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History of Atomic Theory

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1 History of Atomic Theory
Democritus B.C. ancient Greek philosopher believed all matter consisted of extremely small particles that could not be divided atoms, from Greek word atomos, means “uncut” or “indivisible” Aristotle believed all matter came from only four elements—earth, air, fire and water

2 Who Was Right? Greek society was slave based No experiments
It was all a thought game Settled disagreements by argument Aristotle was more famous so he won His ideas carried through to the middle ages.

3 John Dalton (Late 1700’s) School teacher in England
Based his conclusions on experimentation and observations. Combined ideas of elements with that of atoms

4 Dalton’s Atomic Theory
All elements are composed of submicroscopic indivisible parts called atoms. Atoms of the same element are identical, those of different atoms are different. Atoms of different elements combine in whole number ratios to form compounds. Chemical reactions involve the rearrangement of atoms. No new atoms are created or destroyed.

5 Parts of Atoms Most of Dalton’s theory is accepted today.
Except the part about atoms being indivisible

6 J.J. Thomson and the Cathode Ray Tube 1897
English physicist Provided the first evidence that atoms are made of even smaller particles Description of a cathode ray tube and a short video of how it works:

7 Thomson’s Experiment - +

8 Thomson’s Experiment + -

9 Thomson’s Experiment - +

10 Thomson’s Experiment - +
Passing an electric current makes a beam appear to move from the negative to the positive end.

11 Thomson’s Experiment - +
Passing an electric current makes a beam appear to move from the negative to the positive end.

12 Thomson’s Experiment By adding an electric field

13 Thompson’s Experiment
By adding an electric field

14 Thompson’s Experiment
By adding an electric field

15 Thompson’s Experiment
By adding an electric field he found the moving particles were negative

16 Thompson’s Model Found the electron
1 unit of negative charge Mass 1/2000 of hydrogen atom Later refined by Millikan to 1/1840 Concluded that there must be a positive charge since atom was neutral Atom was like plum pudding A bunch of positive stuff, with electrons able to be removed.

17 Other Pieces Proton – positively charged pieces 1,840 times heavier than the electron Neutron – no charge but the same mass as a proton.

18 Ernest Rutherford Former student of J.J. Thomson
Believed in plum pudding Wanted to find out how big they are Fired positively charged alpha particles at a piece of gold foil, which can be made a few atoms thick

19 Rutherford’s Experiment
When alpha particles hit a flourescent screen it will glow. Here’s what it looked like (pg. 90)


21 What he expected to see

22 Alpha particles should pass through without change in direction
Positive charges were spread out evenly. Alone they were not enough to stop an alpha particle


24 What he got

25 How he explained it Atom is mostly empty
Small dense, positive piece at the center Alpha particles are deflected if they get close enough to positive center

26 Niels Bohr ( ) Electrons have orbits about the nucleus (planetary theory) Electrons could only exist at given energy levels An energy level is where an electron is likely to be moving Energy levels were like steps on a ladder An electron can only be at any given step at any given time

27 Modern Atomic Theory Bohr Model—shows electrons in orbit around protons and neutrons Quantum-mechanical model—doesn’t show exact location of electrons, just probable place

28 Structure of the Atom There are two regions The nucleus Electron cloud
Protons and neutrons Positive charge Almost all of the mass Electron cloud Most of the volume of an atom Region where electron can be found

29 Subatomic particles

30 Counting the pieces Atomic number = number of protons
Same as the number of electrons in a neutral atom Mass number = the number of protons + neutrons

31 Atomic Mass Unit AMU Mass of a proton = 1.67 x 10 -27g
A pretty inconvenient number New unit referenced to mass of an isotope of carbon: carbon -12 Carbon-12 has 6 protons and 6 neutrons Has a mass of amu – an atomic mass unit Therefore 1 proton and 1 neutron has a mass of 1 amu.

32 So why not whole numbers for atomic masses in periodic table?
Reported numbers are average atomic mass units, reflecting the abundance of isotopes for any given number. In nature most elements occur as a mixture of two or more isotopes

33 Isotopes Atoms of the same element can have different numbers of neutrons Different mass numbers Called isotopes

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