1. Chapter 11 11.1 Rutherford’s Atom: To describe Rutherford’s model of the atom. 11.2 Energy and Light: To explore the nature of electromagnetic radiation. 11.3: Emission of Energy by Atoms: To see how atoms emit light.

2. 11.1 Rutherford’s Atom Gold Foil experiment: small dense nucleus containing protons and neutrons QUESTIONS: What are the electrons doing? How are the electrons arranged and how do they move? Rutherford suggested that electrons might revolve around the nucleus.

3. 11.2: Energy and Light Electromagnetic Radiation: (energy transmitted as light). Figure 11.4  Rays X rays UV Visible Infrared Micro Radio Wavelength 10 -12 to 10 4 Wave: wavelength( ) frequency(  and speed ( how fast)

4. 11.2: Energy and Light Light can behave as waves or photons (stream of tiney packets of energy). Wave-particle nature of light. Red light- lower energy. Longer the wavelength the lower the energy of its photons.

5. 11.3: Emission of Energy by Atoms Atoms emit light of different energy. Heat from the flame allows atoms to absorb energy-we say that the atoms become excited. Some of this excess energy is released as light.

6. 11.3: Emission of Energy by Atoms Li+ : gives off a beautiful, deep-red color (emits photons of red light) Cu2+: burns green Na+: burns yellow-orange color Excited Li atom Photon of red light emitted Li atom in lower Energy state Reminder: High energy (short Wavelengths) Red light has long wavelength, lower energy

7. 11.4: The Energy Levels of Hydrogen Objective: To understand how the emission spectrum of hydrogen demonstrates the quantized nature of energy.

8. Excited Li atom Photon of red light emitted Li atom in lower Energy state 11.4: The Energy Levels of Hydrogen An atom with excess energy is called the EXCITED STATE The lowest possible enery state of an atom is called the GROUND STATE ENERGY contained in the photon corresponds to the change in energy that The atom experiences in going from the excited state to the lower state.

9. Excited Li atom Photon of red light emitted Li atom in lower Energy state 11.4: The Energy Levels of Hydrogen An atom with excess energy is called the EXCITED STATE The lowest possible enery state of an atom is called the GROUND STATE ENERGY contained in the photon corresponds to the change in energy that The atom experiences in going from the excited state to the lower state.

10. 11.4: The Energy Levels of Hydrogen An atom with excess energy is called the EXCITED STATE The lowest possible enery state of an atom is called the GROUND STATE ENERGY contained in the photon corresponds to the change in energy that The atom experiences in going from the excited state to the lower state. Energy 4 excited states Ground state H atoms only certain types of photons are produced. We only see certain colors. This means H has discrete energy levels (with specific wavelengths NEVER EMIT photons w/ energies in between

11. 11.4: The Energy Levels of Hydrogen Energy levels of hydrogen are QUANTIZED That is only certain values are allowed. Scientists have found that the energy levels of all atoms are quantized. Energy 4 excited states Ground state H atoms only certain types of photons are produced. We only see certain colors. This means H has discrete energy levels (with specific wavelengths NEVER EMIT photons w/ energies in between

12. 11.5: The Bohr Model of the Atom Objective: To learn about Bohr’s model of the hydrogen atom.

13. 11.5: The Bohr Model of the Atom Neils Bohr: at the age of 25 earned his PhD in Physics. Constructed a model of the hydrogen atom with quantized energy levels that agreed with the hydrogen emission results. Current theory of atomic structure Is not the same as the Bohr model. Electrons do not move around the Nucleus in circular orbits. WE DO NOT KNOW EXACTLY How the electrons move in an atom.

14. 11.6: The Wave Mechanical Model of the Atom Objective: To understand how the electron’s position is represented in the wave mechanical model.

15. 11.6: The Wave Mechanical Model of the Atom By the mid 1920’s, Bohr’s model was shown To be incorrect. Victor deBroglie (France) Erwin Schrodinger (Austria) Suggested that because light seems to have both wave and particle characteristics. Electrons also exhibit both of these characteristics.

16. 11.6: The Wave Mechanical Model of the Atom Wave Mechanical Model of the atom -electron states are described by orbitals. Orbitals are nothing like orbits. H atoms can be pictured as a probability of An electron being in a certain position. It does not tell you when the electron occupies this space or how it moves.

17. 11.7: The Hydrogen Orbitals Objective: To learn about the shapes of orbitals designated by s, p and d.

