the
  oxidation ). The 2 occur together, they are
  opposite sides of the same coin.
• A good way to remember LEO the lion goes GER
  losing electrons oxidation..gaining electrons
  reduction OR OIL RIG oxidation is losing,
  reduction is gaining.

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Oxidizing Reducing Agents

• An oxidizing agent is the species that accepts
  the electrons i.e. the H ion above.
• Non-metals tend to be oxidizing agents.
• A reducing agent is the species that donates the
  electrons i.e. the Zn above.
• Metals tend to be reducing agents.

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Oxidation Numbers

• The first step to balancing any redox is
  assigning oxidation numbers to reactants and
  products in the equation please reference pg.
  89 in text
• The oxidation of an element in its elemental
  state is 0.
• The oxidation of an element in a monatomic ion
  is equal to the charge of that ion.
• Certain elements have the same oxidation in all
  or almost all their compounds.
• The sum of the oxidation numbers in a neutral
  species is 0 in a polyatomic ion, it is equal to
  the charge of that ion.

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Assigning Oxidation s Practice

• What is the oxidation number of phosphorus
• in sodium phosphate?
• P 5
• In the dihydrogen phosphate ion?
• P 5

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Balancing Redox Reactions

• Assign oxidation numbers (Rules Pg. 89)
• Identify the oxidation and reduction reactions.
• Split in 2 half reactions.
• Balance the element being oxidized and reduced.
• Balance the elements (that is being reduced or
  oxidized) oxidation by adding electrons.
  Oxidation adds to the right, reduction adds to
  the left.
• Balance the oxygens by adding water molecules.
• Balance the hydrogens by adding H ions.
• If the electrons on both sides are not the same
  you must find the least common multiple between
  the 2 electrons. Multiply each reaction to get
  the same number of electrons on both sides.
• It is important to check that all atoms and
  charges balance at this point. THE MISTAKES ARE
  LIKELY TO HAPPEN HERE!!!
• Combine the reactions and simplify if necessary.
• If in basic solution add OH- ions to both sides
  to produce water molecules on the side with H
  ions. Simplify the water molecules if necessary.

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Zn (s) 2HCl (aq) ? ZnCl2 (aq) H2 (g)

• Oxidation Zn ? Zn2 2e-
• Reduction 2H 2e- ? H2

Net ionic equation Zn (s) 2H (aq) ? H2
(g) Zn2 (aq)
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Balance the following redox

• (a) Fe2(aq) MnO4- (aq) ?Fe3(aq) Mn2(aq)
  (acidic solution)
• (b) Cl2(g) Cr(OH)3(s) ? Cl- (aq) CrO42- (aq)
• (basic solution)

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MC 1

• When 70. milliliter of 3.0-molar Na2CO3 is added
  to 30. milliliters of 1.0-molar NaHCO3 the
  resulting concentration of Na is
• (A) 2.0 M(B) 2.4 M(C) 4.0 M(D) 4.5 M (E) 7.0
  M

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MC 2

• A student wishes to prepare 2.00 liters of
  0.100-molar KIO3 (molecular weight 214 g/mol).
  The proper procedure is to weigh out
• (A) 42.8 grams of KIO3 and add 2.00 kilograms of
  H2O(B) 42.8 grams of KIO3 and add H2O until the
  final homogeneous solution has a volume of 2.00
  liters(C) 21.4 grams of KIO3 and add H2O until
  the final homogeneous solution has a volume of
  2.00 liters(D) 42.8 grams of KIO3 and add 2.00
  liters of H2O(E) 21.4 grams fo KIO3 and add 2.00
  liters of H2O

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MC 3

• A 20.0-milliliter sample of 0.200-molar K2CO3
  solution is added to 30.0 milliliters of
  0.400-molar Ba(NO3)2 solution. Barium carbonate
  precipitates. The concentration of barium ion,
  Ba2, in solution after reaction is
• (A) 0.150 M(B) 0.160 M(C) 0.200 M(D) 0.240
  M(E) 0.267 M

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MC 4

• The weight of H2SO4 (molecular weight 98.1 g/mol)
  in 50.0 milliliters of a 6.00-molar solution is
• (A) 3.10 grams(B) 12.0 grams(C) 29.4 grams(D)
  294 grams(E) 300. grams

