Stoichiometry
Stoichiometry is the
basic chemical calculation that states quantitative relation of chemical
formulas and chemical equations. In chemistry, stoichiometry (sometimes called
stoichiometry of reaction to distinguish it from compositional stoichiometry)
is a science that studies and quantifies quantitative relationships of
reactants and products in chemical reactions (chemical equations). This word
comes from the Greek stoikheion (element) and metriā (size). Stoichiometry is
based on the basic laws of chemistry, namely the law of conservation of mass,
the law of fixed comparison, and the law of multiple comparisons
Early
Stoichiometric Stage
In
early chemistry, the quantitative aspect of chemical change, ie stoichiometry
of chemical reactions, did not receive much attention. Even when attention has
been given, experimental techniques and tools do not produce the correct
results.
One
example involves the theory of flogstones. Flogistonis tried to explain the
phenomenon of burning with the term "combustible substances".
According to the flogitonists, combustion is the release of a substance can be
etrbakar (from a burning substance). This substance is then called
"flogiston". Based on this theory, they define combustion as a
flogiston release of a combustible substance. Changing the mass of wood when
burning fits well with this theory. However, changes in metal mass when calcined
do not match this theory. Nevertheless flogistonis accept that both processes
are essentially identical. Increasing the mass of calcined metals is a fact.
Flogistonis attempts to explain this anomaly by stating that flogiston has a
negative mass.
Philosopher
from Flanders Jan Baptista van Helmont (1579-1644) conducted a famous
"willow" experiment. He grows willow seeds after measuring the mass
of the flowerpot and the soil. Since there is no change in the masses of flower
pots and soil when the seeds grow, he assumes that the mass obtained only by
the water entering the ore. He concludes that "the root of all matter is
water". Based on the current view, the hypothesis and experiments are far
from perfect, but the theory is a good example of the growing demeanor of
quantitative chemical aspects. Helmont recognized the importance of
stoichiometry, and clearly preceded his time.
In
the late 18th century, German chemist Jeremias Benjamin Richter (1762-1807)
discovered the equivalent concept (in terms of modern chemical equivalent
chemistry) with a careful observation of the acid / base reaction, ie the
quantitative relationship between acid and base in the neutralization reaction.
Richter equivalents, or what are now called chemical equivalents, indicate a
certain amount of matter in the reaction. An equivalent in neutralization
relates to the relationship between a number of acids and a number of bases to
neutralize them. Proper knowledge of equivalents is essential to produce good
soap and gunpowder. So, this kind of knowledge is very important in practical
terms.
At
the same time Lavoisier establishes the law of conservation of mass, and
provides the basic concepts equivalent to its accurate and creative
experiments. Thus, the stoichiometry that handles quantitative aspects of
chemical reactions becomes a basic chemical methodology. All fundamental laws
of chemistry, from the law of conservation of mass, the law of comparison
remained until the laws of gas reaction were all grounded stoichiometry. These
fundamental laws form the basis of atomic theory, and are consistently
explained by atomic theory. However, it is interesting to note that, the
equivalent concept was used before the atomic theory was introduced.
A.
Basic Law of Chemistry
1.
Understanding Basic Law of Chemistry
The basic law of chemistry is the basic law
governing the mechanism of the occurrence of a chemical reaction involving
reactants and products.
2.
Basic Chemical Laws
The basic law of chemistry consists of 5 basic laws
such as mass conservation law, fixed comparison law, multiple comparison law,
volume comparison law and avogadro law.
ü The
Basic Law of Conservation of Mass
As the name suggests,
this law was discovered by a scientist named Antonie Lavoiser. In his law, he
states that "The total mass of matter after reaction and before reaction
is the same". The point of his statement is that "the sum of the mass
of a substance acting as a reactant is equal to the sum of the mass of a
substance acting as a product".
The point is → Reactant
Mass = Mass Product
ü Comparative
law remains = law proust
"The ratio of the
mass of the elements in each compound is fixed"
Example:
A.
In the compound NH3 ------ mass N: mass H = 1 Ar. N: 3 Ar. H
=
1 (14): 3 (1)
=
14: 3
B.
On SO3 compound ------ mass S: mass O = 1 Ar. S: 3 Ar. O
=
1 (32): 3 (16)
=
32: 48
=
2: 3
Advantages of Proust
Law:
If known mass of a
compound or mass of one element that make up the compound then the mass of
other elements can be known.
ü Law
of multiple comparisons = dalton's law
"If two elements
can form two or more compounds for the mass of one element equal to the number
then the ratio of the mass of the second element will be proportional to the
integer and the simple".
Example:
When the element of
Nitrogen with compounded oxygen can be formed,
NO
where the mass N: O = 14: 16 = 7: 8
NO2
where the mass N: O = 14: 32 = 7: 16
For the same mass of
Nitrogen the Oxygen mass ratio of the compound
NO: NO2 = 8: 16 = 1: 2
Besides that, another
example that is Nitrogen and Oxygen can form six kinds of compound. The ratio
of the weight of oxygen that reacts with one part of nitrogen is:
0.57:
1.14: 1.74: 2.28: 2.86: 3.42
1:
2: 3: 4: 5: 6
This comparison is an
easy & round number, so it is in accordance with Comparative law.
