Option B 13.1.8-13 Atomic Spectra and atomic energy states

1. Option B13.1.8-13 Atomic Spectra and atomic energy states These notes were typed in association with Physics for use with the IB Diploma Programme by Michael Dickinson For further reading and explanation see: Physics, Tsokos (purple): Ch6.4 Physics, Giancoli (mountain): Ch27

2. 13.1.8 Outline a laboratory procedure for producing and observing atomic spectra. • See notes on section 7.1.4!!! • Easy we’ve already talked about this.

3. 7.1.3 Outline one limitation of the simple model of the nuclear atom.7.1.4 Outline evidence for the existence of atomic energy levels. • If low-pressure gases are heated or current is passed through them they glow. • Different colors correspond to their wavelengths. • Visible spectrum 400nm(violet) to 750nm(red)

4. 7.1.3 Outline one limitation of the simple model of the nuclear atom.7.1.4 Outline evidence for the existence of atomic energy levels. • When single element gases such as hydrogen and helium are excited only specific wave lengths were emitted. • These are called emission line spectra

5. 7.1.3 Outline one limitation of the simple model of the nuclear atom.7.1.4 Outline evidence for the existence of atomic energy levels. • If white light is pass through the gas the emerging light will show dark bands called absorption lines. • They correspond to the emission lines.

6. 13.1.9 Explain how atomic spectra provide evidence for quantization of energy in atoms. Overall • Every time a spectra is created the exact same lines(frequencies) are indicated. • There is NEVER any activity in between those empty spaces (frequencies) • It showed that energy is quantized. It can never be a partial amount.

7. 13.1.0 Calculate wavelengths of spectral lines from energy level differences and visa versa. • Suppose an electron becomes excited(gains energy) and jumps up an orbital level. • Another will drop down and fill the void. Goes from high energy to low energy. • The difference between the high and low energy levels is emitted in the form of electromagnetic radiation.

8. 13.1.10 Calculate wavelengths of spectral lines from energy level differences and visa versa. • The “falling” electron could fall from 2nd, 3rd, 4th, … energy level. • The bigger the “fall” the more energy released as electromagnetic radiation. • The energy emitted is proportional to the frequency of the e.m. radiation. So that means that each energy level difference corresponds to a discrete and specific frequency, meaning different wavelengths too.

9. 13.1.10 Calculate wavelengths of spectral lines from energy level differences and visa versa.

10. 13.1.10 Calculate wavelengths of spectral lines from energy level differences and visa versa. • Electrons on the very outside have no electrostatic forces on them and are called free electrons with an energy of zero. • As it drops, gives off energy and its own energy levels lower to negative values

11. 13.1.10 Calculate wavelengths of spectral lines from energy level differences and visa versa. • Electrons on the very outside have no electrostatic forces on them and are called free electrons with an energy of zero. • As it drops, gives off energy and its own energy levels lower to negative values • So Eemitted = Einitial - Efinal

12. 13.1.10 Calculate wavelengths of spectral lines from energy level differences and visa versa. Example A • An electron drops from n=4 to n=2.What frequency does that correspond to? What wavelength does that correspond to? • Eemitted = Einitial - Efinal Answer f= 6.15 x 1014 Hz Answer λ= 4.88 x 10-7m • This is the second visible line on the emission spectrum for atomic Hydrogen.

13. 13.1.10 Calculate wavelengths of spectral lines from energy level differences and visa versa. Example A • An electron drops from n=4 to n=2.What frequency does that correspond to? What wavelength does that correspond to? • Eemitted = Einitial– Efinal Solution Guide f = E / h E = 2.55 (1.6x10-19) v(c) = f λ

14. 13.1.11 Explain the origin of atomic energy levels in terms of the “electron in a box” model. • An electron is trapped in a box. • If it behaves like a wave then it can create a standing wave. • You might have to adjust the length but that’s ok. • You can then change the frequencies of the electron and create 2, 3 and 4 loop waves. • Each loop then corresponds to a energy level • n = 1 (L = ½ λ) • n = 2 (L = λ) • n = 3 (L = 3/2 λ)

15. 13.1.11 Explain the origin of atomic energy levels in terms of the “electron in a box” model. IB Equation • EK = (n2h2) / (8meL2) n = energy level h = planck’s constant m = mass of electron L = orbital circumference

16. 13.1.12 Outline the Schrodinger model of the hydrogen atom. • De Broglie thought that particles have specific wavelengths and that electrons standing waves have specific boundary (quantum). • Schrodinger associated the electrons standing wave to a wave function,Ψ. • The wave function, Ψ, indicated the probability of finding the electron in a particular location. • High Ψ places the electron close the and orbit • Low Ψ places the electron between orbits. (regions NOT allowed in Bohr’s model) • Schrodinger model looks at the PROBABILITY of finding an electron. It doesn’t treat is as a definite particle. • When this model is used to plot the probability of finding an electron at a distance from the nucleus you get a spherical shape. • This is now called the electron cloud or orbital.

18. 13.1.12 Outline the Schrodinger model of the hydrogen atom. Features of the Schrodinger model • Electrons can be described as a wave • The “electron wave” in the atom has to fit boundary conditions • Only certain wavelengths are allowed, set by the boundary conditions of the standing waves • The wavelength of the electron determines its energy • An electron’s position is undefined • Wave function determines probability of finding a particle

19. 13.1.12 Outline the Schrodinger model of the hydrogen atom. Why was the Schrodinger model an improvement on the Bohr model? • It does not have well defined orbits for the electrons • It does not treat the electron as a localized particle • It assigns a probability wave function to an electron • It was not limited to atomic hydrogen, but could be used to predict the electron probability cloud for any atom. • It can predict the relative intensities of various spectral lines Schrodinger’s Cat http://www.youtube.com/watch?v=IOYyCHGWJq4&list=PL80C5AF536A5A90DF&safety_mode=true&persist_safety_mode=1

20. 13.1.13 Outline the Heisenberg uncertainty principle with regard to position-momentum and time-energy. Heisenberg Uncertainty Theory • Heisenberg built on Schrodinger’s ideas • Investigated the certainty of finding the position and momentum of an electron at a specific moment in time. • To try and find or “see” an electron, you must shine some light on it. The light is then reflected toward you or a piece of lab equipment. • The problem is that the light you was shown in the electron carries energy with it. This means a simply viewing the electron you have changed it’s energy making even more uncertain where it is and what’s its momentum.

21. 13.1.13 Outline the Heisenberg uncertainty principle with regard to position-momentum and time-energy. • Let’s get specific… • Using de Broglie’s wave equation, λ= h / p we can find the exact momentum. • But that means we are now looking at wave that is identical along its entire length. • We have no clue where on the wave it is!?!?!? • The more information we have about it’s momentum the less we know about it’s position and visa versa. IB Definition Heisenberg Uncertainty Principle– It is impossible to simultaneously measure both the position and momentum of a particle without a degree of uncertainty. http://www.youtube.com/watch?v=7vc-Uvp3vwg&list=PL80C5AF536A5A90DF&safety_mode=true&persist_safety_mode=1

22. Review Video Part 1 • http://www.youtube.com/watch?v=1W55RanmJS8&list=PL80C5AF536A5A90DF&safety_mode=true&persist_safety_mode=1 Part 2 • http://www.youtube.com/watch?v=ZmEvF2WetnE&list=PL80C5AF536A5A90DF&safety_mode=true&persist_safety_mode=1