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Modern Atomic Theory and the Periodic Table Chapter 10 :

1 Modern Atomic Theory and the Periodic Table Chapter 10 Hein and Arena Eugene Passer Chemistry Department Bronx Community College © John Wiley and Sons, Inc Version 2.0 12 th Edition

Chapter Outline:

2 Chapter Outline 10.1 A Brief History 10.2 Electromagnetic Radiation 10.3 The Bohr Atom 10.5 Atomic Structures of the First 18 Elements 10.6 Electron Structures and the Periodic Table 10.4 Energy Levels of Electrons

10.1 A Brief History:

3 10.1 A Brief History

Slide 4:

4

10.2 Electromagnetic Radiation:

5 10.2 Electromagnetic Radiation

Energy can travel through space as electromagnetic radiation.:

6 Energy can travel through space as electromagnetic radiation . Examples

Slide 7:

7 light from the sun x-rays microwaves radio waves television waves radiant heat All show wavelike behavior. Each travels at the same speed in a vacuum. 3.00 x 10 8 m/s

Characteristics of a Wave:

8 Characteristics of a Wave

Wavelength (λ):

9 Wavelength ( λ)

Slide 10:

10 wavelength (measured from peak to peak) wavelength (measured from trough to trough) 10.1 Light has the properties of a wave.

Frequency (ν):

11 Frequency ( ν)

Slide 12:

12 Frequency is the number of wavelengths that pass a particular point per second. 10.1

Speed (v):

13 Speed ( v)

Slide 14:

14 Speed is how fast a wave moves through space. 10.1

Slide 15:

15 Light also exhibits the properties of a particle. Light particles are called photons . Both the wave model and the particle model are used to explain the properties of light.

The Electromagnetic Spectrum :

16 The Electromagnetic Spectrum

Slide 17:

17 10.2 visible light is part of the electromagnetic spectrum X-rays are part of the electromagnetic spectrum Infrared light is part of the electromagnetic spectrum

10.3 The Bohr Atom:

18 10.3 The Bohr Atom

Slide 19:

19 At high temperatures or voltages, elements in the gaseous state emit light of different colors. When the light is passed through a prism or diffraction grating a line spectrum results.

Slide 20:

20 Line spectrum of hydrogen . Each line corresponds to the wavelength of the energy emitted when the electron of a hydrogen atom, which has absorbed energy falls back to a lower principal energy level. These colored lines indicate that light is being emitted only at certain wavelengths. Each element has its own unique set of spectral emission lines that distinguish it from other elements. 10.3

Niels Bohr:

21 Niels Bohr

Niels Bohr, a Danish physicist, in 1912-1913 carried out research on the hydrogen atom.:

22 Niels Bohr, a Danish physicist, in 1912-1913 carried out research on the hydrogen atom.

The Bohr Atom:

23 The Bohr Atom

Slide 24:

24 Electrons revolve around the nucleus in orbits that are located at fixed distances from the nucleus. 10.4 An electron has a discrete energy when it occupies an orbit.

Slide 25:

25 When an electron falls from a higher energy level to a lower energy level a quantum of energy in the form of light is emitted by the atom. 10.4 The color of the light emitted corresponds to one of the lines of the hydrogen spectrum.

Slide 26:

26 Different lines of the hydrogen spectrum correspond to different electron energy level shifts. 10.4

Slide 27:

27 Light is not emitted continuously. It is emitted in discrete packets called quanta . 10.4

Slide 28:

28 An electron can have one of several possible energies depending on its orbit. E 2 E 3 E 1 10.4

Slide 29:

29 Bohr’s calculations succeeded very well in correlating the experimentally observed spectral lines with electron energy levels for the hydrogen atom. Bohr’s methods did not succeed for heavier atoms. More theoretical work on atomic structure was needed.

Slide 30:

30 In 1924 Louis De Broglie suggested that all objects have wave properties. De Broglie showed that the wavelength of ordinary sized objects, such as a baseball, are too small to be observed. For objects the size of an electron the wavelength can be detected.

Slide 31:

31 In 1926 Erwin Schr ö edinger created a mathematical model that showed electrons as waves. Schr ö edinger’s work led to a new branch of physics called wave or quantum mechanics . Using Schr ö edinger’s wave mechanics, the probability of finding an electron in a certain region around the atom can be determined. The actual location of an electron within an atom cannot be determined.

Slide 32:

32 Based on wave mechanics it is clear that electrons are not revolving around the nucleus in orbits. Instead of being located in orbits, the electrons are located in orbitals. An orbital is a region around the nucleus where there is a high probability of finding an electron.