18. 11.7: The Hydrogen Orbitals An orbital is a probability map for the location of an electron. Chemists arbitrarily define its shape as 90% of the total electron probability HYDROGEN 1s orbital Lowest energy state (the ground state)

19. 11.7: The Hydrogen Orbitals Principal Energy Levels Sublevels

20. Principal energy level Shape

21. Figure 11.25: The three 2p orbitals.

23. Figure 11.27: Relative sizes of the spherical 1s, 2s, and 3s orbitals of hydrogen.

24. Figure 11.28: The shapes and labels of the five 3d orbitals.

25. Summary 1)The number tells the principal energy level. 2)The letter tells the shape. The letter s means a spherical orbital; the letter p means a two-lobed orbital. The x, y, or z subscript on a p orbital label tells along which of the coordinate axes the two lobes lie.

26. 11.8: The Wave Mechanical Model: Further Development Objective: To review the energy levels and orbitals of the wave mechanical model of the atom. To learn about electron spin.

27. 11.8: The Wave Mechanical Model: Further Development The last property of electrons is spin. Pauli’s exclusion principle says that no orbital can contain more than 2 electrons and those electrons must have opposite spins.

28. Summary of Wave Mechanical Model Electrons are organized in 4 levels These are called the quantum numbers n, l, m l, m s n : Principal quantum number(principal energy level) l: azimuthal quantum number (shape) is the sublevel l=0 for s sublevel, l=1 for p sublevel, l=2 for d sublevel, and l=3 for f sublevel m l =magnetic quantum number When l=0, m l =0 l=1, m l = -1, 0,+1 l=2, m l =-2, -1, 0, 1, 2 l=3, m l =-3, -2, -1, 0, 1, 2, 3 m s =spin quantum number +1/2, -1/2 Add n+l = the higher the number the higher the energy level

29. THREE RULES Aufbau Principle: “building up” fill lower energy levels first Hund’s Rule: every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin.orbitals Pauli exclusion principle: In each energy level, electrons go in as opposite spins

30. 11.9: Electron Arrangements in the First 18 Atoms on the Periodic Table Objective: To understand how the principal energy levels fill with electrons in atoms beyond hydrogen. To learn about valence electrons and core electrons.

31. 11.9: Electron Arrangements in the First 18 Atoms on the Periodic Table Electrons fill the orbitals in the following order: 1s 2s 2p 3s 3p Electron arrangement or Electron configuration 1s 1 orbital diagram(box diagram) 1s

32. 11.9: Electron Arrangements in the First 18 Atoms on the Periodic Table Valence electrons: electrons in the outermost (highest) principal energy level of an atom. These electrons are the most important because they are involved in bonding. Inner electrons are called the core electrons. Atoms of elements in the same group(vertical) have the same number of electrons in a given type of orbital (sublevel) except that the orbitals are in different principal energy levels.

33. 11.9: Electron Arrangements in the First 18 Atoms on the Periodic Table So another way of writing them are: K 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1, or [Ar]4s 1. Ca 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2, or [Ar]4s 2

34. 11.9: Electron Arrangements in the First 18 Atoms on the Periodic Table Atoms of elements in the same group(vertical) have the same number of electrons in a given type of orbital (sublevel) except that the orbitals are in different principal energy levels. Elements with the same valence electron arrangement show very similar chemical behavior.

35. 11.10: Electron Configurations and the Periodic Table Objective: To learn about the electron configuration of atoms with Z greater than 18.

36. 11.10: Electron Configurations and the Periodic Table 1) In a principal energy level that has d orbitals, the s orbital from the next level fills before the d orbitals 2) The lanthanides, fill the seven 4f orbitals. 3) The actinide series fill the seven 5f orbitals. Except for helium, the group numbers indicate the sum of electrons in the ns and np orbitals in the highest principal energy level that contains electrons.

37. The same type of elements occur periodically so that groups of elements show similar chemistry indicate the total # of valence electrons For the atoms in this group. Group 5 ns 2 np 3 Main group elements Or representative elements

38. 11.10: Electron Configurations and the Periodic Table 1)The group labels for Groups 1,2,3,4,5,6,7+8 indicate the total number of valence electrons for the atoms in these groups. 2)The elements in Groups 1,2,3,4,5,6,7,8 are often called the main-group elements, or representative elements.

39. 11.11 Atomic Properties and the Periodic Table Objective: To understand the general properties in the periodic table.

40. 11.11 Atomic Properties and the Periodic Table Objective: To understand the general properties in the periodic table.

42. METALS: lustrous appearance, the ability to change shape without breaking and excellent conductivity of heat and electricity NONMETALS: TEND to GAIN ELECTRONS to form NEGATIVE IONS. METALS: TEND to LOSE ELECTRONS to Form POSITIVE IONS

43. H Li Na K Rb Cs Cs > Rb > K > Na > Li LosesLeast anlikely electronto lose mostan easilyelectron Most chemically active metals Most chemically active nonmetals

45. Ionization Energies Increase I O N I Z A T I O N E N E R G Y DECREASES Ionization Energy of an atom is the energy required to remove an electron from an individual atom in the gas phase. Metals have Low ionization energies