Chemical Reactions - Mr. Fischer

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Chemical Reactions - Mr. Fischer Powered By Docstoc
  Chapter 4

     Balancing Chemical Equations
“Matter is conserved in
 chemical change”
           Antoine Lavoisier, 1789
An equation must be balanced:
It must have the same number
of atoms of each kind on both sides

             Combustion Reactions
   In combustion, a hydrocarbon or C–H–O fuel
    combines with O2 to form CO2 and H2O
       __ CH4 + __ O2  __ CO2 + __ H2O
       1 CH4 + 2 O2  1 CO2 + 2 H2O
    Balanced equation shows 1 C, 4 H, and 4 O on each side
   If N or S are in the formula for the fuel, assume it
    is oxidized to NO2 or SO2

                 Example 4-2
 Write a balanced equation for the complete
  combustion of glycerol, C3H8O3
 Write a balanced equation for the complete
  combustion of thiosalicylic acid, C7H6O2S

 Stoichiometry is chemical accounting
 The heart of stoichiometry is the mole ratio given
  by the coefficients of the balanced equation

 Stoichiometry is chemical accounting
 The heart of stoichiometry is the mole ratio given
  by the coefficients of the balanced equation

                         mole ratio
               moles A                moles B
                          moles B
                          moles A

                  Example 4-3B
   How many moles of Ag are produced in the
    decomposition of 1.00 kg of silver (I) oxide:
       2 Ag2O (s)  4 Ag (s) + O2 (g)

 Example 4-6B
The model problem describes
an Al-Cu alloy composed of
93.7% Al and 6.3% Cu by
mass, with a density of 2.85
g/cm3. The Al (but not the Cu)
reacts with HCl:
  2 Al (s) + 6 HCl (aq)  2 AlCl3 (aq) + 3 H2 (g)
How    many grams of Cu are present in a sample of
alloy that yields 1.31 g H2 when it reacts with HCl?

                  Example 4-7B
   A vinegar contains 4.0% HC2H3O2 by mass and
    has a density of 1.01 g/mL. It reacts with sodium
    hydrogen carbonate:
       HC2H3O2 (aq) + NaHCO3 (s) 
                         NaC2H3O2 (aq) + H2O (l) +
    CO2 (g)
    How many grams of CO2 are produced by the
    reaction of 5.00 mL of this vinegar with NaHCO3?

        Chemical Reactions in Solution
   Most reactions occur in aqueous solution
     SOLUTE is the substance to be dissolved in solution
     SOLVENT is the substance (often a liquid) the solute
      dissolves in
   The concentration of the solution is
       Molarity (M) = moles solute
                         L solution

                  Example 4-8B
   15.0 mL of concentrated acetic acid, HC2H3O2
    (d = 1.048 g/mL), are dissolved in enough water to
    produce 500.0 mL of solution. What is the
    concentration of the solution?

                 Example 4-9B
   How many grams of Na2SO4 • 10 H2O are needed
    to prepare 355 mL of 0.445 M Na2SO4?

              Dilution problems
 It is common to prepare a solution by diluting a
  more concentrated solution (the stock solution).
 The moles of solute taken from the stock solution
  are given by moles solute = volume x molarity
 All the solute taken from the stock appears in the
  diluted solution, so moles solute are constant:
            VstockMstock = VdiluteMdilute

                 Example 4-10A
   15.00 mL of 0.450 M K2CrO4 solution are diluted
    to 100.00 mL. What is the concentration of the
    dilute solution?

                 Example 4-10B
   After being left out in an open beaker, 275 mL of
    0.105 M NaCl has evaporated to only 237 mL.
    What is the concentration of the solution after

              Stoichiometry in Solution
   Stoichiometry in solution is just the same as for
    mass problems, except the conversion into or out
    of moles uses molarity instead of molar mass:
    grams A                                        grams B

                            mole ratio
                  moles A                moles B
                             moles B
                             moles A
      mL A                                         mL B

                Example 4-11B
K2CrO4 (aq) + 2 AgNO3 (aq)  Ag2CrO4 (s) + 2 KNO3 (aq)

   How many mL of 0.150 M AgNO3 must react with
    excess K2CrO4 to produce exactly 1.00 g

               Limiting reactant
 In a given reaction, often there is not enough of
  one reactant to use up the other reactant completely
 The reactant in short supply LIMITS the quantity
  of product that can be formed

               Goldilocks Chemistry
   Imagine reacting different amounts of Zn with
    0.100 mol HCl:
       Zn (s) + 2 HCl (aq)  ZnCl2 (aq) + H2 (g)

                    Rxn 1       Rxn 2        Rxn 3
    Mass Zn         6.54 g      3.27 g       1.31 g
    Moles Zn        0.100 mol   0.0500 mol   0.0200 mol
    Moles HCl       0.100 mol   0.100 mol    0.100 mol
    Ratio mol HCl   1.00        2.00         5.00
          mol Zn

           Limiting reactant problems
   The easiest way to do these is to do two
    stoichiometry calculations
       Find the amount of product possible from each reactant
 The smaller answer is the amount of product you
  can actually make (you just ran out of one reactant)
 The reactant on which that answer was based is the
  limiting reactant

                 Example 4-13A
   When 215 g P4 react with 725 g Cl2
      P4 (s) + 6 Cl2 (g)  4 PCl3 (l) (example 4-12A)
    which reactant is in excess and what mass of that
    reactant remains after the reaction is finished?

