1. The Development of Atomic Models: A Historical Perspective

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

3. atom: (Gk. atomos—indivisible)
Atomic Models
Democritus’ Idea
An object CANNOT be divided indefinitely.
There is a smallest particle.
Democritus – Chemistry textbook p. 73 (BJU Press)
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.
Democritus – Chemistry textbook p. 73 (BJU Press)

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 5

6. Atomic Models John Dalton Framed the first experimental model
Experimented by weighing substances
Assigned relative masses to elements
John Dalton – Chemistry textbook p. 74 (BJU Press)
Dalton’s model – Chemistry textbook p. 74 (BJU Press)

7. Atomic Models John Dalton Believed that:
Atoms of different elements form compounds.
Atoms of different elements have different masses.
John Dalton – Chemistry textbook p. 74 (BJU Press)
Dalton’s model – Chemistry textbook p. 74 (BJU Press)

8. Example: Water’s composition by weight is always an 8:1 ratio.
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.
Thomson’s plum-pudding model – Chemistry textbook p. 75 (BJU Press)

12. Atomic Models Thomson’s Model “Plum-pudding” model
(It looks more like a chocolate chip cookie.)
Thomson’s plum-pudding model – Chemistry textbook p. 75 (BJU Press)

13. Atomic Models Thomson’s Model Shows inaccuracies with Dalton’s model:
Atoms aren’t solid.
Atoms aren’t indestructible.
Thomson’s plum-pudding model – Chemistry textbook p. 75 (BJU Press)

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. Discovery of the Proton
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. Rutherford’s Experiment
Atomic Models
Rutherford’s Experiment
alpha particle slightly deflected
gold foil
zinc sulfide screen
beam of alpha particles
alpha particle greatly deflected
Rutherford’s experiment setup – Chemistry textbook p. 76 (BJU Press)

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.
Rutherford’s model, the nuclear model of the atom – Chemistry textbook p. 78 (BJU Press)

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

22. This is NOT a planetary model.
Rutherford’s Model
Rutherford’s model, the nuclear model of the atom – Chemistry textbook p. 78 (BJU Press)
This is NOT a planetary model.

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

24. Discovery of the Neutron
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. Discovery of the Neutron
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. Energy Levels for Electrons
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. Energy Levels for Electrons
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 continuous spectrum and a bright-line spectrum of hydrogen – Chemistry textbook p. 80 (BJU Press)
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.
Bohr’s model, the “planetary” model of the atom – Chemistry textbook p. 80 (BJU Press)
It is a planetary model.

39. Bohr’s Model
Bohr’s orbits (energy levels) are called principal energy levels.
Seven levels can be measured.
Bohr’s model, the “planetary” model of the atom – Chemistry textbook p. 80 (BJU Press)

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.

43. Level (n)
Total # e− (2n2) 1 2 8 3 18 4 32 5
50* (32) 6
72* (18) 7
98* (8)

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’s Hypothesis
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’s Hypothesis
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.

51. Heisenberg Uncertainty Principle
We can know where an electron is or where it is going; but we cannot know both at the same time.

52. Heisenberg Uncertainty Principle
When an electron is hit by photons (light), it disappears and reappears somewhere else, but it does not travel between.

53. Heisenberg Uncertainty Principle
Bohr’s precise, planet-like orbits were replaced by orbitals.

54. Heisenberg & Electrons
Electrons are in orbitals (areas where electrons are most likely found).
Orbitals look like fuzzy, shaded clouds.
Darker shading indicates a higher probability of an electron existing there.