The Development of Atomic Theory. SCH12U February Mr. Dvorsky
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1 The Development of Atomic Theory SCH12U February Mr. Dvorsky
2 Nearly 2500 years ago Greek philosophers (i.e. Democritus) expressed a belief matter is composed of tiny indivisible particles called atoms (atomos is the Greek word for indivisible ) These conclusions were not based on any evidence; they were derived from philosophical reasoning.
3 Experimentation by many scientists during the 18th and 19th centuries led to the development of 2 Laws: 1. The Law of Conservation of Mass During chemical change no loss or gain of mass occurs. 2. The Law of Definite Proportions Compounds contain elements in fixed proportions by mass.
4 John Dalton, early 19 th century took these findings and developed Atomic Theory. 1. Matter consists of particles called atoms. 2. Atoms are indestructible. In chemical reactions atoms rearrange but are not broken apart. 3. Atoms in one particular element are identical, but differ from atoms of other elements. 4. Compounds are created when atoms of different elements combine in definite proportions.
5 JJ Thompson Discovery of Electrons Source of Electrical Potential Stream of negative particles (electrons) Metal Plate Gasfilled glass tube Metal plate Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 58
6 Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 58 A Cathode Ray Tube
7 Thomson s Experiment voltage source vacuum tube metal disks
8 Thomson s Experiment voltage source vacuum tube metal disks
9 Thomson s Experiment ON OFF voltage source Passing an electric current makes a beam appear to move from the negative to the positive end
10 Thomson s Experiment ON OFF voltage source
11 Thomson s Experiment ON OFF voltage source By adding an electric field he found that the moving pieces were negative.
12 The Effect of an Obstruction on Cathode Rays Dorin, Demmin, Gabel, Chemistry The Study of Matter, 3 rd Edition, 1990, page 117
13 The Effect of an Electric Field on Cathode Rays Dorin, Demmin, Gabel, Chemistry The Study of Matter, 3 rd Edition, 1990, page 118
14 J.J. Thomson J.J. Thomson He proved that atoms of any element can be made to emit tiny negative particles. From this he concluded that ALL atoms must contain these negative particles. He knew that atoms did not have a net negative charge and so there must be balancing the negative charge.
15 William Thomson (Lord Kelvin) In 1910 proposed the Plum Pudding model Negative electrons were embedded into a positively charged spherical cloud. Spherical cloud of Positive charge Electrons Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 56
16 Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 56 PlumPudding Model
17 Thomson Model of the Atom J.J. Thomson discovered the electron and knew that electrons could be emitted from matter (1897). William Thomson proposed that atoms consist of small, negative electrons embedded in a massive, positive sphere. The electrons were like currants in a plum pudding. This is called the plum pudding model of the atom. electrons
18 Rutherford, see animation fired alpha particles at a very thin piece of foil. The alpha particles to pass through without changing direction (very much) Because... The positive charges were spread out evenly. Alone they were not enough to stop the alpha particles
19 Because he thought the mass was evenly distributed in the atom.
20 Because, he thought the mass was evenly distributed in the atom
21 Explanation of AlphaScattering Results Alpha particles Nucleus Plumpudding atom Thomson s model Nuclear atom Rutherford s model
22 Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 323 The Rutherford Atom
23 Bohr suggested the planetary model of the atom could be rescued if one assumption is made: certain special states of motion of the electron corresponding to electron shells would not result in radiation and therefore the electron can exist indefinitely.
24 Bohr was saying, in effect, is that the atom can exist only in certain discrete energy states: the energy of the atom is quantized. Bohr noted that this quantization nicely explained the observed emission spectrum of the hydrogen atom. The electron is normally in its smallest allowed orbit, corresponding to n = 1; upon excitation in an electrical discharge or by ultraviolet light, the atom absorbs energy and the electron gets promoted to higher quantum levels.
25 These higher excited states of the atom are unstable, so after a very short time (around 10 9 sec) the electron falls into lower orbits and finally into the innermost one, which corresponds to the atom's ground state. The energy lost on each jump is given off as a photon, and the frequency of this light provides a direct experimental measurement of the difference in the energies of the two states, according to the PlanckEinstein relationship e = hν.
26
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