Unit Four Chemical Reactions

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Unit Four Chemical Reactions Powered By Docstoc
					    As you come in,
       The Materials:
         Remote control
         Pick up packet.
         Paper, pencil, calculator, periodic table for
          notes
       The Plan:
           Learn about stoichiometry
       The Assessments:
         Stoichiometry Quiz on Thursday
         Limiting Reactants Quiz on Friday


Halloween Pandora Station – Project Eliminate Challenge
Unit 7
Stoichiometry
Chapter 9
Info in a Chemical Equation
       2K + CuSO4 à Cu + K2SO4

A balanced chemical equation shows the
correct ratios required for a chemical
reaction to occur.
Knowing the ratio allows us to provide the
appropriate amount of reactants AND
predict the amount of products.
Mole to Mole Relationships
 Since chemical equations give
  appropriate relationships of moles
  of each compound, mole ratios
  can be written.
 Mole ratio: uses the coefficients in
  the balanced equation to show
  how two compounds are related in
  a reaction
Example 9.2
Goal: Write a mole ratio & solve the problem.
What    number of moles of O2
  will be produced by the
  decomposition of 5.8 mol of
  water?

         the balanced chemical
   First,
    equation must be written for the
    situation described in the prompt.
   2H2O à 2H2 + O2
Example 9.2 (continued)
What number of moles of O2 will
 be produced by the
 decomposition of 5.8 mol of
 water?
            2H2O à 2H2 + O2
Next, the reaction shows that oxygen and
water have a quantitative relationship: a
mole ratio.
           1 mole O2 : 2 moles H2O
Where did the numbers (1 & 2) come from?
Can you write other mole ratios from this
equation?
Example 9.2 (continued)
What  number of moles of O2 will
  be produced by the
  decomposition of 5.8 mol of
  water?
              2H2O à 2H2 + O2
Finally,we can solve this problem using the
mole ratio as a conversion factor in our
dimensional analysis.
   Startwith the given.
   Set up the chart to cancel units AND
    compounds.
   Multiply or divide as usual.
Example 9.3
Goal: Solve the problem with a partner.
Calculate  the number of moles of
 oxygen required to react exactly
 with 4.30 mol of propane, C3H8, in
 the reaction described by the
 following:
 C3H8(g) + 5O2(g) à 3CO2(g) + 4H2O(g)
  Calculate the number of moles of
oxygen required to react exactly with
   4.30 mol of propane, C3H8, in the
 reaction described by the following:
  C3H8(g) + 5O2(g)à 3CO2(g) + 4H2O(g)
A.) 0.860 moles of O2
B.) 7.17 moles of O2
C.) 21.5 moles of O2
D.) none of these
  Calculate the number of moles of
oxygen required to react exactly with
   4.30 mol of propane, C3H8, in the
 reaction described by the following:
  C3H8(g) + 5O2(g)à 3CO2(g) + 4H2O(g)
A.) 0.860 moles of O2
B.) 7.17 moles of O2
C.) 21.5 moles of O2
D.) none of these
Example 9.4
Is another example needed?
   Ammonia is used in huge quantities as a
    fertilizer. It is manufactured by combining
    nitrogen and hydrogen according to the
    following equation:
                N2(g) + 3H2(g) = 2NH3(g)
    Calculate the number of moles of NH3 that
    can be made from 1.30 mol H2(g) reacting
    with excess N2(g).
    Suggested Homework
   Page 287 4 and 5
Mega Mole Map
   Mole to mole relationship
     Mole Ratio = conversion factor (bridge)


 In our examples so far, we’ve been given
  moles of one compound and asked to
  convert to moles of another compound.
 What if I told you that I won’t always start
  out with moles or ask for moles? Starting
  or ending with mass, liters, or particles is
  very common.
 Work with your partner to expand on our
  current Mole Concept Map.
Mega Mole Map
   Mole to mole relationship
     Mole Ratio = conversion factor (bridge)


   Mole to mass relationship
     Molar Mass = conversion factor (bridge)


   Mole to particle relationship
     Avogadro’s # = conversion factor (bridge)


   Mole to volume relationship
     Molar volume = conversion factor (bridge)
Example 9.5
Goal: Use new mole concepts to solve the
problem.
 Consider the reaction of powdered
  aluminum metal and finely ground iodine
  to produce aluminum iodide. The
  balanced equation for this vigorous
  chemical reaction is:
           2Al(s) + 3I2(s) = 2AlI3(s)
  Calculate the mass of I2(s) needed to just react
  with 35.0 g of Al(s).

 First,   write the given and set up your chart.
Example 9.5 (continued)
 Consider   the reaction of powdered
    aluminum metal and finely ground iodine to
    produce aluminum iodide. The balanced
    equation for this vigorous chemical reaction
    is:
               2Al(s) + 3I2(s) = 2AlI3(s)
    Calculate the mass of I2(s) needed to just react with
    35.0 g of Al(s).

 Next, consult your map to determine a problem-
  solving path. How many steps will this problem
  require?
 3 steps
Example 9.5 (continued)
 Consider   the reaction of powdered
    aluminum metal and finely ground iodine to
    produce aluminum iodide. The balanced
    equation for this vigorous chemical reaction
    is:
               2Al(s) + 3I2(s) = 2AlI3(s)
    Calculate the mass of I2(s) needed to just react with
    35.0 g of Al(s).

