1. Electronic Structure of Atoms Chapter 6 Chemistry 100

2. What is light?  Light is obviously real - it is part of our world. Darkness is the absence of light  Light is NOT a solid, a liquid, or even a gas. So what is it? It is a form of energy  A form of radiant energy because it carries energy through space

3. Electromagnetic radiation  Visible light is a type of electromagnetic radiation  Other types include: infra-red, ultra-violet, X-rays, radar waves, microwaves, radio and TV waves  Electromagnetic radiation has wave like properties

4. Waves  Wavelength (lamda)  Frequency (nu)  Speed c (see?)  = c  c = 3.00  10 8 m/s for all types of electromagnetic radiation.  So how is IR different from UV, for example?

6. Electromagnetic Spectrum  Electromagnetic radiation is characterized by a wave length () and a frequency ()  Frequency: number of cycles (vibrations) per second. Unit is second -1 or s -1 or the Hertz (SI unit for frequency). Hence, 82,000 s -1 is the same as 82 kHz (kiloHertz)

7. Units for wavelength UnitSymbolLength (m)Type of Radiation AngstromÅ 10 -10 X-ray Nanometrenm 10 -9 UV & visible Micrometre  m10 -6 IR Millimetremm 10 -3 IR Centimetrecm 10 -2 microwave Metrem1TV, Radio

8. Max Planck and his constant h  Suggested that energy is quantized - comes in small chunks  E = h where n = 1, 2, 3  Compare the potential energy of a brick on a staircase to one on a slope

9. Can this be true?  We do not find that energy is quantized in everyday life - h is very small. Cannot see the difference between 200,000,000h and 200,000,001 h  Einstein used Planck’s idea to explain the photoelectric effect  For electromagnetic radiation, E = h where is the frequency of the radiation. High frequency  more energy

10. What is light?  Examine how light behaves in experiments with lens, mirrors, etc., we are led to believe that light has wave properties  In the photoelectric effect, light appears to consist of particles - which we call photons  Dual nature of electromagnetic radiation

11. Bohr’s Atom  Bohr said: if energy is quantized then the energy of an electron in an atom is quantized  Radius of its orbit cannot be any arbitrary value  Must obey the quantum theory. Only certain orbits are allowed

12. Allowed Orbitals in Bohr’s Atom The quantity n is a quantum number

13. Bohr’s Atom 1913  Electrons move in orbitals with specified radii  Each orbital is associated with a specific energy  This explains why atoms emit (or absorb) light of well-defined frequency. Examples: the yellow sodium street light and the neon tube.

14. Wave Behaviour  Louis de Broglie (1892-1987) If light can have both wave and particle behaviour, why not wave behaviour for all particles? = h/m  He talked about matter waves

15. Matter waves Find for electron moving at 5.97 10 6 m/s Find for baseball moving at 100 km/h

16. Heisenberg  Postulated that there is a limit to how precisely we can measure both position and momentum  The measurement effects the object being measured  Heisenberg’s Uncertainty Principle

17. Schrödinger’s wave equation  In 1926, Schrödinger put de Broglie’s and Heisenberg’s ideas together and came up with the wave equation The quantity  2 provides information about the electron's position when it has energy E!

18. Quantum Numbers  Schrödinger's wave equation has three quantum numbers. Principal quantum number n. Has integer values 1, 2, 3 Azimuthal quantum number, l. Allowed values values of 0, 1... up to n - 1 Magnetic quantum number, m l. Allowed values -l … 0 … +l  There is also the Spin quantum number, m s. It can have a value of -½ or +½

19. Atomic orbitals  The first shell  n = 1The shell nearest the nucleus  l = 0We call this the s subshell (l = 0)  m = 0There is one orbital in the subshell s = -½The orbital can hold two electrons s = + ½ one with spin “up”, one “down”  No two electrons in an atom can have the same value for the four quantum numbers: Pauli’s Exclusion Principle

20. The second shell n = 2l = 0 or 1There are two subshells l = 1The p subshell m = -1, 0, +1Three orbitals in the subshell s = -½ or + ½ Each orbital can hold 2 electrons. p subshell can hold 6 electrons l = 0 The s subshell m = 0One orbital in the subshell s = -½ or + ½ Subshell can hold two electrons The second shell can hold 8 electrons: 2 in s orbitals and 6 in p orbitals

21. If the principal quantum number is n, the shell can hold up to 2n 2 electrons

22. s Orbitals are Spherical

23. p Orbitals are Dumbbell Shaped

24. d Orbitals are Complex

25. Aufbau Principle 1s2s 2p 3s 3p 4s 3d 4p 5s4d 5p 6s 4f 5d 6p 7s5f 6d 6f

26. Let’s do Sodium, Z = 11  Aufbau Principle 1s 2s 2p 3s ….  First 2 electrons1s 2 that’s 2  Next 2 electrons2s 2 that’s 4  Six this time2p 6 that’s 10  1 more to go3s 1 that’s all, folks Electronic configuration of Na is 1s 2 2s 2 2p 6 3s 1

27. Hund’s Rule The configuration with the maximum spin is more stable. Shall we use 1s 2s 2p (  ) (  ) (  )(  ) Or, shall we use 1s 2s 2p (  ) (  ) (  )(  )(  )

28. Shorthand configurations  The configuration of Neon is: 1s 2 2s 2 2p 6  Na is 1s 2 2s 2 2p 6 3s 1, or in short form: [Ne]3s 1  The configuration of Argon:1s 2 2s 2 2p 6 3s 2 3p 6  K is: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1, which in short form becomes [Ar]4s 1  Note the similarity of the two elements from the same group in the periodic table.The incomplete orbitals are 3s 1 and 4s 1.

29. Same group, similar configuration  Fluorine:[He]2s 2 2p 5  Chlorine:[Ne]3s 2 3p 5  Bromine:[Ar]3d 10 4s 2 4p 5  Iodine:[Kr]4d 10 5s 2 5p 5  The outer-shell configuration in each case is s 2 p 5  We need not be concerned with the d electrons here because d 10 is a filled subshell.

30. Electronic Configuration & Periodic Table

31. I’m in a spin!!!  Nitrogen has Atomic Number 7 2  Electronic Configuration: 1s 2 2s 2 2p 3  Let’s draw an orbital diagram: