1. Pre-IB/Pre-AP CHEMISTRY
Chapter 4 – Arrangement of Electrons in Atoms

2. Section 1 Objectives
Be able to define: electromagnetic radiation, electromagnetic spectrum, wavelength, amplitude, frequency, photoelectric effect, quantum(pl. quanta), photon, ground state, excited state, line emission spectrum, continuous spectrum, energy level.

3. Section 1 Objectives
Be able to explain the mathematical relationship between speed, wavelength, and frequency of a wave.
Be able to describe what is meant by the wave-particle duality of light.
Be able to discuss how the photoelectric effect and the line emission spectrum of hydrogen lead to the development of the atomic model.

4. Section 1 Objectives
Be able to describe the Bohr model of the atom.

6. Wave
A wave is a method of transferring energy. This transfer does not require matter as a medium.

7. Wave
Some waves travel through matter (sound, water waves, etc.).

8. Wave
Some waves do not require matter and can travel through empty space (light).

9. Wave Properties
Waves can be described by their wavelength, amplitude, and frequency.

10. Wavelength A crest is the highest point on a wave.
A trough is the lowest point on a wave.
Crest
Trough

11. Wavelength
Wavelength is simply the length of a wave. It is the distance between two crests or two troughs.
Wavelength is measured in m, mm, or nm.
Wavelength
Crest
Trough

12. Amplitude
Amplitude is simply the height of a wave. It is the distance between the crest and trough of a wave.
Amplitude is measured in units of distance.
Amplitude

13. Frequency
Frequency is the number of waves passing a given point in a given time.
Frequency describes the energyof a wave.

14. Frequency
Frequency describes the energy of a wave: the higher the frequency, the greater the energyof that wave.

15. Frequency
Frequency is measured in hertz or cycles per secondor vibrations per second or 1/sec or sec-1 - they all mean the same thing.

16. Frequency
As the wavelength increases, frequency decreases. This is called an inverse relationship.

17. Wave Properties
Wavelength and amplitude give waves their distinctive properties. For example, the loudness of a sound wave is its amplitude, the color of visible light is its wavelength.

18. Types of waves
Electromagnetic waves do not require a medium or matter in order to travel. Light is an example.

19. Light Light is an electromagnetic wave.
Visible light is a small part of the electromagnetic spectrum that humans are able to see.

20. Light
The electromagnetic spectrum consists of different kinds of light of different wavelengths.

21. EM Spectrum

22. EM Spectrum

23. EM Spectrum

24. EM Spectrum

25. EM Spectrum

26. EM Spectrum

27. EM Spectrum

28. Light Interactions
White light is light consisting of all colorsof visible light. These colors are visible in a rainbow or through a prism.

29. Velocity
The velocity of a wave is a product of its frequency and wavelength.
v= fl
v = velocity
f = frequency
l = wavelength

30. Velocity
The velocity of light through a vacuum(c) is about 3.0 x 108 m/sec. It is slightly slower through matter.

31. Photoelectric Effect
Photoelectric effect refers to the emission of electrons from a metal when light shines on the metal.

32. Photoelectric Effect
It was found that light of a certain frequency would cause electrons to be emitted by a particular metal. Light below that frequency had no effect.

33. Emission Spectra
If an object becomes hot enough it will begin to emit light.

34. Emission Spectra
Max Planck suggested that hot objects emit light in specific amounts called quanta (sing. quantum).

35. Equantum= hf Emission Spectra
Planck showed the relationship between a quantum of energy and the frequency of the radiation.
Equantum= hf
Equantum= energy of a quantum in joules
h = Planck’s constant
f = frequency

36. Wave-Particle Duality
Einstein later said that light had a dual nature – it behaved as both a particle and a wave.

37. Wave-Particle Duality
Each particle of light, Einstein said, carries a particular quantum of energy.

38. Wave-Particle Duality
Einstein called the “particles” of light photons which had zero mass and carried a quantum of energy. The energy is described as:
Ephoton= hf

39. Photoelectric Effect
Einstein explained photoelectric effect by saying in order for an electron to be ejected from a metal, the photon striking it must have enough energy to eject it.

40. Attraction
Different metals have stronger attraction for their electrons than other. Therefore, some must absorb more energy than others to emit electrons.

41. Ground State
The lowest energy state of an atom is called its ground state.

42. Excited State
When a current is passed through a gas at low pressure, the atoms become “excited.”

43. Excited State
Atoms in an excited state have a higher potential energy than their ground state.

44. Excited State
An “excited” atom will return to its ground state by releasing energy in the form of electromagnetic radiation.

45. Emission Spectra
Elements will emit radiation of certain frequencies. This reflects the energy states of its electrons and is called a bright-line or emission spectrum.

46. Emission Spectra
The emission spectrum of an element is like its “fingerprint”.
Sodium
Helium
Mercury

47. Energy Levels
Studying the emission spectrum of hydrogen lead Niels Bohr to the idea of energy levels.

48. Energy Levels
The spectrum Bohr and others observed was the result of excited electrons releasing photons as they returned to their ground states.

49. Energy Levels
The difference in the energy of photons was reflected in the different frequencies of light they observed.

50. Section 2 Objectives
Be able to define: diffraction, interference, Heisenberg Uncertainty Principle, Quantum Theory, quantum numbers, principal quantum number, angular momentum quantum number, magnetic quantum number, spin quantum number.
Be able to distinguish between the Bohr model and the quantum model of the atom.

51. Section 2 Objectives
Be able to explain how the Heisenberg Uncertainty Principle and the Schroedinger Wave Equation led to the idea of atomic orbitals.
Be able to list the four quantum numbers that describe each electron in an atom.

52. Section 2 Objectives
Be able to relate the number of sublevels corresponding to each of an atom’s main energy levels, the number of orbitals per sublevel, and the number of orbitals per main energy level.

54. Electrons as Waves
French scientist Louis De Broglie demonstrated that electrons had a dual nature also.

55. E = hf Electrons as Waves
De Broglie showed that electrons behaved as waves confined to the atom. The energy of those electrons could be found like that of waves:
E = hf

56. Electrons as Waves
Electron beams were shown to exhibit the wave properties of diffraction and interference.

57. Heisenberg Uncertainty
Werner Heisenberg tried to find the location and velocity of electrons in the atom.

58. Heisenberg Uncertainty
Heisenberg found that it is impossible to simultaneously determine the position and velocity of an electron or any other particle (The Heisenberg Uncertainty Principle).

59. Schrödinger Wave Equation
Erwin Schrödinger said that electrons had a dual nature(like light) and treated them as waves.

60. Quantum Theory
Schrödinger’s wave equation and Heisenberg’s Uncertainty Principle laid the foundation of modern quantum theory.

61. Quantum Theory
Quantum theory describes mathematically the wave properties of electrons and other very small particles.

62. Quantum Theory
According to the Bohr model we should be able to predict the location and velocity of an electron at any time.

63. Quantum Theory
Quantum theory disagrees with the Bohr model and says that electrons can be found in regions of high probability but cannot be pinpointed.

64. Orbitals
Quantum theory describes electrons as inhabiting a three-dimensional region around the nucleus that indicates their probable locations. These regions are called orbitals.

66. Orbitals
Scientists use quantum numbers to describe orbitals. These numbers describe the properties of the orbitals and the electrons which occupy them.

67. s and p orbitals

68. d orbitals

69. Pg. 110