Chapter 5 Electrons In Atoms

Transcription

1 Chapter 5 Electrons In Atoms 5.1 Revising the Atomic Model 5.2 Electron Arrangement in Atoms 5.3 Atomic Emission Spectra and the Quantum Mechanical Model 1 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

2 CHEMISTRY & YOU What makes the electron configuration of an atom stable? Energy and stability play an important role in determining how electrons are configured in an atom. 2 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

3 Electron Configurations Electron Configurations What are the three rules for writing the electron configurations of elements? Three rules the aufbau principle, the Pauli exclusion principle, and Hund s rule tell you how to find the electron configurations of atoms. 3 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

4 Electron Configurations The ways in which electrons are arranged in various orbitals around the nuclei of atoms are called electron configurations. 4 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

5 Increasing energy 5.2 Electron Arrangement in Atoms > Electron Configurations Aufbau Principle 6s 5s 4s 6p 5p 4p 5d 4d 3d 4f 3p 3s 2s 1s 2p According to the aufbau principle, electrons occupy the orbitals of lowest energy first. In the aufbau diagram, each box represents an atomic orbital. 5 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

6 Increasing energy 5.2 Electron Arrangement in Atoms > Electron Configurations Aufbau Principle 6s 5s 4s 6p 5p 4p 5d 4d 3d 4f 3p 3s 2p 2s 1s The aufbau diagram shows the relative energy levels of the various atomic orbitals. Orbitals of greater energy are higher on the diagram. 6 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

7 Increasing energy 5.2 Electron Arrangement in Atoms > Electron Configurations Aufbau Principle 6s 5s 4s 6p 5p 4p 5d 4d 3d 4f 3p 3s 2p 2s 1s The range of energy levels within a principal energy level can overlap the energy levels of another principal level. 7 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

8 Electron Configurations Pauli Exclusion Principle According to the Pauli exclusion principle, an atomic orbital may describe at most two electrons. To occupy the same orbital, two electrons must have opposite spins; that is, the electron spins must be paired. 8 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

9 Electron Configurations Pauli Exclusion Principle Spin is a quantum mechanical property of electrons and may be thought of as clockwise or counterclockwise. A vertical arrow indicates an electron and its direction of spin ( or ). An orbital containing paired electrons is written as. 9 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

10 Electron Configurations Hund s Rule According to Hund s rule, electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible. 10 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

11 Electron Configurations Hund s Rule Three electrons would occupy three orbitals of equal energy as follows. Electrons then occupy each orbital so that their spins are paired with the first electron in the orbital. 11 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

12 Electron Configurations Look at the orbital filling diagram of the oxygen atom. An oxygen atom contains eight electrons. Electron Configurations of Selected Elements Element 1s 2s 2p x 2p y 2p z 3s H 1s 1 He 1s 2 Electron configuration Li 1s 2 2s 1 C 1s 2 2s 2 2p 2 N 1s 2 2s 2 2p 3 O 1s 2 2s 2 2p 4 F 1s 2 2s 2 2p 5 Ne 1s 2 2s 2 2p 6 Na 1s 2 2s 2 2p 6 3s 1 12 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

13 Electron Configurations Look at the orbital filling diagram of the oxygen atom. The 1s orbital has two electrons of opposite spin. Electron Configurations of Selected Elements Element 1s 2s 2p x 2p y 2p z 3s H 1s 1 He 1s 2 Li 1s 2 2s 1 Electron configuration C 1s 2 2s 2 2p 2 N 1s 2 2s 2 2p 3 O 1s 2 2s 2 2p 4 F 1s 2 2s 2 2p 5 Ne 1s 2 2s 2 2p 6 Na 1s 2 2s 2 2p 6 3s 1 13 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

