Chapter 10. Chemical Calculations II Chemical Reactions

1
Chapter 10.Chemical Calculations IIChemical
Reactions

• Stoichiometry is the calculation of the
  quantities of reactants and products in a
  chemical reaction.
• It's very important that we can do this. We can
  plan the amounts of ingre-dients needed to make a
  compound such as a drug.

2
Conservation of Mass

• Atomic theory states that atoms are not created
  or destroyed in a chem-ical reaction, they just
  trade partners.
• The Law of Conservation of Mass states that mass
  is neither created nor destroyed in a chemical
  reaction.
• The reaction equation for a chemical change uses
  these principles.

3
Conservation of Mass

• CH4 2 O2 ?? CO2 2 H2O
• Each side of the equation has
• 1 carbon, 4 oxygens, 4 hydrogens

4
Conservation of Mass

• CH4 2 O2 ?? CO2 2 H2O
• 16.05 g CH4 44.01 g CO2
• 64.00 g O2 36.04 g H2O
• 80.05 g reactants 80.05 g products

5
Chemical Equations

• A chemical equation is a representation of a
  chemical reaction that uses chemical symbols
  instead of words to describe the changes that
  occur in the reaction.
• It's a sentence that describes something.
• It's a mathematical equation that describes a
  quantitative relationship.

6
Chemical Equations

• Parts of a chemical equation
• (a) Reactants are compounds that are present at
  the start of a reaction. Their formulas are shown
  on the left of the reaction arrow
• (b) Products are compounds that result from the
  reaction. Their formulas are shown on the right
  of the reaction arrow.
• (c) The reaction arrow is read as "to produce"
  and can be treated mathematically like "".
• (d) Coefficients show the number of moles of each
  compound in the reaction.

7
Chemical Equations

• A balanced chemical equation has coefficients
  that show the number of moles of reactants and
  products, such that the same number of each type
  of atoms appear on both sides of the equation.
• A skeleton equation shows reactants and products,
  but no coefficients. It must be balanced to be
  useful in chemical calcu-lations.

8
Chemical Equations

• Balancing chemical equations by inspection
  involves some guesswork and testing.
• Start with the most complex compounds, and work
  toward the simpler compounds and elements.
• Work with coefficients. Do not change
  sub-scripts in chemical formulas!
• Make sure the result shows the lowest
  whole-number ratio of coefficients.

9
Chemical Equations

• Examples
• Balance the following skeleton equations.
• (a) H2 O2 ?? H2O
• (b) Fe2O3 CO ?? Fe CO2
• (c) C3H8 O2 ?? CO2 H2O
• (d) C4H10 O2 ?? CO2 H2O
• Competency II-1

10
Equations and Moles

• 3 H2 N2 ?? 2 NH3
• This balanced equation can be read at the
  molecular level or the molar level.
• The 312 ratio holds for molecules and moles,
  and for any quantities of reactants.
• 9 moles of H2 will react with 3 moles of N2 to
  produce 6 moles of NH3

11
Equations and Moles

• 3 H2 N2 ?? 2 NH3
• We can treat the mole ratios as conversion
  factors and calculate relative moles of reactants
  and products based on these.
• 3 mol H2 2 mol NH3 2 mol NH3
• 1 mol N2 3 mol H2 1 mol N2

12
Equations and Moles

• 3 H2 N2 ?? 2 NH3
• How many moles of NH3 are produced if 4.85 moles
  of H2 are reacted with excess N2?
• And how realistic is that problem? We can't
  usually measure things in moles. We have to use
  mass and convert.

13
Reaction Calculations

• Typical types of calculations that are done for
  chemical reactions
• I have X amount of a reactant. How much
  product can I get?
• (b) I want to make X amount of a product. How
  much of each reactant do I need?
• (c) When I make X amount of the desired product,
  how much of the side products will be produced?

14
Reaction Calculations

• These are called theoretical yield calcula-tions,
  and they allow us to determine how much of a
  product is formed, or a reactant is needed, based
  on given quantities.

15
Reaction Calculations

• In a theoretical yield problem, look for
• (a) "Wants" the quantity requested.
• (b) "Gots" the quantity given. Enough
  information is given to figure out moles
  of "gots".
• (c) "Wants/Gots Factor" the of moles
  of the requested quantity divided by the
  of moles of the given quantity. The balanced
  reaction equation gives these numbers.

16
Reaction Calculations

• Steps
• (a) Calculate moles of Gots
• (b) Multiply that by Wants/Gots factor to
  get moles of Wants
• (c) Multiply moles of Wants by its
  molar mass to get mass of Wants

17
Reaction Calculations

• Example
• Fe2O3 3 CO ?? 2 Fe 3 CO2
• With 500 g of Fe2O3 and excess CO, how much iron
  metal (Fe) can be produced?

18
Reaction Calculations

• Example
• Fe2O3 3 CO ?? 2 Fe 3 CO2
• With 500 g of Fe2O3 and excess CO,
• (a) How much CO will be consumed?
• (b) How much CO2 will be produced?
• Competency II-2

19
Limiting Reactants

• We often don't use exactly the amounts of
  reactants called for in the balanced equation.
  One (usually the cheap-est!) is present in
  excess.
• The limiting reactant is the reactant that is
  entirely consumed when the reaction stops. Other
  reactants are present in excess.

20
Limiting Reactants

• If I have 55 nuts, 48 bolts, and 92 washers, how
  many sets of 1 bolt, 1 nut, 2 washers can I make?
  What part limits me? What is left over?

21
Limiting Reactants

• Chemistry is the same way.
• CH4O NaI ?? CH3I NaOH
• If one starts with 750 g of both reactants, what
  is the theoretical yield of CH3I (the maximum
  that can be obtained)? Which reactant is
  limiting?
• Work this as two theoretical yield problems, the
  lower result is the theoretical yield and the
  reactant that produces it is limiting.

22
Actual and Percent Yield

• We rarely obtain the theoretical yield of a
  prod-uct. The actual yield is the mass of
  product obtained in a reaction. It's usually
  lower than the theoretical yield, sometimes much
  lower! It's determined by experiment.
• The percent yield is the ratio
• actual yield x 100
• theoretical yield

23
Actual and Percent Yield

• Why do we care?
• Percent yield shows how efficient the reaction
  is, and/or how well the scientist or student did
  the reaction.
• Also, if a reaction routinely gives 50 yield,
  one had better start with twice the reactant as
  needed for the calculated theoretical yield!

24
Actual and Percent Yield

• Example
• If my theoretical yield for CH3I was 710 g, but I
  only isolated 650 g, what was my yield?
• Competency II-3