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Chapter 21. Atomic Physics. Table of Contents. Section 1 Quantization of Energy Section 2 Models of the Atom Section 3 Quantum Mechanics. Chapter 21. Section 1 Quantization of Energy. Objectives. Explain how Planck resolved the ultraviolet catastrophe in blackbody radiation.

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table of contents

Chapter 21

Atomic Physics

Table of Contents

Section 1 Quantization of Energy

Section 2 Models of the Atom

Section 3 Quantum Mechanics

objectives

Chapter 21

Section 1 Quantization of Energy

Objectives
  • Explainhow Planck resolved the ultraviolet catastrophe in blackbody radiation.
  • Calculateenergy of quanta using Planck’s equation.
  • Solveproblems involving maximum kinetic energy, work function, and threshold frequency in the photoelectric effect.
blackbody radiation

Chapter 21

Section 1 Quantization of Energy

Blackbody Radiation
  • Physicists studyblackbody radiationby observing a hollow object with a small opening, as shown in the diagram.
  • Ablackbodyis a perfect radiator and absorber and emits radiation based only on its temperature.

Light enters this hollow object through the small opening and strikes the interior wall. Some of the energy is absorbed, but some is reflected at a random angle. After many reflections, essentially all of the incoming energy is absorbed by the cavity wall.

blackbody radiation continued

Chapter 21

Section 1 Quantization of Energy

Blackbody Radiation, continued
  • Theultraviolet catastropheis the failed prediction of classical physics that the energy radiated by a blackbody at extremely short wavelengths is extremely large and that the total energy radiated is infinite.
  • Max Planck (1858–1947) developed a formula for blackbody radiation that was in complete agreement with experimental data at all wavelengths by assuming thatenergy comes in discrete units,or isquantized.
blackbody radiation1

Chapter 21

Section 1 Quantization of Energy

Blackbody Radiation

The graph on the left shows the intensity of blackbody radiation at three different temperatures. Classical theory’s prediction for blackbody radiation (the blue curve) did not correspond to the experimental data (the red data points) at all wavelengths, whereas Planck’s theory (the red curve) did.

quantum energy

Chapter 21

Section 1 Quantization of Energy

Quantum Energy
  • Einstein later applied the concept of quantized energy to light. The units of light energy calledquanta(now calledphotons)are absorbed or given off as a result of electrons “jumping” from one quantum state to another.
  • Theenergy of a light quantum,which corresponds to the energy difference between two adjacent levels, is given by the following equation:

E = hf

energy of a quantum = Planck’s constant  frequency

Planck’s constant (h) ≈ 6.63  10–34 J•s

quantum energy1

Chapter 21

Section 1 Quantization of Energy

Quantum Energy
  • If Planck’s constant is expressed in units of J•s, the equation E = hf gives the energy in joules.
  • However, in atomic physics,energy is often expressed in units of theelectron volt, eV.
  • An electron voltis defined as the energy that an electron or proton gains when it is accelerated through a potential difference of 1 V.
  • The relation between the electron volt and the joule is as follows:

1 eV = 1.60  10–19 J

energy of a photon

Chapter 21

Section 1 Quantization of Energy

Energy of a Photon
the photoelectric effect

Chapter 21

Section 1 Quantization of Energy

The Photoelectric Effect
  • Thephotoelectric effectis the emission of electrons from a material surface that occurs when light of certain frequencies shines on the surface of the material.
  • Classical physics cannot explain the photoelectric effect.
  • Einstein assumed that an electromagnetic wave can be viewed as a stream of particles calledphotons.Photon theory accounts for observations of the photoelectric effect.
the photoelectric effect1

Chapter 21

Section 1 Quantization of Energy

The Photoelectric Effect
the photoelectric effect2

Chapter 21

Section 1 Quantization of Energy

The Photoelectric Effect
the photoelectric effect continued

Chapter 21

Section 1 Quantization of Energy

The Photoelectric Effect, continued
  • No electrons are emitted if the frequency of the incoming light falls below a certain frequency, called thethreshold frequency (ft).
  • The smallest amount of energy the electron must have to escape the surface of a metal is thework functionof the metal.
  • The work function is equal tohft.
the photoelectric effect continued1

