Courtesy: www.lab-initio.com
•400 B.C. – Matter was composed of four fundamental
substances: fire, earth, water, and air.
•Greeks considered whether matter was
continuous (infinitely divisible into smaller
pieces) or composed of small indivisible
particles.
•Democritus (460 – 370 B.C.) used the term
atomos (which later became atoms) to describe
the small indivisible particles.
•No experiments to test theories so no
conclusion reached about divisibility of matter.
•Next 2000 years were dominated by a pseudoscience called
alchemy.
•Some alchemists were fakes who were obsessed with
turning base metals into gold.
•Others were serious scientists who made several advances
such as the discovery of several elements.
•The first “chemist” to perform true
quantitative experiments was Robert
Boyle (1627-1691).
•He measured the relationship
between pressure and volume of air.
•Published The Skeptical Chymist in
1661.
•German chemist Georg Stahl (1660-1734)
suggested a substance burning in a closed
container stopped when the air became
saturated with “phlogiston” ( a substance
that flowed out of a burning material).
•Joseph Priestley (1733-1804) discovered
oxygen gas.
•Oxygen gas, he found, supported
combustion and was supposed to be low in
phlogiston. Oxygen was originally called
“dephlogisticated air”.
•Antoine Lavoisier (1743-1794), a French
chemist, explained the true nature of
combustion.
•He also verified the law of conservation of
mass.
•Law of conservation of mass – mass is
neither created nor destroyed.
•Lavoisier was also a tax collector and was executed on the
guillotine in 1794.
•After 1800, chemists used experimentation to
study chemical reactions and determine
composition of chemical compounds.
•Joseph Proust (1754-1826), a French chemist,
showed that a given compound always contains
exactly the same proportion of elements by mass.
•This principle is known as the law of definite
proportions.
•For example, oxygen makes up 8/9 of the mass of
any sample of water, while hydrogen makes up the
remaining 1/9 of the mass.
•John Dalton (1766-1844), an
English schoolteacher, proposed
that atoms are the particles that
compose elements.
•Dalton discovered a principle
that became known as the law
of multiple proportions.
•Law of multiple proportions:
when two elements form a series
of compounds, the ratios of the
masses of the second element
that combine with 1 gram of the
first element can always be
reduced to small whole numbers.
Example: The Law of Multiple Proportions
The following data was collected for two compounds of
hydrogen and oxygen: water (H2O) and hydrogen peroxide,
H2O2.
Mass of Oxygen That Combines with
1 gram of Hydrogen
__________________________________________________
Water, H2O
7.92 g
Hydrogen Peroxide, H2O2
15.84 g
For the law of multiple proportions to be true, the ratios of
the masses of oxygen combining with 1 gram of hydrogen
should be a small whole number:
15.84 g
 2.00
7.92 g
This is a “small whole number” that supports the
law of multiple proportions.
•In 1808 Dalton published A New System of Chemical
Philosophy, in which he presented his theory of atoms:
How does this relate to chemistry?
1. Each element is made up
of tiny particles called
atoms.
a. There are 6.02 x 1023 atoms in
55.85 g of iron.
2. The atoms of a given
element are identical;
the atoms of different
elements are different in
some fundamental way
or ways.
b. Although graphite and diamond
have different properties, they
are both composed of carbon.
The carbon atoms are identical.
3. Chemical compounds are
formed when atoms
combine with each other. A
given compound always has
the same relative numbers
and types of atoms.
c. C + O2 → CO2; CO2 is not CO,
CO3 or Fe2O3.
4. Chemical reactions involve
reorganization of the atoms
– changes in the way they
are bound together. The
atoms themselves are not
changed in a chemical
reaction.
d. 2H2 + O2 → 2H2O, not CS2 or
NaCl.
•Dalton also prepared the first table of atomic masses.
•Many of the masses were proved to be wrong but the
construction of the table was an important step.
•Joseph Gay-Lussac (1778-1850) and Amadeo Avogadro
(1776-1856) provided the keys to the determination of
absolute formulas for compounds.
Gay-Lussac
Avogadro
•Based on the work of Dalton, Gay-Lussac, and Avogadro,
the concept of the atom was becoming generally accepted.
•The next questions to be addressed were:
•What is an atom made of?
•How do the atoms of the various elements differ?
•The first important experiments on the
composition of the atom were conducted by
J. J. Thomson (1856-1940) from 1898 to 1903.
•J. J. Thomson (1898-1903)
•Postulated the existence of electrons using cathode-ray
tubes.
•Determined the charge-to-mass ratio of an electron.
•The atom must also contain positive particles that
balance the exact negative charge carried by particles we
now know as electrons.
•The cathode ray was deflected by the negative pole of an
applied electrical field, Thomson postulated the ray was a
stream of negatively charged particles, now called electrons.
•Thomson knew atoms were electrically neutral so he
assumed atoms contain some positive charge to balance
the negative charge.
•Developed the “plum pudding” model.
Robert Millikan (1909)
•Performed experiments involving charged oil drops.
•Determined the magnitude of the charge on a single
electron.
•Calculated the mass of the electron.
Millikan put a charge on a drop of oil and measured how strong an applied electrical
field had to be in order to stop the drop from falling. Determined the mass and the
force of gravity on the drop and from this calculated the charge.
•Ernest Rutherford (1911)
•Explained the nuclear atom.
•Atom has a dense center of the positive charge called the
nucleus.
•Performed “gold foil experiment”.
•If Thomson’s “plum pudding” model were correct all alpha particles
would pass through.
•Rutherford found some particles were deflected.
•Postulated must have encountered some positive charge in center of
atom.
•Called this center of positive charge the “nucleus”.