Preview. Atomic Physics Section 1. Section 1 Quantization of Energy. Section 2 Models of the Atom. Section 3 Quantum Mechanics

Transcription

1 Atomic Physics Section 1 Preview Section 1 Quantization of Energy Section 2 Models of the Atom Section 3 Quantum Mechanics

2 Atomic Physics Section 1 TEKS The student is expected to: 8A describe the photoelectric effect and the dual nature of light

3 Atomic Physics Section 1 What do you think? Albert Einstein, while working as a Swiss patent clerk in 1905, submitted three papers for publication. The topics for the three papers were: (1) the special theory of relativity, including E = mc 2 (2) an explanation of Brownian motion (3) a quantum theory of light used to explain the photoelectric effect Sixteen years later he was awarded the Nobel Prize in Physics for one of these papers. Which paper earned him this award?

4 Atomic Physics Section 1 Blackbody Radiation A blackbody is a perfect absorber and perfect emitter of radiation. All radiation passing through the hole is trapped by multiple reflections and very little is emitted. Called a blackbody When you look into the hole, it is black because no light escapes.

5 Atomic Physics Section 1 Blackbody Radiation The radiation that escapes a blackbody has a wide range of wavelengths. At higher temperatures, the shorter wavelengths such as visible and UV light become more prominent.

6 Atomic Physics Section 1 Failure of Classical Wave Theory The wavelength distribution predicted by classical wave theory was wrong. Max Planck proposed a theory that matched the radiation curve. Energy is given off in discrete amounts of energy. Each unit of light energy is called a quanta.

7 Atomic Physics Section 1 Planck s Quantum Theory Energy can only be absorbed or released in multiples of a specific value: E = nhf h is Planck s constant ( J s) f is the frequency n represents the multiple (1,2,3,.)

8 Atomic Physics Section 1 Planck s Quantum Theory The energy absorbed or released must always be a whole number multiple of hf. A unit of energy used for extremely small values is the electron volt (ev). The amount of energy an electron gains when changing its potential difference by 1.0 V 1 ev = J

9 Atomic Physics Section 1 Blackbody Radiation and the Ultraviolet Catastrophe Click below to watch the Visual Concept. Visual Concept

10 Atomic Physics Section 1 Classroom Practice Problem Find the energy of a photon of green light with a frequency of Hz. Give your answer in joules and in electron volts. Answer: J or 2.28 ev

11 Atomic Physics Section 1 The Photoelectric Effect When light strikes a metal surface, electrons may be ejected. To escape from the binding forces of the atoms, electrons must absorb enough energy from the light. Classical wave theory did not predict the experimental results.

12 Atomic Physics Section 1 Photoelectric Effect

13 Atomic Physics Section 1 Photoelectric Effect Einstein theorized that light was quantized. Photons, not waves E = hf Based on this theory, either the frequency (or energy) of the photons was high enough to eject the electrons, or it wasn t. If energy (hf) was not high enough, no electrons would escape no matter how intense the light or how much time the light shined on the metal.

14 Atomic Physics Section 1 Photoelectric Effect Click below to watch the Visual Concept. Visual Concept

15 Atomic Physics Section 1 Photoelectric Effect Work function is the minimum energy needed to escape the metal s surface. Some electrons require more than this minimum. Any additional energy above hf t is the KE the electron has after ejection.

16 Atomic Physics Section 1 Photoelectric Effect Einstein s photon theory explains the photoelectric effect. No matter how bright the light (more photons), no electrons are emitted unless the photons have enough energy (hf t ). A brighter light produces more electrons because more photons strike the surface, but the maximum KE of the electrons does not change. Electrons are emitted almost instantaneously because each photon ejects a single electron.

17 Atomic Physics Section 1 The Dual Nature of Light Click below to watch the Visual Concept. Visual Concept

18 Atomic Physics Section 1 Compton Scattering If light has particle-like properties, it could collide with electrons and lose some energy. Like billiard balls The scattered photon has less energy (E = hf), so the wavelength is greater. The wavelength change is called the Compton shift.

