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Chapter 13 Electrons in Atoms

Chapter 13 Electrons in Atoms. Section 13.1 Models of the Atom. OBJECTIVES: Summarize the development of atomic theory. Section 13.1 Models of the Atom. OBJECTIVES: Explain the significance of quantized energies of electrons as they relate to the quantum mechanical model of the atom.

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Chapter 13 Electrons in Atoms

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  1. Chapter 13Electrons in Atoms

  2. Section 13.1Models of the Atom • OBJECTIVES: • Summarize the development of atomic theory.

  3. Section 13.1Models of the Atom • OBJECTIVES: • Explain the significance of quantized energies of electrons as they relate to the quantum mechanical model of the atom.

  4. Greek Idea • Democritus • Matter is made up of solid indivisible particles • John Dalton - one type of atom for each element

  5. J. J. Thomson’s Model • Discovered electrons • Atoms were made of positive stuff • Negative electron floating around • “Plum-Pudding” model

  6. Ernest Rutherford’s Model • Discovered dense positive piece at the center of the atom- nucleus • Electrons would surround it • Mostly empty space • “Nuclear model”

  7. Niels Bohr’s Model • He had a question: Why don’t the electrons fall into the nucleus? • Move like planets around the sun. • In circular orbits at different levels. • Amounts of energy separate one level from another. • “Planetary model”

  8. Bohr’s planetary model • Energy level of an electron • analogous to the rungs of a ladder • electron cannot exist between energy levels, just like you can’t stand between rungs on ladder • Quantum of energy required to move to the next highest level

  9. The Quantum Mechanical Model • Energy is quantized. It comes in chunks. • A quanta is the amount of energy needed to move from one energy level to another. • Since the energy of an atom is never “in between” there must be a quantum leap in energy. • Erwin Schrodinger derived an equation that described the energy and position of the electrons in an atom

  10. The Quantum Mechanical Model • Things that are very small behave differently from things big enough to see. • The quantum mechanical model is a mathematical solution • It is not like anything you can see.

  11. The Quantum Mechanical Model • Has energy levels for electrons. • Orbits are not circular. • It can only tell us the probability of finding an electron a certain distance from the nucleus.

  12. The Quantum Mechanical Model • The atom is found inside a blurry “electron cloud” • A area where there is a chance of finding an electron. • Draw a line at 90 % • Think of fan blades

  13. Atomic Orbitals • Principal Quantum Number (n) = the energy level of the electron. • Within each energy level, the complex math of Schrodinger’s equation describes several shapes. • These are called atomic orbitals - regions where there is a high probability of finding an electron. • Sublevels- like theater seats arranged in sections

  14. Two representations of the hydrogen 1s, 2s, and 3s orbitals.

  15. Representation of the 2p orbitals. (a) The electron probability distributed for a 2p orbital. (b) The boundary surface representations of all three 2p orbitals.

  16. Representation of the 3d orbitals.

  17. Representation of the 4f orbitals in terms of their boundary surface.

  18. Summary # of shapes Max electrons Starts at energy level s 1 2 1 p 3 6 2 5 10 3 d 7 14 4 f

  19. First Energy Level only s orbital only 2 electrons 1s2 Second Energy Level s and p orbitals are available 2 in s, 6 in p 2s22p6 8 total electrons By Energy Level

  20. Third energy level s, p, and d orbitals 2 in s, 6 in p, and 10 in d 3s23p63d10 18 total electrons Fourth energy level s,p,d, and f orbitals 2 in s, 6 in p, 10 in d, ahd 14 in f 4s24p64d104f14 32 total electrons By Energy Level

  21. Any more than the fourth and not all the orbitals will fill up. You simply run out of electrons The orbitals do not fill up in a neat order. The energy levels overlap Lowest energy fill first. By Energy Level

  22. Section 13.2Electron Arrangement in Atoms • OBJECTIVES: • Apply the aufbau principle, the Pauli exclusion principle, and Hund’s rule in writing the electron configurations of elements.