14 Electron Configurations Look at the orbital filling diagram of the oxygen atom. The 1s orbital has two electrons of opposite spin. The 2s orbital also has two electrons of opposite spin. Electron Configurations of Selected Elements Element 1s 2s 2p x 2p y 2p z 3s H 1s 1 He 1s 2 Li 1s 2 2s 1 Electron configuration C 1s 2 2s 2 2p 2 N 1s 2 2s 2 2p 3 O 1s 2 2s 2 2p 4 F 1s 2 2s 2 2p 5 Ne 1s 2 2s 2 2p 6 Na 1s 2 2s 2 2p 6 3s 1 14 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

15 Electron Configurations Look at the orbital filling diagram of the oxygen atom. Each of the three 2p orbitals has one electron. The remaining electron now pairs with an electron occupying one of the 2p orbitals. Electron Configurations of Selected Elements Element 1s 2s 2p x 2p y 2p z 3s H 1s 1 He 1s 2 Li 1s 2 2s 1 Electron configuration C 1s 2 2s 2 2p 2 N 1s 2 2s 2 2p 3 O 1s 2 2s 2 2p 4 F 1s 2 2s 2 2p 5 Ne 1s 2 2s 2 2p 6 Na 1s 2 2s 2 2p 6 3s 1 15 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

16 Electron Configurations A convenient shorthand method for showing the electron configuration of an atom involves writing the energy level and the symbol for every sublevel occupied by an electron. You indicate the number of electrons occupying that sublevel with a superscript. 16 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

17 Electron Configurations For hydrogen, with one electron in a 1s orbital, the electron configuration is written 1s 1. For oxygen, with two electrons in a 1s orbital, two electrons in a 2s orbital, and four electrons in 2p orbitals, the electron configuration is 1s 2 2s 2 2p 4. Note that the sum of the superscripts equals the number of electrons in the atom. 17 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

18 CHEMISTRY & YOU Explain why the correct electron configuration of oxygen is 1s 2 2s 2 2p 4 and not 1s 2 2s 2 2p 3 3s Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

19 Sample Problem 5.1 Writing Electron Configurations The atomic number of phosphorus is 15. Write the electron configuration of a phosphorus atom. 19 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

20 Electron Configurations Exceptional Electron Configurations You can obtain correct electron configurations for the elements up to vanadium (atomic number 23) by following the aufbau diagram for orbital filling. If you were to continue in that fashion, however, you would assign chromium and copper the following incorrect configurations. Cr 1s 2 2s 2 2p 6 3s 2 3p 6 3d 4 4s 2 Cu 1s 2 2s 2 2p 6 3s 2 3p 6 3d 9 4s 2 20 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

21 Electron Configurations Exceptional Electron Configurations The correct electron configurations are as follows: Cr 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 1 Cu 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 1 These arrangements give chromium a half-filled d sublevel and copper a filled d sublevel. 21 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

22 Electron Configurations Exceptional Electron Configurations Some actual electron configurations differ from those assigned using the aufbau principle because although half-filled sublevels are not as stable as filled sublevels, they are more stable than other configurations. 22 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

23 What is the correct electron configuration of a sulfur atom? A. 1s 2 2s 2 2p 4 3s 2 3p 6 B. 1s 2 2s 2 2p 6 3s 2 3p 3 C. 1s 2 2s 2 2p 6 3s 2 3p 4 D. 1s 2 2s 2 2p 6 3s 6 3p 2 23 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

24 Key Concept Three rules the aufbau principle, the Pauli exclusion principle, and Hund s rule tell you how to find the electron configurations of atoms. 24 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

25 Glossary Terms electron configuration: the arrangement of electrons of an atom in its ground state into various orbitals around the nuclei of atoms aufbau principle: the rule that electrons occupy the orbitals of lowest energy first Pauli exclusion principle: an atomic orbital may describe at most two electrons, each with opposite spin direction spin: a quantum mechanical property of electrons; it may be thought of as clockwise or counterclockwise Hund s rule: electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible 25 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

26 Chapter 5 Electrons In Atoms 5.1 Revising the Atomic Model 5.2 Electron Arrangement in Atoms 5.3 Atomic Emission Spectra and the Quantum Mechanical Model 26 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

27 CHEMISTRY & YOU What gives gas-filled lights their colors? An electric current passing through the gas in each glass tube makes the gas glow with its own characteristic color. 27 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

