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The Mole
When chemists work with elements in the laboratory, they are not dealing with single _____
____
or
_________
______
and
________
amounts contain huge numbers of
these very tiny particles.
Rather than thinking in terms of these individual small
particles, chemists use a larger quantity of matter called a
______
(symbol
_____
A
mole is just a
____________
of material and is the primary measuring unit of
_______________
One mole of material can be
_____
_________
_____________
(used for ionic
compounds) or
____
depending on the composition of the substance. However, for
ALL
materials, the number of particles in a mole is the same,
________________
This is called
___________________________
and is symbolized by
____
A
For example
A mole of sodium (Na) contains
______________________
of sodium
A mole of water (H
2
O) contains
_______________________
of water
A mole of ammonium ions (NH
4
) is
_______________________
of NH
4
A mole of potassium chloride contains
_______________________
of KCl
A mole of eggs is
_____________
eggs
It is very difficult to imagine the immensity of Avogadro’s Number. Perhaps these sizeanalogies will help:
Avogadro’s number of seconds is about 2 x 10
14
centuries, which is roughly a million times
the best current estimate of the age of the entire universe.
Avogadro’s number of chemistry textbooks would cover the entire surface of the Earthto a height of about 300 km.
If you won 1 mole of dollars in a lottery the day you were born, and spent a billion dollarsa second, you would still have more than 99.999% of the prize money left if you lived tobe 90 years old.
In order to obtain a mole of grains of sand it would be necessary to dig the entiresurface of the Sahara desert (8 million square kilometres) to a depth of 6 meters.
Population of the Earth – 6,000,000,000 people1 mole of people – 602,000,000,000,000,000,000,
Molar Mass
Chemical formulas and equations are expressed using amounts in moles, but in alaboratory, we measure mass. Thus, we constantly need to convert amount in moles intomass and vice versa.The mass of one mole of any element is numerically equivalent to the average atomicmass of the element. The mass of one mole of a substance is called
______ _______
(symbol
____
) and is expressed in grams/mole (
________
i)
Molar mass of Na
ii)
Mass of one mole of fluorine
When atoms combine to form compounds, mass is neither created nor destroyed, thus,the molar mass of a compound can be obtained by adding the atomic masses of eachatom in the compound.
Remember subscripts tell us how many of each atom are
present. iii)
Mass of one mole of Fe
2
O
3
iv)
Molar mass of Mg
3
(PO
4
2
The relationship between number of moles, mass of a substance and molar mass isrepresented by the following equation:
Example #1 Convert a mass of 3.6 grams of table salt to an amount in moles.
Example #
Determine the mass of a 6.5 molar solution of hydrochloric acid.
Mole Conversions
The mole is used to help us “count” atoms, molecules ions etc… The relationship betweenmoles, number of particles and the Avogadro constant is:N = n x N
A
Where:
N = number of particlesn = number of molesN
A
= Avogadro constant
Using the above equation and the flow chart listed below as a guide, you will be able toperform a number of calculations involving moles, atoms, particles and eventually mass. Example #1 A sample of benzene, C
6
H
6
, contains 5.69 mol.
a) How many molecules are present in this sample?b) How many hydrogen atoms are in the sample? Example #2 A sample of zinc oxide contains 3.28 x 10
24
formula units. How many moles of zinc oxide
are in the sample?
Empirical and Molecular Formula of a Compound
An
_______________
shows the
____________
in which atoms combine to form a
molecule. An empirical formula
____________
necessarily provide the correct
information about the
_______________
in a molecule. A
_________________
will
give the
____________
and
________________
bonded to form a molecule.
A comparison of empirical and molecular formula
Name of compound
Molecular
formula
Empirical
formula
Lowest ratio of
elements
Benzene Glucose Ammonia Hexene Octene It is possible for a molecule to have the
______
empirical formula and molecular
formula. For example,
_______________
in the above chart.
It is also possible for two compounds to have the
______
empirical formula but
____________
molecular formula. For example
_________
and
___________
You can determine the empirical formula for a compound if you are given experimentallyderived
_________
or the
________________
for the compound.
If you are given weights, you will use the
___________
to calculate
______
If you
are given percent composition, you assume a
__________________
and then
determine the number of
________
. In both cases you will need to know the
_________________
of your element.
Once moles have been determined, you find the
__________________________
(using the smallest mole value as your
_____________________
Example #1 A sample of a compound contains 2.40 g of carbon and 0.800 g of hydrogen.
