A few years after Boyle
came two French chemists who extended scientific understanding of the elements.
Antoine Lavoisier (1743-1794) affirmed the definition of an element as a simple
substance that could not be broken down into a simpler substance, and noted
that elements always react with one another in the same proportions.
Law of Conservation of Mass
The
first of these Fundamental Laws to be discovered was the Law of Conservation of
Mass.
The
total mass of material present after a chemical reaction is the same as before
the reaction.
This
Law was discovered by Antoine Lavoisier in about 1789. In a turn about of the
Scientific Method, Lavoisier had always assumed this Law was true, and ought out experiments which would verify his
assumptions. As a result of numerous combustion experiments conducted on
systems in closed containers, so as to retain any gases present, Lavoisier was
able to unambiguously verify his assumptions and formally state the Law of
Conservation of Mass. For an example, consider our combustion reactions of
elemental Carbon. If the mass of the gasses are accounted for, it is found:
Before
Rxn: 1.00g 2.66g
= 3.66g
After
Rxn: 0.00g 0.00g 3.66g = 3.66g
Before
Rxn: 1.00g 1.66g
= 2.66g
After
Rxn: 0.00g 0.00g 2.66g
= 2.66g
Of
course, these results require that each reactant be present in perfectly
balanced amounts, such that the full quantity of each is consumed completely
during the reaction. If this is not the case, some of the reagent in excess
will remain at the conclusion of the
reaction. However, the Law of Conservation of Mass will still apply.
Before
Rxn: 2.00g 10.00g
= 12.00g
After
Rxn: 0.74g 0.00g 11.26g = 12.00g
From
this example, we see a total of 12.00g of material is present both before and
after the chemical reaction occurs, with some of the hydrogen reagent remaining
as excess. Further, we can also note that oxygen is the Limiting Reagent in
carrying out this reaction; it limits the production of water. If more oxygen
were present, a greater amount of water would be produced.
Finally,
once this Law is accepted, it can be used to predict the amount of an
"unseen" reactant consumed or produced without direct measurement.
For instance, when iron burns in the air, its mass is seen to increase:
Before
Rxn: 5.00g ?g
After
Rxn: 0.00g 0.00g 7.15g
From
these results we can calculate the mass of oxygen needed to carry-out the
complete combustion of 5.00g of iron:
mass
Oxygen = 7.15g - 5.00g = 2.15g
Finally,
it must be noted the Law of Conservation of Mass, though a undamental Law of Chemistry, is not a fundamental
law of nature. When an energy difference occurs during a reaction, minute
amounts of mass are either gained or lost. Mass is either converted to energy
or energy is converted to mass. The energy-mass equivalence was first
postulated by Einstein in his famous formula; E = mc2. While these mass
differences are not detectable by the chemist, they are important in nuclear
reactions.
Law of Definite Proportions
A chemical compound, no matter what its
origin or its method of preparation, always has the same composition;
i.e., the same proportions by mass of constituent elements.
This
Law, sometimes known as the Law of Definite Composition, was first enunciated
by Joseph Proust in 1799. Proust discovered this law while analyzing samples of
Cupric Carbonate. He found two samples, one prepared via synthetic methods, and
the other mined naturally (Malachite Green), possessed the same composition of
elemental Carbon, Oxygen and Copper:
% Copper % Oxygen %
Carbon
Synthetic
Sample: 51.35% 9.74% 38.91%
Natural
Sample: 51.35% 9.74% 38.91%
So,
for example, if we decompose water by electrolysis and we recover the elemental
gases hydrogen and oxygen (not a difficult task experimentally), and
subsequently measure the masses of each gas respectively, we can determine the
composition of this compound:
Before
Rxn: 10.00g
After
Rxn:
1.12g 8.88g
This
data yields an elemental composition of:
%
Oxygen =
x 100 %
=
%
Hydrogen =
x 100 %
=
In
a similar manner, from the data presented above, we can determine the elemental
composition of the two Oxides of Carbon:
Carbonic Acid Carbonic Oxide
%
Carbon 27.29 % 42.88 %
%
Oxygen 72.71 % 57.12 %
Thus,
we begin to see how these two compounds of carbon and oxygen differ, they
differ in their relative proportions of the two constituent elements. Each
compound has a definite, well defined composition, but different compounds of
the same elements will have different compositions.
The
validity of the Law of Definite Proportions was firmly established by a number
experiments conducted by Jons Jakob Berzelius. To cite one example, he heated
powdered elemental lead with powdered elemental sufur, in different
proportions, to form the compound lead sulfide. In each case he found the
elemental composition of the lead sulfide remained unchanged.