28 Light and Atomic Emission Spectra Light and Atomic Emission Spectra What causes atomic emission spectra? 28 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

29 Light and Atomic Emission Spectra The Nature of Light By the year 1900, there was enough experimental evidence to convince scientists that light consisted of waves. The amplitude of a wave is the wave s height from zero to the crest. The wavelength, represented by (the Greek letter lambda), is the distance between the crests. 29 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

30 Light and Atomic Emission Spectra The Nature of Light The frequency, represented by (the Greek letter nu), is the number of wave cycles to pass a given point per unit of time. The SI unit of cycles per second is called the hertz (Hz). 30 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

31 Light and Atomic Emission Spectra The Nature of Light The product of frequency and wavelength equals a constant (c), the speed of light. c = 31 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

32 Light and Atomic Emission Spectra The frequency ( ) and wavelength ( ) of light are inversely proportional to each other. As the wavelength increases, the frequency decreases. 32 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

33 Light and Atomic Emission Spectra The Nature of Light According to the wave model, light consists of electromagnetic waves. Electromagnetic radiation includes radio waves, microwaves, infrared waves, visible light, ultraviolet waves, X-rays, and gamma rays. All electromagnetic waves travel in a vacuum at a speed of m/s. 33 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

34 Light and Atomic Emission Spectra The Nature of Light The sun and incandescent light bulbs emit white light, which consists of light with a continuous range of wavelengths and frequencies. When sunlight passes through a prism, the different wavelengths separate into a spectrum of colors. In the visible spectrum, red light has the longest wavelength and the lowest frequency. 34 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

35 Light and Atomic Emission Spectra The electromagnetic spectrum consists of radiation over a broad range of wavelengths. Low energy ( = 700 nm) High energy ( = 380 nm) Frequency (s -1 ) 3 x x x Wavelength (m) 35 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

36 Light and Atomic Emission Spectra Atomic Emission Spectra When atoms absorb energy, their electrons move to higher energy levels. These electrons lose energy by emitting light when they return to lower energy levels. 36 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

37 Light and Atomic Emission Spectra Atomic Emission Spectra A prism separates light into the colors it contains. White light produces a rainbow of colors. Light bulb Screen Slit Prism 37 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

38 Light and Atomic Emission Spectra Atomic Emission Spectra Light from a helium lamp produces discrete lines. Helium lamp Screen Slit Prism 38 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

39 Light and Atomic Emission Spectra Atomic Emission Spectra The energy absorbed by an electron for it to move from its current energy level to a higher energy level is identical to the energy of the light emitted by the electron as it drops back to its original energy level. The wavelengths of the spectral lines are characteristic of the element, and they make up the atomic emission spectrum of the element. No two elements have the same emission spectrum. 39 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

40 Sample Problem 5.2 Calculating the Wavelength of Light Calculate the wavelength of the yellow light emitted by a sodium lamp if the frequency of the radiation is Hz ( /s). 40 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

41 Sample Problem Analyze List the knowns and the unknown. Use the equation c = to solve for the unknown wavelength. KNOWNS frequency ( ) = /s c = m/s UNKNOWN wavelength ( ) =? m 41 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

42 Sample Problem Calculate Solve for the unknown. Write the expression that relates the frequency and wavelength of light. c = 42 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

43 Sample Problem Calculate Solve for the unknown. Rearrange the equation to solve for. c = c = Solve for by dividing both sides by : c = 43 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

44 Sample Problem Calculate Solve for the unknown. Substitute the known values for and c into the equation and solve. c = = 8 m/s = m /s 44 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

45 Sample Problem Evaluate Does the answer make sense? The magnitude of the frequency is much larger than the numerical value of the speed of light, so the answer should be much less than 1. The answer should have 3 significant figures. 45 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

46 The Quantum Concept and Photons The Quantum Concept and Photons How did Einstein explain the photoelectric effect? 46 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