Chapter 21

Section 1 Quantization of Energy

The Photoelectric Effect, continued

Because energy must be conserved, the maximum kinetic energy (of photoelectrons ejected from the surface) is the difference between the photon energy and the work function of the metal.

maximum kinetic energy of a photoelectron

KEmax = hf – hft

maximum kinetic energy = (Planck’s constant  frequency of incoming photon) – work function

compton shift

Chapter 21

Section 1 Quantization of Energy

Compton Shift
  • If light behaves like a particle, thenphotons should have momentumas well as energy; both quantities should be conserved in elastic collisions.
  • The American physicist Arthur Compton directed X rays toward a block of graphite to test this theory.
  • He found that the scattered waves hadless energyandlonger wavelengthsthan the incoming waves, just as he had predicted.
  • This change in wavelength, known as theCompton shift,supports Einstein’s photon theory of light.
compton shift1

Chapter 21

Section 1 Quantization of Energy

Compton Shift
objectives1

Chapter 21

Section 2 Models of the Atom

Objectives
  • Explainthe strengths and weaknesses of Rutherford’s model of the atom.
  • Recognizethat each element has a unique emission and absorption spectrum.
  • Explainatomic spectra using Bohr’s model of the atom.
  • Interpretenergy-level diagrams.
early models of the atom

Chapter 21

Section 2 Models of the Atom

Early Models of the Atom
  • The model of the atom in the days of Newton was that of a tiny, hard, indestructible sphere.
  • The discovery of the electron in 1897 promptedJ. J. Thomson(1856–1940) to suggest a new model of the atom.
  • In Thomson’s model, electrons are embedded in a spherical volume of positive charge like seeds in a watermelon.
early models of the atom continued

Chapter 21

Section 2 Models of the Atom

Early Models of the Atom, continued
  • Ernest Rutherford(1871–1937) later proved that Thomson’s model could not be correct.
  • In his experiment, a beam ofpositively charged alpha particleswas projected against a thin metal foil.
  • Most of the alpha particles passed through the foil. Some were deflected through very large angles.
early models of the atom continued1

Chapter 21

Section 2 Models of the Atom

Early Models of the Atom, continued
  • Rutherford concluded thatall of the positive chargeinan atom andmost of the atom’s massare found in a region that is small compared to the size of the atom.
  • He called this region the thenucleusof the atom.
  • Anyelectronsin the atom were assumed to be in the relatively large volume outside the nucleus.
early models of the atom continued2

Chapter 21

Section 2 Models of the Atom

Early Models of the Atom, continued
  • To explain why electrons were not pulled into the nucleus, Rutherford viewed theelectrons as movingin orbits about the nucleus.
  • However, accelerated charges should radiate electromagnetic waves, losing energy. This would lead to a rapid collapse of the atom.
  • This difficulty led scientists to continue searching for a new model of the atom.
atomic spectra

Chapter 21

Section 2 Models of the Atom

Atomic Spectra

When the light given off by an atomic gas is passed through a prism, a series of distinct bright lines is seen. Each line corresponds to a different wavelength, or color.

  • A diagram or graph that indicates the wavelengths of radiant energy that a substance emits is called anemission spectrum.
  • Every element has a distinct emission spectrum.
atomic spectra continued

Chapter 21

Section 2 Models of the Atom

Atomic Spectra, continued
  • An element can also absorb light at specific wavelengths.
  • The spectral lines corresponding to this process form what is known as anabsorption spectrum.
  • An absorption spectrum can be seen by passing light containing all wavelengths through a vapor of the element being analyzed.
  • Each line in theabsorption spectrumof a given element coincides with a line in theemission spectrumof that element.
the bohr model of the hydrogen atom

