1 / 36

Atomic Structure and the Periodic Table

Atomic Structure and the Periodic Table. The electronic structure of an atom determines its characteristics. Studying atoms by analyzing light emissions/ absorbtions. Spectroscopy: analysis of light emitted or absorbed from a sample Instrument used = spectrometer

reilly
Download Presentation

Atomic Structure and the Periodic Table

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript


  1. Atomic Structure and the Periodic Table

  2. The electronic structure of an atom determines its characteristics

  3. Studying atoms by analyzing light emissions/absorbtions • Spectroscopy: analysis of light emitted or absorbed from a sample • Instrument used = spectrometer • Light passes through a slit to become a narrow beam • Beam is separated into different colors using a prism (or other device) • Individual colors are recorded as spectral lines

  4. Electromagnetic radiation • Light energy • A wave of electric and magnetic fields • Speed = 3.0 x 108 m/s • Wavelength () = distance between adjacent peaks • Unit = any length unit • Frequency () = number of cycles per second • Unit = hertz (Hz)

  5. Relationship between properties of EM waves • Wavelength x frequency = speed of light ·v = c Calculate the frequency of light that has a wavelength of 6.0 x 107 m. Calculate the wavelength of light that has a frequency of 3.7 x 1014 s-1

  6. Visible Light • Wavelengths from 700 nm (red) to 400 nm (violet) • No other wavelengths are visible to humans

  7. Quanta and Photons • Quanta: discrete amounts • Energy is quantized – restricted to discrete values • Only quantum mechanics can explain electron behavior • Analogy: Water flow

  8. Another analogy for quanta • A person walking up steps – his potential energy increases in a quantized manner

  9. Photons • Packets of electromagnetic energy • Travel in waves • Brighter light = more photons passing a point per second • Higher energy photons have a higher frequency of radiation • Planck constant h = 6.63 x 10-34Js E = hv The energy of a photon is directly proportional to its frequency

  10. Deriving Planck’s constant In a laboratory, the energy of a photon of blue light with a frequency of 6.4 x 1014 Hz was measured to have an energy of 4.2 x 1019 J. Use Planck’s constant to show this: E = (6.63 x 10-34 J·s) x (6.4 x 10141/s) = 4.2 x 1019 J

  11. Evidence for photons • Photoelectric effect – the ejection of electrons from a metal when exposed to EM radiation • Each substance has its own “threshold” frequency of light needed to eject electrons

  12. Determining the energy of a photon E = hv • Use Planck’s constant! • What is the energy of a photon of radiation with a frequency of 5.2 x 1014 waves per second?

  13. Another problem involving photon energy • What is the energy of a photon of radiation with a wavelength of 486 nm?

  14. Louis de Broglie – proposed that matter and radiation have properties of both waves and particles (Nobel Prize 1929) • Calculate the wavelength of a hydrogen atom moving at 7.00 x 102 cm/sec • = h • m m = mass h = Planck’s constant  = velocity

  15. Hydrogen spectral lines Balmer series: n1 = 2 and n2 = 3, 4, … Lyman series (UV lines): n1= 1 and n2 = 2, 3, …

  16. Atomic Spectra and Energy Levels • Observe the hydrogen gas tube, use the prism to see the frequencies of EM radiation emitted • Johann Balmer– noticed that the lines in the visible region of hydrogen’s spectrum fit this expression: v= (3.29 x 1015 Hz) x 1 - 1 4 n2 n = 3, 4, …

  17. Rydberg equation: works for all lines in hydrogen’s spectrum v= RH x 1 - 1 n12 n22 RH = 3.29 x 1015 s-1 Rydberg Constant

  18. Energy associated with electrons in each principal energy level • Energy of an electron in a hydrogen atom -2.178 x 10-18 joule E = n2 n= principal quantum number

  19. Differences in Energy Levels of the hydrogen atom Use the Rydberg Equation OR Use the expression for each energy level’s energy in the following equation: E = Efinal – Einitial

  20. Niels Bohr’s contribution Assumed e- move in circular orbits about the nucleus Only certain orbits of definite energies are permitted An electron in a specific orbit has a specific energy that keeps it from spiraling into the nucleus Energy is emitted or absorbed ONLY as the electron changes from one energy level to another – this energy is emitted or absorbed as a photon

  21. Summary of spectral lines When an e- makes a transition from one energy level to another, the difference in energy is carried away by a photon Different excited hydrogen atoms undergo different energy transitions and contribute to different spectral lines

  22. The Uncertainty Principle – Werner Heisenberg • The dual nature of matter limits how precisely we can simultaneously measure location and momentum of small particles • It is IMPOSSIBLE to know both the location and momentum at the same time

  23. Atomic Orbitals – more than just principal energy levels • Erwin Schrodinger (Austrian) • Calculated the shape of the wave associated with any particle • Schrodinger equation – found mathematical expressions for the shapes of the waves, called wavefunctions(psi) 

  24. Born’s contribution • Max Born (German) • The probability of finding the electron in space is proportional to 2 Called the “probability density” or “electron density”

  25. Atomic Orbital – the wavefunction for an electron in an atom • s – high probability of e- being near or at nucleus ELECTRON IS NEVER AT THE NUCLEUS IN THE FOLLOWING ORBITALS: • p – 2 lobes separated by a nodal plane • d – clover shaped • f – flower shaped

  26. More about orbitals • Each orbital can hold 2 electrons • Orbitals in the same subshell have equal energies

  27. Quantum numbers – like an “address” for an electron n = principal quantum number As n increases * orbitals become larger • electron is • farther from nucleus more often • higher in energy • less tightly bound to nucleus

  28. Quantum numbers • l = angular momentum quantum number • Values: 0 to n – 1 • Defines the shape of the orbital

  29. Quantum numbers Example: for d orbitals, m can be -2, -1, 0, 1, or 2 For p orbitals, m can be -1, 0, or 1 • ml=the magnetic quantum number • Orientation of orbital in space (i.e. pxpy or pz) • Values: between – l and l, including 0

  30. Quantum numbers • ms = the spin number • When looking at line spectra, scientists noticed that each line was really a closely-spaced pair of lines! • Why? Each electron has a SPIN – it behaves as if it were a tiny sphere spinning upon its own axis • Spin can be + ½ or -1/2 • Each represents the direction of the magnetic field the electron creates

  31. Describe the electron that has the following quantum numbers: Principal level 4 4p orbital px orbital spin up n = 4, l = 1, ml = -1, ms = +1/2

  32. Are these sets of quantum numbers valid? • 3, 2, 0, -1/2 • 2, 2, 0, 1/2 NO! Level 2 2d orbital – does not exist! YES! Level 3 3d orbital 3dxz Spin down

  33. Electron configuration: rules • Aufbau principle – electrons fill lowest energy levels first • Pauli exclusion principle – only 2 electrons may occupy each orbital, must have opposite spins • Hund’s rule – the lowest energy is attained when the number of electrons with the same spin is maximized (because electrons repel each other)

  34. Energy level specifics 4s • s and d orbitals are close in energy • Example • 4s electrons have slightly lower energy than 3d electrons • The s electrons can penetrate to get closer to the nucleus, giving them slightly lower energy 3d

  35. Noble Gas Configuration • A shorter electron configuration • Write the symbol for the noble gas BEFORE the element in brackets • Write the remainder of the configuration • Examples: • Cl • Cs

  36. Special rules • One electron can move from an s orbital to the d orbital that is closest in energy • Only happens to create half or whole-filled d orbitals • Examples: Cr, Cu

More Related