47 The Quantum Concept and Photons The Quantization of Energy German physicist Max Planck ( ) showed mathematically that the amount of radiant energy (E) of a single quantum absorbed or emitted by a body is proportional to the frequency of radiation ( ). E or E = h 47 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

48 The Quantum Concept and Photons The Quantization of Energy The constant (h), which has a value of J s (J is the joule, the SI unit of energy), is called Planck s constant. E or E = h 48 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

49 The Quantum Concept and Photons The Photoelectric Effect Albert Einstein used Planck s quantum theory to explain the photoelectric effect. In the photoelectric effect, electrons are ejected when light shines on a metal. 49 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

50 The Quantum Concept and Photons The Photoelectric Effect Not just any frequency of light will cause the photoelectric effect. Red light will not cause potassium to eject electrons, no matter how intense the light. Yet a very weak yellow light shining on potassium begins the effect. 50 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

51 The Quantum Concept and Photons The Photoelectric Effect The photoelectric effect could not be explained by classical physics. Classical physics correctly described light as a form of energy. But, it assumed that under weak light of any wavelength, an electron in a metal should eventually collect enough energy to be ejected. 51 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

52 The Quantum Concept and Photons The Photoelectric Effect To explain the photoelectric effect, Einstein proposed that light could be described as quanta of energy that behave as if they were particles. 52 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

53 The Quantum Concept and Photons The Photoelectric Effect These light quanta are called photons. Einstein s theory that light behaves as a stream of particles explains the photoelectric effect and many other observations. Light behaves as waves in other situations; we must consider that light possesses both wavelike and particle-like properties. 53 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

54 The Quantum Concept and Photons The Photoelectric Effect No electrons are ejected because the frequency of the light is below the threshold frequency. If the light is at or above the threshold frequency, electrons are ejected. If the frequency is increased, the ejected electrons will travel faster. 54 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

55 Sample Problem 5.3 Calculating the Energy of a Photon What is the energy of a photon of microwave radiation with a frequency of /s? 55 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

56 Sample Problem Analyze List the knowns and the unknown. Use the equation E = h to calculate the energy of the photon. KNOWNS frequency ( ) = /s h = J s UNKNOWN energy (E) =? J 56 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

57 Sample Problem Calculate Solve for the unknown. Write the expression that relates the energy of a photon of radiation and the frequency of the radiation. E = h 57 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

58 Sample Problem Calculate Solve for the unknown. Substitute the known values for and h into the equation and solve. E = h = ( J s) ( /s) = J 58 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

59 Sample Problem Evaluate Does the result make sense? Individual photons have very small energies, so the answer seems reasonable. 59 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

60 What is the frequency of a photon whose energy is J? E = h = h E E = = 34 J h J s = Hz 60 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

61 An Explanation of Atomic Spectra An Explanation of Atomic Spectra How are the frequencies of light emitted by an atom related to changes of electron energies? 61 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

62 An Explanation of Atomic Spectra When an electron has its lowest possible energy, the atom is in its ground state. In the ground state, the principal quantum number (n) is 1. Excitation of the electron by absorbing energy raises the atom to an excited state with n = 2, 3, 4, 5, or 6, and so forth. A quantum of energy in the form of light is emitted when the electron drops back to a lower energy level. 62 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

63 An Explanation of Atomic Spectra The light emitted by an electron moving from a higher to a lower energy level has a frequency directly proportional to the energy change of the electron. 63 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

64 An Explanation of Atomic Spectra The three groups of lines in the hydrogen spectrum correspond to the transition of electrons from higher energy levels to lower energy levels. 64 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

65 CHEMISTRY & YOU The glass tubes in lighted signs contain helium, neon, argon, krypton, or xenon gas, or a mixture of these gases. Why do the colors of the light depend on the gases that are used? Each different gas has its own characteristic emission spectrum, creating different colors of light when excited electrons return to the ground state. 65 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

66 In the hydrogen spectrum, which of the following transitions produces a spectral line of the greatest energy? A. n = 2 to n = 1 B. n = 3 to n = 2 C. n = 4 to n = 3 66 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

