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Title: Chapter 4 Aqueous Reactions and Solution Stoichiometry


1
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.

2
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

3
Aqueous Solutions
  • Water is the dissolving medium, or solvent.

4
Figure 4.1 The Water Molecule is Polar
5
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.

6
Solutions
  • Homogeneous mixtures of two or more pure
    substances.
  • The solvent is present in greatest abundance.
  • All other substances are solutes.

7
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)

8
A Solvent
  • retains its phase (if different from the solute)
  • is present in greater amount (if the same phase
    as the solute)

9
Polar Water Molecules Interact with the Positive
and Negative Ions of a Salt
10
BaCI2 Dissolving
11
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.

12
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.

13
Electrolytes and Nonelectrolytes
  • Soluble ionic compounds tend to be electrolytes.

14
Electrolytes and Nonelectrolytes
  • Molecular compounds tend to be nonelectrolytes,
    except for acids and bases.

15
Electrolytes
  • A strong electrolyte dissociates completely when
    dissolved in water.
  • A weak electrolyte only dissociates partially
    when dissolved in water.

16
Strong Electrolytes Are
  • Strong acids

17
Strong Electrolytes Are
  • Strong acids
  • Strong bases

18
Strong Electrolytes Are
  • Strong acids
  • Strong bases
  • Soluble ionic salts

19
Electrolytes
  • Strong - conduct current efficiently
  • NaCl, HNO3
  • Weak - conduct only a small current
  • vinegar, tap water
  • Non - no current flows
  • pure water, sugar solution

20
Strong Electrolytes
  1. Soluble salts- eg. NaCl, Pb(ClO3)2
  2. Strong acids- eg. HCl, H2SO4
  3. Strong bases- eg. NaOH, KOH

21
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

22
September 23-Section 4.2Precipitation Reactions
  • SOLUBILITY
  • PRECIPITATION REACTIONS
  • DOUBLE REPLACEMENT (METATHESIS EXCHANGE )
    REACTION
  • NET IONIC EQUATIONS
  • SPECTATOR IONS

23
Precipitation Reactions
  • When one mixes ions that form compounds that are
    insoluble (as could be predicted by the
    solubility guidelines), a precipitate is formed.

24
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.

25
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.
26
  • 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.

27
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28
  • 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.

29
Molecular Equation
  • The molecular equation lists the reactants and
    products in their molecular form.
  • AgNO3 (aq) KCl (aq) ?? AgCl (s) KNO3 (aq)

30
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)

31
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)

32
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)

33
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)

34
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.

35
  • 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.

36
  • 2. Aqueous solutions of barium chloride and
    potassium sulfate are mixed.

37
  • 3. Aqueous solutions of silver nitrate and
    ammonium chloride are mixed.

38
Homework
  • Page 157 answer 4.7 to 4.9
  • Page 158 4.19 to 4.28 odd
  • Read Section 4.3

39
September 24- Section 4.3 Acid Bases Reactions
  • Strong and weak Acid and Bases
  • Neutralization Reactions

40
Classify as strong, weak or nonelectrolyte
  • CaCl2
  • HNO3
  • C2H5OH
  • HC2H3O2

41
Rank the following solutions in order of
increasing electrical conductivity
  • Ca(NO3)2
  • C6 H12 O6
  • NaC2 H3O2
  • HC2 H3O2

42
Acids
  • Substances that increase the concentration of H
    when dissolved in water (Arrhenius).
  • Proton donors (BrønstedLowry).

43
Acids
  • There are only seven strong acids
  • Hydrochloric (HCl)
  • Hydrobromic (HBr)
  • Hydroiodic (HI)
  • Nitric (HNO3)
  • Sulfuric (H2SO4)
  • Chloric (HClO3)
  • Perchloric (HClO4)

44
  • 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.

45
Bases
  • Substances that increase the concentration of OH-
    when dissolved in water (Arrhenius).
  • Proton acceptors (BrønstedLowry).

46
  • Bases
  • Bases substances that react with the H ions
    formed by acids (e.g. NH3, Drano, Milk of
    Magnesia). Proton acceptors (Bronsted-Lowry).

47
  • 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.

48
Acid-Base Reactions
  • In an acid-base reaction, the acid donates a
    proton (H) to the base.

49
  • 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)

50
Neutralization Reactions
  • Observe the reaction between Milk of Magnesia,
    Mg(OH)2, and HCl.

51
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.

52
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
53
  • 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)

54
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)

55
Sulfite ions produce SO2 and H2O
  • SrSO3 (s) 2 HI (aq) ??SrI2 (aq) SO2 (g) H2O
    (l)

56
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)

57
LEO GER
  • LOSING
  • ELECTRONS
  • OXIDATION
  • O.N. increases in oxidation
  • GAINING
  • ELECTRONS
  • REDUCTION
  • O.N. decreases in reduction

58
Oxidation-Reduction Reactions
  • An oxidation occurs when an atom or ion loses
    electrons.
  • A reduction occurs when an atom or ion gains
    electrons.

59
Oxidation-Reduction Reactions
  • One cannot occur without the other.

60
Oxidizing and Reducing Agents
  • Oxidizing agents cause oxidation to occur. How?
  • Reducing agents cause reduction to occur. How?

61
A Summary of an Oxidation-Reduction Process
62
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

63
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.

64
  • 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

65
  • 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

66
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.

67
  • Examples Determine the oxidation numbers of
    each element in each of the following
  • H2SO4
  • KMnO4
  • NO3-
  • C2H6
  • CH3OH

68
Displacement ReactionsSingle Replacement(always
redox reactions!)
  • In displacement reactions, ions oxidize an
    element.
  • The ions, then, are reduced.

69
Displacement Reactions
  • In this reaction,
  • silver ions oxidize
  • copper metal.
  • Cu (s) 2 Ag (aq) ?? Cu2 (aq) 2 Ag (s)

70
Displacement Reactions
  • The reverse reaction,
  • however, does not
  • occur.
  • Cu2 (aq) 2 Ag (s) ?? Cu (s) 2 Ag (aq)

x
71
  • 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.

72
  • 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.

73
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74
  • 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.

75
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

76
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.

77
Example
  • Balance the following oxidation reduction using
    the half-reaction method
  • MnO2 Cl- ? Mn2 Cl2

78
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.

79
Example
  • Balance the following reaction which occurs in
    alkaline solution.
  • Ag (s) CN-(aq) O2(g) ? Ag(CN)2-(aq)

80
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

81
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

82
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.

83
  • 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

84
  • 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

85
  • 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

86
Mixing a Solution
87
Dilution
88
  • Dilution
  • We recognize that the number of moles are the
    same in dilute and concentrated solutions.
  • So
  • MdiluteVdilute moles MconcentratedVconcentrate
    d
  • M1V1 M2V2

89
  • Find the molarity of the solution if 1 ml of the
    solution are diluted to 500 ml
  • Step 1 M V M V

90
Example What volume of 18.0 M H2SO4 solution is
required to prepare 1000.0 mL of 1.00 M H2SO4?
91
October 1 -Section 4.6Solution Stoichiometry and
Chemical analysis
  • Titration
  • Problems

92
  • 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.

93
  • Use the stoichiometric coefficients to move
    between reactants and product.

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95
  • 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.

96
  • 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.

97
  • 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?

98
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?
99
Titration
  • The analytical technique in which one can
    calculate the concentration of a solute in a
    solution.

100
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.

101
Titration
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