1. Chapter 4 Arrangement of Electrons in Atoms

2. Section 4-1 The Development of a New Atomic Model

3. A new Atomic Model  The Rutherford model was created in 1913, but did not accurately describe the behavior of electrons.  A new model was developed in the early twentieth century, which arose as a result of the study of light.

4. Light as a Wave  Prior to 1900, scientists believed light behaved as a wave.  This is because light is a type of electromagnetic radiation.  This is a form of energy that exhibits wave behavior as it travels through space.

5. Electromagnetic Radiation  All of the forms of electromagnetic radiation make up the electromagnetic spectrum. (p. 92)  All electromagnetic radiation moves through space at a speed of 3.0 * 10 8 meters per second.

6. Characteristics of Waves  Waves are repetitive.  This allows them to have two characteristic properties.  Wavelength – ( ג ) the distance from crest to crest or trough to trough on adjacent waves.  Frequency – (v) The number of waves to pass a specific point in a given amount of time, usually a second.

7. Frequency  Measured in waves per second.  One wave per second is a Hertz (Hz).  Frequency and wavelength are inversely proportional.  This is shown by the formula c = ג v, where c = speed of light.

8. There are five parts to a wave  Wavelength  Amplitude  Origin  Crest  Trough

9. The Photoelectric Effect  In the early 1900’s the wave theory of light began to be criticized.  Two experiments could not be explained by it.  The first of these was the photoelectric effect.

10. What was the Photoelectric Effect?  Electrons called photoelectrons are emitted from some metals, especially group IA metals, when light shines on them.  Light of too low a frequency does not do this, even if it is extremely bright.  Light of a higher frequency always does this, even if it is very dull.

11. What was the Photoelectric Effect?  Scientists theorized that frequency of light had some impact on the photoelectric effect.  Wave theory of light predicts that eventually any wave will produce the effect.  Scientists could not explain this phenomena.

12. The Photoelectric Effect  Coincidentally, the photoelectric effect is used commercially today.  Every solar panel, from the large ones to power homes to the small ones in calculators use this phenomena to generate power.

13. Max Planck (1858 – 1947)  Studied the emission of light by hot objects.  Hot objects do not give off light continuously, as we would expect if it were a wave.  Light is given off in small, discrete units called quanta.

14. Max Planck (1858 – 1947)  A quantum is the minimum energy that can be lost or gained by an atom.  Planck described this relationship by the formula E = hv, where h is the value of Planck’s Constant, 6.626 * 10 -34 J.

15. Albert Einstein (1879 – 1955)  Expanded the work of Planck.  Decided that all electromagnetic radiation existed as both a wave and a particle.  Each particle contains a quantum of energy.  Einstein called these particles photons.

16. Photons  A photon is a particle of electromagnetic radiation having no mass and carrying a quantum of energy.  The energy of a photon depends on the frequency of the radiation.  Einstein used this theory to explain the photoelectric effect.

17. Explanation of the Photoelectric Effect  Electromagnetic radiation is absorbed only in whole numbers of photons.  A photon must have a minimum amount of energy to remove an electron.  According to Planck’s formula, E = hv, the energy of a photon is determined by frequency.  If the frequency of a photon is too low, no electron is emitted.

18. Atomic Emission Spectra  When a current is passed through a gas at a low pressure, the potential energy of some of the gas atoms increases.  The lowest energy state of an atom is its ground state.  The state in which it has a higher potential energy is its excited state.

19. Atomic Emission Spectra  When an atom in the excited state returns to the ground state, it gives off a photon of energy.  The eye sees this as colored light.  When light is passed through a prism, it is separated into bands of color.

20. Atomic Emission Spectra When colored light is passed through a prism, only specific frequencies of color are produced. These are part of the element’s atomic emission spectrum.

21. Atomic Emission Spectra  Classical Theories of physics predict that atoms would be excited by any amount of energy, and that all atoms would produce a full rainbow of colors.  Since they were not, a new theory had to be developed that would describe atoms. This new theory was called Quantum Theory.

22. Quantum Theory  Quantum Theory predicts that when an atom returns to its ground state from the excited state, it releases a photon of light.  The energy of a photon is equal to the difference between the atom’s ground state and its excited state.  The fact that atoms only give off specific bands of visible, ultraviolet, and infrared light, indicates that electrons have constant energy states.

23. Niels Bohr (1885 – 1962)  Developed a new model of the atom in 1913.  Electrons circle the nucleus only in allowed paths or orbits.  The electron can move from one orbit to the next, but cannot exist between orbits.  The farther away from the nucleus an electron is, the more energy it has.

24. Bohr Model of the Atom

25. Section 4-2 The Quantum Model of the Atom

26. Louis de Broglie (1892 – 1987)  Suggested in 1924, that electrons should be considered as waves confined to the area around the nucleus of an atom.  This was soon confirmed, but left the question, if electrons are waves, then where are they in the atom?

27. Werner Heisenberg (1901 – 1976)  Tried to find the answer in 1927.  Electrons are detected by their interaction with photons.  An attempt to locate an electron with a photon knocks the electron off course.

