Masses of both reactants are given Determine the grams or moles of both

Document Sample
Masses of both reactants are given Determine the grams or moles of both Powered By Docstoc
					Stoichiometry
   Chapter 12
What is stoichiometry?
   The study of quantitative relationships
    between amounts used and products
    formed by a chemical reaction
       Based on the law of conservation of mass
            Chemical bonds in reactants break and new chemical
             bonds form to produce products, but the amount of
             matter present at the end of the reaction is the same
             as was present in the beginning
            Mass of reactants = mass of products
Mass of reactants = Mass of products
Example: CH4 + 2O2  CO2 + 2H20
Mole Mass
 CH4


   2O2

   CO2

   2H20
Mole Ratios
   Mole Ratio- a ratio between the number of
    moles of any substances in a BALANCED
    chemical equation
    Example: CH4 + 2O2  CO2 + 2H20
        1 mol CH4 : 2 mole O2


         2   mol H20 : 2 mol O2
   Your Turn  Determine all possible mole ratios
    for the following chemical equation:
       3Fe(s) + 4 H20(l)  Fe3O4(s) + 4 H2(g)
3Fe(s) + 4 H20(l)  Fe3O4(s) + 4 H2(g)
Practice Problems
   Write a balanced chemical equation for
    each reaction and determine the possible
    mole ratios.
    1.   Nitrogen reacts with hydrogen to produce
         ammonia (NH3)
    2.   Hydrogen peroxide (H2O2) decomposes to
         produce water and oxygen.
    3.   Pieces of zinc react with a phosphoric acid
         solution to produce solid zinc phosphate and
         hydrogen gas.
N2(g) + 3H2(g)  2 NH3(g)



2H2O2  2H2O + O2



3Zn(s) + 2 H3PO4(aq)  Zn3(PO4)2(aq) + 3 H2   (g)
Stoichiometric Calculations
   Mole Mole Calculations
       Plan to solve:
        1.   Balance Equation
        2.   Identify given quantity (in mol)
        3.   Conversion factor (mole ratio)
        4.   Calculate unknown quantity (in mol)

Example: The elements lithium and oxygen react explosively
   to form lithium oxide, Li2O. How many moles of lithium
   oxide will form if 2 mol of lithium react?
   Mole Mass Calculations
       Plan of solve:
        1.   Balance Equation
        2.   Identify given quantity (in mol)
        3.   Conversion factor (mole ratio)
        4.   Molar mass of unknown quantity (in g/mol)
        5.   Calculate unknown quantity (in grams)
   What mass, in grams, of glucose is produced
    when 3.00 mol of water reacts with carbon
    dioxide?
    __CO2 (g) + __H2O(l) __C6H12O6 (s) + __O2(g)
Mass Mole Calculations
       Plan of attack:
        1. Balance Equation

        2. Identify given quantity (in grams)

        3. Molar mass of given quantity (in g/mol)

        4. Conversion factor (mole ratio)

        5. Calculate unknown quantity (in mol)

   The first step in the industrial manufacture of
    nitric acid is the catalytic oxidation of ammonia.
    The reaction is run using 824g NH3 and excess
    oxygen. How many moles of NO are formed?
    __NH3(g) + __O2 (g)  __NO(g) + __H2O (g)
Mass Mass Calculations
       Plan of attack:
        1. Balance Equation

        2. Identify given quantity (in grams)

        3. Molar mass of given quantity (in g/mol)

        4. Conversion factor (mole ratio)

        5. Molar mass of unknown quantity (in g/mol)

        6. Calculate unknown quantity (in grams)

   How many grams of SnF2 are produced from the
    reaction of 30.00g of HF with Sn?
Practice Problem 
   Sulfuric acid is formed when sulfur
    dioxide reacts with oxygen and water.
    Write the balanced chemical equation for
    the reaction.
       If 12.5 mol SO2 reacts, how many moles of
        H2SO4 can be produced?
       If 2.50 g SO2 react with excess oxygen and
        water, how many grams of H2SO4 are
        produced?
2SO2(g) + O2(g) + 2H2O(l)  2H2SO4(aq)
Limiting Reactants
   Why do reactions stop?




   How many complete cars can you make?
   Limiting Reactant - The reactant in a chemical reaction
    that limits the amount of product that can be formed. The
    reaction will stop when all of the limiting reactant is
    consumed.
   Excess Reactant - The reactant in a chemical reaction
    that remains when a reaction stops when the limiting
    reactant is completely consumed. The excess reactant
    remains because there is nothing with which it can react.
The “Have and Need” Method
   How can you determine which reactant is
    limited?
       Masses of both reactants are given
       Determine the grams or moles of both
        reactants given (“Have”)
       Determine the grams or moles “needed” of
        each reactant by relating them
Practice Problem
   When 36.0 g of H2O is mixed with 167g of
    Fe, which is the limiting reactant?
           What mass in grams of black iron oxide is produced?
           What mass in grams of excess reactant remains
            when the reaction is completed?
       3Fe   (s)   + 4H2O(g)  Fe3O4   (s)   + 4H2 (g)
           Determine the grams of both reactants given
            (“Have”)
           Determine the grams “needed” of each reactant by
            relating them with their “Have and Need”
3Fe (s) + 4H2O(g)  Fe3O4 (s) + 4H2 (g)

   36.0 g H20


   167 g Fe




       Use the limiting reactant to calculate the
        product mass
Addition Practice
6 Na(s) + Fe2O3(s)  3Na2O (s) + 2 Fe (s)
If 100.0g Na and 100.0g Fe2O3 are used in the reaction,
   determine
      The limiting reactant
      The reactant in excess
      The mass of solid iron produced
      The mass of excess reactant that remains after the reaction is
       complete
6 Na(s) + Fe2O3(s)  3Na2 (s) + 2 Fe (s)
100.0g Na and 100.0g Fe2O3
Percent Yield
   The theoretical yield is the maximum amount of
    product from a given amount of reactant
   The actual yield is the amount of product actually
    produced when the chemical reaction is carried
    out in an experiment
   The percent yield of product is the ratio of the
    actual yield to the theoretical yield expressed as
    a percent

    % Yield =   actual yield (from an experiment)      x 100
                theoretical yield (from calculation)
Example
   If 75.0g of CO reacts to produce 68.4g of
    methanol (CH3OH), what is the percent
    yield of CH3OH?
        CO(g)   + 2H2 (g)  CH3OH   (l)
2Al(s) + 3CuSO4(aq) 3Cu (s) + Al2 (SO4)3 (aq)

   Aluminum reacts with an aqueous solution containing
    excess copper (II) sulfate. If 1.85 g Al reacts and the
    percentage yield of Cu is 56.6%, what mass of Cu is
    produced?
Use the data to determine the percent
yield of the following reaction.
   2Mg   (s)   + O2   (g)    2MgO   (s)   (Oxygen is in excess)


                  Reaction Data
Mass of crucible                  35.67 g
Mass of cruicble + Mg             38.06 g
Mass of Mg
Mass of Crucible + MgO            39.15 g
Mass of MgO

				
DOCUMENT INFO
Shared By:
Categories:
Tags:
Stats:
views:222
posted:2/25/2012
language:English
pages:24