4/9/2013
Stoichiometry
Chemistry 1010 Review Tutorial
Stoichiometry and Lewis Structures
April 9th, 2013
Stoichiometry
Types of Problems
Stoichiometry involves MOLES
Elements/compounds can only be compared
side by side using moles
1 mol = 6.022 × 1023 individual particles
(Avogadro’s #) – applies to anything
We cannot compare masses (or other
properties) as they differ for each element
or compound
Mass Percent and Percent Composition
Empirical and Molecular Formulas
Combustion Analysis
Dry Stoichiometry
Percent Yield
Limiting Reagent
Titration (solution stoichiometry)
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Mass Percent
Mass Percent Example
To find the mass percent of a particular
element in a compound:
Find the mass percent of each element in
C6H12O6
Find the molar mass of the compound
Find the TOTAL mass of the element in question
in the compound
Divide the mass of element by the mass of
compound and multiply by 100
Mass Element × 100% = Percent Composition
Mass Compound
Empirical Formula from %
Empirical Formula from %
To find the empirical formula given percent
composition:
Divide each number of moles by the
smallest number of moles to get a simple
whole number ratio
If you get a decimal, helpful to convert to a
fraction
Assume 100g of sample (MUST state this!)

With this assumption, percentages for each
element equate to masses in grams
Convert each mass to number of moles by
diving by molar mass

Then you can see what number you have to
multiply each fraction by
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Empirical Formula Example
Molecular Formulas
Butanoate is 58.8% carbon, 9.87%
hydrogen, and 31.3% oxygen by mass.
What is its empirical formula?
Normally, you will use the empirical
formula provided, or one determined in a
previous question
The molecular mass of the compound will
be provided
Find the molar mass of the empirical
formula: how many empirical formulas
make up the molecular formula?
Don’t forget – you can’t compare masses!
Must look at moles!
Molecular Formulas
Molecular Mass
= Whole Number Ratio
Empirical Mass
(Empirical) × whole # = Molecular
If you don’t find a whole number, CHECK
your work!
Molecular Formula Example
A compound has the empirical formula
CH2O. Its molar mass was determined to be
60.06 g/mol. What is its molecular formula?
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Combustion Analysis
You will be given the mass of an unknown
compound (generally containing only C, H,
and O) and told it is burned
The masses of H2O and CO2 (which are
always products of combustion) produced
will also be provided
Since all C and H from compound form CO2
and H2O, you can determine their amounts
Combustion Analysis
Mass Unknown = Mass C + Mass H + Mass O
Mass of O is easy to find! Must convert to
MOLES to find empirical formula.
Use moles of C, H, O to determine empirical
formula as before (divide by smallest)
Can be tangly, but remember – you’re just
finding the number of moles of each element
Combustion Analysis
To find the empirical formula of unknown:
Determine moles of CO2 and H2O produced

Moles = Mass/Molar Mass
Find moles of C and H alone

Molar ratios (eg. 1 mol C per mol CO2)
Turn moles of C and H into masses
To find O - Law of Conservation of Mass
Combustion Analysis Example
Vitamin C (ascorbic acid) is composed of
carbon, hydrogen, and oxygen. Combustion
of a 3.87 g sample yields 5.80 g of CO2 and
1.58 g of H2O

What is the empirical formula of Vitamin C?

If the molar mass of Vitamin C is 176.1256
g/mol, what is its molecular formula?
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Dry Stoichiometry
Dry Stoichiometry Example
For given a reaction, you will be provided
with the amount of one substance and asked
to find the amount of another – Reaction
MUST be balanced!
Convert given mass of substance into moles
Relate moles of what you want to moles of
what you have (mole ratio)
Convert moles of desired substance back to
the required unit (e.g. mass in grams)
Given the unbalanced reaction:
Percent Yield
Percent Yield Example
Normally given as a secondary part of a
longer question
You will be given the mass of product
actually obtained for a particular reaction
To find the % yield, divide this value by the
calculated (theoretical) mass and multiply
by 100%
If 2.12 g of sodium carbonate was produced
in the previous reaction, what would the
percent yield be?
NaHCO3(s)  Na2CO2(s) + CO2(s) + H2O(l)
What mass of sodium carbonate is produced
when 4.00 g of sodium bicarbonate is
decomposed?
NaHCO3(s)  Na2CO2(s) + CO2(s) + H2O(l)
Actual Mass × 100% = Percent Yield
Theoretical Mass
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Limiting Reagent
Limiting Reagent
Similar to a regular stoichiometry problem,
except you will be given amounts of two or
more reactants and asked to determine the
amount of product that would be produced
To begin, convert all reactant masses into
moles (n = m/MM)
Determine the amount of product that
would result if each reactant was used up
completely (multiply by mole ratio – comes
from BALANCED equation)

