1. Concept #4 “Electrons in the Atom” Honors Chemistry 1

2. The story continues, Rutherford discovered the nucleus and proposed the “plum-pudding” model of the atom. Scientists knew that the model was flawed and worked on improving it. Their investigations shed “light” on the problem.

3. Different elements give different colors.

4. EM Waves Electromagnetic radiation is a form of energy that exhibits wavelike behavior as it travels through space. Together, all the forms of electromagnetic radiation form the electromagnetic spectrum.

6. The Nature of Light Electromagnetic radiation, including light, is described using wave terms. Wavelength ( ) Amplitude Crest Trough origin

7. Frequency and Wavelength Frequency ( ) is the number of waves that pass a point in a given amount of time. Hertz (Hz) are the units and 1 Hz is one wavelength/second. Wavelength (  is the distance from corresponding points on adjacent waves. Light has a speed of 3.00 x 10 8 m /s (c). For light, c = , this also applies to all electromagnetic radiation.  and are inversely proportional.

8. Particle Nature of Light The photoelectric effect refers to the emission of electrons from a metal when light shines on the metal. This phenomena did not fit with the wave description of light.

10. Light can be a particle also. Max Planck explained a “problem” that scientists had with glowing hot objects by suggesting that light could be a particle. Matter could lose or gain energy in small packets of energy called “quanta”. A quantum is the smallest amount of energy that an atom can lose or gain.

11. Planck’s Equation By applying a constant (Planck’s constant) Planck found the relationship between the energy of a photon (a quantum of light) was: E = h, where h is Planck’s constant and E is the energy measured in Joules. h = 6.626 x 10 -34 J.s

12. Planck’s Equation A photon is a particle of electromagnetic radiation having zero mass and carrying a quantum of energy. The energy of a particular photon depends on the frequency of the radiation. E photon = hv

13. Rainbows are “continuous spectra”. The refraction and reflection of sunlight by water droplets produces light in the full or continuous spectrum.

14. Atomic Emission Spectra Because energy must be released only in certain size packages, the light given off by exciting gases can be resolved into a series of lines with only certain colors of light. Each acts like a “fingerprint”.

15. Different gases produce different lights.

16. The Bohr Model Niels Bohr attempted to “fix” Rutherford’s model by having the electrons move in energy levels around the nucleus. Each energy level could only hold a certain number of electrons. The larger the energy level, the further from the nucleus it was.

17. Bohr Model of the Atom 1234 e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- The numbers represent the principal quantum numbers.

18. Ground State Electrons When an electron is in the lowest energy level possible it is said to be in the ground state. When an electron doesn’t occupy the lowest energy level possible due to outside energy it is said to be in the excited state.

19. Changing Energy Levels When ever an electron moves to a higher or lower energy level an energy change is required. If the right amount of energy is added, the electron moves up. To move down, a certain amount of energy must be released.

20. The Quantum Mechanical Model Louis De Broglie combined Planck’s equation E=h and Einstein’s E=mc 2 to produce an equation that explains why the electrons can only occupy certain energy levels. = h /m

21. Heisenberg Werner Heisenberg proposed that we could never tell the position and momentum (speed) of an electron at the same time. This uncertainty is the “Heisenberg Uncertainty Principle”.

22. Schrodinger Erwin Schrodinger derived a formula that described the space that an electron of a certain energy would occupy. Instead of an orbit, he described an “electron cloud” where you would have the greatest probability finding the electron.

23. Four Quantum numbers for electrons in an atom. The principal quantum number (n) describes the size and energy of the electron orbital. Angular momentum quantum number ( l ) describe the shape of orbitals. The number of sublevels = n The magnetic quantum number(m) describes orientation of the sublevels. Spin quantum number (s) refers to how an electron spins.

24. Shapes of Orbitals

25. Quantum numbers

27. Electron configuration The arrangement of electrons in an atom is known as the atom’s electron configuration. The lowest energy levels are filled first, this is the “Aufbau Principle”. Once an energy level is full electrons can then fill the next highest energy level. “Pauli Exclusion Principle” states that no 2 electrons in the same atom can have the same 4 quantum numbers.

28. Electron Sublevels s sublevels can have up to 2 e - and is spherical. p sublevels can have up to 6 e - and looks like 3 “dumbbells” in the x, y and z axis d sublevels can have up to 10 e - and the shape is more complicated f sublevels can have up to 14 e -.

29. p and d orbitals

30. Electron orbitals Electron orbitals can contain 0, 1 or 2 e -. s has 1 orbital, p has 3, d has 5, and f has 7. We show orbitals as a box with electrons represented as arrows. Spin is represented by the direction of the arrows.

31. Hund’s Rule Electrons will seek an unoccupied orbital within a sublevel before it will pair up with another electron in an orbital. NOYES P sublevel

33. Orbital Filling The e - must fill the lowest energy level first and only 2 e - can be placed per orbital. The first e - goes into the lowest principle energy level, n=1. Level 1 only has 1 sublevel, s An s sublevel can only have 1 orbital

34. Orbital filling continued When the first energy level is filled, the e- can occupy the second energy level, n=2. Level 2 can have 2 sublevels, s and p After sublevel s is filled then p is filled according to Hund’s rule.

35. The Flow Chart of Science Electron filling follows the following pattern: 1s 2sp 3sp 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p Notice that each p is followed by the next higher level’s s.

36. Electron configuration An easy way to represent the e - in an atom is through electron configuration. 2 p 5 The principle energy level is 2 and there are 5 e - in the p sublevel.

37. Example: The electron configuration for the sulfur (Z= 16) atom is: 1s 2 2s 2 p 6 3s 2 p 4 This says that the total number of e - is 16 and the highest energy level is n=3.

38. Valence Electrons The chemical characteristics of atoms are based upon the number of e - that are in the outer energy levels. These are called valance electrons. They are only s and p sublevel e -. For the sulfur atom, 1s 2 2s 2 p 6 3s 2 p 4, there are 6 valence e -.

39. Electron Configuration Noble-Gas Notation The Group 18 elements (helium, neon, argon, krypton, xenon, and radon) are called the noble gases. A noble-gas configuration refers to an outer main energy level occupied, in most cases, by eight electrons.