Atomic Theory

1. Atomic Theory You too can be as smart as Einstein (almost)

2. History - Greeks • The elements • Earth – dry, heavy • Water – wet, heavy • Air – cool, light • Fire – warm, light • The composition of a substance could be estimated from its properties.

3. History - Greeks • These ideas were based on observation, logic and reason, but not experimentation. • Democritus (460 B.C. - 370 B.C.)

4. History - Greeks • Matter is made of small, hard indivisible particles called atoms, which exist in the void. • These atoms differ in size and shape, but not in any other way. Quantitative differences (how much) vs. Qualitative differences (what kind)

5. History - French • Antoine Lavoisier (1743-1794) • Discoverer of Oxygen (disputed) • His work refuted the phlogiston theory • Responsible for the law of conservation of matter.

6. History – French 1800 • Claude Louis Berthollet Joseph Louis Proust

7. History - French • Berthollet – “compounds do not have a fixed composition”. Cu + S  CuxSy • Every time he tried the experiment he got a different result.

8. History - French • Proust - compounds have a fixed composition. 2H2 + O2 2H2O • He always got the same result. • Proust’s argument is called The Law of Definite Proportions. He was proved to be right.

9. History - English • John Dalton (ca. 1804) • The father of modern atomic theory • Schoolteacher • Colorblind – studied colorblindness

10. Dalton’s Atomic Theory • The points of Dalton’s theory • All matter is made of atoms • Atoms are indivisible and indestructible • All atoms of one element are exactly alike, and atoms of different elements are different. • Atoms combine in small whole number ratios to form compounds.

11. Dalton’s Atomic Theory • The Law of Multiple Proportions: If two elements combine to make two different compounds, the ratios of the elements involved are small whole numbers. •  Examples: CO and CO2 CuS and Cu2S H2O and H2O2

12. Inside the Atom • J. J. Thomson and the Electron (1897)

13. The Electron • Thomson discovered the electron - he called it a “corpuscle”. • He used an instrument called a Crookes tube. Cathode (-) Evacuated tube Anode (+)

14. The Electron • He noticed a stream of charged particles coming from the cathode, called cathode rays. • Thomson proposed the "plum pudding" atomic model - negatively charged corpuscles swarm inside a cloud of massless positive charge.

15. Ernest Rutherford and the Nucleus

16. The Nucleus • The gold foil experiment (1909)

17. Gold Foil Experiment • Most of the alpha particles went straight through, and a few were bounced straight back. •  Rutherford’s interpretation: The atom has a small, hard, dense and positively charged nucleus. The electrons are outside the nucleus.

18. The Proton and the Neutron • Discovery of the proton: Henry Moseley (1913) • Moseley bombarded metals with x-rays • Each successive element had one more positive charge – called “atomic number” • Rutherford proved that the nucleus of nitrogen contains hydrogen nuclei – a “proton” (1918-19) • Discovery of the neutron – James Chadwick (1932)

19. Parts of the Atom

20. Isotopes • Atomic number = number of protons in the nucleus • Atomic number determines the identity of the element • Mass number = protons + neutrons • Number of electrons = number of protons • Isotopes: two atoms of the same element with different numbers of neutrons • C-12 and C-13 are isotopes of carbon

21. Nomenclature and symbols • Nuclear symbols 13C • Write the nuclear symbol for lead-206. 206Pb 6 82

22. Periodic table Atomic number 20 Ca Calcium 40.078 Symbol Name Average atomic mass

23. Average atomic mass • Average mass of all the isotopes of an element • Average is weighted • Example: Boron has two isotopes, B-10 and B-11 B-10: 19.9% B-11: 80.1% • Average atomic mass of boron: 10x0.199 = 1.99 11x0.801 = 8.811 Average atomic mass = 1.99 + 8.811 = 10.8amu

24. Outside the Nucleus • Niels Bohr and the stepwise atom (ca. 1918)

25. Rutherford-Bohr Model of the Atom (1911-1913) • Rutherford suggested that electrons orbit around the nucleus like planets around the sun. • This did not explain emission spectra, which gave sharp lines. • He theorized that electrons could only travel in certain sized orbits, and not anywhere in between.

26. Rutherford-Bohr Model of the Atom (1911-1913)

27. Bohr Model of the Atom • The orbits were called energy levels. Each orbit has a specific energy. •  Electrons can jump from one level to another; as they do, they absorb or emit energy.

28. Quantum Mechanics • Erwin Schrödinger and probable cause (ca. 1935)

29. Quantum Mechanics • Schrödinger’s work showed that electrons do not move in actual “orbits”. • Electrons move randomly and form “probability clouds”. The shape of these clouds is similar to the shape of Bohr’s orbits. • The position and momentum of an electron cannot be determined simultaneously (Heisenberg Uncertainty Principle)

30. Quantum Mechanics • Schrödinger’s “electron cloud”

31. Electron Energy Level Populations • Bohr suggested that electrons inhabit energy levels around the nucleus. • Each level has a specific energy associated with it. • The outermost (highest energy) level is called the “valence shell”. • The electrons in the valence shell are called the “valence electrons”. • The valence electrons are the most important electrons in the chemistry of the atom.

32. Electron Energy Level Populations

33. Electron Energy Level Populations • The number of levels depends on the number of electrons. • The first level (K) holds two electrons. • The second level holds eight electrons. • The third level holds 18, and the fourth 32. • No atom can have more than eight electrons in its valence shell. • When the valence shell reaches eight electrons, the next two electrons are put in a higher level. Then the lower level can be filled.

34. Lewis Electron Dot Structures • Lewis dot structures show how many electrons are in the valence shell of an atom. Lewis dot structure for sodium • The first electron always goes to the right of the symbol. • The second is paired with the first.

35. Lewis Dot Structures Lewis dot structure of magnesium • The third goes on top. Lewis dot structure of aluminum

36. Lewis Dot Structures • The fourth goes on the left, and is not paired. The fifth goes on the bottom, and successive electrons are paired until a total of eight is reached. Lewis dot structure of silicon

37. Lewis Dot Structures Lewis dot structure of oxygen

38. Atomic Spectra • Bohr’s model based on atomic spectra • Obtaining emission atomic spectra • Energy is applied to a gas or liquid sample. • Flame test (for samples in solution) • Gas discharge tube • The energy makes an electron or two jump to a higher energy level. • The electrons fall back down to a lower level, and give off energy in the form of light – bright lines against a dark background.

39. Atomic Spectra • Absorption spectra – light is passed through a sample and analyzed – looks like a rainbow with dark lines • Interpreting atomic spectra • The light given off is viewed through a spectroscope. • The spectroscope has either a prism or a grating, which splits the light into its component colors.

40. Atomic Spectra

41. Atomic Spectra • Only a few sharp lines appear in the spectrum. • Each line corresponds to a specific electron transition. • Transition = jump from one energy level to another

42. Light Energy • Light energy travels in the form of waves.

43. Light Energy • Color depends on frequency. • High frequency = violet end of spectrum • Low frequency = red end of spectrum • Energy also depends on frequency, so each color has its own energy. Blue or violet is higher energy than red or green. • When a specific color line is seen in a spectrum, the energy of the electron transition responsible can be calculated.

44. Light Energy and Bohr’s Model • Bohr reasoned that since only certain lines are seen in atomic spectra, only certain energies must be allowed in electron orbits.