28. The Heisenberg Uncertainty Principle  If you locate an electron, you can only do it by changing its velocity.  The Heisenberg Uncertainty Principle states that it is impossible to know the location and velocity of an electron at the same time.  This is currently being expanded to all aspects of research. Basically, the results don’t exist until you look for them.

29. Erwin Schrödinger (1887 – 1961)  Used Heisenberg’s Principle to develop an equation that treated electrons as waves.  Result of his work was quantized energy.  This lead to quantum theory.

30. Quantum Theory  Describes mathematically, the wave properties of electrons and other small particles.  Does not support Bohr’s idea of circular orbits.  Describes that electrons exist around the nucleus in three dimensional regions called orbitals.

31. The Quantum Mechanical Model  This lead to a new model of the atom.  Instead of circular orbits, electrons exist in an electron cloud.  The electron cloud has a region within it, where electrons can be found 90% of the time.

32. Linus Pauling (1901 – 1994)  First proposed the existence of orbitals.  Described them as a three dimensional region around the nucleus that indicates the probable location of an electron.  Orbitals are not physical shapes that can be seen or measured.

33. Atomic Orbitals  Orbitals are a numeric answer to a probability formula that are graphed as physical models.  Pauling helped develop these shapes, and the quantum numbers that describe them.

34. Other contributions by Pauling  Described the bonding of organic molecules.  Determined the structure of hemoglobin and its role in sickle cell anemia.  Only man to win the Nobel prize twice.

35. Quantum Numbers  Determine the properties of atomic orbitals and of the electrons in those orbitals.  There are four quantum numbers.  Principle Quantum Number  Angular Momentum Quantum Number  Magnetic Quantum Number  Spin Quantum Number

36. Principle Quantum Number  Indicates the main energy level occupied by the electron.  Symbolized by the letter n.  Must be a positive integer.  As n increases, the distance from the electron to the nucleus increases.  If n=1, the electron is in the lowest energy level.  Energy levels are also called shells.

37. Angular Momentum Quantum Number  Indicates the shape of the orbital.  Every energy level except the first contains groups of orbitals with different shapes known as sublevels.  The numbers range from 0 – 4, but we commonly use the letters s, p, d, and f to represent them.  The first energy level has only an s sublevel, the second level has an s and a p sublevel. The third energy level has s, p, and d sublevels. All others have all four of them.

38. Magnetic Quantum Number  Indicates the orbital within the sublevel.  S sublevels contain only 1 orbital, p sublevels contain 3, d sublevels contain 5, and f sublevels contain 7.  Usually use letters to represent this, normally these are x, y, and z to correspond to a Cartesian plane.  Usually used only for p orbitals.

39. Spin Quantum Number  Indicates the individual electron in the orbital.  Only two electrons may occupy an orbital.  The two electrons must spin in opposite directions.  There are only two spin quantum numbers, +1/2 or -1/2.  +1/2 indicates a clockwise spin, -1/2 is counterclockwise.

40. S and P orbitals

41. D orbitals

43. F orbitals

45. Section 4-3 Electron Configurations

46.  The four quantum numbers allow us to distinguish any electron in an atom. This arrangement is the electron configuration.  Electron Configurations are only done for the ground state.  Three rules determine how an electron configuration is done.

47. Rules for Electron Configurations  Aufbau Principle  Pauli Exclusion Principle  Hund’s Rule

48. Aufbau Principle  An electron must occupy the lowest energy orbital available.

49. Pauli Exclusion Principle  No two electrons in the same atom can have the same four quantum numbers.

50. Hund’s Rule  Orbitals of equal energy are each occupied by one electron before two electrons can occupy any orbital in the same sublevel.

51. Electron Configurations  There are two types of electron configurations.  Orbital Notation – Uses lines to show orbitals and arrows to show electrons.  Electron Configuration Notation – Uses superscript numbers to show the number of electrons in an orbital.

52. Three Terms are Important for Electron Configurations  Highest Occupied Energy Level – The main electron containing energy level with the highest principle quantum number.  Inner Shell Electrons – Electrons that are not in the highest occupied energy level.  Valence Electrons – Electrons that are in the highest occupied energy level.

53. Noble Gas Configurations  Noble Gases are group VIIIA elements.  All noble gases have electron configurations that end in p 6.  To put an electron configuration into noble gas configuration, select the first noble gas with an atomic number lower than the element you are working with. Put its symbol in brackets, and then fill in the rest of the electron configuration after the noble gas.

54. Exceptions to Noble Gas Configurations  Half full orbitals are more stable than orbitals that are more than half full.  Because of orbital overlap, the elements Cr, Cu, Mo, and Ag steal electrons from one orbital to fill or half – fill another.  As we move into the rare earth elements, there are several exceptions, and these will not be covered on a test.

55. Shortcut to Electron Configurations  The periodic table can be used to determine how an electron configuration should end. An element in the s block ends in an s orbital, the p block in a p orbital, etc.  The number of columns in an element is tells how many electrons in that block it has.  The number of periods down an element is indicates its highest occupied energy level.  For d – block elements, subtract one from the period number. For f – block elements, subtract two.