Empirical Formulas from Analysis

When iron reacts with oxygen in the air a compound (rust) is formed
that is 69.9% Fe and 30.1% O. Determine the empirical formula of
rust and write a balanced chemical equation for the process.

The mole ratios should be rounded to the nearest whole number
OR

to the nearest common fraction (1/4, 1/3, ½) and multiplied by the least
common denominator.
Combustion analysis
Experimental method for determining empirical formulas



When a compound containing carbon and hydrogen is
completely combusted, all the carbon in the compound is
converted to CO2 and all the hydrogen is converted to H2O.
The CO2 and H2O that is produced is weighed to determine
the empirical formula of the compound.
“Empirical” means “based on observation and experiment”.
Combustion analysis
Experimental method for determining empirical formulas

Isopropyl alcohol is composed of C, H, and O. Combustion
of 0.255 g of isopropyl alcohol produces 0.561 g CO2 and
0.306 g H2O. Determine the empirical formula.
Combustion analysis
Experimental method for determining empirical formulas

Caproic acid (stinky socks) is composed of C, H, and O
atoms. Combustion of a 0.225-g sample of this compound
produces 0.512 g CO2 and 0.209 g H2O.


What is empirical formula of caproic acid?
Caproic acid has a molar mass of 116 g/mol. What is its molecular
formula?
Quantitative Information from
Balanced Equations

The coefficients in a balanced chemical equation tell us the
relative number of molecules (or relative number of moles)
involved in the reaction.
2H2(g)
2 molecules
2(6.022x1023 molecules)
2 mol
+
O2(g)

1 molecule
1(6.022x1023 molecules)
1 mol
2H2O(l)
2 molecules
2(6.022x1023 molecules)
2 mol
Quantitative Information from
Balanced Equations
2H2(g)
2 molecules
2(6.022x1023 molecules)
2 mol


+
O2(g)

1 molecule
1(6.022x1023 molecules)
1 mol
2H2O(l)
2 molecules
2(6.022x1023 molecules)
2 mol
2 mol H2, 1 mol O2, and 2 mol H2O are stoichiometrically
equivalent quantities.
2 mol H2  1 mol O2  2 mol H2O
These stoichiometric relations can be used to convert
between quantities of reactants and products.
Stoichiometric Calculations
2H2(g)

+
O2(g)

2H2O(l)
How many moles of H2O are produced from 1.57 mol of O2?
 2 mol H 2O 
  3.14 mol H 2O
(1.57 mol O2 )
 1 mol O2 
From stoichiometry
1 mol O2  2 mol H2O
Stoichiometric Calculations

Calculate the mass of CO2 produced when 1.50 g of
butane (C4H10) is burned.
1.
Obtain balanced chemical equation for reaction.
2C4H10(l) + 13O2(g)  8CO2(g) + 10H2O(g)
Stoichiometric Calculations

Calculate the mass of CO2 produced when 1.50 g of butane (C4H10)
is burned.
2.
Use molar mass to determine the moles of butane consumed.
 1 mol C4 H10 
  0.0258 mol C4 H10
1.50 g C4 H10 
 58.1 g C4 H10 
Molar Mass C 4 H10  4(12.0 g/mol)  10(1.01 g/mol)
 58.1 g/mol
Stoichiometric Calculations

Calculate the mass of CO2 produced when 1.50 g of butane (C4H10)
is burned.
3.
Use stoichiometry from the balanced chemical equation to find the
moles of CO2.
2C4H10(l) + 13O2(g)  8CO2(g) + 10H2O(g)
2 mol C4H10  8 mol CO2
 8 mol CO2 
  0.1032 mol CO2
 2 mol C4 H10 
0.0258 mol C4 H10 
Stoichiometric Calculations

Calculate the mass of CO2 produced when 1.50 g of butane (C4H10)
is burned.
4.
Use molar mass to convert moles of CO2 to grams.


0.1032 mol CO2  44.0 g CO2   4.54 g CO2
 1 mol CO2 
Molar Mass CO 2  1(12.0 g/mol)  2(16.0 g/mol)
 44.0 g/mol
Stoichiometric Calculations
Summary of procedure
AB
Stoichiometric Calculations

The decomposition of KClO3 is commonly used to prepare small
amounts of O2 in the laboratory. How many grams of O2 can be
prepared from 4.50 g of KClO3?
2KClO3(s)  2KCl(s) + 3O2(g)
Stoichiometric Calculations

Propane, C3H8, is commonly used for cooking and home heating.


What mass of O2 is consumed when 5.00 g of propane is burned?
What mass of CO2 is produced when 5.00 g of propane is burned?
Stoichiometric Calculations

The complete combustion of octane (C8H18), a component of
gasoline, proceeds as follows:
2C8H18(l) + 25O2(g)  16CO2(g) + 18 H2O(g)

How many grams of CO2 are produced when 1.00 gallon of octane
(density = 0.692 g/mL) is burned?
1 gallon =3.7854 L

How many gallons of octane must be burned to produce 1.00 gallons of
water (density = 1.00 g/mL)?