2. 1 Mr. ShieldsRegents Chemistry U06 L02

3. 2 H2H2H2H2 We saw that Bohr Was able to equate orbits With Energy levels And secondly he could Then equate energy Levels with specific Wavelengths of light. So how did he arrive at this model? To see how his maodel came about we need to step back a little bit in time and look at some significant Contributions made by other contemporary scientists Development of the Bohr Model

4. 3 The Quantum Max Plank (1858 – 1947) 1900 – Max Plank Studied the light (electromagnetic Radiation) emitted by Heated Objects. He concluded that matter can only Gain or lose energy in small specific multiples. He called the minimum Energy that matter can gain or lose the “QUANTUM”

5. 4 Quantum Max Plank (1858 – 1947) Equantum = h x frequency h = plank’s constant So matter can only emit radiation in multiples of 1(h x freq), or 2(h x freq) etc. Energies between these values DO NOT exist. (Sound familiar? Yep … Bohr’s assumption about where Electrons can reside i.e. quantized orbits) Furthermore, the energy of a quantum Is equal to the frequency of the emitted Radiation times a constant.

6. 5 Photoelectric effect Recall that Frequency is also related to energy E quantum = Plank’s Constant (h) x Frequency This is known as the Light directed at metal will release electrons from the Surface. However, only light having a certain minimum Frequency will do this PHOTOELECTRIC EFFECT eV = unit of Energy (electron volts) e-e- But if light has no mass how does it knock an electron out of the atom?

7. 6 Einstein & the Photoelectric effect In 1905 Einstein explained how this happened. He said Light could behave either as a wave or a particle. As a Particle light is called a PHOTON.

8. 7 Einstein & the Photoelectric effect And, extending Plank’s idea: E quantum = h x freq Einstein equated the photon with the quantum, so: Einstein equated the photon with the quantum, so: E photon = h x freq Therefore E photon = E quantum

9. 8 The Bohr model proposed electrons in specific energy levels. These were designated n=1, 2, 3 etc. Rule 1: The number Of electrons Allowed in Each Orbit is 2n 2 2n 2 2 electrons 8 electrons 18 32

10. 9 Bohr Electron Configurations Rule 2: When electrons are added to atoms they are placed in the lowest energy levels first. - What value of “n” represents the lowest energy level? Look at Neon on the Periodic Table. - How many electrons it have? How do you know? - Into what orbits can I put these electrons? So, Neon’s “electron configuration” is designated 2-8 - Look at the electron configs in your reference table

11. 10 Bohr Electron Configuration Try these – How many electrons does Mg have? 2-8-2 What is the electron configuration? In what orbits are these electrons located? 12 n= 1, 2, and 3 Now answer these same questions for Br …

12. 11 Electrons could only exist in specific orbits.Electrons could only exist in specific orbits. Each orbit has a definite energy. The orbit closest to the nucleus has the lowest energy.Each orbit has a definite energy. The orbit closest to the nucleus has the lowest energy. When electrons absorb energy they moves to higher orbits. The Atom is said to be in aWhen electrons absorb energy they moves to higher orbits. The Atom is said to be in a “EXCITED STATE” “EXCITED STATE” Bohr’s Model stated: When all excited state electrons return to their original energy levels the atom is said to be in the “GROUND STATE” E5 E3 E1 E6 E4 E2

13. 12 NIELS BOHR Electrons can not reside between these orbits. To move from one orbit - to the next e - must absorb or release specific amounts of energy called “quantums“ x

14. 13 Excited State Ground State Remember – a photon is electromagnetic energy that Is behaving as a particle but has no mass and carries a quantum of energy.

15. 14 NOTE Since energy is quantized the energy lost in the transition from 3 to 2 and then 2 to 1 is exactly equal to the energy of transition from 3 to 1 or even 1 to 3.

16. 15 Hydrogen e - Transition Series Balmer: Visible light nx to n=2 Paschen: IR nx to n=3 Lyman: UV nx to n=1

17. 16 When excited electrons drop back to n=2 from n= 6, 5, 4 and 3, photons of an energy equal to that Transition are emitted producing hydrogen’s Emission spectrum. For example the E3 – E2 transition produces the red Emission line at 656.2 nm And the E4 – E2 transition produces the blue-green Emission line at 486.1 nm Hydrogen’s emission spectra

18. 17 Excited State Electron configuration We saw previously that the electron configuration Of Bromine is 2-8-18-7 This is know as the “ground state electron configuration” Why? What energy levels are they in And how many e - are in each ? n = 1 2 3 4

19. 18 Excited State Electron configuration We’ve said that an excited state of an atom is one in which any number of electrons have moved to higher Energy levels. Even orbits not previously occupied by electrons! Let’s look at some examples: The ground state e- config of Bromine is 2-8-18-7 One possible excited state electron config might be 2-7-18-8 ( (from which orbit was the e- excited and to what orbit did it go ?)

20. 19 Excited State Electron configurations Bromines electron config is 2-8-18-7 Can one possible excited state be Sure! All the electrons in any given level can be excited to higher energy levels having 0 or more e - Can one possible excited state for Br be 2-7-19-7? NO! I can’t have more electrons than are allowed in Any given level. 0-8-18-8-1 ?

21. 20 Bohr model - Shortcomings The Bohr model correctly explained the discontinuous lines seen in the Emission spectra and even allowed one To Calculate which lines you should get! It also took care of the problems in Rutherford’s model but… It couldn’t explain the spectra for ANY OTHER ELEMENT! To solve this problem required one more step to be made in the development of the Atomic model – The development of the Quantum Mechanical Concept