1. Chapter 5 Light Electrons in atoms

2. Models of the atom Rutherford’s model of the atom did not show or explain chemical properties of elements Needed some modifications Along comes Bohr….

3. THE BOHR MODEL A Danish physicist, Niels Bohr, proposed that an electron is found only in specific circular paths, or orbits, around the nucleus  This is a combination of Rutherford’s model and Einstein’s famous E=mc 2 Each orbital has a fixed energy, known as energy levels  These levels are similar to rungs on a ladder. The lowest rung has the least amount of energy. The highest rung has the most energy

4. A quantum of energy is the amount of energy needed to move an electron from one orbital to another  When electrons jump from one orbital to another, they release light of different frequencies  These different frequencies result in different colors of light

5. Light behaves two ways: 1. As a wave 2. As a particle

6. LIGHT: What Is It? Light Energy Atoms As atoms absorb energy, electrons jump out to a higher energy level. Electrons release photons (bundles of energy released) when falling down to the lower energy level.

7. Light Waves Light pieces are called = PHOTONS Light travels as a wave and has a QUANTUM of energy Wavelength (λ)= distance from crest to crest, meters crest trough λ Frequency (ν or f)= # of waves that pass a point in 1 sec., Hertz or Hz = 1/sec ν Speed of Light (c)= 3.00 x 10 8 m/sec c = λ f

8. ELECTROMAGNETIC SPECTRUM

9. Light as a Particle* photoelectric effect Emission of electrons from a metal when light is shined onto the metal. This helped us understand it also acts as a particle!

11. Light Energy Quantum = minimum amount of Energy (in joules) that can be lost or gained by an atom E=energy (J); f=frequency (1/sec, or Hz) E = h f Planck’s constant = h = 6.626 x 10 -34 J sec

12. Light Spectra Ground State – lowest energy state of an atom Excited State – an atom in a higher energy state Line Emission Spectrum –is the spectrum of frequencies of electromagnetic radiation emitted by atoms or molecules when they are returned to their ground state. Each element's emission spectrum is unique.

13. Line Emission Spectrum

14. Quantum Model Notes We don’t know exactly where the electrons are, but we can have a MODEL showing the most likely place for an electron.

15. Quantum Model- 3D model to identify where electrons exist. ENERGY LEVELS= a possible location for electrons found around the nucleus There are seven (maybe more!) energy levels.

16. SUBLEVEL= For each level, (EXCEPT LEVEL #1!) there can be orbitals of different shapes, called sublevels. We name these sublevels s,p,d,f

17. Long Periodic Table

18. Orbitals a single, allowed location for the atomic electron. Each orbital can hold at most 2 electrons can be empty or only half-full

19. Sublevel# OrbitalsTotal # of e- (# of orbitals*2) s1 2 p3 6 d5 10 f7 14

20. 3.What does the quantum mechanical model determine about electrons in atoms? 4.How do 2 sublevels of the same principal energy level differ from each other? 5.How can electrons move from one energy level to another? 6.How many orbitals are in the following sublevels? 1.3p 2.3d 3.2s 4.4f 5.4p

21. Aufbau Principle: Electrons fill the lower energy levels first Example: Climbing a staircase. You have to start at the bottom and move your way up Pauli Exclusion Principle: Each orbital can hold only 2 electrons of opposite spin Example: Shoes in shoe box

22. QUESTION FOR DISCUSSION Modeling the Pauli Exclusion Principle: How are 2 bar magnets similar to electrons’ spins in the same orbitals?

23. Rules for electron Configurations Hund’s Rule: Electrons fill orbitals of equal energy before sharing orbitals. Ex: In a house, usually all kids bedrooms are filled with one kid before their parents make them share a room.

24. The Energy Levels of an atom Sublevels = s,p,d,forbitals = Energy nucleus n=1 n=2 n=3 n=4 1s 2s 3s 4s 2p 3p 4p 3d 4d4f