19 Atomic Physics Section 1 Compton Shift Click below to watch the Visual Concept. Visual Concept

20 Atomic Physics Section 1 Now what do you think? In science, theories are developed and tested. When a theory fails to match the results of experiments, the theory must be modified. In what ways did the wave theory fail to explain the ejection of electrons from a metal surface when it was illuminated with light of various wavelengths? How did Einstein s explanation of the photoelectric effect change our ideas about light and also explain the experimental results?

21 Atomic Physics Section 2 TEKS The student is expected to: 2C know that scientific theories are based on natural and physical phenomena and are capable of being tested by multiple independent researchers. Unlike hypotheses, scientific theories are well-established and highly-reliable explanations, but may be subject to change as new areas of science and new technologies are developed 3D explain the impacts of the scientific contributions of a variety of historical and contemporary scientists on scientific thought and society 8B compare and explain the emission spectra produced by various atoms

22 Atomic Physics Section 2 What do you think? Describe an atom. Consider the following in your explanation: What are the component parts of an atom? How do they compare in size? How does the charge on each compare? Where are these parts located in the atom? Do the component parts of an atom move about or remain in fixed positions? If they move, describe the motion.

23 Atomic Physics Section 2 Models of the Atom The earliest models described atoms as tiny, indestructible, neutral spheres. J.J. Thomson discovered electrons in Since atoms had negatively charged electrons, they must have an equal positive charge of some sort. Thomson s model is like seeds (electrons) in watermelon (positive charge).

24 Atomic Physics Section 2 Ernest Rutherford Rutherford fired positive alpha particles at a thin sheet of gold foil. Most passed straight through the foil. Some were deflected as shown. Thomson s spheres of + charge would not cause this much deflection.

25 Atomic Physics Section 2 Rutherford's Gold Foil Experiment Click below to watch the Visual Concept. Visual Concept

26 Atomic Physics Section 2 Rutherford s Atom A small, positively-charged core called the nucleus Mostly empty space outside the nucleus Electrons circle the nucleus like planets around the sun. There was a problem with the electron motion in this model: A circling charged particle is accelerating and would radiate energy, and would spiral into the nucleus as it lost energy. Atoms could not exist in this model for more than one-billionth of a second.

27 Atomic Physics Section 2 Atomic Spectra The specific wavelengths of light given off when an electric current passes through a gas Every element produces different spectral lines. Rutherford s model could not explain this phenomenon.

28 Atomic Physics Section 2 Atomic Spectra

29 Atomic Physics Section 2 Atomic Spectra Emission spectra are the colors given off by the gas. When white light passes through hydrogen, it absorbs the same frequencies. Absorption spectra are used to study gases surrounding the stars.

30 Atomic Physics Section 2 Niels Bohr s Model of the Atom Bohr s model assumed that the orbits of electrons were quantized. Electrons existed only in certain energy levels. Ground state is the lowest level. Electrons absorb energy and move to excited states. They release the energy when returning to ground state.

31 Atomic Physics Section 2 Hydrogen s Spectrum Explained The visible lines in hydrogen s spectrum result from jumps to level 2 from levels 6, 5, 4 and 3. Would drops to level 1 produce photons with more or less energy? Where would these spectral lines appear in the spectrum?

32 Atomic Physics Section 2 Bohr Model of the Atom Click below to watch the Visual Concept. Visual Concept

33 Atomic Physics Section 2 Classroom Practice Problems The five lowest energy levels for Hg vapor are shown above. Assume an electron falls from E 5 to E 2. What amount of energy is given off in ev and in joules? What is the frequency of this photon? Is this photon in the visible spectrum? Answers: 2.01 ev, J Hz, visible

34 Atomic Physics Section 2 Bohr s Model of the Atom Bohr s model was the first to explain spectral lines, but it raised some questions as well. It could not be extended well to multi-electron atoms. It did not explain why electrons did not radiate energy as they circled the nucleus. It did not explain why electrons had only certain stable orbits and other orbits did not exist.

