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Welcome to Quarter 2!

Explore the arrangement of electrons within atoms, electromagnetic radiation, quantum concepts, atomic emission spectra, Bohr model, and more. Engage in calculations and concepts of energy levels, wavelengths, frequencies, and electron configurations.

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Welcome to Quarter 2!

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  1. Welcome to Quarter 2! • Make sure you have an electron note packet, your reference packet, and a calculator • Remember ALL makeup work is due today at 3:00pm – NO EXCEPTIONS

  2. Electrons in Atoms

  3. Electrons • Scientists pursue an understanding of how electrons are arranged within atoms • Electron arrangement plays a role in chemical behavior • Early 1900s- scientists observed that certain elements emit visible light when heated in a flame

  4. Wave Nature of Light • Electromagnetic Radiation- form of energy that exhibits wavelike behavior as it travels through space

  5. Wavelength • Wavelength- shortest distance between equivalent points on a wave • Symbol- λ • Unit- meters, centimeters, or nanometers (1 nm= 1x10-9m)

  6. Frequency • Frequency- the number of waves that pass a given point per second • Symbol- ν • Unit- Hertz (SI Unit)= (1/s)= (s-1) cycle per second

  7. Electromagnetic Spectrum • ALL electromagnetic waves, including visible light, travel at the speed of light c = 3.00x108 m/s • Wavelength must be in meters! C= λν

  8. Electromagnetic Spectrum • Encompasses all forms of electromagnetic radiation • The only differences in the types of radiation being their wavelengths and frequencies

  9. As the wavelength increases, the frequency decreases. • As the frequency increases, the energy increases. Electromagnetic Spectrum

  10. Calculations  1. Microwaves are used to transmit information. What is the wavelengthof a microwave having a frequency of 3.44 x109 Hz? • C= λν • C= 3.00x108 m/s • ν = 3.44 x109 Hz • λ = ???

  11. 3.00x108 = λ(3.44x109 Hz) λ= 8.72 x10-2 m

  12. 2. Yellow light has a wavelength of 589 nm. What is the frequency? 5.09x1014 Hz

  13. Particle Nature of Light (Honors) • Quantum Concept • Explained why colors of heated matter correspond to different frequencies and wavelengths • Max Plank- “matter can gain or lose only in small, specific amounts called quanta” • Quantum- the minimum amount of energy that can be gained or lost by an atom

  14. Particle Nature of Light (Honors) • Energy of a quantum is related to the frequency of the emitted radiation by the equation: Equantum= hv • E = energy • h = Plank’s Constant (6.63x10-34Js) • v = frequency • Joule (J)= SI unit for energy

  15. Particle Nature of Light (Honors) • Photon- a particle of EM radiation with no mass that carries a quantum of energy Ephoton= hv

  16. Example • Calculate the quantum of energy that an object can absorb from light with a wavelength of 477 nm. 4.17x10-19 J

  17. Atomic Emission Spectra • Set of frequencies of the electromagnetic waves emitted by atoms of the element • Example- The light of neon sign is produced by passing electricity through a tube filled with neon gas. Neon atoms release energy by emitting light.

  18. An atomic emission spectrum is characteristic of the element being examined and can be used to identify that element

  19. Bohr Model of the Atom • Proposed that the hydrogen atom has only certain allowable energy states • Ground State - lowest energy state of an atom • Excited State - higher energy state

  20. Bohr Model of the Atom • An electron must absorb energy to move from a lower energy level to a higher level. • Electrons do not stay in the excited state. When the electrons return to lower energy levels, energy is emitted.

