1. Unit 3
History of the Atom

2. Democritus about 460 bc stated that the “ultimate particle” is atmos
He believed that atoms were indivisible and indestructible
was not based on the scientific method – but just philosophy

3. John Dalton - 1803 English school teacher
the atom was a small, indivisible particle like a “small marble”

4. Dalton’s Atomic Theory (experiment based!)
All matter is made of tiny indivisible particles called atoms
Atoms of the same element are chemically the same. Atoms of different elements are chemically different.
John Dalton
(1766 – 1844)
Individual atoms of an element have slightly different masses. We use average mass.
Different elements have different ave masses
Atoms can’t be divided in normal chemical reactions.

5. Discovery of the Electron
In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle: the electron

6. Modern Cathode Ray Tubes
Television
Computer Monitor
Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

7. Thomson’s Atomic Model
J. J. Thomson
Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.

8. Ernest Rutherford’s Gold Foil Experiment - 1911
Alpha particles are helium nuclei - The alpha particles were fired at a thin sheet of gold foil
Particles that hit on the detecting screen (film) are recorded

9. Rutherford’s Findings
Most of the particles passed right through
A few particles were deflected
VERY FEW were greatly deflected
“Like howitzer shells bouncing off of tissue paper!”
Conclusions:
The nucleus is small
The nucleus is dense
The nucleus is positively charged

10. The Rutherford Atomic Model
Based on his experimental evidence:
atom - mostly empty space
all positive charge, and most mass is in the center called a “nucleus”
nucleus = protons and neutrons (neutrons confirmed by Chadwick in 1932)
electrons distributed around the nucleus, and occupy most of the volume
His model was called a “nuclear model”

11. Niels Bohr
Niels Bohr wondered why horseshoes heated in a forge changed color, but did not react
Nucleus?
Electrons?

12. Neils Bohr
All atoms contain energy – the energy has to do with the electrons
1913 Bohr said that electrons are in definite areas with definite amounts of energy

13. Planetary Model
Electrons are in the “ground state” – normal lowest energy state
*excited state electrons gain energy in fixed amounts (photons) and go to higher levels

14. So, Energy = Planck’s constant x frequency
E = hf
E = energy
h = Planck’s constant
f = frequency
So, Energy = Planck’s constant x frequency

15. Electrons don’t stay in the excited state, but fall back to the ground state and give off energy in the form of light called a spectrum (which is unique for each element)

16. http://www. mhhe. com/physsci/astronomy/applets/Bohr/applet_files/Bohr

17. Speed of light (constant) = 2.998 x 108 m/sec
c = f λ
wavelength, lambda
speed of light
frequency

18. c = f λ
frequency=speed/wavelength
2.998 x 108 m/s divided by 7.6 x 10-7
= 3.9 x /s = 3.9 x 1014 Hz
Energy of an electron = x J/n2

19. Schrodinger
Wave Mechanical Model – electrons exist in areas of probability called “space orbitals” (Rutherford and Bohr support this)

20. Quantum Mechanics – each electron is identified by 4 quantum numbers:
1. Principle quantum # - level
2. Orbital quantum # - shape
3. Magnetic quantum # - orientation
4. Spin quantum # - clockwise or counter

21. Werner Karl Heisenberg -1927
Uncertainty Principle – you can’t locate the exact position of an electron at any given time (too small, too fast)