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Chapter 13

Chapter 13. Electron Configuration. Evolution of Atomic Models. First Model: John Dalton: The atom is a solid, indivisible mass. Evolution of Atomic Models, cont. Second Model: JJ Thompson: He discovered the electron.

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Chapter 13

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  1. Chapter 13 Electron Configuration

  2. Evolution of Atomic Models • First Model: • John Dalton: • The atom is a solid, indivisible mass

  3. Evolution of Atomic Models, cont. • Second Model: • JJ Thompson: • He discovered the electron. • He developed the plum pudding model. This was a positive material (pudding) with negative charges (raisins) stuck on it.

  4. Evolution of Atomic Models, cont. • Third Model: • Ernest Rutherford: • Discovered the nucleus • His model had electrons surrounding a dense nucleus. • The rest of the atom is empty space.

  5. Evolution of Atomic Models, cont. • Fourth Model: • Niels Bohr: • Electrons are arranged around concentric paths known as orbits. • Electrons in a particular path have a fixed amount of energy

  6. Niels Bohr Model, cont. • Energy Levels: • The region around the nucleus where the electron is likely to be moving. • The lowest energy level has the lowest amount of energy. (Think of ladders.) • An electron cannot be between energy levels. • Quantum: • the amount of energy required to move an electron from its present energy level to a higher one. • Quantum Leap = abrupt change

  7. Niels Bohr Model, cont. • Energy Levels, cont.: • The energy levels are not equally spaced. • The further away from the nucleus you are, the closer together the levels are. • Why is it easier to pull electrons away from the atom the further out you are?

  8. Quantum Mechanical Model • Erwin Schrödinger: • Used the quantum theory to write and solve a mathematical equation describing the location and energy of an electron in a Hydrogen atom. • This became known as the Quantum Mechanical Model.

  9. Quantum Mechanical Model, cont. • Quantum Mechanical Model: • Modern description, primarily mathematical of the behavior of electrons in atoms • The energy of electrons is restricted to certain values. • The electrons do not have exact paths around the nucleus. • This strictly ESTIMATES the probability of finding an electron in a certain position.

  10. Quantum Mechanical Model, cont. • The probability of finding an electron is represented by a fuzzy cloud. • Theoretically this area could go on forever, so they show the area where the electron is found 90% of the time. • From this, we can start to see the shapes of the different clouds.

  11. Atomic Orbitals • Energy Levels are also known as principal quantum numbers, represented with the letter n. • n=1, n=2, n=3, n=4, n=5, n=6, n=7 • Hint: This equals the row number on the periodic table. 

  12. Atomic Orbitals, cont.

  13. Atomic Orbitals, cont. • Within each energy level (principle quantum number) there are energy sublevels. • n=1 1 sublevel 1s • n=2 2 sublevels 2s, 2p • n=3 3 sublevels 3s, 3p, 3d • n=4 4 sublevels 4s, 4p, 4d, 4f • n=5 5 sublevels 5s, 5p, 5d, 5f, 5g (These are sublevels)

  14. Sublevels • The s orbital: • Spherical • One orbital • Always the first orbital filled for any energy level. • 1s, 2s, 3s, 4s, 5s, 6s, 7s 3s 2s 1s

  15. Sublevels, cont. • The p orbitals: • Dumbbell shaped • Three orbitals: x, y, and z • Electrons go into the p orbital after the s orbital is filled. pz px py

  16. Sublevels, cont. • The d orbitals: • Cloverleaf shaped • There are 5 orbitals: d1, d2, d3, d4, d5 • Electrons fill the d orbitals after the s and p orbitals are completely filled. d2 d4 d1 d3 d5

  17. Sublevels, cont. • The f orbitals: • These are complexly shaped. • There are 7 orbitals: f1, f2, f3, f4, f5, f6, f7 • These fill after the s, p, and d orbitals are completely filled. f3 f1 f2