May be followed by a percent yield question
You must work with ALL reactants to
determine how much product each would
produce
# mol product
moles reactant × # mol reactant = moles Product
Mole ratio (from coefficients)
Limiting Reagent
Limiting reagent is the reactant that gives
the smallest amount of product
This smallest amount of product is the actual
amount of product produced (in moles)
 Use this value to determine the mass (or other
desired unit) of product produced

Amounts of excess reagents can be
determined by subtracting the moles used
(mole ratio again!) from initial moles of
each excess reactant
Limiting Reagent Example
Ammonia gas can be prepared as follows:
2NH4Cl(s) + CaO(s)  2NH3(g) + H2O(g) + CaCl2(s)
If 112 g of CaO reacts with 224 g of NH4Cl,
what mass of NH3 would be produced?
What mass of excess reagent will remain
after complete reaction?
If 38.7 g of NH3 is obtained, find the % yield
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Titrations
A known amount of a solution with known
concentration is added to a known amount
of a second solution, having unknown
concentration
Identical to other stoichiometry problems,
except we now determine moles using
concentrations and volumes instead of
masses and molar masses
Lewis Structures
Titration Example
A 25.00 mL aliquot of Ba(OH)2 solution is
titrated with standardized 0.3500 mol/L
HBr(aq). At the equivalence point, 15.39 mL
of HBr solution was added. What is the
concentration of the Ba(OH)2 solution?
Lewis Structures
Useful for showing valence electrons in
atoms, ions, and compounds
Use dots, arranged in pairs, to depict
valence electrons
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Molecular Lewis Structures
Molecular compounds contain two or more
non-metals, held together by covalent bonds

Electron sharing
Lewis structures are constructed a little
differently, but in a step-wise manner that is
the SAME for all covalent compounds and
ions!
Drawing Lewis Structures
1. Count total valence electrons for all atoms
Groups 1 and 2 have 1 and 2 valence
electrons respectively
For other main group elements (Grps 13-18),
use Group # - 10: Group 14 has 4 valence eAny ionic charge MUST be accounted for!


Drawing Lewis Structures
2. Determine the central atom – typically the
least electronegative atom
Exceptions:
• Hydrogen is NEVER the central atom
• Carbon is ALWAYS the central atom
For negatively charged ions, add a number of
electrons corresponding to the charge
For positively charged ions, subtract a number of
electrons corresponding to the charge
Drawing Lewis Structures
4. Subtract bonding pair electrons from total
valence electrons (2 e- per single bond)
5. Assign remaining electrons to terminal
atoms, fulfilling octet (8e- total each)
6. Assign any leftover electrons to the central
3. Attach remaining atoms to central atom
via single bonds (bonding pairs of
electrons)
atom (in pairs)
Don’t forget to indicate overall charge when
drawing ions!
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Drawing Lewis Structures
7. Calculate formal charge for each atom:
FC = (valence e-) – (unbonded e-) – (½ bonding e-)
[FC = (valence
e-)
– (unbonded
e-)
Lewis Structure Examples
Draw Lewis Structures for the following:
– (bonds)]
8. Reduce formal charges by making double
or triple bonds wherever possible
If you have a choice, take e- from the less EN
atom
Dichloromethane (CH2Cl2)
Nitrite ion (NO2-)
Chlorine difluoride cation (ClF2+)
9. Check for resonance structures (more than
one possible Lewis Structure)
Preparation for Final
Review old tests, assignments, and tutorial
notes. Do the online practice sheets!
If you’re having trouble with a particular
topic – get some FREE help with it

Chemistry 1010 Help Centre: C-2022

Chemistry Resource Room: C-2010


Mon-Thurs, 9am-4:30pm; Fri, 9am-3:30pm
Open Sat, April 13 and Sun, April 14: 10am-6pm
A Few Pointers…
There are many parts of a question, other than the
‘right answer’, that are worth marks
Show any/all formulas you use
Keep track of sig figs! It’s good to keep an extra digit
or two until the end, but be sure to indicate this!
Always give final answers with proper sig figs
Show all workings!!! (This includes molar mass)
Read and re-read the question – make sure you know
what exactly is expected/required
Show all units, and double check them
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Good Luck on Exams!!!
Have a great summer!
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