Expt.
#1
Before
Rxn: 10.00g 1.56g
After
Rxn: 0.00g 0.00g 11.56g ==>
% Lead = 86.5 %
% Sulfur = 13.5 %
Expt.
#2
Before
Rxn: 10.00g 3.00g
After
Rxn: 0.00g 1.44g 11.56g ==>
% Lead = 86.5 %
% Sulfur = 13.5 %
Expt.
#3
Before
Rxn: 18.00g 1.56g
After
Rxn: 8.00g 0.00g 11.56g ==>
% Lead = 86.5 %
% Sulfur = 13.5 %
From
the results given above, we see attempts to increase the amount of one
elemental substance, without likewise increasing the amount of the other,
simply leads to an excess of that substance remaining after the reaction, and
not a change in the composition of the compound.
It
should be noted that forming a mixture, such as a solution, is a distinctly
different process than the process associated with compound formation. Suppose
copper and zinc are mixed to form brass. This is a physical process rather than
a chemical process. Brass is not a compound because its physical properties
(color, density, melting point, etc.) are not distinct; their exact values
depend on the proportions in which the copper and zinc are mixed. This is very
similar to mixing Table Salt and water to form a Salt-Water solution; the salt
and water can be mixed in different proportions. We would not think of
Salt-Water as a compound. When we combine, or "mix," hydrogen and
oxygen to form water, the result is a compound whose composition is fixed and
whose properties are distinct.
Once
accepted, this Law can be used to predict the amount of product which can be
formed from a given elemental reactant. For example, from the data cited above,
we know that Carbonic Acid is 27.29 % carbon and 72.71 % oxygen. So, how much
Carbonic Acid can be produced from 5.0g of carbon?
%
Carbon =
Or
Mass
Compound =
Further,
this result can be used to determine how much oxygen would be consumed in the reaction
forming this compound:
mass
Oxygen = 18.32g - 5.00g = 13.32g
Law
of Multiple Proportions
The
Law of Multiple Proportions was enunciated by John Dalton at about the same
time he postulated his Atomic Theory of Matter in ~1803. It was experimental
results in the form which suggested the validity of the Law of Multiple
Proportions which provided Dalton with the data needed to formulate the Atomic
Theory. This Law, therefore, is a central linchpin in the development of modern
chemistry.
If two elements form more than a single compound,
the masses of one element combined with a fixed mass of the second are in the
ratio of small whole numbers.
This
Law deals with the relationship between two compounds composed of the same
elements. Our Carbonic Acid - Carbonic Oxide example is a case in point. Both
are composed of the same two elements; carbon and oxygen. Recall, the above
data showed that 1.00g of carbon will combine with 2.66g of oxygen, in the case
of Carbonic Acid, and 1.33g of oxygen in the case of Carbonic Oxide. Thus, the
amount of carbon, in each case, is fixed at 1.00g. We can, in turn, use this
data to illustrate the application of the Law of Multiple Proportions:
As
another example, it is found elemental iron combines with elemental chlorine to
form two different compounds; ferric chloride and ferrous chloride. The
definite composition of these two compounds is:
Ferrous Chloride Ferric Chloride
%
Iron 44.06 % 34.43 %
%
Chlorine 55.94 % 65.57 %
This
data can be used to determine the mass of chlorine per 1.00g of iron needed to
produce these compounds. The results are:
Ferrous Chloride Ferric Chloride
mass
Iron 1.00g 1.00g
mass
Chlorine 1.27g 1.90g
Applying
the Law of Multiple Proportions to these results, we obtain:
In
the case of the Oxides of Carbon, Dalton would interpret the above results to
mean that carbonic acid is composed of 2 Atoms of oxygen for every 1 Atom of
carbon, and carbonic oxide contains 1 Atom of oxygen for every 1 Atom of
carbon. In the case of the Chlorides of Iron, he would conclude ferric chloride
contains 3 Atoms of chlorine for every Atom of iron and ferrous chloride
contains 2 Atoms of chlorine for 1 iron Atom.
It
must be kept in mind, these results are purely experimental in nature. And yet,
they have led directly to an Atomic interpretation for the formation of
compounds. The Law of Multiple Proportions, in conjunction with the other
Fundamental Laws of chemistry, led directly to the postulates of the Atomic
Theory of Matter.
