Rutherford s Gold Foil Experiment. Quantum Theory Max Planck (1910)

Transcription

1

2 Neils Bohr

3 Niels Bohr (1913) developed the planetary model of the atom based upon the following: Rutherford s Gold Foil Experiment E = mc 2 Albert Einstein (1905) Quantum Theory Max Planck (1910)

4 He postulated that the electrons were in specific orbits about the nucleus. That the electrons were spinning so that they would not crash into the nucleus. And he knew the model was very limited and that it was going to be modified as soon as he wrote it down!

5 Bohr stated that the light must be from energy given off from the element Different colors of light must be different energy level transitions This means an element has specific energy level transitions that it can give off light

6 Light can have only discrete amounts of energy Energy is quantized (fixed levels like the steps of a ladder or shelves of a bookshelf) Electrons can have only these values and no others Similar to books on a shelf Can be on the first shelf or the second shelf, but not in between

7

8 Electrons prefer to be in the lowest energy level levels closest to the nucleus Ground state Excited state electron goes from the lowest energy level to a higher energy level when it absorbs energy

9 Excited State Ground State

10 Electrons cannot just jump to a higher state for no reason Something has to make them do it, otherwise they d stay at the ground state If energy is put into the atom, the electron can take that energy and jump to another level This taking in of energy causes the absorption spectra, the releasing of energy causes emission spectra

11 Bohr s idea of the atom worked well for hydrogen Any other gas this was attempted with, the spectra didn t look like they should have Needed something better

12 Neils Bohr I pictured electrons orbiting the nucleus much like planets orbiting the sun. But I was wrong! They re more like bees around a hive. WRONG!!!

13 Rutherford said very little about them Neils Bohr said a lot! But we need to cover more before we get to the Bohr Atom! So. Back to Physics!

14 E = H Equation for probability of an electron being found within a region of space Erwin Schrodinger

15 Schrödinger s model: probability of finding electron in a given volume Orbitals Electron clouds Different shapes for different types of orbitals

16 An orbital is a region within an atom where there is a probability of finding an electron. This is a probability diagram for the s orbital in the first energy level Orbital shapes are defined as the volume that contains 90% of the total electron probability. There are 4 Types of Orbitals, named s, p, d & f

17 The s orbital has a spherical shape centered around the origin of the three axes in space. s orbital shape

18 There are three dumbbell-shaped p orbitals in each energy level above the first, each assigned to its own axis (x, y and z) in space.

19 Things get a bit more complicated with the five d orbitals that are found in the d sublevels beginning with the third energy level. To remember the shapes, think of double dumbbells and a dumbbell with a donut! d orbital shapes

20

21 We know where we might find the electron, but.. Once we find it, it moves! Ok anything else? What really matters to the Chemist? As it happens we are interested in the Energy of the electron, not where it is.

22 One cannot simultaneously determine both the position and momentum of an electron. Werner Heisenberg You can find out where the electron is, but not where it is going. OR You can find out where the electron is going, but not where it is!

23 Since Heisenberg demonstrated that you cannot know both the energy and the position of the electron, Chemists concentrate on the energy of the electron and according to Bohr That means we need to know the energy level the electron occupies.

24 This gives rise to: Electron Configurations or Orbital Notation

25 Aufbau Principle - The electron that distinguishes an element from the previous element enters the lowest energy atomic orbital available. Or: electrons fill up the orbitals from the bottom up lowest energy to highest energy

26 Orbital Notation for carbon 1s 2s 2p Electron configuration for carbon element #6 C - 1s 2 2s 2 2p 2 1s 2s 2p

27 Electrons fill sublevels of an orbital singly before they spin pair. An Orbital can hold a maximum of 2 electrons but those electrons must have opposite spins. 1s 2s 2p Nitrogen

28 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f

29 Get out your Periodic Table! Determine the energy levels used Determine the orbital type Determine the number of electrons in each orbital Continue to fill each higher level until all electrons are accounted for

30 The Orbitals Being Filled for Elements in Various Parts of the Periodic Table

31 Modern View of Atom From past to present