Atomic Structure

1. Atomic Structure Unit 3

2. History of the Atom • 400 B.C. – Democritus & Leucippus • Beach  sand  smaller piece of sand  atomos (indivisible) • Everything is composed of imperishable, indivisible elements called atomos.

3. Aristotle’s views • All substances are made of 4 elements • Fire, Air, Earth, Water • Blend these elements in different proportions to get all substances

4. So who was right? • Greeks settled argument by… • Debating! • Aristotle was a better speaker so he won. • His views carried on through the Middle Ages.

5. 1808 - Dalton • English schoolteacher • Recognized that elements were made of atoms • Combined LOTS of research • Dalton’s Atomic Theory

6. Dalton’s Atomic Theory 1. All matter is made of tiny, indivisible particles called atoms. 2. Atoms of a given element are identical in their physical and chemical properties. 3. Atoms of different elements differ in their physical & chemical properties

7. Dalton’s Atomic Theory 4. Atoms of different elements combine in simple whole # ratios to form compounds 5. In a chemical reaction, atoms are combined, separated, & rearranged, but never created, destroyed, or changed.

8. From Dalton’s Theory… • Law of Definite Proportions (#4) • 2 samples of a compound have the same proportions by mass • 500 kg NaCl = 60.66% Cl & 39.34% Na • 2 mg NaCl = • Law of Conservation of Mass (#5) • Mass of reactants = mass of products 60.66% Cl & 39.34% Na

9. From Dalton’s Theory… • Law of Multiple Proportions • If 2 or more compounds are composed of the same elements, the ratio of the mass of the elements is always a small, whole # • NO 1.14 g O: 1 g N • NO2 2.28 g O: 1 g N • O [NO]: O [NO2] = 1.14 : 2.28 = 1 : 2

10. Another way to look at it… • Water (H2O) has 8 g of oxygen per g of hydrogen. • Hydrogen peroxide (H2O2) is 16 g of oxygen per g of hydrogen. • 16 to 8 is a 2 to 1 ratio of oxygen. • Always whole #s because you have to add a whole atom --- you can’t add a piece of an atom.

11. Dalton’s Atomic Model • The Solid Sphere Model

12. The research continues… • 1897 – JJ Thomson – cathode ray tube • Movie • Pumped air out of glass tube and applied voltage to metal electrodes at either end of the tube • Anode = positive charge • Cathode = negative charge

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

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

16. Voltage source Thomson’s Experiment + - • By adding an electric field he found that the moving pieces were negative • Also placed a paddle wheel in the center – and it turned! So they must have mass! • By adding an electric field

18. Thomson’s Conclusions • The cathode ray consists of particles that have mass & a negative charge • Particles called ELECTRONS • Plum pudding model • Negative electrons in a ball of positive charge

19. The research continues… • We know that an electron has a negative charge and atoms are neutral • Mass of an electron MUCH less than an atom • Something must be missing…

20. Rutherford - 1909 • Beam of small, positively charged particles called ALPHA particles • Aimed at thin gold foil (a few atoms thick) • Measured the angles of deflection • Movie

21. Flourescent Screen Lead block Uranium Gold Foil

22. What Rutherford Expected • The Plum Pudding Model • Alpha particles would pass through without changing direction very much. • Because… • The positive charges were spread out evenly & would not stop the positive alpha particles.

23. What he expected

24. Because

25. Because, he thought the positive charges were evenly distributed in the atom

26. Because, he thought the positive charges were evenly distributed in the atom

27. What he got

28. + Rutherford’s Explanation • Atom is mostly empty. • Small dense, positive piece at center. • Alpha particles are deflected by it if they get hit the dense positive center.

29. +

30. Conclusions from Rutherford • The small dense, positive place at the center is called the NUCLEUS • Radius of nucleus is less than 1/10000 the radius of the atom • PROTON – positive charged particle in nucleus • Charge is exactly equal but opposite to an electron • BUT still not enough mass

31. Still More Research… • NEUTRONS • Found in nucleus with protons • Do not have a charge • Same mass as protons

32. The Nuclear Model - Rutherford • Electrons revolve around nucleus in elliptical orbits • Also called planetary model

33. The Bohr Model • Electrons are found in certain levels or shells around the nucleus

34. More research… • 1924 – de Broglie - Electrons behave like waves around the nucleus • Heisenberg’s Uncertainty Principle • A ceiling fan • Planck & Einstein – quantum theory • Electrons found in clouds instead of strict orbitals

35. Today’s model – Quantum model

36. Historical models of the atom (solid sphere, plum pudding, nuclear model, & planetary model all predicted exact locations of particles. The modern quantum theories combine all of this research with more recent findings that suggest a certain level of unpredictability.

37. Subatomic particles Actual mass (g) Relative mass Name Symbol Charge Electron e- -1 1/1840 9.11 x 10-28 Proton p+ +1 1 1.67 x 10-24 Neutron n0 0 1 1.67 x 10-24 Most of the mass is in the nucleus!

38. The Atom All atoms have protons and electrons Most atoms have neutrons Elements differ from each other in the number of protons in an atom Protons & neutrons made of quarks

39. Atomic Number • The number of protons in an atom • Same in all atoms of an element • Example – Hydrogen=1 • Atomic number also reveals the number of electrons in a neutral atom

40. Mass Number • Number of particles in the nucleus • = # of Protons + # of neutrons • Example – Neon has a mass # of 20 • Can vary among atoms of an element • Different elements can have the same mass number • Not specifically on periodic table

41. Using Atomic Symbols • Each element has a name & a symbol • Examples – • Sulfur = S • Sodium = Na • The subscript to the right tells you how many atoms are present • S8 = 8 sulfur atoms

42. Using Atomic Symbols • Contain the symbol of the element, the mass number, and the atomic number. Mass number X Atomic number

43. Symbols/Notation • Find the … • Atomic number • Mass Number • number of protons • number of neutrons • number of electrons • Name 24 Na 11

44. Symbols/Notation • Find the … • Atomic number • Mass Number • number of protons • number of neutrons • number of electrons • Name 80 Br 35

45. Isotopes • All atoms of an element have the same number of protons but not necessarily the same # of neutrons • ISOTOPE – Atoms of the same element with different numbers of neutrons

46. Naming Isotopes • Name – mass number • Examples • Helium-3, Helium-4 • Symbols/Notation 4 3 He He 2 2

47. Name That Element • if an element has an atomic number of 34 and a mass number of 78 what is the • number of protons • number of neutrons • number of electrons • Complete symbol • Name

48. Name That Element • if an element has 91 protons and 140 neutrons what is the • Atomic number • Mass number • number of electrons • Complete symbol • Name

49. Name That Element • if an element has 78 electrons and 117 neutrons what is the • Atomic number • Mass number • number of protons • Complete symbol • Name

50. Atomic Mass • There are different isotopes of each element that all have different masses • Therefore we look at AVERAGE atomic mass for an element • Based on abundance of each isotope in nature