1. Atoms: Not to Be Cut

2. Dalton’s Theory He deduced that all elements are composed of atoms. He deduced that all elements are composed of atoms. Atoms are indivisible particles. Atoms are indivisible particles. Atoms of the same element are identical. Atoms of the same element are identical. Atoms of different elements are different. Atoms of different elements are different. Compounds are formed by the joining of atoms of two or more elements. Compounds are formed by the joining of atoms of two or more elements.

3. Thomson Model In 1897, J.J. Thomson proposed that an atom is made of even smaller particles with the “Plum Pudding” model. In 1897, J.J. Thomson proposed that an atom is made of even smaller particles with the “Plum Pudding” model. Atoms were made from a positively charged substance with negatively charged electrons scattered about, like raisins in a pudding. Atoms were made from a positively charged substance with negatively charged electrons scattered about, like raisins in a pudding.

4. In 1908, Ernest Rutherford’s experiment Involved firing a stream of tiny positively charged particles at a thin sheet of gold foil In 1908, Ernest Rutherford’s experiment Involved firing a stream of tiny positively charged particles at a thin sheet of gold foil

5. Rutherford Rutherford reasoned that all of an atom’s positively charged particles were contained in the nucleus. The negatively charged particles were scattered outside the nucleus around the atom’s edge. Rutherford reasoned that all of an atom’s positively charged particles were contained in the nucleus. The negatively charged particles were scattered outside the nucleus around the atom’s edge.

6. Bohr Model In 1913, the Danish scientist Niels Bohr proposed that the electrons in an atom existed at specific energy levels. In 1913, the Danish scientist Niels Bohr proposed that the electrons in an atom existed at specific energy levels.

7. Bohr Model According to Bohr’s atomic model, electrons move in definite orbits around the nucleus, much like planets circle the sun. These orbits, or energy levels, are located at certain distances from the nucleus. According to Bohr’s atomic model, electrons move in definite orbits around the nucleus, much like planets circle the sun. These orbits, or energy levels, are located at certain distances from the nucleus.

8. Bohr Atom Model In this model, the nucleus is orbited by electrons, which are in different energy levels. In this model, the nucleus is orbited by electrons, which are in different energy levels.

9. The Bohr Model of the Atom Bohr’s model of the atom –Quantized energy levels –Electron moves in a circular orbit –Electron jumps between levels by absorbing or emitting photon of a particular wavelength

10. The Energy Levels of Hydrogen The Energy Levels of Hydrogen Atomic states Atomic states Excited state – atom with excess energy Excited state – atom with excess energy Ground state – atom in the lowest possible state. Ground state – atom in the lowest possible state. When an atom absorbs energy from an outside source it enters an excited state. When an atom absorbs energy from an outside source it enters an excited state.

11. A. The Energy Levels of Hydrogen Energy level diagram Energy level diagram Energy in the photon corresponds to the energy used by the atom to get to the excited state.

12. A. The Energy Levels of Hydrogen Only certain types of photons are produced when atoms release energy. Why? Only certain types of photons are produced when atoms release energy. Why?

13. A. The Energy Levels of Hydrogen Quantized Energy Levels Quantized Energy Levels –Since only certain energy changes occur the H atom must contain discrete energy levels.

14. B. The Bohr Model of the Atom Bohr’s model of the atom –Quantized energy levels –Electron moves in a circular orbit –Electron jumps between levels by absorbing or emitting photon of a particular wavelength

15. Electrons circle the nucleus due to the Electric force Bohr’s Picture of the Atom Allowed Orbits 1 2 3 4 5 n = Electron in lowest “allowed” energy level (n=1) Electron in excited state (n=5) Before 1 2 3 4 5 Electron falls to the lowest energy level After Radiated photon Note: There are many more energy levels beyond n=5, they are omitted for simplicity