ü Law
of the Gay / Lussac Volume / Legal Comparison
At the same pressure
and temperature, the volume of the reacting gas and the gas volume of the
reaction product are simple and integer ratios. This is the sound of the law
triggered by our legendary scientist named Gay lussac. His theory was not
without evidence or without research. To prove his theory, he conducted a
simple experiment by reacting hydrogen gas with oxygen gas into a container,
then into the container was given an electric flower flow so that oxygen gas
and hydrogen gas can react. After the reaction is complete, water vapor is
generated as a product and residual H2 and O2 gas are not reacted. After that,
the resulting water vapor is directly separated from inside the container. The
experiments were conducted repeatedly at fixed temperature and pressure and the
measurements showed that the ratio of the volume of hydrogen and oxygen gas and
water vapor was always 2: 1: 2.
2 H2 + O2 → 2H2O
Comparison of
coefficient numbers = 2: 1: 2 (the coefficient numbers are integers and simple)
ü Avogadro's
Law
At the same temperature
and pressure, a gas having the same volume also has the same number of
particles, This is the statement of the oldest legendary scientist, Mr.
Avogadro. The point of his statement is that the amount of particles of a gas
is independent of the Mass Or Mr. possessed by the gas, while in the same
volume, temperature and pressure, the amount of particles of a gas will always
be the same. The analogy is this: 1 liter of nitrogen gas and 1 liter of
chlorine gas have the same number of particles while under the same pressure
and temperature. So when a gas is in the same temperature, volume and pressure
the amount of particles from the gas will always be the same.
B.
The Concept of Molar Mol and Mass (})
In
SI systems, one mole is defined as the sum of the material composed of entities
(atoms, molecules, or other particles) a sum of the atoms in 12 grams of
carbon-12. The value of the number of atoms is 6.022 × 1023 called the Avogadro
number, NA.
Dalton
recognizes that it is important to determine the mass of each atom because its
mass varies for each type of atom. Atoms are so small that it is impossible to
determine the mass of one atom. So he focused on the relative mass values and
made the atomic mass table (figure 1.3) for the first time in human history. In
the table, the lightest element mass, hydrogen set one as standard (H = 1). The
atomic mass is a relative value, meaning a dimensionless ratio. Although some
atomic masses differ from modern values, most of the proposed values are in
the range of compatibility with current values. This shows that his ideas and
experiments are correct.
ü Mr. Dan Ar
Understanding Relative
Atomic Mass (Ar)
The relative atomic
mass is the average mass of an atom divided by the mass of the reference atom
C-12 atoms.
ü Relative
Molecular Mass (Mr)
The realmic molecular
mass is the average mass of the molecule divided by the mass of the reference
atom C-12 atoms.
Empirical Formulas And
Molecular Formulas
Molecular formula is a
formula that states the number and type of atoms that make up a compound, for
example water, water has the formula H2O molecule which means the water
compound is composed by two types of atoms H and O atoms and the number of each
atom consists of 2 H atoms And 1 O atom, so this is what is meant by the
molecular formula.
While the empirical
formula is not much different from the molecular formula, since both show the
number and type of atoms that make up a compound, the difference in this
empirical formula comparison of the number of atoms simplified as small as
possible, for example in benzene compounds, benzene has the formula C6H6
molecule, to make The empirical formula then each number of atoms must be
divided by 6 for the ratio of the number of atoms to be more simple, so the
empirical formula of the compound C6H6 is CH.
ü Hydrate
Compounds
Hydrate compounds are
one of the chemical compounds that molecules bind to water molecules.
Analoginya like this, the chemical compound is like a cotton that can absorb
water / moisture around the cotton, the more water is absorbed, the more water
molecules contained in the cotton
The concept of Mol
Mol is the number of
substances of an element that contains some form of elements such as atoms,
molecules, ions or electrons. Here is the relationship of mole with several
categories including:
• The Relation of Moles
With Particles
• The connection of Mol
With Mass
• Relation of Mol With
Volume
• The Relation of Moles
With Chemical Reactions
Problem
No. 1
Determine
the number of moles contained in:
A)
96 grams of oxygen (O2)
B)
88 grams of carbon dioxide (CO2)
(Mr.
O2 = 32; Mr. CO2 = 44)
Discussion
Determine
the known mol mass of the substance.
A)
96 grams of oxygen (O2)
M
= 96 grams
Mr
= 32
N
= ....
N
= m / mr
N
= 96/32 = 3 mol
B)
88 grams of carbon dioxide (CO2)
M
= 84 grams
Mr
= 44
N
= ....
N
= m / mr
N
= 88/44 = 2 mol
Problem
No. 2
Determine
the amount of mass contained in:
A)
0.2 mol of oxygen (O2)
B)
0.04 moles of carbon dioxide (CO2)
(Mr.
O2 = 32; Mr. CO2 = 44)
Discussion
Determining
the mass
A)
0.2 mol of oxygen (O2)
N
= 0.2 mol
Mr.
O2 = 32
M
= ....
M
= n × Mr
M
= 0.2 × 32
M
= 6.4 grams
B)
0.04 moles of carbon dioxide (CO2)
N
= 0.04 mol
Mr.
CO2 = 44
M
= ....
M
= n × Mr
M
= 0.04 × 44
M
= 1.76 grams
Problem
No. 3
Determine
the number of particles contained in
A)
0.1 mol of oxygen (O2)
B)
0.02 mol of carbon dioxide (CO2)
Discussion
Determine
the number of particles
A)
0.1 mol of oxygen (O2)
N
= 0.1 mol
X
= ......
X
= n × 6.02 × 1023
X
= 0.1 × 6.02 × 1023
X
= 6.02 × 1022 particles
B)
0.02 mol of carbon dioxide (CO2)
N
= 0.02 mol
X
= ......
X
= n × 6.02 × 1023
X
= 0.02 × 6.02 × 1023
X
= 1.204 × 1022 particles
