1. Quantum Mechanics Directly observing electrons in the atom is impossible, the electron is so small that observing it changes its behavior The quantum-mechanical model explains the manner electrons exist and behave in atoms The model also helps us understand and predict the properties of atoms that are directly related to the behavior of the electrons

2. Why study EMR? Scientists determined in the early 20 th century that the position and momentum of an electron can be accurately described by wave equations

3. The Nature of Light The electromagnetic radiation composed of perpendicular oscillating waves, one for the electric field and one for the magnetic field All electromagnetic waves move through space at the same constant speed of 3.00  10 8 m/s. This velocity is represented by ‘c’.

4. The Wavelength The peak-to-peak distance is called the wavelength. The wavelength is represented by the symbol . Wavelength is usually expressed in units of meters, centimeters or nanometers (1 nm = 10 – 9 m) (also Angstroms=1x10 -10 m)

5. Blue light (high energy) has shorter wavelengths and red light has longer wavelengths (low energy) The Energy and the Wavelength

6. The Frequency Frequency ( ) is the number of waves that pass any reference point per unit of time or waves/time. Frequency is measured in hertz (Hz). 1 Hz = 1 s -1 The frequency can be calculated from the velocity and the and the wavelength of the radiation.

7. Different colors are attributed to the difference in wavelengths of light, and the magnitude of brightness is directly proportional to the amplitude.

8. The Amplitude Amplitude is the vertical distance from the midline of a wave to the peak or trough. Amplitude affects the intensity or the brightness of the radiation. http://id.mind.net/~zona/mstm/physics/waves/introduction/introductionWaves.html

9. The color of light is determined by its wavelength or frequency. When an object absorbs some of the wavelengths of white light while reflecting others, it appears colored, the observed color is predominantly the colors reflected. An object appears red because it is predominantly reflecting red light while absorbing most other colors.

10. The Electromagnetic spectrum The electromagnetic spectrum is divided into regions according to the wavelengths or the frequency of the radiation.

11. The Nature of Matter End of the 19 th century---Matter was composed of particles and Energy was composed of waves…right?? BUT There was a problem that classical physics couldn’t explain- that a glowing hot object doesn’t emit UV radiation as expected……

12. In 1900 the German scientist Max Planck proposed that the electromagnetic radiation could be viewed as a stream of tiny energy packets or quanta we now call photons. He proposed that there is a minimum amount of energy that can be gained or lost by an atom The Photons Max Plank (1858 – 1947)

13. The Photoelectric effect Another dilemma in classical physics….

14. The Photoelectric Effect It was observed that many metals emit electrons when a light shines on their surface, this effect is called the Photoelectric Effect.

15. The classic wave theory attributed the photoelectric effect to the light energy being transferred to the electron. According to this theory, if the wavelength of light is made shorter (higher energy) more electrons should be ejected If a dim light was used there would be a lag time before electrons were emitted In experiments with the photoelectric effect, it was observed that there was a maximum wavelength for electrons to be emitted called the threshold frequency (regardless of the intensity). It was also observed that high frequency light with a dim source caused electron emission without any lag time

16. Einstein (1879 – 1955) Einstein to the rescue…….. Explaimed Photoelectric Effect-- Radiant energy striking the metal surface behaves not as a wave but as a stream of tiny packets of energy.

17. Einstein (1879 – 1955) The Energy of the Electromagnetic Radiation Planck proposed, and Einstein confirmed, that the energy of a photon is proportional to its frequency. E = h Where is h is the Plank’s constant and it equals = 6.626 × 10 – 34 J s This means that both electrons and electromagnetic radiation can be represented as either waves (E) or particles (h ).

18. Ground State and The Excited State Atoms when exposed to electromagnetic radiation, they emit the absorbed energy in the form of light as electrons return to a lower state.

19. Spectrum of ordinary white light The visible spectrum of white light is called a continuous spectrum, because it contains continuous distribution of all colors.

