Electrons and Quantum Mechanics

1. Electrons and Quantum Mechanics Unit 5

2. Electrons • Rutherford described the dense center of the atom called the nucleus. • But the Electrons spin around the outside of that nucleus. • Provide the chemical properties of the atoms. • Responsible for color and reactivity.

3. Energy • Energy is transmitted from one place to another. • Light carries this energy. • Converted into heat. • Light is called Electromagnetic Radiation.

4. Electromagnetic Spectrum • Radio • Infrared • Visible Light • ROY G BIV • Ultraviolet • X Rays • Gamma Rays

5. Light • Light travels as a wave. • Wave Properties • Wavelength (λ) = distance between two waves (m) • Frequency (f) = number of peaks per second (Hz) • Speed of Light (c) = how fast light moves.

6. Light • Light Equation c= ƒλ • Speed of light is a constant = 3 x 108m/s • Nothing travel faster than the speed of light! • Maybe?!?!?!?!?!?!?!?!?

7. Light • The Dual Nature of Light • Light carries energy through space like a wave. • Light also behaves like a particle?!? • A beam of light is made of tiny packets of energy called PHOTONS! • Which travel in waves!?!

8. Light • The Energy of a photon depends on its frequency. • So is the color of light!!! E = hƒ • ELECTRONS are like photons! • Act as waves and particles. • Orbit the nucleus in a wave-like motion.

9. Blackbody Radiation • Rutherford could never explain why objects change colors when they are heated. • As the object heats, it must give off electrons of certain frequencies and energies.

10. Photoelectric Effect • Similarly, light on a metal object can knock off electrons. • Shine different colors on a metal. • Measure the number of electrons knocked off. • Found that no electrons were knocked off below a certain frequency.

11. The Bohr Model • Proposed the electrons orbit the nucleus with fixed energies. • Called Energy Levels • Much like the rungs of a ladder. • Quantum describes the amount of energy required to move an electron from one level to another.

12. The Bohr Model • Ground State • Lowest possible energy of an electron. • Normal location • Excited State • If electron absorbs energy, it moves up an energy level (absorption) • If an electron gives off energy, it moves down an energy level (emission).

13. The Bohr Model

14. Atomic Spectra • Hydrogen Atom Line Emission Spectrum • Expected continuous spectrum of light • But only specific frequencies were given off. • Red (656.6 nm) • Blue-green (486.1 nm) • Violet (434.1 nm) • Violet (419.2 nm)

15. Atomic Spectra • Shine a light on an Atom • When atoms absorb energy, electrons move to higher energy levels. • When atoms release the energy, electrons return to the lower energy level. • Atomic Spectra • Frequencies of light emitted by a certain element. • No two elements have the same spectrum. http://student.fizika.org/~nnctc/spectra.htm

16. Flame Tests • Because no two atoms produce the same spectrum, elements can be identified by the colors they emit. • Spectral Analysis uses this properties to identify elements.

17. Quantum Mechanics • Max Planck (1900) • Founder of Quantum Mechanics E = hf • Albert Einstein (1905) • Wave-Particle Duality • Electrons are small particles that move like waves.

18. Quantum Mechanics • Neils Bohr (1922) • Electrons orbit in distinct energy levels. • Louis de Brogelie (1923) • Wave Mechanics says that ALL MATTER behaves like waves. mv/λ = h

19. Quantum Mechanics • Werner Heisenberg (1927) • Principle of Indeterminacy • You can’t know both the position and the velocity of an electron. • Erwin Schrödinger (1930) • Used wave mechanics to show the PROBABLE location of an electron. • Electrons exist in 3D clouds of probability!!!

21. Quantum Mechanical Model • Uses Schrodinger’s equation to predict the probable location of an electron. • Determines the energies an electron is allowed to have. • Determines how likely it is to find the electron in various locations around the nucleus.

22. Quantum Numbers • Describes the location and behavior of an electron • Like an electron’s address • No two electrons can have the same quantum numbers. • Four Numbers

23. Quantum Numbers • Principle (1st) Quantum Number (n) • The Energy Level • Describes the size of the cloud and the distance of the cloud from the nucleus. • Shows the number of electrons n = 1 = 2 electrons n = 2 = 8 e- n = 3 = 18 e- n = 4 = 32 e-

25. Quantum Numbers • 2nd Quantum Number (l) • Each energy level has sublevels. • The number of sublevels equals n. • Sublevels are called: s = spherical p= peanut-shaped d = daisy-shaped f = unknown?

27. Quantum Numbers • 3rd Quantum Number (ml) • Divides sublevels into orbitals. • Tells the shape the electron moves in. • Number of orbitals = n2 • Examples s = 1 orbital p = 3 orbitals d = 5 orbitals f = 7 orbitals

28. Quantum Numbers • 4th Quantum Number (ms) • Describes the electron’s spin. • Only two electrons fit in an orbital. • Their charges repel causing them to spin in opposite directions (+½ or –½) • Use up and down arrows.

30. Quantum Numbers • Pauli Exclusion Principle • No two electrons can have the same set of 4 quantum numbers. • The electrons repel each other. • Hund’s Rule • Every orbital must get one electron before doubling up.

31. Quantum Numbers 1s • Diagonal Rule • Electrons fill orbitals in predictable patterns • Some People Do Forget • Electrons dill the lowest energy level possible. 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f

32. Orbital Notation • Draw out the locations of each electron in an atom with arrows.

33. Electron Configuration • Write out the configurations of electrons using superscripts. • Examples: • H = 1s1 • He = 1s2

34. Electron Configurations • Noble Gas Shorthand • Write the Noble Gas just before the element. • Add the remainder of the configuration.

35. Lewis Dot Diagrams • A way to show the number and position of the valence electrons. • Outermost energy level • Look at the column number to get this number. • Use the chemical symbol and number of valence electrons. • All four sides must have a dot before you double up. p orbitals s orbital X p1 s p2 p3