35 Atomic Physics Section 2 Now what do you think? Describe an atom. Consider the following in your explanation. What are the component parts of an atom? How do they compare in size? How does the charge on each compare? Where are these parts located in the atom? Do the component parts of an atom move about or remain in fixed positions? If they move, describe the motion.

36 Atomic Physics Section 3 TEKS The student is expected to: 3D explain the impacts of the scientific contributions of a variety of historical and contemporary scientists on scientific thought and society 8A describe the photoelectric effect and the dual nature of light

37 Atomic Physics Section 3 What do you think? The photoelectric effect shows that electromagnetic waves can act like particles. Does it seem possible that particles such as bullets or electrons can show wavelike behavior? If hundreds of baseballs pass through a doorway, will they spread out like a sound wave? Will a beam of electrons going through a small opening diffract? After splitting a beam of electrons, could an electron from one beam interfere constructively with one from the other beam?

38 Atomic Physics Section 3 Wave-Particle Duality Low-frequency photons (radio waves) have very little energy. Energy values (E = hf) are about J. Too small to detect single photons A large number of photons/second are needed Behaves more like a wave of photons Visible light photons have enough energy to be detected in some experiments, like the photoelectric effect. High-frequency photons such as gamma rays behave more like particles. Wave properties (diffraction, interference) are hard to observe.

39 Atomic Physics Section 3 Matter Waves Louis de Broglie hypothesized that, if light could behave like particles, then particles could behave like waves. All forms of matter have wavelike properties. He derived an equation for the wavelength (λ).

40 Atomic Physics Section 3 Matter Waves De Broglie postulated that Planck s equation could be used to find the frequency of matter waves. De Broglie wavelengths and frequencies for matter waves depend on momentum (mv) and energy (E).

41 Atomic Physics Section 3 Matter Waves Light waves passing through an opening exhibit interference. Interference, a wave property, was demonstrated with electrons passing through a crystal of nickel years after De Broglie s prediction. The spacing between nickel atoms was about the same as the wavelength of the electrons.

42 Atomic Physics Section 3 Classroom Practice Problems Calculate the De Broglie wavelength of an electron moving at a speed of m/s. Answer: m Calculate the De Broglie wavelength for a kg fastball thrown at 40.0 m/s. Answer: m

43 Atomic Physics Section 3 De Broglie Waves in Atoms Electrons in atomic orbits behave like waves and set up standing waves around each orbital. Only certain multiples are possible. Just like standing waves on a string

44 Atomic Physics Section 3 De Broglie and the Wave-Particle Nature of Electrons Click below to watch the Visual Concept. Visual Concept

45 Atomic Physics Section 3 The Uncertainty Principle Heisenberg s uncertainty principle states that it is impossible to simultaneously determine a particle s position and momentum with infinite accuracy. Measuring one of the quantities alters the other one. For example, measuring the position of an electron requires that photons bounce off the electron so we can see it. However, the photon striking the electron will change the electron s momentum.

46 Atomic Physics Section 3 Electron Probability Since it is impossible to measure the exact position of an electron, we can only calculate the probability of finding an electron at various locations. Erwin Schrödinger developed a mathematical model of the electron using a wave function to describe it. The wave function is used to calculate the probability of finding electrons at various distances from the nucleus.

47 Atomic Physics Section 3 Electron Cloud The most likely position of the electron in the ground state, according to the wave function, is the same as the Bohr radius. However, the electron may be closer or farther away as shown by other probabilities. An electron cloud is a picture of the probabilities, where denser means more probable.

48 Atomic Physics Section 3 Heisenberg Uncertainty Principle Click below to watch the Visual Concept. Visual Concept

49 Atomic Physics Section 3 Now what do you think? The photoelectric effect shows that electromagnetic waves can act like particles. Does it seem possible that particles such as bullets or electrons can show wavelike behavior? If hundreds of baseballs pass through a doorway, will they spread out like a sound wave? Will a beam of electrons going through a small opening diffract? After splitting a beam of electrons, could an electron from one beam interfere constructively with one from the other beam?