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MC 5

• Given that a solution is 5 percent sucrose by
  mass, what additional information is necessary to
  calculate the molarity of the solution?
• I. The density of waterII. The density of the
  solutionIII. The molar mass of sucrose
• (A) I only(B) II only(C) III only(D) I and
  III(E) II and III

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MC 6

• A yellow precipitate forms when 0.5 M NaI(aq) is
  added to a 0.5 M solution of which of the
  following ions?
• A) Pb2 (aq)B) Zn2 (aq) C) CrO42 (aq) D)
  SO42 (aq) E) OH (aq)

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MC 7

• When 100 mL of 1.0 M Na3PO4 is mixed with 100 mL
  of 1.0 M AgNO3, a yellow precipitate forms and
  Ag becomes negligibly small. Which of the
  following is a correct listing of the ions
  remaining in solution in order of increasing
  concentration?
• A) PO43 lt NO3 lt NaB) PO43 lt Na
  lt NO3C) NO3 lt PO43 lt NaD) Na lt
  NO3 lt PO43E) Na lt PO43 lt NO3

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MC 8

• The volume of distilled water that should be
  added to 10.0 mL of 6.00 M HCl(aq) in order to
  prepare a 0.500 M HCl(aq) solution is
  approximately
• A) 50.0 mLB) 60.0 mLC) 100. mLD) 110. mLE)
  120. mL

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MC 9

• The net ionic equation for the reaction between
  silver carbonate and hydrochloric acid is
• (A) Ag2CO3 2H 2 Cl ---gt 2 AgCl H2O CO2
• (B) 2Ag CO32 2 H 2 Cl ---gt 2 AgCl
  H2O CO2(C) CO32 2 H ---gt H2O CO2(D) Ag
  Cl ---gt AgCl(E) Ag2CO3 2H ---gt 2Ag
  H2CO3

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MC 10

• 5 Fe2 MnO4 8 H ltgt 5 Fe3 Mn2 4H2O
• In a titration experiment based on the equation
  above, 25.0 milliliters of an acidified Fe2
  solution requires 14.0 milliliters of standard
  0.050-molar MnO4 solution to reach the
  equivalence point. The concentration of Fe2 in
  the original solution is
• (A) 0.0010 M(B) 0.0056 M(C) 0.028 M(D) 0.090
  M(E) 0.14 M

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FRQ 1

• A 1.2516 gram sample of a mixture of CaCO3 and
  Na2SO4 was analyzed by dissolving the sample and
  completely precipitating the Ca2 as CaC2O4. The
  CaC2O4 was dissolved in sulfuric acid and the
  resulting H2C2O4 was titrated with a standard
  KMnO4 solution.
• (a) Write the balanced equation for the titration
  reaction, shown unbalanced below
• MnO4- H2C2O4 H ? Mn2 CO2 H2O
• (i) Indicate which substance is the oxidizing
  agent and which substance is the reducing agent.
• (b) The titration of the H2C2O4 obtained required
  35.62 milliliters of 0.1092 molar MnO4- solution.
  Calculate the number of moles of H2C2O4 that
  reacted with the MnO4-
• (c) Calculate the number of moles of CaCO3 in the
  original sample.
• (d) Calculate the percentage by weight of CaCO3
  in the original sample.

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FRQ 2

• Permanganate ion, MnO4-, oxidizes sulfite ions to
  sulfate ion. The manganese product depends upon
  the pH of the reaction mixture. The mole ratio of
  oxidizing to reducing agent is two to five at pH
  1 (acidic), and is two to one at pH 13 (basic).
  For each of these cases, write a balanced
  equation for the reaction, and indicate the
  oxidation state of the manganese in the product
  containing manganese.

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FRQ 3

• A 0.150 g sample of solid lead(II) nitrate is
  added to 125 mL of 0.100 M sodium iodide
  solution. Assume no change in volume of the
  solution. The chemical reaction that takes place
  is represented by the following equation
• Pb(NO3)2(s) 2 NaI(aq) ? PbI2(s) 2NaNO3(aq)
• (a) List an appropriate observation that
  provides evidence of a chemical reaction between
  the two compounds.
• (b) Calculate the number of moles of each
  reactant.
• (c) Identify the limiting reactant. Show
  calculations to support your identification.
• (d) Calculate the molar concentration of
  NO3(aq) in the mixture after the reaction is
  complete.

