Chapter 7. The Quantum Mechanical Model of the Atom
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1 Chapter 7 The Quantum Mechanical Model of the Atom
2 The Nature of Light:Its Wave Nature Light is a form of 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 x 10 8 m/s in a vacuum = the speed of light, c
3 The Relationship Between Wavelength and Frequency ν = c λ ν = frequency c = speed of light λ = wavelength
4 Characterizing Waves The amplitude is the height of the wave. The amplitude is a measure of how intense the light is the larger the amplitude, the brighter the light. The wavelength (λ) is a measure of the distance covered by the wave.
5 The Wave Nature of Light Amplitude (intensity) of a wave
6 Characterizing Waves The frequency (ν) is the number of waves that pass a point in a given period of time. The number of waves = number of cycles units are hertz (Hz) or cycles/s = s 1 1 Hz = 1 s 1 or 1/s The total energy is proportional to the amplitude of the waves and the frequency. The larger the amplitude, the more force it has. The more frequently the waves strike, the more total force.
7 Characterizing Waves low frequency low amplitude Increasing energy higher frequency higher amplitude higher frequency higher amplitude
8 The Electromagnetic Spectrum RedOrangeYellowGreenBlueViolet Shorter wavelengths of energy have higher energy than longer wavelengths: Radiowaves have the lowest energy. Gamma rays have the highest energy.
9 White Light Produces a Continuous Spectrum
10 19th Century-Atomic Spectra When atoms or molecules absorb energy, that energy is often released as light energy. When that emitted light is passed through a prism, a pattern of particular wavelengths of light is seen that is unique to that type of atom or molecule the pattern is called an emission spectrum. non-continuous can be used to identify the material
11 Identifying Elements with Flame Tests Na K Li Ba
12
13 Gas Atoms Can Be Excited with Electrical Energy to Emit Light
14 Emission Spectrum Violet line λ = nm Blue line λ = 434 nm Green line λ = 486 nm Red line λ = nm
15 Emission vs. Absorption Spectra of Mercury Vapor
16 The Beginnings of Quantum Mechanics The field of quantum mechanics began with the studies of physicists in the early the 20th century. Max Planck (1918) Albert Einstein (1921) Neils Bohr (1922) Arthur Compton (1927) Louis de Broglie (1929) Werner Heisenberg (1932) P. A. M. Dirac (1933) Erwin Schrödinger (1933)
17 Quantum Mechanics The Behavior of the Very Small Electrons are incredibly small. Electron behavior determines much of the behavior of atoms. Directly observing electrons in the atom is impossible, the electron is so small that observing it changes its behavior.
18 The Path from Classical Theory to Quantum Theory
19 Classical Physical Theory Matter particulate massive Energy continuous wavelike Quantum Theory Could energy be discontinuous and quantized? Could matter have wavelike properties? Observation Theory Conclusion Black-body radiation Planck Energy is quantized Photoelectric effect Einstein Light has particulate behavior Atomic line spectra Bohr Energy of atoms is quantized
20 The Electromagnetic Spectrum RedOrangeYellowGreenBlueViolet Shorter wavelengths of energy have higher energy than longer wavelengths: Radiowaves have the lowest energy. Gamma rays have the highest energy.
21 The Relationship Between Wavelength and Frequency ν = c λ ν = frequency c = speed of light λ = wavelength
22 The Electromagnetic Spectrum RedOrangeYellowGreenBlueViolet Energy Increases
23 Blackbody Radiation Intensity increases from right to left on the curve as wavelength decreases. As the wavelength continues to decrease, however,intensity reaches a maximum and then drops off to zero. Max Planck
24 Good Grief! It s an ultraviolet catastrophe!
25 Blackbody Radiation The dependence of the intensity of blackbody radiation on wavelength at two different temperatures. Intensity increases from right to left on the curve as wavelength decreases. As the wavelength continues to decrease, intensity reaches a maximum and then drops off to zero.
26 Planck: ν = c λ Ephoton = ν C Ephoton= λ Ephoton= ν = x Jᐧs
27 The Photoelectric Effect Observation: Many metals emit electrons when a light shines on their surface. Classic wave theory: Light energy is transferred to the electron, which is then ejected from the metal. Prediction: If the wavelength of light is made shorter, or the light waves intensity made brighter, more electrons should be ejected.
28 The Photoelectric Effect A beam of white light is dispersed into its wavelength components by a quartz prism and falls on a metal sample (potassium, in this case). Light of the highest frequencies (violet and ultraviolet) produces the most energetic photoelectrons (longest arrows). Light of lower frequencies (for example, orange) results in less energetic photoelectrons (shorter arrows).
29 The Photoelectric Effect Light with a frequency lower than 4.23 x s-1 (710 nm) produces no photoelectric effect at all on potassium, regardless of how bright (intense) the light is.
30 The Bohr Model of the Atom Neils Bohr ( ) The energy of the atom is quantized. The amount of energy in the atom is related to the electron s position in the atom. Quantized means that the atom could only have very specific amounts of energy. Bohr correlated these allowed energy levels with allowed radii of electron orbits.
31 The Bohr Model-Allowed Radii of Electron Orbits Z = atomic number ao = a constant derived from the mass and charge of the electron, Planck s constant, and Coulomb s constant n = state of the atom n can only be an integer!!
32 The Bohr Model-The Hydrogen Atom En = ev/n n=1 n=2 n=3 n=4 n=5 radius n = n 2 rb 1rb 4rb 9rb 16rb 25rb nm nm nm nm 1.32 nm
33 The Bohr Model-The Hydrogen Atom
34 Bohr s Model The electrons travel in orbits that are at a fixed distance from the nucleus (stationary states). The energy of the electron is proportional to the distance the orbit was from the nucleus. Electrons emit radiation when they jump from an orbit with higher energy down to an orbit with lower energy. Emitted radiation is a photon of light. The distance between the orbits determines the energy of the photon of light produced.
35 Bohr s Model
36 Excitation and Emission When an atom absorbs energy, an electron is excited to a higherenergy orbit. The electron relaxes to a lower energy level, emitting a photon of light. n=1 Energy is absorbed!! n=2 n=3 n=4 n=5 Energy is emitted!!
37 The Bohr Model-Potential Energy of Electrons Energy Radius
38 The Bohr Model-Potential Energy of Electrons Energy Radius
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