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Chapter 4 Arrangement of Electrons in Atoms

Chapter 4 Arrangement of Electrons in Atoms. 4.1 The Development of a New Atomic Model. A. Properties of Light. Electromagnetic Radiation: EM radiation are forms of energy which move through space as waves There are many different types of EM waves visible light – colors of the spectrum

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Chapter 4 Arrangement of Electrons in Atoms

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  1. Chapter 4 Arrangement of Electrons in Atoms 4.1 The Development of a New Atomic Model

  2. A. Properties of Light • Electromagnetic Radiation: EM radiation are forms of energy which move through space as waves • There are many different types of EM waves • visible light – colors of the spectrum • x-rays- used by doctors • ultraviolet light -insects use to find nectar • infrared light -used in heat lamps • radio waves - disturb electromagnetic fields • gamma -short wavelengths and high frequency

  3. B. EM Waves • Move at speed of light: 3.00 x 108 m/s • Speed is equal to the frequency times the wavelength c = v • Frequency (v) is the number of waves passing a given point in one second • Wavelength () is the distance between peaks of adjacent waves • Speed of light is a constant, so vis also a constant; vandmust be inversely proportional

  4. Quantum Effect energy absorption spectrum p+ e no ground state e-

  5. Photoelectric Effect energy photon emission spectrum! p+ e no When the “excited” electron returns to lower energy levels, it releases energy in the form of light travels at the speed of light (3.00 x 108 m/s)

  6. A photon is a particle of energy having a rest mass of zero and carrying a quantum of energy • A quantum is the minimum amount of energy that can be lost or gained by an atom

  7. When a specific or quantized amount of energy is exposed to the atom, the electron jumps from its “ground” or original state to an “excited” state • When the “excited” electron returns to lower energy levels, it releases energy in the form of light.

  8. Electromagnetic Spectrum • Wavelength increases→ • Frequency decreases→ • Energy decreases→

  9. ROY G BIV • Red: The longest wavelength Least amount of energy Violet: Shortest wavelength Greatest amount of energy

  10. Emissions Spectrum • Bright line spectrum: Light is given off by excited atoms as they return to lower energy states • Light is given off in very definite wavelengths • It is like a “finger print” to the element • A spectroscope reveals lines of particular colors- light passed through a prism; specific frequencies given off.

  11. Bright Line Emission Spectrum of Elements

  12. Comparison of Spectrums

  13. 4.2 The Quantum Model of The Atom

  14. Bohr Model Niels Bohr e- Electrons circle around the nucleus on their energy level 1st ring = 2e- 2nd ring = 8e- 3rd ring= 18e- 4th ring = 32e- 5th ring = 32e- p+ no Energy levels

  15. The Bohr Model of the Atom • Electron Orbits, or Energy Levels • Electrons can circle the nucleus only in allowed paths or orbits • The energy of the electron is greater when it is in orbits farther from the nucleus • The atom achieves the ground state when atoms occupy the closest possible positions around the nucleus • Electromagnetic radiation is emitted when electrons move closer to the nucleus.

  16. The Bohr Atomic Model

  17. 4.3 Electron Configurations

  18. Energy Levels

  19. A. Writing Electron Configurations Aufbau Principle An electron occupies the lowest-energy orbital that can receive it. Each p and d sublevel must have 1e before you double them up. Pauli Exclusion Principle The maximum number of electrons per orbital is 2 Rules: Hund's Rule The first electron must face up and the second electron must face down.

  20. B. Principal Energy Level (n) The n value gives you two things: • The principal energy level (PEL) • The number of sublevels. n=1 n= 2 n=3 n=4 1s 2sp 3spd 4spdf 2s2p 3s3p3d 2s4p4d4f

  21. Orbital Filling Order

  22. C. Electrons in each subshell • The maximum number of electrons in each subshell is as follows: • s = 2e- • p = 6e- • d =10e- • f =14e-

  23. D. Practice with Electron Configuration LithiumCarbon 1s2s1 1s22s22p2 MagnesiumIron 1s22s22p63s2 1s22s22p63s2 Silver 1s22s22p63s23p64s23d104p65s24d9

  24. E. Shortcut Notation • The configuration begins with the preceding noble gas’s symbol in brackets and is followed by the rest of the configuration for the particular element. • [Ne] 3s23p5

  25. F. Terms • Valence shell: Electrons that are in the outermost energy level Inner shell electrons: Electrons that are not in the highest energy level • Octet Rule: The valence shell has no more than 8 electrons. ( 8 is great!). It gives the electron stability.

  26. G. Regents way for Electron Configuration • States the total number of electrons in the entire principal energy level. • Copy the electron configuration directly from the periodic tablein the packet. Calcium Zinc Lead Ca Zn Pb 2-8-8-2 2-8-18-2 2-8-18-32-18-4

  27. H. Magnetic Fields & Electrons • Paramagnetic: When an atom has unpaired electrons, it will be attracted into a magnetic field 1s22s22p4 • Dimagnetic: When an atom has only paired electrons, it will be slightly repelled by a magnetic field 1s22s22p63s2

  28. I. Shapes of Orbitals • S orbitals are Spherical in shape

  29. P Orbitals • P orbitals are peanut in shape • They exist in all three dimensional planes

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