Dalton's Atomic Theory of Matter
The
enunciation of the Law of Multiple Proportions and the Atomic Theory of Matter
by John Dalton occurred, historically, at the same point in time. The
experimental data led to both conclusions simultaneously. Dalton's Atomic
Theory consisted of five basic postulates:
Elements
are composed of indivisible Atoms.
Atoms
are alike for a given element.
Atoms
for different elements differ in size and mass and other properties.
Compounds
are formed from two or more Atoms of different elements.
Atoms
combine in simple numerical ratios to form these compounds. These ratios are different
for different compounds.
As
has been mentioned, the Oxides of Carbon are found to form from 2 Atoms versus
1 Atom of oxygen to every Atom of carbon in the compound. In Dalton's view,
these compounds could be represented as:
Carbonic Acid Carbonic Oxide
This
would clearly explain the results of the Law of Multiple Proportions. Further,
since the atoms combine without changing, losing or gaining parts, but instead
by merely recombining, this would explain the Law of Conservation of Mass. And,
since each "molecule" of Carbonic Acid contains 2 Atoms oxygen to 1
Atom carbon, the proportions of each are always the same; explaining the Law of
Definite Proportions. A similar conclusion can be drawn for the case of Carbonic
Oxide In modern parlance, we would name the compounds Carbonic Acid and Carbonic
Oxide, Carbon Dioxide and Carbon Monoxide, respectively. Finally, we would represent
these compounds with the chemical formulas CO2 and CO, respectively; a
convention introduced by Berzelius, and strongly opposed by Dalton.
The
compounds Ferric Chloride and Ferrous Chloride are represented by the chemical
formulas FeCl3 and FeCl2; where the number of iron atoms is denoted by the
symbol Fe (iron is ferrum in latin) and the number of chlorine
atoms is denoted by the symbol Cl.
Even
at this point in the historical development of chemistry, many questions
remained. For instance, the application of the Law of Multiple Proportions is
not unambiguous. How do we know the order in which to form the ratios? Also,
could we interpret the results for the Oxides of Carbon to mean Carbon Dioxide
is really 4 Atoms oxygen to 1 Atom carbon and Carbon Monoxide is 2 Atoms oxygen
to 1 Atom carbon? This interpretation is in fact consistent with the experimental
data and the Law of Multiple Proportions. Thus, the determination of exact chemical
formulas for compounds cannot be done on the basis of the Law of Multiple Proportions
alone. Much more data is required to perform this task. Further, we are now
left with the question of how the atoms bind together to form
"molecules?" What accounts for the bonding of these atoms?
Lastly,
not all of Dalton's postulates have withstood the test of time. Although his
theory is correct in its broad outlines, and is based directly on solid
experimental data, several of his postulates have to be modified in order to
conform with modern experimental results. We will now turn to an examination of
those results.
Charles-Gay-Lussac's
law
Charles-Gay-Lussac's Law tells us that at conditions
of constant pressure and constant amount of gas, the change of volume is proportional
to the change of temperature. In its simplest way this law can be expressed
with the formula V=constant * T, where V is the volume and T is the
thermodynamic temperature.
The above mentioned formula would mean that the volume
of a ideal gas tends to reach zero when the temperature approaches the amount
of 0 kelvins. This would mean that the density of gas must have an infinite
value at the temperature of 0 kelvins, which is impossible, and therefore it is
also not possible to reach absolute zero, but extremly low temperatures have
been successfully reached (even as low as 0.001 kelvins).
This law can be visualized with the help of this
graph. As we can see, volume is proportional to the temperature, which means
that by increasing temperature (at constant pressure and amount of gas) the
volume will also increase, and vice versa.
Law
of combining volumes
The law of combining volumes says that the volumes of
gases that react with each other or are formed by a chemical reaction, are in a
ratio of small whole numbers when the measurements are performed at constant
pressure and temperature.
For instance, one liter of oxygen reacts with exactly
two liters of hydrogen, giving two liters of water, i.e. water vapor. Similarly
1 liter of oxygen reacts with 2 liters of carbon monoxide, giving 2 liters of
carbon dioxide.
After the discovery of this law, some scientists have
proposed that there must be a simple relationship between the number of atoms
present in the same volumes of different gases at the same conditions.
Initially it was thought that all gases of equal
volume at the same conditions contain the same number of atoms. However this
has been proven wrong. More about this can be found in the tutorial concerning
Avogadro's law.
Experiment and theory in
Fundamental laws of chemical reactions and chemical
equations
Mark E. Tuckerman 2011-09-03
