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Lecture 20: Periodic Trends

Lecture 20: Periodic Trends. Reading: Zumdahl 12.14-12.15 Outline Periodic Trends Ionization Energy, Electron Affinity, and Radii A Case Example. Periodic Trends. The valence electron structure of atoms can be used to explain various properties of atoms.

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Lecture 20: Periodic Trends

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  1. Lecture 20: Periodic Trends • Reading: Zumdahl 12.14-12.15 • Outline • Periodic Trends • Ionization Energy, Electron Affinity, and Radii • A Case Example

  2. Periodic Trends • The valence electron structure of atoms can be used to explain various properties of atoms. • In general, properties correlate down a group of elements. • A warning: such discussions are by nature very generalized…exceptions do occur.

  3. Periodic Trends: Ionization • If we put in enough energy, we can remove an electron from an atom. • The electron is completely “removed” from the atom (potential energy = 0).

  4. Periodic Trends: Ionization • Generally done using photons, with energy measured in eV (1 eV = 1.6 x 10-19 J). • The greater the propensity for an atom to “hold on” to its electrons, the higher the ionization potential will be. • Koopmans’ Theorem: The ionization energy of an electron is equal to the energy of the orbital from where the electron came.

  5. Periodic Trends: Ionization • One can perform multiple ionizations: first Al(g) Al+(g) + e- I1 = 580 kJ/mol second Al+(g) Al2+(g) + e- I2 = 1815 kJ/mol third Al2+(g) Al3+(g) + e- I3 = 2740 kJ/mol fourth Al3+(g) Al4+(g) + e- I4 = 11,600 kJ/mol

  6. Which reaction represents the ionization of F? A. 1s22s22p5 1s22s22p6 B. 1s22s22p5 1s22s22p43s1 C. 1s22s22p5 1s22s22p4 D. 1s22s22p5 1s22s12p6

  7. Periodic Trends: Ionization • First Ionization Potentials: Column 8 Column 1

  8. Periodic Trends: Ionization • First Ionization Potentials: • Increases as one goes from left to right. • Reason: increased Z+ • Decrease as one goes down a group. • Reason: increased distance from nucleus

  9. Periodic Trends: Ionization • Removal of valence versus core electrons Na(g) Na+(g) + e- I1 = 495 kJ/mol (removing “valence” electron) [Ne]3s1 [Ne] Na+(g) Na2+(g) + e- I2 = 4560 kJ/mol [Ne] 1s22s22p5 (removing “core” electron) • Takes significantly more energy to remove a core electron….tendancy for core configurations to be energetically stable.

  10. Which atom would you expect to have the lowest ionization energy? 1s22s22p3 B. 1s22s22p63s23p5 C. 1s22s22p63s23p64s2 D. 1s22s22p63s23p64s23d104p65s1

  11. Periodic Trends: Electron Affinity • Electron Affinity: the energy change associated with the addition of an electron to a gaseous atom.

  12. Periodic Trends: Electron Affinity • We will stick with our thermodynamic definition, with energy released being a negative quantity. Wow!

  13. Periodic Trends: Electron Affinity • Elements that have high electron affinity: • Group 7 (the halogens) and Group 6 (O and S specifically).

  14. Periodic Trends: Electron Affinity • Some elements will not form ions: N? • Orbital configurations can explain both observations.

  15. Periodic Trends: Electron Affinity • Why is EA so great for the halogens? F(g) + e- F-(g) EA = -327.8 kJ/mol [Ne] 1s22s22p5 1s22s22p6 • Why is EA so poor for nitrogen? N(g) + e- N-(g) EA > 0 (unstable) 1s22s22p3 1s22s22p4 (e- must go into occupied orbital)

  16. Periodic Trends: Electron Affinity • How do these arguments do for O? O(g) + e- O-(g) EA = -140 kJ/mol 1s22s22p4 1s22s22p5 Bigger Z+ overcomes e- repulsion. • What about the second EA for O? O-(g) + e- O2-(g) EA > 0 (unstable) 1s22s22p5 1s22s22p6 [Ne] configuration, but electron repulsion is just too great.

  17. Which diagram indicates the evolution in electron affinity from high affinity to low affinity?

  18. Atomic Radii • Atomic Radii are defined as the covalent radii, and are obtained by taking 1/2 the distance of a bond: r = atomic radius

  19. Atomic Radii • Decrease to right due due increase in Z+ • Increase down column due to population of orbitals of greater n.

  20. Looking Ahead • We can partition the periodic table into general types of elements. Metals: tend to give up e- non-Metals: tend to gain e- Metalloids: can do either

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