Chapter 21

Section 2 Models of the Atom

The Bohr Model of the Hydrogen Atom
  • In 1913, the Danish physicistNiels Bohr(1885–1962) proposed a new model of the hydrogen atom that explained atomic spectra.
  • In Bohr’s model,only certain orbits are allowed.The electron is never found between these orbits; instead, it is said to “jump” instantly from one orbit to another.
  • In Bohr’s model,transitionsbetween stable orbits with different energy levels account for thediscrete spectral lines.
the bohr model continued

Chapter 21

Section 2 Models of the Atom

The Bohr Model, continued
  • When light of a continuous spectrum shines on the atom, only the photons whose energy (hf) matches the energy separation between two levels can be absorbed by the atom.
  • When this occurs, an electron jumps from a lower energy state to ahigherenergy state, which corresponds to an orbitfartherfrom the nucleus.
  • This is called anexcited state. The absorbed photons account for the dark lines in the absorption spectrum.
the bohr model continued1

Chapter 21

Section 2 Models of the Atom

The Bohr Model, continued
  • Once an electron is in an excited state, there is a certain probability that it will jump back to a lowerenergy level byemitting a photon.
  • This process is calledspontaneous emission.
  • The emitted photons are responsible for the bright lines in the emission spectrum.
  • In both cases, there is a correlation between the“size”of an electron’s jump and theenergy of the photon.
the bohr model of the atom

Chapter 21

Section 2 Models of the Atom

The Bohr Model of the Atom
sample problem

Chapter 21

Section 2 Models of the Atom

Sample Problem

Interpreting Energy-Level Diagrams

An electron in a hydrogen atom drops from energy level E4 to energy level E2. What is the frequency of the emitted photon, and which line in the emission spectrum corresponds to this event?

sample problem continued

Chapter 21

Section 2 Models of the Atom

Sample Problem, continued
  • Find the energy of the photon.

The energy of the photon is equal to the change in the energy of the electron. The electron’s initial energy level was E4, and the electron’s final energy level was E2. Using the values from the energy-level diagram gives the following:

E = Einitial – Efinal =E4 – E2

E = (–0.850 eV) – (–3.40 eV) = 2.55 eV

sample problem continued1

Chapter 21

Section 2 Models of the Atom

Sample Problem, continued

Tip: Note that the energies for each energy level are negative. The reason is that the energy of an electron in an atom is defined with respect to the amount of work required to remove the electron from the atom. In some energy-level diagrams, the energy of E1 is defined as zero, and the higher energy levels are positive.

In either case, the difference between a higher energy level and a lower one is always positive, indicating that the electron loses energy when it drops to a lower level.

sample problem continued2

Chapter 21

Section 2 Models of the Atom

Sample Problem, continued

Tip:Note that electron volts were converted to joules so that the units cancel properly.

2. Use Planck’s equation to find the frequency.

sample problem continued3

Chapter 21

Section 2 Models of the Atom

Sample Problem, continued
  • Find the corresponding line in the emission spectrum.

Examination of the diagram shows that the electron’s jump from energy level E4 to energy level E2 corresponds to Line 3 in the emission spectrum.

sample problem continued4

Chapter 21

Section 2 Models of the Atom

Sample Problem, continued

4. Evaluate your answer.

Line 3 is in the visible part of the electromagnetic spectrum and appears to be blue. The frequency f = 6.15  1014 Hz lies within the range of the visible spectrum and is toward the violet end, so it is reasonable that light of this frequency would be visible blue light.

the bohr model continued2

Chapter 21

Section 2 Models of the Atom

The Bohr Model, continued
  • Bohr’s modelwas not considered to be a complete picture of the structure of the atom.
    • Bohr assumed that electrons do not radiate energy when they are in a stable orbit, but his model offered no explanation for this.
    • Another problem with Bohr’s model was that it could not explain why electrons always have certain stable orbits
  • For these reasons, scientists continued to search for a new model of the atom.
objectives2

Chapter 21

Section 3 Quantum Mechanics

Objectives
  • Recognizethe dual nature of light and matter.
  • Calculatethe de Broglie wavelength of matter waves.
  • Distinguishbetween classical ideas of measurement and Heisenberg’s uncertainty principle.
  • Describethe quantum-mechanical picture of the atom, including the electron cloud and probability waves.
the dual nature of light

