The Development of Atomic Models: A Historical Perspective

1. The Development of Atomic Models: A Historical Perspective

2. Model of an Atom An IDEA of what it looks like (a working representation)

3. Atomic Models Democritus’ Idea • An object CANNOT be divided indefinitely. • There is a smallest particle. atom: (Gk. atomos—indivisible)

4. Atomic Models Democritus’ Idea was CORRECT! • There is a basic unit of matter—the atom. • Chemists found this • out by looking at the mass ratio of substances. • But the atom is NOT indivisible.

5. 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

6. Atomic Models John Dalton • Framed the first experimental model • Experimented by weighing substances • Assigned relative masses to elements

7. Atomic Models John Dalton • Believed that: • Atoms of different elements form compounds. • Atoms of different elements have different masses.

8. law of definite composition Every compound has a definite composition by weight. Example: Water’s composition by weight is always an 8:1 ratio.

9. Atomic Models Joseph Thomson • Created a vacuum in a tube • Light came from the “cathode” end. • Noted that cathode rays are affected by magnets and electricity, but not by gravity

10. Atomic Models Joseph Thomson • Decided that the cathode rays were negatively (−) charged and calculated their charge-to-mass ratio • Found that every element he tested emitted these cathode rays, later called electrons

11. Atomic Models Thomson’s Model • The electrons (e−) are in a positive jell-like substance. • The electrons and the • positive substance cancel out. • Electrons can be removed from the atom.

12. Atomic Models Thomson’s Model • “Plum-pudding” model • (It looks more like a chocolate chip cookie.)

13. Atomic Models Thomson’s Model • Shows inaccuracies with Dalton’s model: • Atoms aren’t solid. • Atoms aren’t indestructible.

14. Question What do we now know is part of the atom, but is missing from Thomson’s model? • Cathode rays • Electrons • Nucleus

15. Atomic Models 2 Theories • Continuous: Matter can be subdivided forever. • Particulate: A smallest particle exists.

16. Atomic Models Ernest Rutherford Discovery of the Proton • 1910 Alpha Particle Experiment: He aimed alpha particles at a thin gold foil. • Alpha particles are relatively heavy and positively (+) charged.

17. Atomic Models Ernest Rutherford Exper. Fact Conclusion A few alpha particles bounced back completely. Atoms have a dense positive charge—the proton (p+).

18. Atomic Models Ernest Rutherford Exper. Fact Conclusion Most alpha particles passed through the gold foil. Atoms are mostly empty space. The nucleus is only 1/100,000th the size of the atom.

19. alpha particle slightly deflected gold foil zinc sulfide screen beam of alpha particles alpha particle greatly deflected Atomic Models Rutherford’s Experiment

20. Atomic Models Rutherford’s Model • Protons form a small, central nucleus. • A proton is 1800 times as massive as an electron. • Electrons are somewhere away from the nucleus.

21. Question What do we now know is in the atom, but was missing from Rutherford’s model? • Neutrons • Protons • Electrons • Nucleus

22. Rutherford’s Model This is NOT a planetary model.

23. atomic number (Z) the number of protons in the nucleus of an atom

24. Atomic Models James Chadwick Discovery of the Neutron • According to Rutherford’s model, there is not enough mass with just a proton and an electron.

25. Atomic Models James Chadwick Discovery of the Neutron • In 1932, Chadwick discovered a neutral particle with almost the same mass as a proton. • He called it a neutron (N).

26. Atomic Models Niels Bohr Energy Levels for Electrons • Where are the electrons? • Are they moving? • If so, in what manner? • Why don’t the electrons fall into the nucleus? Don’t protons attract them?

27. Atomic Models Niels Bohr Energy Levels for Electrons • 1913 spectroscopy experiment

28. Spectroscopy Heated atoms absorb electromagnetic radiation. They emit energy as light, which is part of the spectrum.

29. Electromagnetic Spectrum • It contains different wavelengths of energy— electricity, radio, microwave, visible, cosmic rays. • As the wavelength decreases, there is more energy in each wave.

30. Electromagnetic Spectrum • A continuous spectrum comes from a source that emits all energies (colors) visible to our eyes. • A line spectrum comes from a source that emits only exact (quantized) energy.

31. Electromagnetic Spectrum a continuous spectrum a bright-line spectrum of hydrogen

32. Electron Motion • How do the e− make the line spectra? • By only giving off exact amounts of energy. • Bohr said they do this by moving only in restricted steps—from level to level (orbit to orbit). • These levels are therefore said to be quantized.

33. Question What causes the line spectra? • e− orbiting • e− moving between levels • p+ orbiting • p+ moving between levels

34. Electron Levels • Electrons can move from orbit to orbit. • A more distant orbit is higher in energy.

35. Electron Levels • So, an electron absorbs energy, “jumps” to a higher orbit, and then “falls” back down, emitting energy. • The lowest level is called the ground state.

36. Electron Orbits • Since the electron falls back down from one exact orbit to another, it releases an exact amount of energy (quanta). • This exact energy packet produces a line of the line spectra.

37. Question Why do the e− give off line spectra and not continuous spectra? • They move too fast. • They are too far from the nucleus. • They move only from one exact level to another. • They can only give off energy bursts, not continuously.

38. Bohr’s Model It is a planetary model.

39. Bohr’s Model • Bohr’s orbits (energy levels) are called principal energy levels. • Seven levels can be measured.

40. Spectroscopy • Bunsen and Kirchoff invented the spectroscope. • Different elements, when heated, produce different sets of wavelengths.

41. Spectroscopy • It acts like a fingerprint to identify elements. • It gives evidence that different elements have different numbers of electrons in different locations.

42. Principal Energy Levels • There are 7. • Each can hold more than 1 electron. • Bohr’s model works well only for atoms with one electron.

44. Where is the Electron? • In Bohr’s model: The electron is in an exact location, orbiting like a planet.

45. Where is the Electron? • Problems with Bohr’s model: • It only works for atoms with 1 electron. • Electrons are not planets made up of many atoms; they are a part of an atom. • Forces between atoms are not the same as forces within atoms.

46. Where is the Electron? • Conclusion: • Electrons are NOT in exact orbits.

47. Light • Has properties of a wave • Refraction • Reflection • Diffraction

48. Light • Einstein suggested that light consisted of massless particles called photons. • De Broglie suggested the opposite: that particles could act like waves.

49. de Broglie’sHypothesis • All particles act like waves. • The bigger the particle, the smaller the wave. • The waves are only significant for very small particles.

50. de Broglie’sHypothesis • An exact number (from 1–7) of de Broglie wavelengths can fit in a Bohr orbit around an atom, depending on the size of the orbit.