Download Chem 1A Lecture 1

4/15/2010
Chapter 3 Objectives
• Understand 4 general reaction types
• Be able to write and balance chemical equations
• Understand the concepts of formula weight and the mole
as a counting number for particles (atoms or molecules)
• Be able to use balanced chemical equations to covert
between particles/mole/mass of one reactant or product
to particles/mol/mass of another
• Understand the concepts of limiting reactant, theoretical
yield and percent yields and be able use balanced
equations and given information to calculate each.
Find the Limiting Reactant: Mass
2 Na + Cl2 g 2 NaCl
We have 30.0 g Na and 40.0 g of Cl2. What is the limiting
reactant? Start with mole ratios:
• 1 mol of Na makes 1 mol NaCl
– 30.0 g Na x 1 mol Na/22.99 g Na = 1.30 mol Na g 1.30 mol NaCl
• 1 mol of Cl2 makes 2 mol NaCl
– 40.0 g Cl2 x 1 mol Cl2/70.90 g Cl2 = 0.564 mol Cl2 g 1.13 mol NaCl
• Cl2 is the limiting reactant because it yields the smaller
amount of NaCl
• The theoretical yield is 1.13 mol NaCl; it’s the max we can
make if we use up all of the Cl2
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Percent Yield
• The measured amount of product made in a chemical
reaction is called the experimental yield.
• We can determine the percent yield of a reaction or
process by comparing the amount we actually got to the
maximum or theoretical yield:
Experimental Yield
100% Percent Yield
Theoretical Yield
• The maximum or theoretical yield is calculated based on
the limiting reactant.
• Because of both controllable and uncontrollable factors,
the experimental yield of product is always be less than the
theoretical yield (unless something has gone wrong….)
Percent Composition
•
•
Percentage of each element in a compound by mass
Can be determined from either:
o
o
•
The formula of the compound.
The experimental mass analysis of the compound.
The percentages may not always total to 100% due to
rounding.
Percentage
mass of element X in 1 mol
100%
mass of 1 mol of the compound
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Combustion Analysis
• Compounds containing C, H and O are routinely
analyzed through combustion in a chamber like this.
 C is determined from the mass of CO2 produced.
 H is determined from the mass of H2O produced.
 O is determined by difference after the C and H have been
determined.
Elemental Analyses
Compounds
containing other
elements are
analyzed using
methods analogous
to those used for C, H
and O.
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Mass Percent as a
Conversion Factor
• The mass percent tells you the mass of a constituent
element in 100 g of the compound.
o The fact that NaCl is 39% Na by mass means that
100 g of NaCl contains 39 g Na.
• This can be used as a conversion factor.
o 100 g NaCl g 39 g Na
g NaCl
39 g Na
g Na
100 g NaCl
g Na
100 g NaCl
g NaCl
39 g Na
Empirical Formulas
• Empirical means determined by experiment.
o Finding the ratios of the weights of one element to another in
a compound is the first step in figuring out the formula of an
unknown substance
• The empirical formula of a chemical compound is a simple
expression of the relative numbers of each type of atom in it
• The empirical formula only tells you the simplest, wholenumber ratio of atoms in a molecule.
o Can be determined from percent composition or
combining masses.
• More than one molecule can have the same empirical
formula
• The molecular formula is a multiple of the empirical formula.
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Empirical Formulas
Hydrogen Peroxide
Molecular formula = H2O2
Empirical formula = HO
Benzene
Molecular formula = C6H6
Empirical formula = CH
Glucose, Fructose, Galactose
Molecular formula = C6H12O6
Empirical formula = CH2O
Determine the Empirical Formula of Benzopyrene, C20H12,
•
Find the greatest common factor (GCF) of the subscripts.
20 factors = (10 x 2), (5 x 4)
12 factors = (6 x 2), (4 x 3)
GCF = 4
•
Divide each subscript by the GCF to get the empirical
formula.
C20H12 = (C5H3)4
Empirical formula = C5H3
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Finding an Empirical Formula
from Experimental Data
gC
mol C
gH
mol H
gO
mol O
1.
2.
Mole ratio
formula
mole
ratio
whole
number
ratio
empirical
formula
Convert the percentages to grams, if necessary.
Convert grams to moles.
a.
Use molar mass of each element.
Write a mole ratio formula using moles as subscripts.
Divide all by smallest number of moles.
Multiply all mole ratios by a whole number to make all
whole numbers, if necessary.
3.
4.
5.
o
If ratio is 0.5, multiply all by 2; if ratio 0.33 or 0.67, multiply all by
3, etc.
Empirical to Molecular Formulas
• The molecular formula is a multiple of the empirical
formula.
• To determine the molecular formula, you need to know
the empirical formula and the molar mass of the
compound.
Molar massreal formula = Factor used to multiply subscripts
Molar massempirical formula
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For the Coming Week
• Review Chapters 1 - 3
• Work on text problems for both chapters
Mastering Chemistry Chap2 &3 due: Thursday 4/22/10 6 pm
• PreLab for 7 Up Lab due at beginning of Lab Lecture
• Tuesday: Nomenclature Worksheet Due
Finish 7-Up lab
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