10.4 Energy Levels of Electrons:

33 10.4 Energy Levels of Electrons

According to Bohr the energies of electrons in an atom are quantized.:

34 According to Bohr the energies of electrons in an atom are quantized . The wave-mechanical model of the atom also predicts discrete principal energy levels within the atom

Slide 35:

35 The first four principal energy levels of the hydrogen atom. As n increases, the energy of the electron increases. 10.7 Each level is assigned a principal quantum number n .

Slide 36:

36 Each principal energy level is subdivided into sublevels. 10.7, 10.8

Slide 37:

37 Within sublevels the electrons are found in orbitals. An s orbital is spherical in shape. The spherical surface encloses a space where there is a 90% probability that the electron may be found. 10.9

Slide 38:

38 An electron can spin in one of two possible directions represented by ↑ or ↓. The two electrons that occupy an atomic orbital must have opposite spins. This is known as the Pauli Exclusion Principal. 10.10 An atomic orbital can hold a maximum of two electrons.

Slide 39:

39 Each p orbital has two lobes. Each p orbital can hold a maximum of two electrons. A p sublevel can hold a maximum of 6 electrons. 10.9 A p sublevel is made up of three orbitals.

Slide 40:

40 The three p orbitals share a common center. 10.10 p x p y p z The three p orbitals point in different directions.

Slide 41:

41 The five d orbitals all point in different directions. Each d orbital can hold a maximum of two electrons. A d sublevel can hold a maximum of 10 electrons. 10.11 A d sublevel is made up of five orbitals.

Slide 42:

42 Number of Orbitals in a Sublevel 10.8 10.10 10.11

Slide 43:

43 n=1 1s n=2 2s 2p 2p 2p n = 3 3s 3p 3p 3p 3d 3d 3d 3d 3d n = 4 4s 4p 4p 4p 4d 4d 4d 4d 4d 4f 4f 4f 4f 4f 4f 4f Distribution of Subshells by Principal Energy Level

The Hydrogen Atom:

44 The Hydrogen Atom In the ground state hydrogen’s single electron lies in the 1s orbital. Hydrogen can absorb energy and the electron will move to excited states. 10.12 The diameter of hydrogen’s nucleus is about 10 -13 cm. The diameter of hydrogen’s electron cloud is about 10 -8 cm. The diameter of hydrogen’s electron cloud is about 100,000 times greater than the diameter of its nucleus.

10.5 Atomic Structure of the First 18 Elements:

45 10.5 Atomic Structure of the First 18 Elements

To determine the electronic structures of atoms, the following guidelines are used.:

46 To determine the electronic structures of atoms, the following guidelines are used.

Slide 47:

47 No more than two electrons can occupy one orbital 10.10

Slide 48:

48 1 s orbital Electrons occupy the lowest energy orbitals available. They enter a higher energy orbital only after the lower orbitals are filled. For the atoms beyond hydrogen, orbital energies vary as s<p<d<f for a given value of n. 2 s orbital 10.10

Slide 49:

49 Each orbital in a sublevel is occupied by a single electron before a second electron enters. For example, all three p orbitals must contain one electron before a second electron enters a p orbital. 10.10

Nuclear makeup and electronic structure of each principal energy level of an atom. :

50 Nuclear makeup and electronic structure of each principal energy level of an atom. number of electrons in each sublevel number of protons and neutrons in the nucleus 10.13

Electron Configuration:

51 Electron Configuration Arrangement of electrons within their respective sublevels. 2 p 6 Principal energy level Type of orbital Number of electrons in sublevel orbitals

Orbital Filling:

52 Orbital Filling

Slide 53:

53 In the following diagrams boxes represent orbitals. Electrons are indicated by arrows: ↑ or ↓. Each arrow direction represents one of the two possible electron spin states.

Filling the 1s Sublevel:

54 Filling the 1 s Sublevel

Slide 55:

55 ↑ 1 s 2 ↓ H ↑ 1 s 1 Hydrogen has 1 electron. It will occupy the orbital of lowest energy which is the 1s. He Helium has two electrons. Both helium electrons occupy the 1s orbital with opposite spins.

