Chapter 3
Stoichiometry
http://www.shodor.org/UNChem/basic/stoic/
THINK

Why do the masses on the periodic table
have decimals if we can’t have part of a
proton or a neutron?
http://richardbowles.tripod.com/chemistry/structure/structure.htm
1961: Scientists decided to have a
standard of masses, they chose Carbon.
They decided it has a mass of 12 atomic
mass units (amu) and base all other
elements on this.
 Example: He is 1/3 the mass of Carbon.
What is Heliums atomic mass?

http://richardbowles.tripod.com/chemistry/structure/structure.htm
Calculating the Mass of a
Compound
Compound
Table salt
Potassium
Bromide
Aluminum
Hydroxide
Magnesium
Iodide
Formula
Formula Mass
Calculating the Mass of a Molecule
Molecule
Water
Dinitrogen
Pentoxide
Simple Sugar
Sulfur Dioxide
Formula
MolecularMass
THINK

What is the difference between a
compound and a molecule?
A.M.U?!?!?!

Define:

Units:

Examples and Calculations:
http://en.wikipedia.org/wiki/Atomic_mass_unit
What is Stoichiometry?

Define:
http://www.chem4kids.com/files/react_stoichio.html
What is a mole?
Definition: (and citation)
 Value:
 Another name for this value:
 What does it mean? (in your own words!)

http://www.youtube.com/watch?v=1R7NiIum2TI
Conversions
1 mole= atomic mass of substance
1 mole= 6.022 X 1023 atoms
1 mole= 22.4L
How many grams, liters and atoms are
in the following?
1 mole of Ag=
1 mole of H2O=
1 mole of CO2
Mole-Gram Conversions

Use Dimensional Analysis to calculate the
answer to the following:
How many moles are in 11.5 grams of C2H5OH?
http://dbhs.wvusd.k12.ca.us/webdocs/Mole/Moles-to-Grams.html

How many moles of water are in
1.20X1025 atoms?

If the volume of Nitrogen gas is 75.0 L,
how many grams are present?

How many atoms are in 16.2 grams of N2?

How much volume of ammonia will be
present in 45.2 grams?
 If
there are 44.5L of Hydrogen gas, how many
grams are present?
Percent composition from Formula
Formula

The subscripts in a
formula represent
not only the atom
ration in which the
different elements
are combined, but
also the
______________
____________.
Atom
Ratio
Mole Ratio
H2O
KNO3
C12H22O11
http://www.chemcool.com/regents/molesstoichiometry/aim4.htm
Examples

Sodium hydrogen carbonate, commonly
called “bicarbonate of soda”, is used in
many commercial products to relieve an
upset stomach. It has the formula
NaHCO3. What are the mass percents of
Na, H, C, and O in sodium hydrogen
carbonate?
How to Solve

% of element = ____________________ x 100
Element
Number of
Moles
Atomic Mass
Compound
mass =
Molar Mass

Use these numbers to determine the
percent of each component:
Example 2

An iron containing mineral responsible
for the red color of soil in many parts of
the country is limonite, which has a
formula Fe2O3 • 3/2 H2O. What mass of
iron in grams can be obtained from a
metric ton (103 kg = 106 g) of limonite?
Simplest Formula from Chemical
Analysis

Simplest formula: Gives the simplest
_____________________
___________ of the atoms present.
http://www.carlton.srsd119.ca/chemical/Molemass/empirical_formula.htm
Example

A 25.0 g sample of an orange compound
contains 6.64 g of potassium, 8.84 g of
chromium, and 9.52 g of oxygen. Find the
simplest formula.
How To Solve




Change all grams to moles for comparisons.
Calculate the mole ratio.
Make them all whole numbers.
Because the mole ratio is the same as the atom
ratio, you are done!! Just write the final answer and
box it!
Example 2

When a sample of ethyl alcohol is burned
in air it is found that 5.00 grams of ethyl
alcohol convert to 9.55 grams of carbon
dioxide, and 5.87 grams of water. What
is the simplest formula of ethyl alcohol?
(Hint: Find the mass of each ELEMENT
first!! Then continue the problem)
Molecular formula from Simplest
Formula

The molecular formula is a
_______________________
________________ of the simplest
formula.
Example

Vitamin C’s formula is found to be
C3H4O3. From an experiment the
molecular mass is found to be 180 g/mol.
What is the molecular formula of vitamin
C?
How To Solve It


Find the simplest formula
Figure out what to multiply by determining the ration
of the simplest to the molecular masses.
Writing and Balancing Equations

word:

formula:
Balancing Equations

Ca
+

Mg
+

AgNO3
O2
N 2
+ Cu
CaO
Mg3N2
Cu(NO3)2
http://richardbowles.tripod.com/chemistry/balance.htm
+ Ag
Mass Relations from Equations

The coefficients of a balanced equation
represent the number of
_______________ of reactant and
products
Example

Ammonia is used in fertilizer and is made by
reacting nitrogen of the air with hydrogen. (a)
How much ammonia (in grams) is formed from
1.34 mol of nitrogen? (b) how much nitrogen
(in grams) is required to form 1.00 kg of
ammonia? (c) How much hydrogen (in grams)
is needed to react with 6.00 grams of nitrogen?
How To Solve It



Write a balanced equation for this reaction
Use the mole rations (coefficients of the reaction) to
relate moles of one substance to moles of another.
Change from moles to grams.
Limiting Reactant

Limiting reactant: The reactant that
_________________ the amount of
_____________.
http://www.chem.tamu.edu/class/majors/tutorialnotefiles/limiting.htm
Example
If 3.5 grams of copper is added to 6.0 grams of
AgNO3, which one is the limiting reactant?

Cu +
AgNO3  Cu(NO3)2 +
Ag
How To
Solve It
Take grams of reactant #1 and find the mass of the
product.
Take grams of reactant #2 and find the mass of the
product.
The reactant that produces less product is the limiting reactant
Example 2
Identify the limiting reactant when 1.7 grams of
sodium reacts with 2.6 L of chlorine to produce
salt.

Na + Cl2  NaCl
Experimental Yield; Percent Yield
Percent Yield:
Actual Yield:
Expected Yield:
http://www.800mainstreet.com/6/0006-007-percent-yield.html
Example
A piece of copper with a mass of 5.00 grams is
placed in a solution of silver nitrate. The silver
metal that is produced has a mass of 15.2 grams.
What is the percent yield?
[Hint: 5.00 g Cu -> ? g Ag (theoretical yield)]
Cu + AgNO3  Cu(NO3)2 + Ag
Example 2
Determine the percent yield for the reaction between
2.80 grams of Al(NO3)3 and excess NaOH if 0.966
grams of Al(OH)3 is recovered.
Al(NO3)3 + NaOH  NaNO3 + Al(OH)3