                  Example 4-13B
   12.2 g H2 and 154 g O2 are allowed to react.
    Identify the limiting reactant, which gas remains
    after the reaction, and what mass of it is left over.
              2 H2 (g) + O2 (g)  2 H2O (l)

                     Percent Yield
   In real experiments we often do not get the amount
    of product we calculate we should, because
     the reactants may participate in other reactions (side
      reactions) that produce other products (by-products)
     The reaction often does not go to completion.

   Percent yield tells the ratio of actual to theoretical
    amount formed.

                  Percent Yield
 Suppose you calculate that a reaction will produce
  50.0 g of product. This is the theoretical yield.
 The reaction actually produces only 45.0 g of
  product . This is the actual yield.

   Percent yield = 45.0 g (actual) x 100 = 90.0%
                   50.0 g (theoretical)

                  Example 4-14B
   What is the percent yield if 25.0 g P4 reacts with
    91.5 g Cl2 to produce 104 g PCl3:
       P4 (s) + 6 Cl2 (g)  4 PCl3 (l)

                 Example 4-15B
   What mass of C6H11OH should you start with to
    produce 45.0 g C6H10 if the reaction has 86.2%
    yield and the C6H11OH is 92.3% pure:
        C6H11OH (l)  C6H10 + H2O (l)

                    Exercise 26
   Balance these equations by inspection
     (NH4)2Cr2O7 (s)  Cr2O3 (s) + N2 (g) + H2O (g)
     NO2 (g) + H2O (l)  HNO3 (aq) + NO (g)
     H2S (g) + SO2 (g)  S (g) + H2O (g)
     SO2Cl2 + HI  H2S + H2O + HCl + I2

                     Exercise 30
   Write balanced equations for these reactions:
     Sulfur dioxide gas with oxygen gas to produce sulfur
      trioxide gas
     Solid calcium carbonate with water and dissolved
      carbon dioxide to produce aqueous calcium hydrogen
     Ammonia gas and nitrogen monoxide gas to produce
      nitrogen gas and water vapor

                    Exercise 32
       3 Fe (s) + 4 H2O (g)  Fe3O4 (s) + H2 (g)
     How many moles of H2 can be produced from 42.7 g Fe
      and excess steam?
     How many grams of H2O are consumed in the
      conversion of 63.5 g Fe to Fe3O4?
     If 7.36 mol H2 are produced, how many grams of Fe3O4
      must also be produced?

                   Exercise 36
   Silver oxide decomposes above 300 °C to yield
    metallic silver and oxygen gas. 3.13 g impure
    silver oxide yields 0.187 g O2. Assuming there is
    no other source of O2, what is the % Ag2O by mass
    in the original sample?

                    Exercise 42
   How many grams of CO2 are produced in the
    complete combustion of 406 g of a bottled gas that
    consists of 72.7% C3H8 (propane) and 27.3%
    C4H10 (butane), by mass?

                    Exercise 45
   What are the molarities of these solutes?
     150.0 g sucrose (C12H22O11) in 250.0 mL aqueous
     98.3 mg of 97.9% pure urea, CO(NH2)2, in 5.00 mL
      aqueous solution
     12.5.0 mL methanol (CH3OH, density = 0.792 g/mL) in
      15.0 L aqueous solution

                    Exercise 52
   After 25.0 mL of aqueous HCl solution is diluted
    to 500.0 mL, the concentration of the diluted
    solution is found to be 0.085 M HCl. What was
    the concentration of the original HCl solution?

                       Exercise 56
 Ca(OH)2 (s) + 2 HCl (aq)  CaCl2 (aq) + 2 H2O (l)
     How many grams of Ca(OH)2 will react completely with 415 mL
      of 0.477 M HCl?
     How many kilograms of Ca(OH)2 will react with 324 L of an
      HCl solution that is 24.28% HCl by mass, density = 1.12 g/mL?

                   Exercise 63
   0.3126 g oxalic acid, H2C2O4, is exactly
    neutralized by 26.21 mL of a NaOH solution.
    What is the concentration of the NaOH solution?
    + 2 NaOH  Na2C2O4 + 2 H2O

                    Exercise 70
   Chlorine can be generated by heating calcium
    hypochlorite and hydrochloric acid to form
    chlorine gas, calcium chloride, and water. If 50.0 g
    Ca(OCl)2 and 275 mL 6.00 M HCl react, how
    many grams of Cl2 gas form? Which reactant is
    left over, and how much (in grams)?

                       Exercise 72
2 C6H5NO2 + 4 C6H14O4  (C6H5N)2 + 4 C6H12O4 + 4 H2O
nitrobenzene   triethylene   azobenzene
   If 0.10 L nitrobenzene (d = 1.20 g/mL) react with
    0.30 L triethylene glycol (d = 1.12 g/mL) to form
    55 g azobenzene, find
     Theoretical yield
     Actual yield
     Percent yield


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