What is its
empirical formula? Example #2 A compound was determined experimentally to be 29.0% sodium, 40.6% sulfur and30.4% oxygen mass.
What is its empirical formula?
If the
_______________
of the
_______________
is known, one can determine the
_______________
of the compound.
To accomplish this, first determine the
________
formula of the compound. Then compare the
______________
of the
substance with the
______________
derived from the
___________________
. To
determine how much greater your molecular formula is you can use the followingequation:
Example #3 A compound is 82.8% carbon and 17.2% hydrogen by mass. Its molar mass is 58.0g/mol.What is the molecular formula of this compound?
Hydrated Ionic Compounds
A
______
is an
_______
compound that has
___________
from
a water solution and has
_____
molecules incorporated into their
crystal structure. Hydrates have a
________________
of water
molecules
___________
bonded to each
_________________
Compounds that have no water molecules incorporated into themare called
_________________
to distinguish them from their
hydrated forms.For example, you can have different forms of sodium carbonate:
Note
: the raised dot in the chemical formula does not mean
multiplication, but rather a weak bond between an ionic compoundand one or more water molecules.The molar mass of a hydrated compound must include the mass ofany water molecules that are in the compound.What is the molar mass of Na
2
CO
3
10H
2
O?
Na
C
O
H
We can use the coefficients in a balanced chemical equation to determine a number of quantities:
i.
How many molecules react
ii.
How many molecules are produced
Example
:
___Al + ___Fe
2
O
3
→
___Al
2
O
3
According to the balanced equation, ___ atoms of Al react with ___ formula units of Fe
2
O
3
What if you had 8 atoms of Al?You would need ___ formula units of Fe
2
O
3
to react, and you would produce ___ formula units of aluminum
oxide and ___ atoms of iron.This is because in the above equation, aluminium and iron (III) oxide interact in a __________ ratio.Balance the following equation and state the ratio of reactants and products:
___ C
3
H
8
2
→
___ CO
2
2
O
How many molecules of oxygen are required to generate 100 water molecules? The coefficients in balanced equations actually represent moles and can be used to determine molar relationships or moleratios. Balance the following equation and state the mole ratio of reactants and products:
___ N
2
2
→
___ NH
3
How many moles of nitrogen and hydrogen would you need to generate 8.6 mol of ammonia?
In order for scientists to apply stoichiometry, chemists must always find the moles of reactants beforethey set-up their ratio.
Often it is then useful to convert amounts of products in moles to a more useful
measurement - mass. To solve stoichiometry problems involving mass:
i.
Begin with a balanced chemical equation and write the mass of reactant or product beneath thecorresponding formula
ii.
Determine the molar mass of the necessary reactants and products
iii.
Determine the number of moles from the given mass
iv.
Use the mole ratio to predict the amount in moles of desired substance
v.
Convert the predicted amount in moles into mass or molecules
Examples
:
Ammonium sulfate is used as a source of nitrogen in some fertilizers.
It reacts with sodium hydroxide to
produce sodium sulfate, water and ammonia.
What mass of sodium hydroxide is required to react
completely with 15.4 g of ammonium sulfate?Powdered zinc reacts rapidly with powdered sulfur in a highly exothermic reaction.
What mass of zinc
sulfide is expected when 32.0 g of sulfur reacts with sufficient zinc?
___ Zn (s) + ___ S
8
(s)
→
___ ZnS (s)
The Limiting Reactant
In a balanced chemical equation, the stoichiometric coefficientspredict the mole ratio for a reaction that has gone to completion.In practice, however, there are often reactants left over.The reactant that is completely used up in a chemical reaction iscalled the
limiting reactant
. The limiting reactant will determine
how much product is produced. When the limiting reactant is usedup, the reaction stops.A reactant that remains after a reaction is over is called the excess reactant
Example
Powdered zinc and sulfur react in an extremely rapid, exothermic reaction. If a 6.00 g sample of zinc isallowed to react with 3.35 g of sulfur, which is the limiting reactant?
Zn
(s)
+ S
(s)
ZnS
(s)
You can use your knowledge about finding limiting reactants to predict the amount ofproduct that is expected in a reaction. This type of prediction is a routine part of achemist’s job.
Example
Iron can be produced when iron (III) oxide reacts with carbon monoxidegas. If 11.5 g of Fe
2
O
3
reacts with 2.
×
24
molecules of CO, what mass
of Fe is expected?
Fe
2
O
3(s)
+ 3CO
(g)
2Fe
(s)
+ 3CO
2(g)
