1. Electronic Configurations of Atoms

2. The Development of Atomic Models
5.1
The Development of Atomic Models
The Development of Atomic Models
What was inadequate about Rutherford’s atomic model?

3. The Development of Atomic Models
5.1
The Development of Atomic Models
Rutherford’s atomic model could not explain the chemical properties of elements.
Rutherford’s atomic model could not explain why objects change color when heated.
Rutherford’s model fails to explain why objects change color when heated. As the temperature of this horseshoe is increased, it first appears black, then red, then yellow, and then white. The observed behavior could be explained only if the atoms in the iron gave off light in specific amounts of energy. A better atomic model was needed to explain this observation.

4. The Development of Atomic Models
5.1
The Development of Atomic Models
The timeline shoes the development of atomic models from 1803 to 1911.
These illustrations show how the atomic model has changed as scientists learned more about the atom’s structure.

5. The Development of Atomic Models
5.1
The Development of Atomic Models
The timeline shows the development of atomic models from 1913 to 1932.
These illustrations show how the atomic model has changed as scientists learned more about the atom’s structure.

6. 5.1
The Bohr Model
The Bohr Model
What was the new proposal in the Bohr model of the atom?

7. 5.1
The Bohr Model
Bohr proposed that an electron is found only in specific circular paths, or orbits, around the nucleus.

8. Each possible electron orbit in Bohr’s model has a fixed energy.
5.1
The Bohr Model
Each possible electron orbit in Bohr’s model has a fixed energy.
The fixed energies an electron can have are called energy levels.
A quantum of energy is the amount of energy required to move an electron from one energy level to another energy level.

9. 5.1
The Bohr Model
Like the rungs of the strange ladder, the energy levels in an atom are not equally spaced.
The higher the energy level occupied by an electron, the less energy it takes to move from that energy level to the next higher energy level.
These ladder steps are somewhat like energy levels. In an ordinary ladder, the rungs are equally spaced. The energy levels in atoms are unequally spaced, like the rungs in this ladder. The higher energy levels are closer together.

10. The Quantum Mechanical Model
5.1
The Quantum Mechanical Model
The Quantum Mechanical Model
What does the quantum mechanical model determine about the electrons in an atom?

11. The Quantum Mechanical Model
5.1
The Quantum Mechanical Model
The quantum mechanical model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus.

12. The Quantum Mechanical Model
5.1
The Quantum Mechanical Model
Austrian physicist Erwin Schrödinger (1887– 1961) used new theoretical calculations and results to devise and solve a mathematical equation describing the behavior of the electron in a hydrogen atom.
The modern description of the electrons in atoms, the quantum mechanical model, comes from the mathematical solutions to the Schrödinger equation.

13. The Quantum Mechanical Model
5.1
The Quantum Mechanical Model
In the quantum mechanical model, the probability of finding an electron within a certain volume of space surrounding the nucleus can be represented as a fuzzy cloud. The cloud is more dense where the probability of finding the electron is high.
The electron cloud of an atom is compared here to photographs of a spinning airplane propeller. a) The airplane propeller is somewhere in the blurry region it produces in this picture, but the picture does not tell you its exact position at any instant. b) Similarly, the electron cloud of an atom represents the locations where an electron is likely to be found.

14. Atomic Orbitals How do sublevels of principal energy levels differ?
5.1
Atomic Orbitals
Atomic Orbitals
How do sublevels of principal energy levels differ?

15. 5.1
Atomic Orbitals
An atomic orbital is often thought of as a region of space in which there is a high probability of finding an electron.
Each energy sublevel corresponds to an orbital of a different shape, which describes where the electron is likely to be found.

16. 5.1
Atomic Orbitals
Different atomic orbitals are denoted by letters. The s orbitals are spherical, and p orbitals are dumbbell-shaped.
The electron clouds for the s orbital and the p orbitals are shown here.

17. 5.1
Atomic Orbitals
Four of the five d orbitals have the same shape but different orientations in space.
The d orbitals are illustrated here. Four of the five d orbitals have the same shape but different orientations in space. Interpreting Diagrams How are the orientations of the dxy and dx2 – y2 orbitals similar? How are they different?

18. 5.1
Atomic Orbitals
The numbers and kinds of atomic orbitals depend on the energy sublevel.

19. 5.1
Atomic Orbitals
The number of electrons allowed in each of the first four energy levels are shown here.

20. Electron Configurations
5.2
Electron Configurations
Electron Configurations
What are the three rules for writing the electron configurations of elements?

21. Electron Configurations
5.2
Electron Configurations
The ways in which electrons are arranged in various orbitals around the nuclei of atoms are called electron configurations.
Three rules—the aufbau principle, the Pauli exclusion principle, and Hund’s rule—tell you how to find the electron configurations of atoms.

22. Electron Configurations
5.2
Electron Configurations
Aufbau Principle
According to the aufbau principle, electrons occupy the orbitals of lowest energy first. In the aufbau diagram below, each box represents an atomic orbital.
This aufbau diagram shows the energy levels of the various atomic orbitals. Orbitals of greater energy are higher on the diagram. Using Tables Which is of higher energy, a 4d or a 5s orbital?

23. Electron Configurations
5.2
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.

24. Electron Configurations
5.2
Electron Configurations
Hund’s Rule
Hund’s rule states that 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. Electron Configurations
5.2
Electron Configurations
Orbital Filling Diagram

26. Electron Configurations
Simulation 2
Fill atomic orbitals to build the ground state of several atoms.

30. for Conceptual Problem 1.1
Problem Solving 5.9 Solve Problem 9 with the help of an interactive guided tutorial.