Chapter 21

Section 3 Quantum Mechanics

The Dual Nature of Light
  • As seen earlier, there is considerable evidence for thephotontheory of light. In this theory, all electromagnetic waves consist of photons, particle-like pulses that have energy and momentum.
  • On the other hand, light and other electromagnetic waves exhibit interference and diffraction effects that are considered to bewavebehaviors.
  • So, which model is correct?
the dual nature of light continued

Chapter 21

Section 3 Quantum Mechanics

The Dual Nature of Light, continued
  • Some experiments can be better explained or only explained by thephotonconcept, whereas others require a wave model.
  • Most physicistsaccept both modelsand believe that the true nature of light is not describable in terms of a single classical picture.
    • At one extreme, the electromagnetic wave description suits the overall interference pattern formed by a large number of photons.
    • At the other extreme, the particle description is more suitable for dealing with highly energetic photons of very short wavelengths.
the dual nature of light1

Chapter 21

Section 3 Quantum Mechanics

The Dual Nature of Light
matter waves

Chapter 21

Section 3 Quantum Mechanics

Matter Waves
  • In 1924, the French physicistLouis de Broglie(1892–1987) extended the wave-particle duality. De Broglie proposed that all forms of matter may have both wave properties and particle properties.
  • Three years after de Broglie’s proposal, C. J. Davisson and L. Germer, of the United States, discovered that electrons can be diffracted by a single crystal of nickel. This important discovery provided thefirst experimental confirmationof de Broglie’s theory.
matter waves continued

Chapter 21

Section 3 Quantum Mechanics

Matter Waves, continued
  • Thewavelengthof a photon is equal toPlanck’s constant (h) divided by the photon’smomentum (p). De Broglie speculated that this relationship might also hold for matter waves, as follows:
  • As seen by this equation, thelarger the momentumof an object, thesmaller its wavelength.
matter waves continued1

Chapter 21

Section 3 Quantum Mechanics

Matter Waves, continued
  • In an analogy with photons, de Broglie postulated that thefrequencyof a matter wave can be found withPlanck’s equation,as illustrated below:
  • Thedual nature of mattersuggested by de Broglie is quite apparent in the wavelength and frequency equations, both of which containparticle concepts (E and mv) and wave concepts (l and f).
matter waves continued2

Chapter 21

Section 3 Quantum Mechanics

Matter Waves, continued
  • He assumed that an electron orbit would be stable only if it contained an integral (whole) number of electron wavelengths.
  • De Broglie saw a connection between histheory of matter waves and the stable orbits in the Bohr model.
the uncertainty principle

Chapter 21

Section 3 Quantum Mechanics

The Uncertainty Principle
  • In 1927, Werner Heisenberg argued that it is fundamentally impossible to make simultaneous measurements of a particle’s position and momentum with infinite accuracy.
  • In fact, the more we learn about a particle’s momentum, the less we know of its position, and the reverse is also true.
  • This principle is known asHeisenberg’s uncertainty principle.
the uncertainty principle1

Chapter 21

Section 3 Quantum Mechanics

The Uncertainty Principle
the electron cloud continued

Chapter 21

Section 3 Quantum Mechanics

The Electron Cloud, continued
  • Quantum mechanics also predicts that the electron can be found in aspherical regionsurrounding the nucleus.
  • This result is often interpreted by viewing the electron as acloudsurrounding the nucleus.
  • Analysis of each of the energy levels of hydrogen reveals that themost probable electron locationin each case is in agreement with each of theradii predicted by the Bohr theory.
the electron cloud

Chapter 21

Section 3 Quantum Mechanics

The Electron Cloud
  • Because the electron’s location cannot be precisely determined, it is useful to discuss theprobabilityof finding the electron at different locations.
  • The diagram shows the probability per unit distance of finding the electron at various distances from the nucleus in the ground state of hydrogen.
atomic spectra1

Chapter 21

Section 2 Models of the Atom

Atomic Spectra