67 Quantum Mechanics Quantum Mechanics How does quantum mechanics differ from classical mechanics? 67 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

68 Quantum Mechanics Given that light behaves as waves and particles, can particles of matter behave as waves? Louis de Broglie referred to the wavelike behavior of particles as matter waves. His reasoning led him to a mathematical expression for the wavelength of a moving particle. 68 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

69 Quantum Mechanics The Wavelike Nature of Matter Today, the wavelike properties of beams of electrons are useful in viewing objects that cannot be viewed with an optical microscope. The electrons in an electron microscope have much smaller wavelengths than visible light. These smaller wavelengths allow a much clearer enlarged image of a very small object, such as this pollen grain, than is possible with an ordinary microscope. 69 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

70 Quantum Mechanics Classical mechanics adequately describes the motions of bodies much larger than atoms, while quantum mechanics describes the motions of subatomic particles and atoms as waves. 70 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

71 Quantum Mechanics The Heisenberg Uncertainty Principle The Heisenberg uncertainty principle states that it is impossible to know both the velocity and the position of a particle at the same time. 71 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

72 Quantum Mechanics The Heisenberg Uncertainty Principle The Heisenberg uncertainty principle states that it is impossible to know both the velocity and the position of a particle at the same time. This limitation is critical when dealing with small particles such as electrons. But it does not matter for ordinary-sized objects such as cars or airplanes. 72 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

73 Quantum Mechanics To locate an electron, you might strike it with a photon. The electron has such a small mass that striking it with a photon affects its motion in a way that cannot be predicted accurately. The very act of measuring the position of the electron changes its velocity, making its velocity uncertain. Before collision: A photon strikes an electron during an attempt to observe the electron s position. After collision: The impact changes the electron s velocity, making it uncertain. 73 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

74 The Heisenberg uncertainty principle states that it is impossible to simultaneously know which two attributes of a particle? velocity and position 74 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

75 Key Concepts and Key Equations When atoms absorb energy, their electrons move to higher energy levels. These electrons lose energy by emitting light when they return to lower energy levels. To explain the photoelectric effect, Einstein proposed that light could be described as quanta of energy that behave as if they were particles. The light emitted by an electron moving from a higher to a lower energy level has a frequency directly proportional to the energy change of the electron. 75 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

76 Key Concepts and Key Equations Classical mechanics adequately describes the motions of bodies much larger than atoms, while quantum mechanics describes the motions of subatomic particles and atoms as waves. C = E = h 76 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

77 Glossary Terms amplitude: the height of a wave s crest wavelength: the distance between adjacent crests of a wave frequency: the number of wave cycles that pass a given point per unit of time; frequency and wavelength are inversely proportional to each other hertz: the unit of frequency, equal to one cycle per second 77 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

78 Glossary Terms electromagnetic radiation: energy waves that travel in a vacuum at a speed of m/s; includes radio waves, microwaves, infrared waves, visible light, ultraviolet waves, X-rays, and gamma rays spectrum: wavelengths of visible light that are separated when a beam of light passes through a prism; range of wavelengths of electromagnetic radiation 78 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

79 Glossary Terms atomic emission spectrum: the pattern formed when light passes through a prism or diffraction grating to separate it into the different frequencies of light it contains Planck s constant: the constant (h) by which the amount of radiant energy (E) is proportional to the frequency of the radiation ( ) photoelectric effect: the phenomenon in which electrons are ejected when light shines on a metal 79 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

80 Glossary Terms photon: a quantum of light; a discrete bundle of electromagnetic energy that interacts with matter similarly to particles ground state: the lowest possible energy of an atom described by quantum mechanics Heisenberg uncertainty principle: it is impossible to know both the velocity and the position of a particle at the same time 80 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.

81 BIG IDEA Electrons and the Structure of Atoms Electrons can absorb energy to move from one energy level to a higher energy level. When an electron moves from a higher energy level back down to a lower energy level, light is emitted. 81 Copyright Pearson Education, Inc., or its affiliates. All Rights Reserved.