Download atomic structure discoveries/experiments conclusions

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3 ESO-D. IES AMES 06-07
EUROPEAN SECTION- CHEMISTRY & PHYSICS
ATOMIC STRUCTURE
DISCOVERIES/EXPERIMENTS
CONCLUSIONS
Early Chemical Discoveries:
Law of Conservation of Mass (1744, Antoine Lavoisier): “The total mass of substances present
after a chemical reaction is the same as the total mass of substances before the reaction”
Law of Constant Composition (1799, Joseph Proust): “All samples of a compound have the
same composition-the same proportions by mass of the constituent elements”
Dalton’s Atomic Theory (1803, John Dalton)
Dalton proposed an "atomic theory" with spherical solid atoms based
upon measurable properties of mass. Dalton's theory was based on the
premise that the atoms of different elements could be distinguished by
differences in their weights.
1. Each chemical element is composed of minute, indestructible
particles called atoms. Atoms can neither be created nor
destroyed during a chemical reaction.
2. All atoms of an element are alike in mass and other properties,
but the atoms of one element are different from those of all
other elements.
3. In each of their compounds, different elements combine in a simple numerical ratio.
Electrons and other discoveries:
Electric charges: static electricity
Electrolysis: Faraday's work on the chemical reaction produced when an electric current passes
through a liquid resulted in the laws of electrolysis.
The discovery of Electrons: Cathode ray tube (Thomson, 1897)
On April 30, 1897, Joseph John Thomson announced that cathode rays were negatively
charged particles which he called 'corpuscles.' He also announced that they had a mass about
1000 times smaller than a hydrogen atom, and he claimed that these corpuscles were the things
from which atoms were built up.
The “Plum pudding” model (1904, J.J. Thomson)
In this model, the atom is composed of electrons (which Thomson still called
"corpuscles," though Stoney had proposed that atoms of electricity be called
electrons in 1894), surrounded by a soup of positive charge to balance the
electron's negative charge, like plums surrounded by pudding.
http://www.sciencemuseum.org.uk/on-line/electron/section2/shockwave2.asp
rd
3 ESO-D. IES AMES 06-07
EUROPEAN SECTION- CHEMISTRY & PHYSICS
X-rays and Radioactivity
9 Wilhelm Roetgen (1895) noticed that when cathode ray tubes were operating, certain
materials outside the tubes would glow or fluoresce. Because of the unknown nature of
this radiation, Roentgen coined the term X-ray. X-ray are a form of high-energy
electromagnetic radiation.
9 By the early 1900s Henry Becquerel, Marie and Pierre Curie discovered radioactivity.
9 Ernest Rutherford identified three types of radiation from radioactive materials:
Rutherford`s atomic model: The nuclear atom (1909, E. Rutherford)
1. Most of the mass and all of the positive charge of an atom are
centred in a very small region called the nucleus. The atom is
mostly empty space.
2. There exist as many electrons outside the nucleus as there are
units of positive charge on the nucleus. The atom as a whole is
electrically neutral.
3. The protons and neutrons in the nucleus take up very little volume
but contain essentially all the atom’s mass. A number of electrons
equal to the number of protons move around the nucleus and
account for most of the atom’s volume.
Rutherford's gold foil experiment
The Rutherford scattering experiment. (a) When a beam of alpha particles is directed at a thin
gold foil, most particles pass through the foil undeflected, but a small number are deflected at
large angles and a few bounce back toward the particle source. (b) A closeup view shows how
most of an atom is empty space and only the alpha particles that strike a nucleus are deflected.
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/ruther14.swf
Protons and neutrons
Rutherford discovered the protons in 1919, and James Chadwick discovered the neutrons in
1932.
ATOMIC NUMBER (Z)
The number of protons in the nucleus of an atom determines the atomic number (Z) and
indicates the element's identity. For a neutral atom, the atomic number also describes the
number of electrons around the nucleus.
MASS NUMBER (A)
The mass number is the number of neutrons added to the number of protons. A= Z + N
ISOTOPES
Atoms of the same element can have different numbers of neutrons; the different possible
versions of each element are called isotopes.
ATOMIC MASS
The atomic mass of a chemical element is the mass of an atom at rest, most often expressed in
unified atomic mass units. The atomic mass is often synonymous with relative atomic mass
(RAM), average atomic mass and atomic weight; however, it is subtly different in that it can
either be the abundance-weighted average of isotope masses of an element or the mass of a
single isotope.
rd
3 ESO-D. IES AMES 06-07
EUROPEAN SECTION- CHEMISTRY & PHYSICS
Quantum theory (Max Planck, 1900)
Electromagnetic spectrum
Black body radiation
Niels Bohr (1913) applies quantum theory to Rutherford's atomic structure by
assuming that electrons travel in stationary orbits defined by their angular momentum. This
led to the calculation of possible energy levels for these orbits and the postulation that the
emission of light occurs when an electron moves into a lower energy orbit.
Photoelectric effect (Einstein, 1905)
Uncertainty Principle- Heisemberg(1927) Described atoms by means of formula connected to
the frequencies of spectral lines. Proposed Principle of Uncertainty- you can not know both the
position and velocity of a particle.
Wave Mechanics- Schrödinger(1930) viewed electrons as continuous clouds and introduced
"wave mechanics" as a mathematical model of the atom.
Particle/Wave duality- De Broglie(1923) Discovered that electrons had a dual nature-similar to
both particles and waves. Particle/wave duality. Supported Einstein.
Quantum Mechanical Model for an atom
According to the Principles of Quantum Mechanics electrons are distributed around the
nucleus in "probability regions". These probability regions are called "atomic orbitals".