16. Atomic Radiation The difference in energy,  E, is given by:  E = E 5 – E 1 = h  photon h = Planck’s constant = 6.6x10 -34 [J s] = frequency of light [hz] The energy of the light is DIRECTLY PROPORTIONAL to the frequency,. Recall that the frequency,, is related to the wavelength by: c =  c  So, higher frequency  higher energy  lower wavelength This is why UV radiation browns your skin but visible light does not ! It is now “known” that when an electron is in an “excited state”, it spontaneously decays to a lower-energy stable state. Before n = 1 n = 2 n = 3 n = 4 n = 5 Energy Electron in excited state (higher PE) E5E5 E4E4 E2E2 E3E3 E1E1 E 5 > E 4 > E 3 > E 2 > E 1 After n = 1 n = 2 n = 3 n = 4 n = 5 Energy Electron in lowest state (lower PE) E5E5 E4E4 E2E2 E3E3 E1E1 One example could be:

17. Light has a dual nature ___________________________ Wave (electromagnetic) - Interference Wave (electromagnetic) - Interference - Diffraction - Diffraction Particle (photons) - Photoelectric effect Particle (photons) - Photoelectric effect - Compton effect - Compton effect Wave - Particle Duality for light

18. What about Matter? _______________________________ If light, which was traditionally understood as a wave also turns out to have a particle nature, might matter, which is traditionally understood as particles, also have a wave nature? If light, which was traditionally understood as a wave also turns out to have a particle nature, might matter, which is traditionally understood as particles, also have a wave nature?Yes!

19. Louis de Broglie’s hypothesis ____________________________ The dual nature of matter In 1924, Louis de Broglie proposed that matter particles have wave nature. In 1924, Louis de Broglie proposed that matter particles have wave nature. A particle with momentum p has a matter wave associated with it, whose wavelength is given by A particle with momentum p has a matter wave associated with it, whose wavelength is given by λ = h/p = h/mv λ = h/p = h/mv

20. The connecting link – Planck’s constant _______________________________ Dual Nature Radiation E = hf Radiation E = hf Matter λ = h/p Matter λ = h/p

21. Matter Waves ? Particles, like photons, also have a wavelength given by: = h/p = h / mv That is, the wavelength of a particle depends on its momentum, just like a photon! The main difference is that matter particles have mass, and photons don’t !

22. Why isn’t the wave nature of matter more apparent to us…? ___________________________________ Planck’s constant is so small that we don’t observe the wave behaviour of ordinary objects – their de Broglie wavelengths could be many orders of magnitude smaller than the size of a nucleus!

23. Real photographs of an electron interference pattern… 70 [s] 70,000 electrons Notice the clear interference fringes. Clear indication of wave phenomenon.

24. Are matter waves for real?! __________________________________  In 1927 Davisson and Germer showed that electrons can diffract – they act like waves  Big application – Electron Microscopes

25. Matter Waves (cont) Compute the wavelength of a 1 [kg] block moving at 1000 [m/s]. = h/mv = 6.6x10 -34 [J s] / (1 [kg])(1000 [m/s]) = 6.6x10 -37 [m]. This is immeasureably small  For ordinary “everyday objects”, we don’t experience that matter can behave as a wave.

26. But, what about small particles ? Compute the wavelength of an electron (m = 9.1x10 -31 [kg]) moving at 1x10 7 [m/s]. = h/mv = 6.6x10 -34 [J s]/(9.1x10 -31 [kg])(1x10 7 [m/s]) = 7.3x10 -11 [m]. = 0.073 [nm] These electrons have a wavelength in the region of X-rays

27. How do we see ? Light reflects (scatters) from a surface and reaches our eye. Our eye forms an image of the object.

28. Wavelength versus Size Even with a visible light microscope, we are limited to being able to resolve objects which are at least about 10 -6 [m] = 1 [  m] = 1000 [nm] in size. This is because visible light, with a wavelength of ~500 [nm] cannot resolve objects whose size is smaller than it’s wavelength. Bacteria, as viewed using visible light Bacteria, as viewed using electrons !

29. Electron Microscope This image was taken with a Scanning Electron Microscope (SEM). These devices can resolve features down to about 1 [nm]. This is about 100 times better than can be done with visible light microscopes! This image was taken with a Scanning Electron Microscope (SEM). These devices can resolve features down to about 1 [nm]. This is about 100 times better than can be done with visible light microscopes!  The electron microscope is a device which uses the wave behavior of electrons to make images  which are otherwise too small for visible light! IMPORTANT POINT HERE: High energy particles can be used to reveal the structure of matter !