History of Atomic Theory

1. History of Atomic Theory Democritus • 460-371 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. Thomsonand the Cathode Ray Tube1897 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:http://www.chem.uiuc.edu/clcwebsite/cathode.html

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 1871-1937

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 http://micro.magnet.fsu.edu/electromag/java/rutherford/

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 (1885-1862) • 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 • 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 12.00000 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