  23. Section 13.2Electron Arrangement in Atoms • OBJECTIVES: • Explain why the electron configurations for some elements differ from those assigned using the aufbau principle.

  24. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s Aufbau diagram - page 367

  25. Electron Configurations • The way electrons are arranged in atoms. • Aufbau principle- electrons enter the lowest energy first. • This causes difficulties because of the overlap of orbitals of different energies. • Pauli Exclusion Principle- at most 2 electrons per orbital - different spins

  26. Electron Configuration • Hund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until they have to. • Let’s determine the electron configuration for Phosphorus • Need to account for 15 electrons

  27. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The first two electrons go into the 1s orbital • Notice the opposite spins • only 13 more to go...

  28. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next electrons go into the 2s orbital • only 11 more...

  29. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next electrons go into the 2p orbital • only 5 more...

  30. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The next electrons go into the 3s orbital • only 3 more...

  31. 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p Increasing energy 3s 2p 2s 1s • The last three electrons go into the 3p orbitals. • They each go into separate shapes • 3 unpaired electrons • = 1s22s22p63s23p3

  32. Exceptional Electron Configurations

  33. Orbitals fill in order • Lowest energy to higher energy. • Adding electrons can change the energy of the orbital. • Half filled orbitals have a lower energy. • Makes them more stable. • Changes the filling order

  34. Write these electron configurations • Titanium - 22 electrons • 1s22s22p63s23p64s23d2 • Vanadium - 23 electrons • 1s22s22p63s23p64s23d3 • Chromium - 24 electrons • 1s22s22p63s23p64s23d4 expected • But this is wrong!!

  35. Chromium is actually: • 1s22s22p63s23p64s13d5 • Why? • This gives us two half filled orbitals. • Slightly lower in energy. • The same principal applies to copper.

  36. Copper’s electron configuration • Copper has 29 electrons so we expect: 1s22s22p63s23p64s23d9 • But the actual configuration is: • 1s22s22p63s23p64s13d10 • This gives one filled orbital and one half filled orbital. • Remember these exceptions: d4, d9

  37. Section 13.3Chemical Physics and the Quantum Mechanical Model • OBJECTIVES: • Calculate the wavelength, frequency, or energy of light, given two of these values.

  38. Section 13.3Chemical Physics and the Quantum Mechanical Model • OBJECTIVES: • Explain the origin of the atomic emission spectrum of an element.

  39. Light • The study of light led to the development of the quantum mechanical model. • Light is a kind of electromagnetic radiation. • Electromagnetic radiation includes many kinds of waves • All move at 3.00 x 108 m/s = c

  40. Crest Wavelength Amplitude Trough Parts of a wave Origin

  41. Parts of Wave - p.372 • Origin - the base line of the energy. • Crest - high point on a wave • Trough - Low point on a wave • Amplitude - distance from origin to trough (-) or crest (+) • Wavelength - distance from crest to crest • Wavelength is abbreviated by the Greek letter lambda = l

  42. Frequency • The number of waves that pass a given point per second. • Units: cycles/sec or hertz (Hz or sec-1) • Abbreviated by Greek letter nu = n c = ln

  43. Frequency and wavelength • Are inversely related • As one goes up the other goes down. • Different frequencies of light are different colors of light. • There are a wide variety of frequencies • The whole range is called a spectrum, Fig. 13.10, page 373

  44. High energy Low energy Low Frequency High Frequency X-Rays Radiowaves Microwaves Ultra-violet GammaRays Infrared . Long Wavelength Short Wavelength Visible Light

  45. Prism • White light is made up of all the colors of the visible spectrum. • Passing it through a prism separates it.

  46. If the light is not white • By heating a gas with electricity we can get it to give off colors. • Passing this light through a prism does something different.

  47. Atomic Spectrum • Each element gives off its own characteristic colors. • Can be used to identify the atom. • How we know what stars are made of.

  48. These are called discontinuous spectra, or line spectra • unique to each element. • These are emission spectra • The light is emitted given off • Sample 13-2 p.375

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