Filling the 2s Sublevel:

56 Filling the 2 s Sublevel

Slide 57:

57 Li 1 s 2 2 s 2 The 1s orbital is filled. Lithium’s third electron will enter the 2s orbital. ↑ ↓ ↑ ↓ 1 s 2 s ↑ 1 s 2 2 s 1 Be ↑ ↓ The 2s orbital fills upon the addition of beryllium’s third and fourth electrons. 1 s 2 s

Filling the 2p Sublevel:

58 Filling the 2 p Sublevel

Slide 59:

59 B 1 s 2 2 s 2 2 p 1 1 s 2 s 2 p ↑ ↓ ↑ ↓ ↑ Boron has the first p electron. The three 2p orbitals have the same energy. It does not matter which orbital fills first. C 1 s 2 s 2 p ↑ ↓ ↑ ↓ The second p electron of carbon enters a different p orbital than the first p electron so as to give carbon the lowest possible energy. 1 s 2 2 s 2 2 p 2 ↑ ↑ ↑ N 1 s 2 s 2 p ↑ ↓ ↓ ↑ The third p electron of nitrogen enters a different p orbital than its first two p electrons to give nitrogen the lowest possible energy . 1 s 2 2 s 2 2 p 3 ↑ ↑

Slide 60:

60 ↑ 1 s 2 2 s 2 2 p 4 ↑ ↑ 1 s 2 s 2 p ↑ ↓ ↑ ↓ O There are four electrons in the 2p sublevel of oxygen. One of the 2p orbitals is now occupied by a second electron, which has a spin opposite to that of the first electron already in the orbital. ↑ ↓ ↑ ↓ ↑ 2 p F 1 s 2 s ↑ ↓ ↑ ↓ There are five electrons in the 2p sublevel of fluorine. Two of the 2p orbitals are now occupied by a second electron, which has a spin opposite to that of the first electron already in the orbital. 1 s 2 2 s 2 2 p 5 ↓

Slide 61:

61 ↑ ↓ ↑ ↓ ↑ ↓ 2 p Ne 1 s 2 s ↑ ↓ ↑ ↓ There are 6 electrons in the 2p sublevel of neon, which fills the sublevel. 1 s 2 2 s 2 2 p 6

Filling the 3s Sublevel:

62 Filling the 3 s Sublevel

Slide 63:

63 Na 1 s 2 2 s 2 2 p 6 3 s 1 1 s 2 s 2 p 3 s ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ The 2s and 2p sublevels are filled. The next electron enters the 3s sublevel of sodium. Mg 1 s 2 s 2 p 3 s ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ The 3s orbital fills upon the addition of magnesium’s twelfth electron. 1 s 2 2 s 2 2 p 6 3 s 2 ↓

Slide 64:

64

Slide 65:

65

10.6 Electron Structures and the Periodic Table:

66 10.6 Electron Structures and the Periodic Table

In 1869 Dimitri Mendeleev of Russia and Lothar Meyer of Germany independently published periodic arrangements of the elements based on increasing atomic masses.:

67 In 1869 Dimitri Mendeleev of Russia and Lothar Meyer of Germany independently published periodic arrangements of the elements based on increasing atomic masses. Mendeleev’s arrangement is the precursor to the modern periodic table.

Slide 68:

68 10.14 Horizontal rows are called periods Period numbers correspond to the highest occupied energy level.

Slide 69:

69 10.14 Elements with similar properties are organized in groups or families.

Slide 70:

70 10.14 Elements in the A groups are designated representative elements

Slide 71:

71 10.14 Elements in the B groups are designated transition elements

Slide 72:

72 10.15 For A family elements the valence electron configuration is the same in each column. The chemical behavior and properties of elements in a family are associated with the electron configuration of its elements.

Slide 73:

73 10.15 With the exception of helium which has a filled s orbital, the nobles gases have filled p orbitals.

The electron configuration of any of the noble gas elements can be represented by the symbol of the element enclosed in square brackets.:

74 The electron configuration of any of the noble gas elements can be represented by the symbol of the element enclosed in square brackets. B 1s 2 2s 2 2p 1 [He] 2s 2 2p 1 Cl 1s 2 2s 2 2p 6 3s 2 3p 5 [Ne] 3s 2 3p 5 Na 1s 2 2s 2 2p 6 3s 1 [Ne] 3s 1

The electron configuration of argon is :

75 The electron configuration of argon is Ar 1s 2 2s 2 2p 6 3s 2 3p 6 Ca 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 [Ar] 4s 2 K 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 [Ar] 4s 1 The elements after argon are potassium and calcium Instead of entering a 3d orbital, the valence electrons of these elements enter the 4s orbital.

Slide 76:

76 10.16 d orbital filling The number of a d orbital is 1 less than its period number Arrangement of electrons according to sublevel being filled.

Slide 77:

77 10.16 f orbital filling The number of an f orbital is 2 less than its period number Arrangement of electrons according to sublevel being filled.

Slide 78:

78 10.17 A period number corresponds to the highest energy level occupied by electrons in the period.

Slide 79:

79 10.17 The group numbers for the representative elements are equal to the total number of outermost electrons in the atoms of the group. The elements of a family have the same outermost electron configuration except that the electrons are in different energy levels.

Slide 80:

80 The End

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