20. The origin of atomic line spectra is the movement of electrons between quantized energy levels The visible line spectrum of excited hydrogen atoms consists of four lines, from indigo at 410 nm to red at 656 nm (not a continuous spectrum)> Each of these wavelengths represents a specific energy transition as excited electrons move from excited states to lower energy states

21. Oxygen spectrum More examples of atomic spectra Neon spectrum

22. The Energy Levels Are Quantized Atomic line spectra tell us that when an excited atom loses energy, not just any arbitrary amount can be lost. This is possible if the electron is restricted to certain energy levels. The energy of the electron is said to be quantized. (a) Continuous energy level (b) The energy level are quantized.

23. Connected the spectra of hydrogen, and the quantum ideas of Einstein and Planck, to explain that single electron of hydrogen could occupy only certain energy states The Bohr Model Niels Bohr (1885 -1962) AN electron would remain in its lowest energy state until otherwise disturbed.

24. The first theoretical model that successfully accounted for the Rydberg equation was proposed in 1913 by the Danish physicist Niels Bohr. The Bohr Model Niels Bohr (1885 -1962) Bohr proposed that the electrons moved around the nucleus is fixed paths or orbits much like the planets move around the sun.

25. Neils Bohr proposed that the electrons could only have very specific amounts of energy-- fixed amounts, quantized

26. The electrons traveled in orbits that were a fixed distance from the nucleus therefore the energy of the electron was proportional the distance the orbital was from the nucleus.

27. When energy is added to an electron, it is promoted to a orbit further away from the nucleus

28. Electrons emit radiation when they “jump” from an orbit with higher energy down to an orbit with lower energy. For example they give off violet EMR when jumping from the 5 th energy level to the more stable 2 nd.

31. Bohr’s model for hydrogen atom

32. Bohr explained the line spectrum of hydrogen. The energies of emitted light are equal to the differences between the energy state of an electron jumping from an outer orbits and an inner orbit into which it can move

34. Rydberg equation The Rydberg equation is used to calculate the energy changes when electrons are promoted to higher energy levels and subsequently fall back to the lower energy levels

36. In general, the line spectrum of an element is rather complicated. The line spectrum of hydrogen, with a single electron, is the simplest. The Rydberg equation can be used to calculated energy levels associated with the spectral lines of hydrogen.: E= -2.178 x 10 -18 J Z 2 n 2 Bohr’s calculation of the energy levels of a hydrogen atom  –  energy of a particular energy level n = Energy level The Rydberg constant, R, is an empirical constant with a value of -2.178 x 10 -18 J/atom Z= nuclear charge aka the number of protons

37. Rydberg equation Energy changes during transitions are proportional to (atomic #). This means that if an electron is promoted from for example level 1 to level 5 in a species that has less protons in the nucleus the same transition for a species with more protons would be more difficult. This is because the protons in the nucleus are attracting the electron to the lower energy level and more energy is required to promote them.

38. Calculate the energy of the n=3 state of the H atom in joules per Atom -2.421x10-19J

39. Calculating the delta energy between two quantized orbits

40. Example Problem 1: Calculate the energy involved when an electron transitions from n=1 to n=3. +1.936 x 10 ─18 J NOTE: energy is positive, therefore process is endothermic, energy is required for the transition  transition an absorption process!

41. Calculate the energy involved for an electron to transition from n=5 to n=2 for an atom of hydrogen. Does this represent absorption or emission?

42. Calculate the wavelength of light observed for the transition in the previous problem.

43. The Lyman series of spectral lines for the H atom occurs in the ultraviolet region. They arise from transitions from higher levels to n=1. Calculate the frequency and wavelength of the least energetic line in this series. Answer: E = -Rhc (1/1 2 1/2 2 ) = -2.179x10-18J/atom (1/1 – 1/4) = -1.634x10-18J E = hν ν = 1.634x10-18 / 6.626x10-34Js = 2.466x1015Hz λ = c/ν = 3.00x108m/s / 2.466x1015Hz = 1.216x10-7m = 121.6nm