Transcript Chapter 3: Chemical Reactions and the Earth’s Composition

Chapter 3: Chemical Reactions
and the Earth’s Composition
Problems: 3.1-3.3, 3.5, 3.11-3.86, 3.95-3.115,
3.119-3.120, 3.122, 3.125-3.128, 3.132, 3.134,
3.136-3.138-3.141
3.2 The Mole
Stoichiometry (STOY-key-OM-e-tree): quantitative study of
reactants and products in a chemical reaction
Interpreting a Chemical Equation
H2 (g) +
Cl2 (g)
 2 HCl (g)
1 molecule 1 molecule
2 molecules
It follows that any multiples of these coefficients will be in same ratio!
2 H2 (g)
+
O2 (g)

2 H2O(g)
1000
_____ molecule(s) _____ molecule(s) _____ molecule(s)
N
_____ molecule(s) _____ molecule(s) _____ molecule(s)
Since N = Avogadro’s # = 6.0221023 molecules = 1 mole
2 H2 (g)
+
O2 (g)

2 H2O(g)
_____ mole(s)
_____ mole(s)
_____ mole(s)
Thus, the coefficients in a chemical equation give the mole ratios of
reactants and products.
Example Problem
Consider the following:
2 C2H6 (g) + 7 O2 (g)  4 CO2 (g) + 6 H2O (g)
1. How many moles of O2 will react with 2.50 moles of C2H6?
2. How many moles of CO2 form when 3.50 moles of O2
completely react?
Stoichiometric Calculations
and the Carbon cycle
Mass-Mass Stoichiometry Problems
MASS OF
KNOWN
Molar
Mass
MOLES OF MOLE-MOLE
Ratio
KNOWN
MOLES OF
UNKNOWN
Molar
Mass
MASS OF
UNKNOWN
Example 1: Photosynthesis is the process of energy from
sunlight being used to convert carbon dioxide into organic
compounds, especially sugars like glucose, C6H12O6:
6 CO2(g)
+
6 H2O(g)

C6H12O6(aq) + 6 O2(g)
What mass (in g) of glucose is produced via photosynthesis
when 25.0 kg of carbon dioxide react with excess steam?
6 CO2(g)
+
6 H2O(g)

C6H12O6(aq) + 6 O2(g)
What mass (in g) of glucose is produced via
photosynthesis when 25.0 kg of carbon dioxide react
with excess steam?
Stoichiometric Calculations
and the Carbon cycle
Example 2:
In biological systems, the reverse reaction,
C6H12O6(aq) + 6 O2(g)

6 CO2(g)
+
6 H2O(g)
is called respiration and is the major source of energy for all
livings things.
a. Calculate the mass (in g) of carbon dioxide produced
when 5.00 lb. of glucose reacts completely. (1 lb. = 453.6 g)
b. How many pounds of carbon dioxide are produced in
problem a. above?
C6H12O6(aq) + 6 O2(g)

