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(Electron Configurations).  Electromagnetic Radiation-form of energy that exhibits wave-like behavior as it travels through space.  Electromagnetic.

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Presentation on theme: "(Electron Configurations).  Electromagnetic Radiation-form of energy that exhibits wave-like behavior as it travels through space.  Electromagnetic."— Presentation transcript:

1 (Electron Configurations)

2  Electromagnetic Radiation-form of energy that exhibits wave-like behavior as it travels through space.  Electromagnetic Spectrum-ordered arrangement by wavelength or frequency for all forms of electromagnetic radiation.

3  Wavelength-lambda (λ) The distance between corresponding points on adjacent waves. Units: m, nm, cm, or Å  Frequency-nu (ν) The number of waves passing a given point in a definite amount of time. Units: hertz (Hz) or cycles/sec = 1/sec = sec -1

4  When an electric field changes, so does the magnetic field. The changing magnetic field causes the electric field to change. When one field vibrates—so does the other.  RESULT-An electromagnetic wave.

5 Waves or Particles  Electromagnetic radiation has properties of waves but also can be thought of as a stream of particles.  Example: Light  Light as a wave: Light behaves as a transverse wave which we can filter using polarized lenses.  Light as particles (photons)  When directed at a substance light can knock electrons off of a substance (Photoelectric effect)

6  c = λ∙ν  λ = wavelength (m)  ν = frequency (Hz)  c = speed of light= 3.0 x 10 8 m/sec (constant)  λ and ν are _______________ related.

7 Truck-mounted helium-neon laser produces red light whose wavelength (λ ) is 633 nanometers. Determine the frequency (v). *Remember that c=3.0x10 8 m/s. *Use the formula v= c λ

8 c =3.0x10 8 m/s c= λ. v v=c / λ λ = 633nm= 6.33x10 -7 m v = 3.0x10 8 m/s = 0.47x s -1 = 4.7x10 14 s x10 -7 m Frequency = 4.7x10 14 Hz (cycles per second)

9 GIVEN: = ? = 434 nm = 4.34  m c = 3.00  10 8 m/s WORK : = c = 3.00  10 8 m/s 4.34  m = 6.91  Hz  EX: Find the frequency of a photon with a wavelength of 434 nm.

10  2 problems that could not be explained if light only acted as a wave.  1.) Emission of Light by Hot bodies: Characteristic color given off as bodies are heated: red  yellow  white If light were a wave, energy would be given off continually in the infrared (IR) region of the spectrum.

11  2.) Absorption of Light by Matter = Photoelectric Effect Light can only cause electrons to be ejected from a metallic surface if that light is at least a minimum threshold frequency. The intensity is not important. If light were only a wave intensity would be the determining factor, not the frequency!

12  When an object loses energy, it doesn’t happen continuously but in small packages called “quanta”. “Quantum”-a definite amount of energy either lost or gained by an atom. “Photon”-a quantum of light or a particle of radiation.

13  Calculate the frequency for the yellow- orange light of sodium.  Calculate the frequency for violet light.

14  Calculate the frequency for the yellow- orange light of sodium.  Calculate the frequency for violet light.

15  E = h∙ν  E = energy (joule)  h = Planck’s constant = 6.63 x j∙sec  ν = frequency (Hz)  E and ν are ______________ related.  Calculate the energy for the yellow-orange light for sodium.  Calculate the energy for the violet light.

16  Excited State: Higher energy state than the atom normally exists in.  Ground State: Lowest energy state “happy state”  Line Spectrum: Discrete wavelengths of light emitted.  2 Types:  1.) Emission Spectrum: All wavelengths of light emitted by an atom.  2.) Absorption Spectrum: All wavelengths of light that are not absorbed by an atom. This is a continuous spectrum with wavelengths removed that are absorbed by the atom. These are shown as black lines for absorbed light.  Continuous Spectrum: All wavelengths of a region of the spectrum are represented (i.e. visible light)

17  Hydrogen’s spectrum can be explained with the wave-particle theory of light.  Niel’s Bohr (1913)  1.) The electron travels in orbits (energy levels) around the nucleus.  2.) The orbits closest to the nucleus are lowest in energy, those further out are higher in energy.  3.) When energy is absorbed by the atom, the electron moves into a higher energy orbit. This energy is released when the electron falls back to a lower energy orbit. A photon of light is emitted.

18  Lyman Series-electrons falling to the 1 st orbit, these are highest energy, _____ region.  Balmer Series- electrons falling to the 2 nd orbit, intermediate energy, _______ region.  Paschen Series-electrons falling to the 3 rd orbit, smallest energy, ______ region.

19  E n = (-R H ) 1/n 2  E n = energy of an electron in an allowed orbit (n=1, n=2, n=3, etc.)  n = principal quantum number (1-7)  R H = Rydberg constant (2.18 x J)  When an electron jumps between energy levels: ΔE =E f – E i  By substitution: ΔE = hν = R H (1/n i 2 - 1/n f 2 )  When n f > n i then ΔE = (+)  When n f < n i then ΔE = (-)

20  DeBroglie (1924)-Wave properties of the electron was observed from the diffraction pattern created by a stream of electrons.  Schrodinger (1926)-Developed an equation that correctly accounts for the wave property of the electron and all spectra of atoms. (very complex)

21  Rather than orbits  we refer to orbitals. These are 3-dimensional regions of space where there is a high probability of locating the electron.  Heisenberg Uncertainty Principle-it is not possible to know the exact location and momentum (speed) of an electron at the same time.  Quantum Numbers-4 numbers that are used to identify the highest probability location for the electron.

22  1.) Principal Quantum Number (n)  States the main energy level of the electron and also identifies the number of sublevels that are possible.  n=1, n=2, n=3, etc. to n=7  2.) Orbital Quantum Number  Identifies the shape of the orbital s (2 electrons) sphere1 orbital P (6 electrons) dumbbell3 orbitals d ( 10 electrons)4-4 leaf clovers & 1-dumbbell w/doughnut 5 orbitals f (14 electrons) very complex7 orbitals

23  3.) Magnetic Quantum Number  Identifies the orientation in space (x, y, z) s  1 orientation p  3 orientations d  5 orientations f  7 orientations 4.) Spin Quantum Number States the spin of the electron. Each orbital can hold at most 2 electrons with opposite spin.

24  1.) Principal Quantum Number (n)  States the main energy level of the electron and also identifies the number of sublevels that are possible.  n=1, n=2, n=3, etc. to n=7  2.) Azimuthal Quantum Number (l)  Values from 0 to n-1  Identifies the shape of the orbital l = 0ssphere1 orbital l = 1pdumbbell3 orbitals l = 2d 4-4 leaf clovers & 1-dumbbell w/doughnut 5 orbitals l = 3fvery complex7 orbitals

25  3.) Magnetic Quantum Number (m l )  Values from –l  l  States the orientation in space (x, y, z) m l = 0sonly 1 orientation m l = -1, 0, +1p3 orientations m l = -2,-1,0,+1,+2d5 orientations m l = -3,-2,-1,0,+1+2,+3f7 orientations 4.) Spin Quantum Number (m s ) Values of +1/2 to -1/2 States the spin of the electron. Each orbital can hold at most 2 electrons with opposite spin.


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