Chapter 15: The Nuclear Atom

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Our purpose in trying to understand light is indirectly to understand the atom. In this chapter, we concentrate on the atom. In the two following chapters, we bring light and the atom together into a final model. How do we know anything about an atom that is so small that even with the most sophicsticated of microscopes, it just looks like a tiny smudge? How do we know that the atom has charged parts? Experiments with gas discharge tubes in the 19th century demonstrated that atoms have charged parts: negative and positive fragments. How do we know the mass of an atom? The "mass spectrometer" measures the masses of atom. Atoms are bombarded until they lose an electron and become a charged object (called an ion). The ion is given a known amount of (gravitational potential, kinetic, radiant?) energy. The (speed, charge?) of the ion is measured separately. From the definition of kinetic energy, Kinetic Energy = 1/2 (mass) x (speed)2, we can solve for the unknown mass. How do we know what the structure of an atom is like? The Thomson Model of the atom imagined the negative electrons stuck in the mass of positive charge like plums in pudding or raisins in bread. Experiments bombarding a gold foil with charged particles by (last name?) in 1911 showed that the plum-pudding model was wrong. The experiments showed that the atoms was mostly space with a very tiny at the center. The matter of the nucleus was unlike matter that anyone had ever imagined. If a pin were made entirely of nuclear matter, it would weight as much as a battleship! Thus, the solar-system model (Rutherford) replaced the plum-pudding model (Thomson). There are two terrible things wrong with a solar-system model of the atom: the electrons cannot be stationary (because they would be pulled into the ) and they cannot be NOT stationary (because they would radiate away their energy as radiation) and, it is virtually impossible to explain the (continuous, discrete?) spectrum of light emitted by atoms if they truly are little solar systems. Neils Bohr developed a modified solar-system model of the atom that explained the discrete spectrum. He assumed: electrons moving around an undisturbed atom ordinarily (do, do not) radiate energy, and only certain discrete orbits are possible and the electron is characterized by a (same, different?) energy in each different orbit. An emission spectrum is a discrete spectrum of colors produced when many atoms of the same kind are bumped so that electrons of the atoms are bumped into orbits of (higher, lower?) energy. The electron gets the extra energy from the thing that bumped it. When the electron spontaneously jumps back down to an orbit of lower energy, it must get rid of the extra energy. It does so by creating a to carry the energy out of the atom. Photon energy = E2 - E1 = (Planck's constant) x ( ). An absorption spectrum is a spectrum created when light with a continuous spectrum is passed through a gas. Only photons with exactly the right energy can be absorbed by in the atom to make them jump from one allowed orbit up to an orbit of higher energy. The gas subtracts just these photons (colors) from the original continuous spectrum as the light passes through, yielding a continuous spectrum with a discrete set of (colored, black?) lines.

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