1. Goal 5.01 Atomic Theory

2. Basic Laws of Chemistry Law of conservation of massLaw of conservation of mass: mass is neither created or destroyed during chemical reactions or physical changes Law of conservation of mass Law of definite proportionsLaw of definite proportions: compounds contain the same proportions of elements by mass regardless of the size of the sample Law of definite proportions Law of multiple proportions: different compounds made from the same elements result from the combination of different whole numbers of atoms

3. Dalton Dalton turned Democritus' idea of a indivisible particle of matter into a theory that could be tested. He called it the atomic theory. It explained the laws of chemistry that all atoms followed. atoms Not all aspect’s of Dalton’s theory have been proven correct.

4. Atomic Theory All elements are composed of atoms that cannot be divided. Atoms of the same element have the same mass and other properties. Atoms cannot be created or destroyed. Compounds contain the atoms of more than one element. In chemical reactions, atoms of the same element always combine, separate, or rearrange in the same way.

5. Where He Went Wrong Parts of Dalton’s theory have been shown to be untrue… Atoms can be divided in to smaller parts. Atoms of the same element can have different masses.

6. Thomson’s Experimental Setup Thomson preformed an experiment using a cathode ray tube. It is a sealed tube that contains a gas at a low pressure. Current is passed through the gas from the cathode (+) to anode (-) sides of the tube. When this is done, a glowing beam of what looks like light is created. This glowing beam is called the cathode ray.

7. Thomson’s Experimental Results When a paddle wheel was placed in the cathode ray, the wheel turned and moved. When a negative magnetic field was placed near the cathode ray, the ray bent away from the charge. When a positive magnetic field was placed near the cathode ray, the ray bent toward the charge.

9. Thomson’s Conclusion Because the paddle wheel turned, the beam must actually be particles. The mass of these particles was smaller than an atom’s mass and therefore the particle must be subatomic. Because of how it was deflected by charges, the particles must be negative. Thompson named these negative, subatomic particles electrons.

10. Inferences This led to other inferences about atomic structure. Because atoms are neutral, there must be a positive charge to balance the negatives. Because electrons weigh so little their must be something else to account for the weight.

11. Rutherford’s Experimental Setup A thin piece of gold foil, full of tiny gold atoms, was bombarded with fast moving alpha particles. Alpha particles are a positive type of radiation with four times the mass of a hydrogen. It was thought that all alpha particles would pass through the gold foil because nothing would be big or strong enough to deflect them.

12. Rutherford’s Experimental Results BUTBUT… 1 in every 8,000 alpha particles were deflected back. BUT This deflection was caused by the powerful repulsive force of a positively charged space within the atom and the density that the space had.

14. Rutherford’s Conclusion Rutherford concluded that the space was small due to how few alpha particles were deflected. Rutherford named the small, positively charged space a nucleus and the space surrounding it an electron cloud. small electron cloudsmall electron cloud He named the particles inside the nucleus that gave it it’s positive charge protons. This positively charged protons were what was canceling out the negative charge of the electrons.

15. Bohr’s Experimental Setup Bohr made electrons absorb energy and jump up in orbit. When this occurs, the atom is said to be excited. As the electron falls from the excited state to the ground state, energy is given back off in the form of a photon of radiation. photon The photon of energy was then sent through a prism where it formed a spectrum.

16. Bohr’s Experimental Results The spectrum that was produced was a line emission spectrum. spectrum This indicated that only fixed amounts of energy were being released/absorbed as electrons moved between orbits.

18. Bohr’s Conclusion Bohr discovered that electrons don’t just move randomly around a nucleus because if they did the spectrum would have been continuous. Instead the electrons follow distinct paths called orbits and these orbits are at set distances from one another. orbits Bohr used the wavelengths of the lines in the spectrum to calculate the exact energy that the electrons have at each level.

19. Electron Probability It is now known that electrons don’t have to follow these paths. Electrons are just likely to follow paths called orbits. They have a probability to. probability