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Atomic Structure

Atomic Structure. Emphasis on electron arrangement. Recall:. Rutherford’s gold foil experiment showed that atoms have very dense, small positive nuclei and a relatively large empty space where electrons exist. But his experiment did not shed light on the arrangement of electrons.

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Atomic Structure

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  1. Atomic Structure Emphasis on electron arrangement

  2. Recall: • Rutherford’s gold foil experiment showed that atoms have very dense, small positive nuclei and a relatively large empty space where electrons exist. • But his experiment did not shed light on the arrangement of electrons

  3. Atomic Spectra • When elements are given energy such as heat or electricity they emit specific colours of light which can be divided using a spectroscope, called a line or emission spectrum:

  4. Neils Bohr • Interpreted these emission spectrum as having something to do with electron arrangement. • Before we get into specifics we have to go over some background fundamentals on light and waves.

  5. Waves • Waves transfer energy in a repeating pattern. • The distance from one point on a wave to the next identical point is the wavelength given the symbol ‘λ’ (called lambda). • The number of cycles that occur at one position over a time of one second is called the frequency, symbol ‘f’. • The speed of a wave is dependent on the medium it travels through, symbol ‘v’. • The relationship between these three variables is called the universal wave equation: v = fλ

  6. Waves can Diffract • The diffraction phenomenon is described as the apparent bending of waves around small obstacles and the spreading out of waves past small openings.

  7. Light • Light also is found to diffract and thus proving it has wave properties. • Light is made up of electromagnetic waves which are perpendicularly oscillating electric and magnetic fields.

  8. Light... • Notice that light has a wavelength, and frequency, and this is what determines the type of light observed:

  9. Speed of light • Light travels at an amazing 300 000 km in one second through space and roughly that for air. • It is given the symbol ‘c’ and it is equal to 3.0 x 108 m/s. • So instead of v= fλ, we can write: c = fλ

  10. Max Planck • Discovered that light energy was quantized as photons having energy related to the frequency. • E = hf • h is called planck’s constant = 6.63 x 10-34 J.s • f is the frequency of light • Since f=c/λ, we can sub this into planck’s equation: • E = hc/λ

  11. Putting it together • Going back to the emission spectrum, Bohr realized that • electrons can exist in specific energy levels and that they can absorb energy to become “excited” • These excited electrons emit energy to return to their natural lower energy level called their “ground state”. • The colour of light in the line spectrum has a specific energy which corresponds to the difference in energy between the excited energy level and the ground state energy level.

  12. Electron Energy • The energy of an electron can be calculated using a variation of the Rydberg equation: • E = 2.18 x 10-18 (1/n2) where ‘n’ is the energy level (only can be used for Hydrogen) • And the change in energy when an electron falls from an excited state back to a lower state: ΔE = 2.18 x 10-18 (1/nf2 – 1/ni2) • And this amount of energy corresponds directly to the energy of the photon released by the electron: where E = hc/λ

  13. Try it! • Calculate the expected wavelength of light if an electron of Hydrogen, falls from the 5th to the 2nd energy level. • ΔE = 2.18 x 10-18 (1/nf2 – 1/ni2) • ΔE = 2.18 x 10-18 (1/22 – 1/52) • ΔE = 2.18 x 10-18 (0.25-0.04) • ΔE = 4.58 x 10-19 J • 4.58 x 10-19 J = hc/λ • λ = 6.63 x 10-34 x 3 x 108 / 4.58 x 10-19 • λ = 4.34 x 10-7 m = 434 nm Violet!

  14. Bottom Line • Electrons exist in energy levels and can be excited into higher levels by absorbing energy in the form of heat or light. • When these electrons return to their lower energy states they emit this difference in energy as light. • Some of the light is visible and can be observed using a spectroscope although some is ultraviolet or infrared. • Bohr’s model worked well for hydrogen but not for others.

  15. What about other elements • Bohr’s model did not work for elements with more than one electron. • The emission spectrum for sodium or helium was more complex with lines closer together • This suggested that there were smaller changes in energy and perhaps sublevels within a level. • There was also the discovery that electrons behaved like waves as well as particles. • After many years of scientific investigation and debate the problems above led us to the quantum mechanical model of the atom.

  16. Quantum Mechanical Model • Developed by Erwin Shrodinger • Based on the idea that electrons have wave properties (diffract) and they exist in energy levels. • Shrodinger developed a mathematical function that describes the electron as a wave having a particular shape and distance away from the nucleus, called an orbital.

  17. Orbitals • The orbital describes the probability that an electron can be found at a particular point at a particular time, shown as an electron probability density graph. • The following EPDG shows is for Hydrogen having only 1 electron in the lowest possible energy level:

  18. Quantum Numbers • To use Shrodinger’s equation an electron is assigned four quantum numbers representing various aspects of the orbital; n, l, ml, ms • n – the principle quantum number, represents the energy level (n = 1,2,3...∞) of the electron. • These correspond with the period number on the periodic table (except the transition metals) • The total number of electrons in an energy level equals 2n2.

