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6.1 The Atom

6.1 The Atom

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6.1 The Atom

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  1. 6.1 The Atom Atomic Structure Nuclear Structure

  2. 6.1.1 • Describe a model of the atom that features a small nucleus surrounded by electrons. • This is a simplified view of the atom known as the Rutherford model

  3. History • Atom – the term atomos comes from Greek and means “indivisable” • Originally thought to be the smallest parts of matter • The spheres to the right are the top layer of a gold sheet • Each sphere is one atom

  4. History • 1803 – John Dalton proposed that atoms reacted with each other • 1897 – JJ Thompson discovered the electron with the invention of the cathod ray tube • Thompson's model was the plumb pudding version, where negative "corpuscles" were distributed though out positive “pudding”

  5. History • The Rutherford model had a positively charged nucleus, with the negatively charge electron orbiting the nucleus • In 1913 Niels Bohr added Quantum energy levels saying the electrons could only exist in certain orbits with definite energy • These models didn’t adequately explain the atom so in 1926 Schrödinger, citing DeBroglie’s wave model for mass, created his Model using the uncertainty principle

  6. The Schrödinger Model • The nucleus contain the nucleons (Protons and Neutrons) with are made up of quarks • The electrons exist in electron clouds around the nucleus • All the components of the atom are made from elementary particles • Try this Atom Builder

  7. 6.1.2 • Outline the evidence that supports a nuclear model of the atom. • A qualitative explanation of the Geiger–Marsden experiment and its results is all that is required. • Also know as the Gold foil experiment, It was performed by Hans Geiger and Ernest Marsden in 1909 under the supervision of Rutherford

  8. Evidence of the Nuclear model • The expected path for the “Plum pudding model” would be for an Alpha particle (He nucleus) to travel straight through or get absorbed • In reality some a-particles where reflected back • Rutherford said of the results "It was almost as incredible as if you fired a fifteen-inch shell at a piece of tissue paper and it came back and hit you". • Click here for a picture

  9. 6.1.3 • Outline evidence for the existence of atomic energy levels. • Students should be familiar with emission and absorption spectra, but the details of atomic models are not required. • Take a trip back to the Bohr model

  10. If I only had an orbit • Imagine you are at the science museum and you see that yellow funnel thing, the spiral wishing well. • You know, the one you drop the penny in and watch it swirl around until in goes into the hole at the center • That is the fundamental problem with the Rutherford model of the atom, when the electron give off energy the loose momentum and thus should spiral into the nucleus

  11. Bohr’s Solution • Bohr solved this by “allowing” the electrons to have discrete energies instead of continuous energies • The electron loses energy when it “drops” to a lower energy state • The emitted energy is the difference between the two states • This is evident in the emission and absorption spectra

  12. Energy Levels • When an electron goes down in energy, the excess energy is released as a photon • In order to go up in energy, a photon must be absorbed • The energy associated with the transition is E = hf • h = 6.63 x 10-34 Js Photon out Photon in

  13. Emission and absorption • Since certain wavelengths with be absorbed or emitted, a given atom will only produce or take in certain frequencies of light • Here are both spectrums for Hydrogen

  14. 6.1.4 Isotopes and neutrons • We already know there are protons and neutrons (nucleons) in the nucleus of the atom, but how do we know this? • The discovery of isotopes, and some very ingenious experiments brought the existence of the neutron to light

  15. 6.1.5 Explanation of terms • Nuclide - This name is given to a nucleus with a specific number of protons and neutrons • Here, A is the total number of nucleons, Z is the number of proton, and X is the atomic symbol • A nuclide is represented the following way:

  16. Notation • This notation can be applied to proton, neutron, and electron as well • Isotopes are easy to identify using this notation • Common isotopes would be Carbon 14, Hydrogen 3, and Oxygen 18

  17. Isotopes Hydrogen Deuterium Tritium • An isotope will have the same number of protons and a different number of neutrons • Isotopes will have the same chemical properties, but their physical properties will be slightly different

  18. 6.1.6 Atomic mass & number • Z denotes the number of protons in a nucleus, also called the atomic number • A denotes the total number of nucleons for an atom, so (Z + N = A) • N is the number of neutrons • If the atomic mass number changes the element changes

  19. 6.1.7 More forces • Obviously something has to hold together the nucleus • We learn in 5.1.7 positive charges repel each other, yet the protons in the nucleus are “glued” together • This “glue” comes from the strong nuclear force, which works at very short distances (10-15 m or less) • The atomic radius is written as R = 1.2 x A1/3 x 10-15, where A is the mass number.

  20. Weak nuclear force • Since Neutrons and protons are made up of quarks there must be a force holding the quarks together as well • This force is the Weak nuclear force • A neutron will decay into a proton, an electron, and an anti-neutrino • This is called Beta decay • Click here to see more