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FRQ 4

• Answer the following questions about
  acetylsalicylic acid, the active ingredient in
  aspirin.
• (a) The amount of acetylsalicylic acid in a
  single aspirin tablet is 325 mg, yet the tablet
  has a mass of 2.00 g. Calculate the mass percent
  of acetylsalicylic acid in the tablet.
• (b) A student dissolved 1.625 g of pure
  acetylsalicylic acid in distilled water and
  titrated the resulting solution to the
  equivalence point using 88.43 mL of 0.102 M
  NaOH(aq). Assuming that acetylsalicylic acid has
  only one ionizable hydrogen, calculate the molar
  mass of the acid.

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FRQ 5 (part I)

• 5 Fe2(aq) MnO4(aq) 8 H(aq) ? 5 Fe3(aq)
  Mn2(aq) 4H2O(l)
• The mass percent of iron in a soluble iron(II)
  compound is measured using a titration based on
  the balanced equation above.
• (a) What is the oxidation number of manganese in
  the permanganate ion, MnO4(aq)?
• (b) Identify the reducing agent in the reaction
  represented above. Explain your reasoning.
• The mass of a sample of the iron(II) compound is
  carefully measured before the sample is dissolved
  in distilled water. The resulting solution is
  acidified with H2SO4(aq). The solution is then
  titrated with MnO4(aq) until the end point is
  reached.
• (c) Describe the color change that occurs in the
  flask when the end point of the titration has
  been reached. Explain why the color of the
  solution changes at the end point.

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FRQ 5 (part II)

• (d) Let the variables g, M, and V be defined as
  follows
• g the mass, in grams, of the sample of the
  iron(II) compound
• M the molarity of the MnO4(aq) used as the
  titrant
• V the volume, in liters, of MnO4(aq) added
  to reach the end point
• In terms of these variables, the number of moles
  of MnO4(aq) added to reach the end point of the
  titration is expressed as M x V. Using the
  variables defined above, the molar mass of iron
  (55.85 g mol-1), and the coefficients in the
  balanced chemical equation, write the expression
  for each of the following quantities
• (i) The number of moles of iron in the sample
• (ii) The mass of iron in the sample, in grams
• (iii) The mass percent of iron in the compound
• (e) What effect will adding too much titrant
  have on the experimentally determined value of
  the mass percent of iron in the compound? Justify
  your answer.

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Equations 1

• (a) A sample of solid iron(III) oxide is reduced
  completely with solid carbon.
• (i) Balanced equation
• (ii) What is the oxidation number of carbon
  before the reaction, and what is the oxidation
  number of carbon after the reaction is complete
• (b) Equal volumes of equimolar solutions of
  ammonia and hydrochloric acid are combined.
• (i) Balanced equation
• (ii) Indicate whether the resulting solution is
  acidic, basic, or neutral. Explain.
• (c) Solid mercury(II) oxide decomposes as it is
  heated in an open test tube in a fume hood.
• (i) Balanced equation
• (ii) After the reaction is complete, is the mass
  of the material in the test tube greater than,
  less than, or equal to the mass of the original
  sample? Explain.

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Equations 2

• (a) A small piece of sodium is placed in a beaker
  of distilled water.
• (i) Balanced equation
• (ii) The reaction is exothermic, and sometimes
  small flames are observed as the sodium reacts
  with the water. Identify the product of the
  reaction that burns to produce the flames
• (b) Hydrogen chloride gas is oxidized by oxygen
  gas.
• (i) Balanced equation
• (ii) If three moles of hydrogen chloride gas and
  three moles of oxygen gas react as completely as
  possible, which reactant, if any, is present in
  excess? Justify your answer.
• (c) Solid potassium oxide is added to water.
• (i) Balanced equation
• (ii) If a few drops of phenolphthalein are added
  to the resulting solution, what would be
  observed? Explain.