  19. Subshells ‘l’ • Recall that Bohr’s model did not work for elements with more than one electron. • And that emission spectrum from these other elements showed a possibility of sublevels within the energy level. • This is what the second quantum number, ‘l’ represents, the subshell or sublevel, which also describes the shape of the orbital. • l values can range in value from 0 to n-1 • So if an electron is in the 4th energy level it could have l values of 0, 1, 2, and 3

  20. Orbital shapes ‘l=0’ • When l = 0, it is called an ‘s’ orbital and has the shape: • Notice that as the energy level increases the electron density gets further away and more spread out. Spherical

  21. Orbital shapes ‘l = 1’ • When l = 1, it is called a ‘p’ orbital and has the shapes: • There are 3 orientations, described by the third quantum number ml. Dumbell shaped

  22. Orbital shapes ‘l = 2’ • When l = 2, it is called an ‘d’ orbital and has the shapes: • There are 5 different shapes described by the third quantum number, ml. 4 flower shaped and one double baby soother

  23. This will help you remember! • http://www.youtube.com/watch?v=K-jNgq16jEY

  24. Orbital shapes ‘l = 3’ • When l = 3, it is called an ‘f’ orbital and has the shapes: • There are 7 different shapes described by the third quantum number, ml.

  25. Orbital Orientation ‘ml’ • These values can range from +l through to –l. • if l = 0, then ml values can only be 0, hence ‘s’ orbitals are only spherical. • if l = 1, then ml values can be +1, 0, -1, hence there are three different p orbital orientations (pxy, pxz, pyz). • if l = 2, then ml values can be +2,+1,0, -1,-2, hence there are 5 different orientations of d orbitals. • if l = 3, then ml can be +3,+2,+1,0, -1,-2,-3 hence there are 7 different orientations of f orbitals.

  26. Magnetic Spin ‘ms’ • The fourth quantum number represents the magnetic spin of the electron (moving electric charges have circular magnetic fields) • The values can either be + ½ or – ½ • Meaning either spin up or spin down. • Two electrons can occupy one orbital but they must have opposite spins.

  27. SUmmary • The first energy level, n=1, can have a maximum of 2n2 electrons, which is 2(1)2 =2, which would be in the l= 0 or ‘s’ orbital. • The second energy level, n=2, can have a maximum of 2n2 electrons, which is 2(2)2 =8, which would be in the l= 0 and 1, being the s and the p orbitals. Two would go in the s orbital first then the remaining six electrons would go in the three p orbitals.

  28. Summary continued • The third energy level, n=3, can have a maximum of 2n2 electrons, which is 2(3)2=18, which would be in the l= 0,1 and 2, being the s, p and d orbitals. Two would go in the s orbital first then six electrons would go in the three p orbitals, then the remaining ten electrons would go into the five d orbitals. • But wait!, the 4s orbital has a lower energy than the 3d orbitals, this is why the Bohr-Rutherford model we taught you worked well for the first 20 elements!

  29. Summary • See the concept organizer on page 135 of the text book. • Answer # 1-5 on page 136, # 2,3,5,6 on p.138 • Read pages 139 -143

  30. Electron Configuration • Electrons will “occupy” the lowest energy levels possible • page 16 of your manual, p. 142 of text

  31. “Filling” orbitals • Use three principles when determining the electron configuration of an atom. • 1. The Aufbau principle – (to build up in german) fill in the lowest sublevel one at a time before proceeding to the next sublevel. • 2. Hund’s Rule – each orbital is “filled” with one electron first before doubling up. When electrons are added singly to orbitals they would all have the same spin.

  32. 3. Pauli Exclusion principle- no two electrons in an atom can have the same quantum numbers. Which means that if two electrons occupy the same orbital they must have opposite spins, indicated by the arrow direction. • Refer to page 144 of the textbook. The orbital diagram shows the electrons as arrows in boxes, and the electron configuration is shown on the right.

  33. Condensed or Abbreviated electron configuration • It can be very lengthy to write the entire e.c. for say an element like Sn, so we use a condensed version. • Write the noble gas preceding the element, on the periodic table, in square brackets, for Tin this would be [Kr] • Then looking at the periodic write in the remaining electrons to get to tin: • [Kr]5s2 4d10 5p2

  34. Try it • P.145-146 # 6,7, 9 • P. 150 # 10- 13 • Summarize the periodic trends in the periodic table in terms of Zeff and electron configuration for: • atomic radius • Ionization energy • Reactivity • Electron affinity See pages 150 -157

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