 Finally, plug in your conversion factors. Multiply
  and divide as usual to give your final answer.
 493.89 g I2
 Section 9.2 Review Questions
 Goal: Solve the problem with a
 partner.
2.   Solutions of sodium hydroxide cannot
     be kept for very long because they
     absorb carbon dioxide from the air,
     forming sodium carbonate. The
     unbalanced equation is:
NaOH(aq) + CO2(g) à Na2CO3(aq) + H2O(l)
     Calculate the number of grams of
     carbon dioxide that can be absorbed
     by complete reaction with a solution
     that contains 5.00 g of NaOH.
  NaOH(aq) + CO2(g) à Na2CO3(aq) + H2O(l)

  Calculate the number of grams of carbon
 dioxide that can be absorbed by complete
 reaction with a solution that contains 5.00 g
                  of NaOH.
A.) 2.75 g CO2

B.) 5.50 g CO2

C.) 9.09 g CO2

D.) none of these
NaOH(aq) + CO2(g) à Na2CO3(aq) + H2O(l)
    Calculate the number of grams of
carbon dioxide that can be absorbed by
complete reaction with a solution that
contains 5.00 g of NaOH.


A.) 2.75 g CO2

B.) 5.50 g CO2

C.) 9.09 g CO2

D.) none of these
Suggested Homework
   11-2 Practice Problems – (1-11…skip 9)

NOTE: We’ll use iRespond TODAY to check
 your answers, and we’ll make colored
 pencil corrections on your work.
Limiting Reactants
 As  you know, a balanced chemical
  equation gives the perfect ratio of
  reactants needed to perform the
  reaction.
 In reality, we rarely have the “perfect”
  ratio.
 With an imperfect ratio, one reactant will
  run out before the other. The reactant
  running out will stop the reaction...limit
  the products.
 The reactant that runs out = Limiting
  reactant
 The reactant that remains = Excess
  reactant
Limiting Reactant Calculations
   Using stoichiometry, we can calculate the
    following predictions:
     Identify the limiting reactant
     Identify the excess reactant
     Predict the amount of product to form
     Predict the amount of excess reactant
      remaining
Practice Problem 9.5
   Calculate the mass of AlI3(s) formed by
    the reaction of 35.0 g Al(s) with 495 g I2.
Example 9.7
 Suppose that 25.0 kg of nitrogen
 gas and 5.00 kg of hydrogen gas
 are mixed and reacted to form
 ammonia. Calculate the mass of
 ammonia produced when this
 reaction is run to completion.
    Two “givens” means two calculations
    The unknown is the same in the two
     calculations
    The smallest answer is the best prediction as it
     tells when the limiting reactant will run out.
Follow-up Questions
 Identify the limiting reactant.
 Identify the excess reactant.
 How much H2 is used? Left?
 How much N2 is used? Left?
Limiting Reactant Practice
1.The reaction between solid white phosphorus and
  oxygen produces solid tetraphosphorus decaoxide
  (P4O10). This compound is often called phosphorus
  pentoxide because its empirical formula is P2O5.

                    P4 + 5O2 à P4O10
     a.   Determine the mass of tetraphosphorus decaoxide
          formed if 25.0 g of phosphorus (P4) and 50.0 g of
          oxygen gas are combined.
     b.   Identify the limiting reactant.
     c.   How much of the excess reactant remains after the
          reaction stops?
FYI
   "Stoichiometry" is derived from the Greek
    words στοιχεῖον (stoicheion, meaning
    element]) and μέτρον (metron, meaning
    measure.)
Section 9.2
Review Questions
4.   You react 10.0 g of nitrogen gas
     with hydrogen gas according to
     the following reaction.
           N2(g) + 3H2(g) = 2NH3(g)
a.   What mass of hydrogen is required
     to completely react with 10.0 g
     sample of nitrogen gas?
b.   What mass of ammonia is
     produced from 10.0 g of nitrogen
     gas and sufficient hydrogen gas?
Practice Problem 9.8
   Lithium nitride, an ionic compound
    containing the Li+ and N3- ions, is
    prepared by the reaction of lithium metal
    and nitrogen gas. Calculate the mass of
    lithium nitride formed from 56.0 g of
    nitrogen gas and 56.0 g of lithium.
Follow-up Questions
 Identify the limiting reactant.
 Identify the excess reactant.
 How much of each reactant is used?
  Left?
Percent Yield
 Yield means product.
 Calculating % yield is calculating what %
  of the product your experiment actually
  produced.
 You are comparing your prediction to
  your action.
% Yield
 (Actual yield / theoretical yield) 100
 Actual yield = either given in the problem
  or carried out in the lab
 Theoretical yield = ALWAYS calculated by
  stoichiometry
Stoichiometry Test Tips
1.   Know how to write and balance a
     combustion reaction. (See #1 on
     your Stoichiometry Quiz.)
2.   Remember, “a liter” or “a single
     gram” means 1 L or 1 gram. The
     given can be written this way.
3.   “How many grams of product are
     produced?” - refers to theoretical
     yield (smallest answer) in a limiting
     reactant problem
4. Metal products are solids.
5. STP = Standard temperature and
    pressure (doesn’t affect the
    calculation)
6. % yield = (ACTUAL/THEORETICAL) 100
7. Excess reactant - Use dimensional
    analysis starting with given amount
    of LR to calculate the amount
    USED; Subtract from ORIGINAL
    amount to calculate excess LEFT
    Look closely at the
    wording…
   Labeling the equation will help!
       If you are given info about BOTH reactants, then
        you have a limiting reactant problem.
           Example: What mass of NaCl will be produced by the
            reaction of 58.7g of NaI with 29.4g of Cl2 gas if the
            products are NaCl and I2?
       If you are given info about ONLY ONE substance,
        then you have a simple stoichiometry problem.
More wording advice…
   LABELS!
       If you are given info about BOTH reactants
        AND a product, then you’ve probably got
        a percent yield problem to solve. The
        product info is the actual yield.
           Determine the percent yield for a reaction
            between 6.92g K and 4.28g of O2 if 7.36g of
            K2O is produced.

				
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