Atomos: Not to Be Cut

1. Atomos: Not to Be Cut The History of Atomic Theory

2. Atomic Models • A model uses familiar ideas to explain unfamiliar facts observed in nature. • A model can be changed as new information is collected.

3. The atomic model has changed throughout the centuries, starting in 400 BC, when it looked like a billiard ball →

4. Democritus 400 BC • This is the Greek philosopher Democritus - • His theory: Matter could not be divided into smaller and smaller pieces forever, eventually the smallest possible piece would be obtained.

5. Atomos • This piece would be indivisible. • He named the smallest piece of matter “atomos,” meaning “not to be cut.”

6. This theory was ignored and forgotten for more than 2000 years!

7. Why? • The eminent philosophers of the time, Aristotle and Plato, had a more respected, (and ultimately wrong) theory. Aristotle and Plato favored the earth,fire, air and waterapproach to the nature of matter. Their ideas held sway because of their eminence as philosophers. The atomos idea was buried for approximately 2000 years.

9. Dalton’s Model • In the early 1800s, the English Chemist John Dalton performed a number of experiments that eventually led to the acceptance of the idea of atoms.

10. Dalton’s Theory • All elements are made of tiny, indivisible particles called atoms. • All atoms of the same element are identical. All atoms of different elements are different (unique). • Atoms combine with different atoms to form compounds. Atoms unite in whole number ratios (H2O, CO2). • Atoms are not created or destroyed, simply rearranged in chemical reactions.

11. . • This theory became one of the foundations of modern chemistry.

12. Thomson’s Plum Pudding Model • In 1897, the English scientist J.J. Thomson provided the first hint that an atom is made of even smaller particles.

13. Thomson Model • Thomson studied the passage of an electric current through a gas. • As the current passed through the gas, it gave off rays of negatively charged particles.

15. Thomson Model • He proposed -“PlumPudding” model. • Atoms were made from a positively chargedsubstance with negatively charged electrons scattered about, like raisins in a pudding.

16. Thomson Model Where did they come from? • This surprised Thomson, because the atoms of the gas were uncharged. Where had the negative charges come from?

17. Rutherford’s Gold Foil Experiment • In 1908, the English physicist Ernest Rutherford. • Rutherford’s experiment Involved firing a stream of tiny positively charged particles at a thin sheet of gold foil (2000 atoms thick)

18. Gold Foil Experiment

19. Gold Foil Experiment • Most of the positively charged “bullets” passed right through the gold atoms in the sheet of gold foil without changing course at all. • Some of the positively charged “bullets,” however, did bounce away from the gold sheet as if they had hit something solid. He knew that positive charges repel positive charges.

20. This could only mean atoms were mostly open space. • Rutherford concluded that an atom had a small, dense, positively charged center that repelled his positively charged “bullets.” • He called the center of the atom the “nucleus” • The nucleus is tiny compared to the atom as a whole.

22. Bohr Model • In 1913, the Danish scientist Niels Bohr proposed that electrons were in a specific energy level.

23. Bohr Model • electrons move in definite orbits around the nucleus, much like planets circle the sun. • These orbits, or energy levels, are located at certaindistances from the nucleus.

24. Wave Model

25. The Wave Model • Today’s atomic model is based on the principles of wavemechanics. • electrons do not move about an atom in a definite path, like the planets around the sun. • In fact, it is impossible to determine the exact location of an electron. The probable location of an electron is based on how much energy the electron has.

26. The Wave Model • According to the modern atomic model, at atom has a small positively charged nucleus surrounded by a large region (electron cloud) in which there are enough electrons to make an atom neutral.

27. Electron Cloud: • Depending on their energy they are locked into a certain area in the cloud. • Electrons with the lowest energy are found in the energy level closest to the nucleus • Electrons with the highest energy are found in the outermost energy levels, farther from the nucleus.

28. The Atom CHAPTER 4, PART I

29. http://www.gpb.org/chemistry-physics/chemistry/302

30. Relative Sizes What is the relative distance from the nucleus of an atom to the electrons? If the nucleus of an atom were the size of a marble how far away would the first electron be?

31. Answer

32. Subatomic Particles • Proton • Neutron • Electron FYI: • Quarks • Up / Down • Charm / Strange • Top / Bottom

33. Proton • Symbol = P+ • Charge = Positive • Mass = 1 atomic mass unit (amu) • Location = In the nucleus

34. Neutron • Symbol = n0 • Charge = Neutral • Mass = 1 atomic mass unit (amu) • Location = In the nucleus

35. Electron • Symbol = e- • Charge = Negative • Mass = 0 atomic mass unit (amu) • Location = Orbiting the nucleus

37. Atomic Number, (Z) = # of P+ in an atom’s nucleus • Mass Number, (A) = Sum of the # of N0 and # of P+ in a given nucleus. A Z A = mass #, Z = atomic #, X = symbol X

38. Atomic Calculations 1) Atomic Number = Number of Protons Z = Atomic Number = # p+ 2) If neutral, then protons must equal electrons 3) Atomic Mass = Number of Protons + Neutrons A = Atomic Mass A = Z + n0

39. Atomic Calculations Game http://education.jlab.org/elementmath/index.html

40. Isotopes – atoms of the same element with the same # of P+, but different # of N0 and different mass numbers. • Ex. • Nuclide – a particular atom containing a definite number of protons and neutrons. • Ex. Carbon-12 Iron-56 • Nucleons – particles that make up the atomic nucleus • Ex. protons and neutrons.

41. Atoms vs Ions B-1 C+1 C

42. Ion – an atom that has lost or gained electrons and now has a charge Number of protons = Z (atomic number) Number of neutrons = A – Z (atomic mass – atomic number) Number of electrons = Number of protons - charge

43. Average Atomic Mass • AMU – atomic mass unit, 1/12 the mass of a carbon-12 atom. • Mass on periodic table is a weighted average mass of all isotopes of an element.

44. Example Neon-12 relative abundance – 98% Carbon-14, relative abundance – 2% 0.98 (12 amu) + 0.02 (14 amu) = 12.04 amu

45. Example 2 • What is the atomic mass of silicon if 92.21% of its atoms have mass of 27.977 amu, 4.70% have mass of 28.976 amu, and 3.09% have a mass of 29.974 amu? • 28.09 amu