  18. f Orbitals, cont. f5 f6 f4 f7

  19. Energy Levels with Orbitals • n=1 1s • n=2 2s, 2px, 2py, 2pz • n=3 3s, 3px, 3py, 3pz, 3d1, 3d2, 3d3, 3d4, 3d5 • n=4 4s, 4px, 4py, 4pz, 4d1, 4d2, 4d3, 4d4, 4d5, 4f1, 4f2, 4f3, 4f4, 4f5, 4f6, 4f7 • How do you remember how many orbitals there are in an energy level? • n2

  20. Electrons • In any orbital, there is a maximum of 2 electrons. • How do you figure out how many electrons are in an energy level? • 2n2 • Example: • n=1, there is a max of 2 electrons • n=2, there is a max of 8 electrons

  21. Periodic Table and All of This • Okay, so where does everything fit on the periodic table? • Group 1A and 2A: s orbitals • Group 3A – 0: p orbitals • Transition Metals: d orbitals • Inner transition Metals: f orbitals

  22. Let’s Try  • Give the electron configuration for Be: • 1s22s2 • Give the electron configuration for H: • 1s1 • Give the electron configuration for S: • 1s22s22p63s23p4

  23. Okay, Now the Harder Ones • Give the electron configuration for Mn: • 1s22s22p63s23p64s23d5 • Give the electron configuration for Au: • 1s22s22p63s23p64s23d104p65s24d105p66s24f145d9 • Give the electron configuration for Ra: • 1s22s22p63s23p64s23d104p65s24d105p66s24f14 5d10 6p67s2

  24. So… • When you get to the transition metals, the number in front of the d, the energy level, is ONE less then the row you are on. • When you get to the inner transition metals, the number in front of the f, the energy level, is TWO less then the row you are on.

  25. Section 13.2 • Electron Configuration: • The ways in which electrons are arranged around the nuclei of atoms • There 3 rules to finding the electron configuration: • Aufbau Principle • Pauli Exclusion Principle • Hund’s Rule

  26. Aufbau Principle • This principle states that electrons enter orbitals of lowest energy first • Various orbitals within the sublevel are equal in energy • For each energy level the s sublevel is lowest in energy.

  27. Pauli Exclusion Principle • This states that at most 2 electrons per atomic orbital. • If one electron is currently in an orbital, the 2nd electron can enter only if it spins in the opposite direction. One spins clockwise, one spins counterclockwise. • Shown:

  28. Hund’s Rule • This states that when electrons occupy an orbital of equal energy, one electron enters each orbital until all orbitals contain one electron with parallel spins. • Example: 2px 2py 2pz

  29. Practice • Write the orbital configurations and electron configurations for the following elements: • H • C • Na • Ti

  30. Exceptional Electron Configurations • You can obtain correct electron configurations up to #23, V, after that it may change. • Reason: half-filled orbitals are more stable than “off balance” orbitals. • Example: Cu has 29 electrons • 1s22s22p63s23p64s23d9 is not as stable as 1s22s22p63s23p64s13d10

  31. Section 13.3 • Electromagnetic radiation: • Includes radio waves, microwaves, infrared waves, visible light, ultraviolet waves, x-rays, and gamma rays. • Sunlight (white light) consists of a continuous range of wavelengths and frequencies. • When passed through a prism the different wavelengths separate into a spectrumof colors. • For example: a rainbow

  32. Atomic Emission Spectrum • Every element emits light when it is excited by the passage of an electricity through its gas or vapor. • The atoms first absorb energy, then lose the energy as they emit light. • Passing this light emitted by and element through a prism gives the atomic emission spectrum. • Each line in the emission spectra corresponds to one exact frequency of light emitted by the atom.

  33. Explanation of Atomic Emission Spectrum • Electrons start off in their own ground state (lowest energy level). • Excitation of the electron raises it from its ground state to an excited state (a higher energy level). • In order for the electron to jump up a level, it had to absorb a quantum of energy. When it falls back down, it releases this same amount of energy (as a photon) and we see the light.

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