Atomic Structure

1. Atomic Structure How do we know that atoms exist?

2. Thought experiment • What’s the smallest possible piece that is still considered sand? • Look at the beach • It’s made of sand • Cut the sand particles in your minds eye – smaller and smaller sand particles

3. History of the atom • Not the history of the atom, but the idea of the atom. • Original idea came from Ancient Greece (400 B.C.) • Democritus and Leucippus- Greek philosophers. • Pondered the fundamental nature of matter. • Matter is made up of indivisible particles • Atomos – indivisible, not to be cut

4. Another Greek • Aristotle - Famous philosopher • All substances are made of 4 elements • Fire - Hot • Air - light • Earth - cool, heavy • Water - wet • Blend these in different proportions to get all substances • All matter is continuous

5. Who Was Right? • Neither view was supported by experimentation. • Greeks settled disagreements by debate. • Aristotle was a better debater - He won. • His ideas carried through middle ages.

6. Foundations of Atomic Theory • Late 1700’s – most chemists accepted the modern definition of an element as a substance that cannot be broken down further by ordinary chemical means. • Elements combine to form compounds whose properties are different from the elements that form them. • Great controversy over whether elements always combine in the same ratio when forming a particular compound.

7. Foundations of Atomic Theory • Significant improvements to the available technologies • Led to a more quantitative study of elements, compounds, and chemical reactions • Led to the discovery of several basic laws

8. Law of Conservation of Mass • Mass is neither created nor destroyed during ordinary chemical reactions or physical changes.

9. Law of Definite Proportions • Regardless of where or how a pure chemical compound is prepared, it is composed of a fixed proportion of. elements. • It is a ratio by mass. • All salt crystals, NaCl, regardless of sample size contains exactly 39.34% sodium and 60.66% chlorine.

10. Foundations of Atomic Theory • In 1808 - John Dalton- England. • Teacher- proposed an explanation for these laws. • Dalton’s Atomic Theory 1) All atoms are composed of extremely small particles called atoms.

11. Dalton’s Atomic Theory 2) Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties. 3) Atoms cannot be subdivided, created, or destroyed. 4) Atoms of different elements combine in simple whole-number ratios to form chemical compounds. 5) In chemical reactions, atoms are combined, separated, or rearranged.

12. Where Dalton Was Wrong • Atoms are divisible into even smaller particles! • A given element can have atoms with different masses! • Atomic theory is an ever evolving concept improving as technology improves. • Unchanged for 200 years: matter is composed of atoms, atoms of one element differ in properties from atoms of another element.

13. The Structure of the Atom • Late 1800’s scientific advances allowed for a deeper exploration into the nature of matter. • Atoms are composed of several basic types of smaller particles. • The number and arrangement of these particles within an atom determines that atom’s chemical properties.

14. The Structure of the Atom • An atom is the smallest particle of an element that retains the chemical properties of that element. • All atoms are composed of two regions: • Nucleus: very small dense region located in the center of the atom. • Made up of at least one positive particle, proton, and usually one or more neutral particles called neutrons.

15. The Structure of the Atom • Surrounding the nucleus is a large region occupied by negatively charged particles called electrons. • This region is called the electron cloud. • Protons, Neutrons, and Electrons are referred to as subatomic particles.

16. Discoveries • J. J. Thomson - English physicist. 1897 • Investigated relationship between electricity and matter. • Made a piece of equipment called a cathode ray tube. • It is a vacuum tube - all the air has been pumped out. • A limited amount of other gases are put in

17. Voltage source Thomson’s Experiment - + Metal Disks

18. Voltage source Thomson’s Experiment - + • Passing an electric current makes a beam appear to move from the negative to the positive end • Called cathode rays

19. Voltage source Thomson’s Experiment + - • By adding a magnetic field

20. Voltage source Thomson’s Experiment + - • By adding a magnetic field he found that the moving pieces were negative. • By adding a magnetic field

21. Thomson’s Experiment • Used many different metals and gases • Beam was always the same • By the amount it bent he could find the ratio of charge to mass • He found the ratio was the same with every material • He concluded that the rays were composed of identical negatively charged particles.

