Week 3/Tu: Lecture Units ‘5 & 6’
Unit 4: Molecules and Compounds
-- nomenclature
-- compounds and formulas
-- polyatomic ions
Unit 5: The Mole
-- definition for counting, molar mass
-- stoichiometry, elemental composition
Unit 6: Stoichiometry for Cpds.
-- Balancing Reactions
-- mass and reactions
Issues: ?
http://www2.chemistry.msu.edu/faculty/morrissey/
© DJMorrissey, 2o12
Week 3/Tu: Counting atoms
We know that atoms have mass and that one isotope (Z,N) is
identical to all other with same (Z,N), e.g., gold is all 197Au
79 protons, 118 neutrons
Modern techniques allow “visualization” of individual
molecules and large atoms like gold.
Can we count them? 1, 2, 3, … 12 (dozen), … 20 (score),
144 (gross), … 500 (ream of paper) …
… 6.0221367x1023 (mole)
How long would it take the fastest computer to count a mole?
Amount of time = 6.022x1023 / CPU-speed
= 6.022x1023 / 16.3 petaFLOPS
= 6.022x1023 / 16.3 x1015 operations/second
= 3.69x107 s à 1.17 years
© DJMorrissey, 2o12
Week 3/Tu: Mass vs. Count
It’s not practical to “count out” atoms in the lab. So we
introduce a definition that connects the mass of an object with
the number of particles using a convenient isotope and weight.
1 mole (Avagadro’s Number) of 12C is defined to be 12 grams
All other masses are then measured relative to that of 12C.
Notice that I was specific about the isotope?
Carbon has two stable isotopes in nature (12C & 13C). Natural
carbon is a mixture of the two so that the number given in the
periodic table reflects the average mass of the isotopes we
have here on earth.
© DJMorrissey, 2o12
Week 3/Tu: Mole à Molecular Masses
If we know the formula for a chemical compound then it is
easy to find the molecular mass of that compound using the
periodic table.
For example, last week we found the names for:
CO Carbon Monoxide 12.01 + 16.00 = 28.01 g/mol
CO2 Carbon Dioxide
12.01 + 2*16.00 = 44.01 g/mol
CCl4 Carbon tetrachloride 12.01 + 4*35.45 = 153.81 g/mol
CsF Cesium Fluoride
132.91 + 19.00 = 151.91 g/mol
Calcium hypochlorite
40.08 + 2*(35.45+16.00) =
(swimming pool chemical) 40.08 + 2*(51.45) =
40.08 + 102.90 = 142.98 g/mol
Ca( ClO)2
© DJMorrissey, 2o12
Week 3/Tu: Molecular Masses à Moles
If we know the formula for a chemical compound then it is
easy to find the number of moles of that compound using the
periodic table.
Other examples:
64 g of methane CH4 12.01 + 4*1.008 = 16.04 g/mol
64 g / 16.08 g/mol = 3.990 moles
= 4.0 moles
10 lbs of
Calcium hypochlorite 10 lb * 453.5924 g / lb = 4535.92 g
Ca( ClO)2
4535.92 g / 142.98 g/mol = 31.724 mol
= 32 moles
© DJMorrissey, 2o12
Week 3/Tu: Elemental Composition -11)  An oxide of copper contains 88.82% copper by mass.
What is the formula (and systematic name) of this
compound?
copper
oxygen
amount
88.82
100 ?– 88.82 = 11.18
Molar
Mass
Moles
Mole
Ratio
63.546 g/mol
16.00 g/mol
=88.82 / 63.546
= 1.398
=11.18 / 16.00
= 0.6988 moles
= 1.398 / 0.6988
= 2.000
© DJMorrissey, 2o12
= 0.6988/0.6988
=1
Cu2O Copper (I) Oxide
Week 3/Tu: Elemental Composition -22) A compound of (only) boron and hydrogen contains
78.14% boron by mass. What is the empirical formula of this
compound? If the molar mass is 27.7 g/mol, what is its
molecular formula?
boron
hydrogen
Amount
78.14
100-78.14 = 21.86
Molar
10.81
1.008 g/mol
Mass
Moles
=78.14/10.81
=21.86/1.008
= 7.23
= 21.69
Mole
Ratio
=7.23/7.23 = 1
= 21.69/7.23 = 3.00
Formula BH3 , Formula Mass = 13.83 g/mol .. Molecule B2H6
© DJMorrissey, 2o12
Week 3/Tu: Elemental Composition -33) What is the % by mass of nitrogen in ammonium nitrate?
This material is used as fertilizer but it also decomposes into
nitrogen, oxygen and water.
Ammonium NH4+ & Nitrate NO3-
à NH4NO3
Molar Mass =[14.01 + (4* 1.008) +14.01 + (3 * 16.00)] g/mol
= 80.05 g/mol
Percent Nitrogen = 100 * ( 28.02 / 80.05 ) = 35.00 %
© DJMorrissey, 2o12
Week 3/Tu: Elemental Composition -44a) What is the average mass of a single molecule of ethanol?
formulas: CH3CH2OH or C2H5OH or C2H6O
molar mass = 2* 12.01 + 6* 1.008 + 16.00 = 46.07 g/mol
molecule mass = molar mass / NA
= 46.07 g/mol / 6.022 x 1023 /mol
= 7.650 x 10-23 g on average
4b) Do all individual molecules of ethanol have exactly the
same mass?
DEMO:
C12H22O11 (s) + excess H2SO4 (l) →12 C + 11 H2O (l) + excess H2SO4 (aq)
Sugar Packet
© DJMorrissey, 2o12
Week 3/Tu: Isotopic Fraction, Mass
Question that requires thinking:
Given that natural carbon has a molar mass of 12.01 g/mol
and has only two stable isotopes (N= 6 & 7). What are (a)
the mass number and (b) the fraction of the heavier isotope?
a)  A (atomic mass number) = Z+N = 6 + 7 =13
b)  Two isotopes, need their masses …
13C = 13.0033 g/mol [in a table], 12C = 12 g/mol [definition]
Fraction = X
(1 – X)
Average mass = 12.01 =
12.01 =
12.01 =
0.01 =
© DJMorrissey, 2o12
X* 13.0033 + (1-X) * 12
13.0033X + 12 – 12X
1.0033X + 12
1.0033 X à X = 0.01 or 1%