Transcript The Mole

FORMULA
MATH
&
THE MOLE
The Mole
Measuring Matter
In Chemistry, we commonly
measure:
Mass
(grams)
Volume (L or mL)
Particles (counted)
We
can relate each of
these measurements to a
single quantity called the
“Mole”.
The
Mole is the SI unit
for the
“amount of something”
Why use it?
To
estimate the number of
particles that are too small
or too numerous to actually
count.
How
big is a Mole?
If you have a mole of
pennies and divide them
equally among the 6 billion
people on Earth, how many
dollars would each person
get?
6.022 X1023 pennies x
6 x 109 people
1$
100
pennies
= $1x1012
person
Each person would get
~ $ 1 trillion!!!
Volume
 1.0
mole = 22.4 Liters of a gas, at
standard temperature and pressure
(STP)
 Std. temp. = 0º C or 273K
 Std. pressure = 1 atmosphere (atm)
or 760 mm Hg (Torr)
Particles
1.0
23
10
mole = 6.022 x
particles of a pure substance.
* Element = atom
* Molecular = molecule (mlc)
*Ionic = formula unit (fmu)
Mass
1.0
mole = ____ grams of a
pure substance.
(“Molar Mass” of that
substance)
Molar Mass –
mass, in grams, of 1 mole
of a pure substance.
What is the Molar Mass of:
Carbon
tetrachloride
Comparing amounts
1
mole CCl4 =
_154.0_ grams
23
6.022 x 10 molecules
Liters is N/A (not a gas)
What is the Molar Mass of:
Oxygen
Comparing amounts
1
mole O2 =
_32.0_ grams
23
6.022 x 10 molecules
22.4 Liters, at STP
What is the Molar Mass of:
Aluminum
carbonate
Comparing amounts
1
mole Al2(CO3)3 =
_234.0_ grams
23
6.022 x 10 formula units
Liters is N/A (not a gas)
Calculating
with
Moles
Remember….
1.0 mole =
23
6.022 x 10 particles =
22.4 liters of gas, at STP =
______ grams (Molar Mass)
Steps to calculate:
READ directions in Lecture
Packet.
Example
What is the mass of 45.8 L of
carbon dioxide gas, at STP?
Example
How many moles are equal to
3.01 x 1022 atoms of magnesium?
Example:
How many molecules are in
38.69 grams of chlorine gas?
Example:
What is the volume, in L, at STP,
occupied by 0.600 moles of
sulfur dioxide gas?
STOP
Percent
Composition
Think about it...
What is the percent
composition of the
number of boys to girls
in this classroom?
Now, think about this...
What is the percent
composition of the
weight of boys to girls
in this classroom?
In Chemistry, it’s calculated
the same way, except we
consider the make-up of
compounds.
What information does %
Composition give us?
The percent, by mass,
of each element in a
compound.
Formula:
grams element
------------------------ X 100 = %
grams compound
2 Ways To Calculate:
1) Chemical Formula is given
2) Masses are given
(no formula given)
If formula is given:
1)
2)
3)
4)
5)
Write chemical formula.
Count atoms of each element in
compound.
Multiply by Mass Number from
P.T.(round all mass # to the 100th).
Add answers.
Plug numbers into formula and
solve.
If actual masses are given:
1) Use Law of Conservation of
Mass to determine elements and
compound mass.
2) Use formula to calculate %.
Example #1
What is the percent
composition of
water?
Example#2
What is the %
composition of calcium
nitrate?
If masses are given:
g element
+ g element
g compound
*Remember the
Law Of Conservation Of Mass
1)
Example#3
What is the percent composition
of a compound if 27.07 grams
of calcium completely reacts
with 47.93 grams of chlorine?
Example#4
Chlorine reacts with 33.86 grams
of phosphorus to form 150.00
grams of a compound. What is
the % composition of the
compound?
STOP
More
Formula
Math!!!
Empirical Formula


The simplest whole number
ratio of elements in a compound.
Subscripts cannot be reduced.
Molecular Formula
 The
actual number of atoms of
each element present in a
substance.
 Is a whole number multiple of the
Empirical Formula.
 Subscripts can be reduced.
From the “mole” perspective…
Ba3N2
3
moles of barium reacts with 2
moles of nitrogen to form
compound.
 Is an Empirical Formula because
subscripts cannot be reduced.
From the “mole” perspective…
Al2(SO4)3
 2 moles of aluminum reacts
with 3 moles of sulfate to form
compound.
 Is an Empirical Formula
because subscripts cannot be
reduced.
From the “mole” perspective…
C6H12O6
6
moles of carbon reacts with 12
moles of hydrogen and 6 moles of
oxygen to form compound.
 Is a Molecular Formula because
subscripts can be reduced.
From the “mole” perspective…
H2O2
2
moles of hydrogen reacts with 2
moles of oxygen to form
compound.
 Is a Molecular Formula because
subscripts can be reduced.
To calculate Empirical Formulas:
READ direction on
CHEAT SHEET!!!
example…
What is the E.F. of a compound
composed of 32.00% Carbon,
42.66% Oxygen, 18.67% Nitrogen,
and 6.67% Hydrogen?
1) Change % to grams
(just change the sign)
2) Convert from grams to moles
(use a 1 step box)
3) Make a “ratio”
(divide all answers by the smallest one)
4) Write formula
example…
What is the E. F. of a compound
composed of 25.9% Nitrogen and
74.1% Oxygen?
example…
What is the E. F. of a compound
composed of 36.6 grams of Carbon
and 9.2 grams of Hydrogen?
To calculate Molecular Formulas:
READ direction on
CHEAT SHEET!!!
example….
What is the M.F. of a compound
with a mass of 92.0 grams and an
E.F. of NO2?
example….
What is the M.F. of a compound
with a mass of 150.0 grams and an
E.F. of CH2O?
What is the M.F. of a compound
with a mass of 150.0 grams and an
E.F. of CH2O?
example….
What is the M.F. of a compound
with a mass of 132 grams that is
composed of 54.6% Carbon, 13.6%
Hydrogen, and 31.8% Nitrogen?