22. Millikan’s Experiment • 1909 Robert Millikan measured the charge of an electron. • Scientists used this information and the charge-to-mass ratio to determine the mass of an electron.

23. Atomizer + - Oil Microscope Metal Plates Millikan’s Experiment

24. Atomizer Oil droplets + - Oil Microscope Millikan’s Experiment

25. Millikan’s Experiment X-rays X-rays give some drops a charge by knocking offelectrons

26. + - Millikan’s Experiment

27. Millikan’s Experiment - - + + They put an electric charge on the plates

28. Millikan’s Experiment - - + + Some drops would hover

29. + Millikan’s Experiment - - - - - - - Some drops would hover + + + + + + +

30. Millikan’s Experiment - - + + From the mass of the drop and the charge on the plates, he calculated the charge on an electron

31. Inferences About Atomic Structure • Atoms are electrically neutral, therefore, they must contain a positive charge to balance the negative electrons. • Electrons have so much less mass than atoms, atoms must contain other particles that account for most of the atoms mass.

32. Thomsom’s Plum Pudding Model • Said the atom was like plum pudding. • For us, like seeds in a watermelon • A bunch of positive stuff, with the electrons able to be removed.

33. Rutherford’s Experiment • Ernest Rutherford English physicist. (1911) • Hans Geiger and Ernest Marsden • Believed the plum pudding model of the atom was correct. • Set out to prove it!

34. Rutherford’s Experiment • Gold Foil Experiment • Bombarded a thin piece of gold foil with fast-moving alpha particles, positively charged particle 4 times the mass of a hydrogen atom. • Expected the particles to go through the gold atoms with only slight deflection. • Most did what was expected but 1 in 8000 was deflected back toward the source.

35. Flourescent Screen Lead block Uranium Gold Foil

36. What he expected

37. Because, he thought the mass was evenly distributed in the atom

39. What he got

40. Rutherford’s Conclusion • When the alpha particles hit a florescent screen, it glows. • The deflected particles experienced a powerful force within the atom. • He figured this force occupied a very small space because most of the particles went through.

41. + Rutherford’sConclusion • Atom is mostly empty space. • Small densely packed bundle of matter with a positive electric charge. • Alpha particles are deflected by it if they get close enough.

42. +

43. What We Know So Far • Except for the simplest hydrogen atom, all nuclei are composed of two kinds of particles • positive Protons • neutral Neutrons • Atoms are electrically neutral, they contain the same number of protons and electrons.

44. What We Know So Far • The nuclei of atoms of different elements differ in their number of protons and therefore in the amount of positive charge. • The number of protons determines the atoms identity.

45. What We Don’t Know Yet • Where are the electrons?

46. Forces In The Nucleus • Like charges repel each other • Except in the nucleus • When like charges are extremely close together there is a strong attraction between them. • More than 100 protons can exist together to help form a nucleus. • These short range forces hold the nuclear particles together, they are called nuclear forces.

47. Subatomic Particles • Proton - positively charged pieces 1836 times heavier than the electron. • Neutron - no charge but the slightly larger than the mass of a proton. • Electron – negatively charged pieces ~1/1837 times the mass of the hydrogen atom. • Where are the pieces?

48. Subatomic particles Actual mass (g) Relative mass Name Symbol Charge Electron e- -1 1/1837 9.109 x 10-28 Proton p+ +1 1 1.673 x 10-24 Neutron n0 0 1 1.675 x 10-24

49. Structure of the Atom • There are two regions. • The nucleus. • With protons and neutrons. • Positive charge. • Almost all the mass. • Electron cloud- most of the volume of an atom. • The region where the electron can be found.

50. Size of an atom • Atoms are small. • Measured in picometers, 10-12 meters. • Hydrogen atom, 32 pm radius. • Nucleus tiny compared to atom. • IF the atom was the size of a stadium, the nucleus would be the size of a marble. • Radius of the nucleus is near 10-15m. • Density near 1014 g/cm3.