31. Exceptional Electron Configurations
5.2
Exceptional Electron Configurations
Exceptional Electron Configurations
Why do actual electron configurations for some elements differ from those assigned using the aufbau principle?

32. Exceptional Electron Configurations
5.2
Exceptional Electron Configurations
Some actual electron configurations differ from those assigned using the aufbau principle because half-filled sublevels are not as stable as filled sublevels, but they are more stable than other configurations.

33. Exceptional Electron Configurations
5.2
Exceptional Electron Configurations
Exceptions to the aufbau principle are due to subtle electron-electron interactions in orbitals with very similar energies.
Copper has an electron configuration that is an exception to the aufbau principle.
Copper is a good conductor of electricity and is commonly used in electrical wiring.

34. 5.2 Section Quiz.
5.2.

35. 5.2 Section Quiz.
1. Identify the element that corresponds to the following electron configuration: 1s22s22p5. F Cl Ne O

36. 5.2 Section Quiz. 2. Write the electron configuration for the atom N.
1s22s22p5
1s22s22p3
1s22s1p2
1s22s22p1

37. 5.2 Section Quiz.
3. The electron configurations for some elements differ from those predicted by the aufbau principle because the
the lowest energy level is completely filled.
none of the energy levels are completely filled.
half-filled sublevels are less stable than filled energy levels.
half-filled sublevels are more stable than some other arrangements. `

38. chemistry

39. Physics and the Quantum Mechanical Model
5.3
Physics and the Quantum Mechanical Model
Neon advertising signs are formed from glass tubes bent in various shapes. An electric current passing through the gas in each glass tube makes the gas glow with its own characteristic color. You will learn why each gas glows with a specific color of light.

40. 5.3
Light
Light
How are the wavelength and frequency of light related?

41. 5.3
Light
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.

42. 5.3
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 a hertz (Hz).

43. 5.3
Light
The wavelength and frequency of light are inversely proportional to each other.
The frequency and wavelength of light waves are inversely related. As the wavelength increases, the frequency decreases.

44. 5.3
Light
The product of the frequency and wavelength always equals a constant (c), the speed of light.

45. 5.3
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  108 m/s.

46. 5.3
Light
Sunlight consists of light with a continuous range of wavelengths and frequencies.
When sunlight passes through a prism, the different frequencies separate into a spectrum of colors.
In the visible spectrum, red light has the longest wavelength and the lowest frequency.

47. 5.3
Light
The electromagnetic spectrum consists of radiation over a broad band of wavelengths.
The electromagnetic spectrum consists of radiation over a broad band of wavelengths. The visible light portion is very small. It is in the 10-7m wavelength range and 1015 Hz (s-1) frequency range. Interpreting Diagrams What types of nonvisible radiation have wavelengths close to those of red light? To those of blue light?

48. Light
Simulation 3
Explore the properties of electromagnetic radiation.

49. 5.1
Sodium vapor lamps produce a yellow glow.

50. 5.1

51. 5.1

52. 5.1

53. for Sample Problem 5.1
Problem-Solving 5.15 Solve Problem 15 with the help of an interactive guided tutorial.

54. 5.3
Atomic Spectra
Atomic Spectra
What causes atomic emission spectra?

55. 5.3
Atomic Spectra
When atoms absorb energy, electrons move into higher energy levels. These electrons then lose energy by emitting light when they return to lower energy levels.

56. 5.3
Atomic Spectra
A prism separates light into the colors it contains. When white light passes through a prism, it produces a rainbow of colors.
A prism separates light into the colors it contains. For white light this produces a rainbow of colors.

57. 5.3
Atomic Spectra
When light from a helium lamp passes through a prism, discrete lines are produced.
A prism separates light into the colors it contains. Light from a helium lamp produces discrete lines. Identifying Which color has the highest frequency?

58. 5.3
Atomic Spectra
The frequencies of light emitted by an element separate into discrete lines to give the atomic emission spectrum of the element.
Mercury
Nitrogen
No two elements have the same emission spectrum. a) Mercury vapor lamps produce a blue glow. b) Nitrogen gas gives off a yellowish-orange light.

59. An Explanation of Atomic Spectra
5.3
An Explanation of Atomic Spectra
An Explanation of Atomic Spectra
How are the frequencies of light an atom emits related to changes of electron energies?

60. An Explanation of Atomic Spectra
5.3
An Explanation of Atomic Spectra
In the Bohr model, the lone electron in the hydrogen atom can have only certain specific energies.
When the electron has its lowest possible energy, the atom is in its ground state.
Excitation of the electron by absorbing energy raises the atom from the ground state to an excited state.
A quantum of energy in the form of light is emitted when the electron drops back to a lower energy level.

61. An Explanation of Atomic Spectra
5.3
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.

62. An Explanation of Atomic Spectra
5.3
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.
The three groups of lines in the hydrogen spectrum correspond to the transition of electrons from higher energy levels to lower energy levels. The Lyman series corresponds to the transition to the n  1 energy level. The Balmer series corresponds to the transition to the n  2 energy level. The Paschen series corresponds to the transition to the n  3 energy level.

63. An Explanation of Atomic Spectra
Animation 6
Learn about atomic emission spectra and how neon lights work.

64. Quantum Mechanics Quantum Mechanics
5.3
Quantum Mechanics
Quantum Mechanics
How does quantum mechanics differ from classical mechanics?

65. 5.3
Quantum Mechanics
In 1905, Albert Einstein successfully explained experimental data by proposing that light could be described as quanta of energy.
The quanta behave as if they were particles.
Light quanta are called photons.

66. The Planck Equation
The energy of electromagnetic radiation is directly related to the frequency
E = hn
h is the Planck constant = x J.s
n is the frequency in Hz

67. END OF SHOW