  21. Heisenberg Uncertainty Principle • The Heisenberg Uncertainty Principle - states that it is impossible to know precisely both the velocity and position of a particle at the same time

  22. The Bohr Model • Using the Bohr Model from your packet, what is the wavelength of energy that is emitted when an electron falls from n= 6 to n=3? wavelength = 1094 nm

  23. B) What is the frequency of this radiation? 2.75x1014 Hz C) What is the energy of a photon of this radiation? (Honors) 1.82x10-19 J

  24. Atomic Orbital-a 3D region around the nucleus describing the electron’s probable location

  25. Atomic Orbitals • Energy Levels (n)- the major energy levels of an atom Ex: n = 1 energy level closest to the nucleus • Energy level → sublevel → orbital • Every orbital can hold up to 2 e-

  26. Sublevels are represented by the letters s, p, d, f lowest energy highest energy

  27. First 4 Principal Energy Levels Energy Level Sublevel Orbital Number of Electrons 1 s 1 2 s 1 2 2 p 3 6 (8 total e-) s 1 2 p 3 6 3 d 5 10 (18 total e-) s 1 2 p 3 6 4 d 5 10 f 7 14 (32 total e-) 2n2 = maximum # of electrons in energy level

  28. Electron Arrangement in Atoms • Electron Configurations-the arrangement of electrons in an atom

  29. Aufbau Principle: Electrons enter orbitals from lowest to highest energy

  30. Writing Electron Configurations • H (1e-) 1s1 energy level sublevel # e-

  31. Writing Electron Configurations • He (2e-) 1s2 • Li (3e-) 1s2 2s1 • Be (4e-) 1s22s2

  32. Writing Electron Configurations • B (5e-) 1s22s22p1 • C (6e-) 1s22s22p2 • Ne (10e-) 1s22s22p6

  33. Writing Electron Configurations • Na (11e-) 1s22s22p63s1 • Si (14e-) 1s22s22p63s23p2 • Cl (17e-) 1s22s22p63s23p5

  34. s p d f

  35. Noble Gas Configuration • Used to shorten electron configurations • Sodium: #11- instead of 1s22s22p63s1 can be shortened to [Ne] 3s1

  36. Examples • Write the shorthand electron configuration of Mn. [Ar]4s23d5 2. At [Xe]6s24f145d106p5

  37. Big Bang – Sheldon • Video

  38. Valence Electrons (V.E.) • Electrons in the atom’s outermost energy level • Determine the chemical properties of an element • V.E. are used in forming chemical bonds

  39. Examples • Write the electron configuration and give the number of valence e-. Mg 2 valence e- Br 7 valence e- V 2 valence e-

  40. Exceptions 1. Cu • not [Ar]4s23d9but [Ar]4s13d10 2. Ag •  [Kr]5s14d10 3.  Au • [Xe]6s14f145d10

  41. Exceptions 4. Cr [Ar]4s13d5 5. Mo [Kr]5s14d5

  42. Ions • Cations (+ ions) –remove e- • Anions (- ions) - add e- • O: • 1s22s22p4 • O2-: • 1s22s22p6 • O2- is isoelectronic with ________. Ne

  43. Examples Write the electron configuration for: • P3-: •  1s22s22p63s23p6 • Al3+: • 1s22s22p6 • Ba2+: • [Xe]

  44. Examples • Pb: • [Xe]6s24f145d106p2 • Pb2+: •  [Xe]6s24f145d10 • Pb4+: • [Xe]4f145d10

  45. Transition Metals • Fe: • [Ar]4s23d6 • Fe2+: •  [Ar]3d6 • Fe3+: • [Ar]3d5

  46. Transition Metals • Mn: •  [Ar]4s23d5 • Mn2+: •  [Ar]3d5 • Mn4+: • [Ar]3d3 • What is the highest possible charge for Mn? • +7

  47. Excited state: e- jumps to higher energy level • Ex: 1s22s22p63p6 • Ground state: normal e- configuration (lowest energy) • Ex: 1s22s22p63s23p1 • Blue Book: pg 358 # 37-39

  48. Orbital Diagrams • Use arrows to represent electrons • Use lines to represent orbitals • Every orbital can hold up to 2 e-

  49. s ____ • p ____ ____ ____ • d ____ ____ ____ ____ ____ • Lines represent orbitals.

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