1. Unit #4 CP Chemistry

2.  Bohr  Quantum Numbers  Quantum Mechanical Model

3.  Bohr worked with the concepts of energy, wavelength and frequency  Each color of light is associated with a different energy  Each atom gives off its own unique color so..  Electrons of different atoms have different energies

4.  Each atom has its own specific electron arrangement  Electrons are in Energy levels in the atom  When an electron goes from an excited state back down to its ground state the atom emits light

5. n = 3 n = 4 n = 2 n = 1

6.  Doesn’t work.  Only works for hydrogen atoms.  Electrons don’t move in circles.  The quantization of energy is right, but not because they are circling like planets.  Back to the drawing board

7.  Valence electrons- the electrons in the outermost energy levels (not d).  Core electrons- the inner electrons  Ground state – all electrons are on the lowest possible energy levels  Excited State – some electrons have more energy than usual, causing some electrons to be in a higher energy level than they should be in

8.  Electromagnetic radiation is a form of energy that exhibits wavelike behavior as it travels through space  Organized on the Electromagnetic Spectrum according to wavelength and frequency

9. Frequency(v) = number of cycles in one second Measured in hertz 1 hz = 1 cycle/second Wavelength ( ) = Length of one wave Measured in unit of distance (m, nm, etc.)

10.  There are many  Different and  Higher Energy = Higher Frequency / Lower Wavelength  Radio waves, microwaves, x rays and gamma rays are all examples.  Visible Light is only the part our eyes can detect. amma Rays Radio waves

12.  As we know, the amount of energy is related to the wavelength  Different wavelengths will show different colors  Each element gives off it own unique set of colors  Therefore each element gives off its own unique amount of energy

13.  Electrons are constantly in motion and give off energy when they move from an excited state to ground state

15.  Planck found energy came in packets  A Packet of energy is called a quantum  A quantum is the minimum amount of energy that can be gained or lost by an atom

16.  Einstein said that light can be viewed as a stream of particles called photons  A Photon is a particle of radiation with zero mass and carrying a quantum of energy

17.  Is energy a wave like light, or a particle?  Yes  Concept is called the Wave -Particle duality.  What about the other way, is matter a wave?  Yes, hence the quantum mechanical model

18.  Each element has a specific number of electrons  In an atom electrons are arranged in a specific arrangement  We know this because each element has its own atomic spectra  Different energy = different colors= different electron arrangement

19.  A totally new approach.  De Broglie said matter could be like a wave, like standing waves.  The vibrations of a stringed instrument.

20.  How we keep track of electrons  Principal quantum number (n) energy level of the electron.  Is a number from 1-7

21.  Angular momentum quantum number (l) gives the shape of the orbital  Has values between 0 and n-1

22. S orbitals l = 0 1 orbital per energy level

23. P orbitals l = 1 3 orbitals per energy level

24. D orbitals l = 2 5 orbitals per energy level

25. F Orbitals l = 3 7 orbitals per energy level

26. F orbitals

27.  Magnetic quantum number (m I )  Takes a guess at what orbital the electron is in  Gives the axis orientation  Tells direction in each shape (x,y,z)  Is a value between – l and + l

28.  Electron spin quantum number (m s )  Can have 2 values.  either +1/2 or -1/2

29. We follow 3 rules to get the correct electron configuration for each atom 1. Aufbau Principle 2. Pauli exclusion Principle 3. Hund’s Rule

30.  Aufbau is German for building up.  As electrons are added to the atom they arrange themselves in orbitals  The orbitals are in order of lowest energy (1s) to the highest energy  The order of the triangle  Fill up in order of energy levels.

31. 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2 electrons 2s 2 4 2p 6 3s 2 12 3p 6 4s 2 20 3d 10 4p 6 5s 2 38 4d 10 5p 6 6s 2 56

32.  Only 2 electrons per orbital  Electrons in the same orbital must have opposite spins  Spin is represented by an arrow

33.  Hund’s Rule- The lowest energy configuration for an atom is the one that has the maximum number of unpaired electrons in the orbital.  C 1s 2 2s 2 2p 2

35. Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f He with 2 electrons