6 CO2(g)
+
6 H2O(g)
Calculate the mass (in g) of carbon dioxide produced when
5.00 lb. of glucose reacts completely.
b. How many pounds of carbon dioxide are produced in
problem a. above?
3.4 Combustion Reactions
CxHy + O2(g)  CO2(g) + H2O(g)
CxHyOz + O2(g)  CO2(g) + H2O(g)
Hydrocarbons (compounds with only C and H) and
hydrocarbon derivatives (compounds with only C, H and
O) burn in O2 to produce CO2 gas and steam, H2O(g).
Combustion Reactions
Example 1: Many home barbecues are fueled with
propane gas (C3H8). Write the balanced equation for the
combustion of propane, then calculate the mass (in kg)
of carbon dioxide produced upon complete combustion
of liquid propane from a 5.0 gal tank.
(The density of liquid propane at 60°F is about 4.2
lbs. per gallon, and 1 lb. = 453.6 g)
Combustion Reactions
Example 2: Everclear is a brand of grain alcohol that
can be as high as 190 proof (or 95% ethanol, C2H5OH,
by volume). Calculate the mass of carbon dioxide
produced upon complete combustion of the ethanol in
a 750 mL bottle of Everclear. Write the balanced
chemical equation for the combustion of ethanol. (The
density of this Everclear is 0.80 g/mL.)
3.9 Limiting Reactants and
Percent Yield
• In practice, reactants will not always be present in the
exact amounts necessary to be converted completely
into products.
• Some reactants (usually the more expensive) are only
present in a limited supply, so these are almost
always completely used up
– “limiting reactant” (or limiting reagent) since it
limits the amount of product made
• Some reactants (usually the less expensive) are
present in larger amounts and are never completely
used up  “reactant(s) in excess”
Guidelines for solving Limiting
Reactant Problems
1. Calculate the mass or the # of moles of the 2nd reactant
needed to completely react with the 1st reactant.
– If the moles needed is greater than the number of
moles present for the 2nd reactant
• That 2nd reactant will run out before the 1st reactant.
• The 2nd reactant = the limiting reactant, and the
1st reactant is in excess.
– If the moles needed is less than the number of moles
present for the reactant,
• The 1st reactant = the limiting reactant, and the
2nd reactant is in excess.
2. Use the amount of the limiting reactant present to solve
for the mass or # of moles of product that can be made.
Limiting Reactant Problems
Consider the reaction to produce ammonia:
N2(g) + 3 H2(g)  2 NH3(g)
Example 1: a) If 50.0 g of N2 react with 10.0 g of H2,
what mass of ammonia is produced?
b) The limiting reactant is _______ and the excess
reactant is _________.
c) What mass of the reactant in excess remains after the
reaction?
Calculating Percent Yield
actual yield
Percent yield =
 100%
theoretical yield
• Theoretical yield: Amount of product one should get based
on the chemical equation and the amount of reactants
present
– One generally calculates this in grams from info given
• Actual yield: Amount of product one actually obtains
– Generally smaller than the theoretical yield because of
impurities and other adverse conditions in the lab
– This is generally determined experimentally in the lab or
given for a problem in lecture.
Calculating Percent Yield
Example 1: N2(g) + 3 H2(g)  2 NH3(g)
a. For the reaction of 50.0 g of N2 with 10.0 g of H2, the
theoretical yield of ammonia was determined to be
what?
theoretical yield =
If 49.6 g of ammonia were actually produced, calculate
the percent yield for the reaction.
percent yield =
Calculating Percent Yield
Example 2:
Consider the following reaction:
2 KClO3(s)  2 KCl(s) + 3 O2(g)
What is the percent yield if 50.0 g of KClO3 decomposes
to produce 16.4 g of oxygen gas?
Calculating Percent Yield
Example 3:
Consider the following reaction:
3 Na2CrO4 (aq) + 2 AlCl3 (aq)

Al2(CrO4)3 (s) +
6 NaCl (aq)
a. What mass of precipitate is produced when 50.0 g of
sodium chromate react with 50.0 g of aluminum
chloride? Which is the limiting reactant and which is the
reactant in excess?
Calculating Percent Yield
Example 3:
Consider the following reaction:
3 Na2CrO4(aq) + 2 AlCl3(aq)
6 NaCl(aq).

Al2(CrO4)3(s) +
b. What mass of the reactant in excess remains after the
reaction?
c. What is the percent yield if 4.32 g of precipitate is
actually produced?
Calculating Percent Yield
Example 4:
Calculate the mass of methane (CH4) that must react to
produce 10.0 kg of carbon dioxide if the percent yield
for the reaction is 88.8%.
Calculating Percent Yield
Example 5:
Consider the thermal decomposition of N2O5:
2 N2O5(g) 4 NO2(g) + O2(g)
If the percent yield for the reaction is 96.8%, and the
density of oxygen gas is 1.31 g/L, calculate the mass of
N2O5 required to produce 50.0 L of oxygen gas.