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The History of Atomic Theory

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1 The History of Atomic Theory
Mr Nelson


3 Democritus 400 BC The Greek philosopher Democritus began the search for a description of matter more than years ago. He asked: Could matter be divided into smaller and smaller pieces forever, or was there a limit to the number of times a piece of matter could be divided?

4 Atomos This piece would be indivisible.
He named the smallest piece of matter “atomos,” meaning “not to be cut.”

5 Why? The eminent philosophers of the time, Aristotle and Plato, had a more respected, theory. Aristotle and Plato favored the earth, fire, air and water approach to the nature of matter.

6 Dalton’s Model In the early 1800s, English Chemist John Dalton performed experiments that verified the existence of atoms.

7 Dalton’s Model 1803 Dalton’s Model was that atoms are indivisible particles.

8 Dalton’s Theory All matter is composed of atoms.
Atoms of the same element are exactly alike. Atoms of different elements are different. Atoms are indestructible and cannot be divided.

9 J. J. Thomson In 1897, the English scientist J.J. Thomson proved the atom is made of even smaller particles.

10 Thomson Model Thomson studied the Cathode Ray Tube.
As the current passed through the gas, it gave off rays of negatively charged particles.

11 Thomson Model the atoms of the gas were uncharged.
Where did they come from? the atoms of the gas were uncharged. Where had the negative charges come from?

12 Thomson Thomson concluded that the negative charges came from within the atom. Thomson called the negatively charged electrons. Since the gas was known to be neutral, he reasoned that there must be positively charged particles in the atom. But he could never find them.

13 Thomson Model “Plum Pudding” model.
Atoms were made from a positively charged substance with negatively charged electrons scattered around

14 Rutherford’s Gold Foil Experiment
In 1908 English physicist Ernest Rutherford, began work on his gold foil experiment.

15 Rutherford’s Hypothesis
Rutherford was trying to verify Thomson’s model. He expected positively charged particles to go straight through a piece of very thin gold.

16 What Happened Most particles passed straight through the gold foil
A small percentage of particles were deflected at large angles or returned to the source


18 Rutherford’s Experiment
There are 2 reasons alpha particles deflected Density of the nucleus Repelling charges

19 Rutherford’s Conclusion
An atom has a small, dense, positively charged center that repelled the positively charged particles. Named the center of the atom the “nucleus” The nucleus is tiny compared to the atom as a whole. The atom is mostly empty space

20 Rutherford’s Nuclear Model
An atom’s positively charged particles were contained in the nucleus. The negatively charged particles were scattered outside the nucleus around the atom’s edge.

21 Neils Bohr Was a Jewish Scientist in Copenhagen during the onset of WWII Hitler was interested in his research of the atom. He was moved to the US to protect his knowledge.

22 Bohr Model Electrons travelled in a specific orbit at a certain distance from the nucleus called an energy level Worked well for Hydrogen and Helium

23 Nuclear symbols In this unit we need to be familiar with this type of symbol A X Z

24 Hyphen Notation Includes an element name a ‘-’ and a number
Example: Sulfur – 32 This sulfur atom has an atomic mass of 32 Since Sulfur has ______ protons & electrons It also has ______ neutrons

25 Particle Symbol Charge Mass (amu) Location Proton Neutron Electron

26 Average Atomic Mass Calculate the average atomic mass of magnesium given the following information: Mg, mass = amu; percent abundance = 78.99% Mg, mass = amu; percent abundance = 10.00% Mg, mass = amu; percent abundance = 11.01%

27 Bohr Model Electrons are a HUGE deal in chemistry
Responsible for light, color & chem reactions

28 Excited atoms

29 Bright-line Spectra Electrons in atoms:
when an atom absorbs energy - electrons jump to higher energy levels. An “EXCITED” electron jumps from its ground state to a higher energy level. When the electron returns it releases the same amount of energy that it absorbed.

30 Pink Floyd

31 Spectral Lines Light is released as photons as electrons return from different energy levels. Some of the photons are visible light Balmer Series

32 The Electromagnetic Spectrum

33 What is light? Wave Particle


35 Some Atomic Emission Spectra
Hydrogen Mercury Argon Helium

36 How many wavelengths are represented in each figure below?
Waves: A Warm up

37 Wave Comparison Red Light Low frequency Long wavelength Violet Light
nm = 1 x 10-9 m Red Light Low frequency Long wavelength Violet Light High frequency Short wavelength

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