1 / 21

Periodic trends

Periodic trends. Chapter 7. Periodic Properties of elements. All atomic properties depend on the energies of the outermost orbitals. The relative energies of these orbitals depend on their size and effective nuclear charge . How much of the nucleus’ charge a valence electron “feels”.

iden
Download Presentation

Periodic trends

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. Periodic trends Chapter 7

  2. Periodic Properties of elements • All atomic properties depend on the energies of the outermost orbitals. • The relative energies of these orbitals depend on their size and effective nuclear charge. • How much of the nucleus’ charge a valence electron “feels”.

  3. orbital energies

  4. across a period Why?

  5. effective nuclear charge • Z* = Z – S • Z* is the effective nuclear charge • Z is the charge of the nucleus • S is the potential energy change caused by electron shielding.

  6. General trend • Highest occupied orbital smaller, lower E larger, higher E

  7. atomic size • ½ of the inter-nuclear distance of a diatomic molecule.

  8. Atomic size trends

  9. Ionization energy • Atom(g) + energy  Atom+(g) + e- • The energy required to remove the least tightly held electron is ionization energy.

  10. Ionization energy trend

  11. overall trend

  12. why the break? Look at what happens to the energy of the outermost electron when you fill a subshell. What about pairing electrons in subshells?

  13. multiple ionizations • Removing the second electron is harder than the first • Effective nuclear charge increases, making the electrons more attracted to the nucleus. • You can tell when an ion is isoelectronic with a noble gas • Look for the huge spike in Ionization Energy

  14. Photoelectron spectroscopy • Fires X-rays to dislodge electrons. • The difference between the energy of the x-ray photon and the ejected electron is the ionization energy. • What can we determine from this graph?

  15. PES Question • Why is the “x” peak for carbon to the left of the peak for boron? • Why is the “z” peak for boron one half the height of the peak for carbon? • Please label each peak with their corresponding orbitals.

  16. electron affinity • Atom(g) + e- Atom-(g) + energy

  17. Why is ea zero sometimes? Look at where the added electron will go. Is there a benefit to put it there?

  18. ion formation • Why do ions form the charges they do?

  19. ions and orbital energy • Metal form cations because you are taking electrons from high energy orbitals, lowering the overall energy. • Nonmetals form anions because adding an electron to a low energy orbital lowers the energy more. • Noble gases can’t lose because their valence electrons are too low energy. They can’t add because you would add to a high energy orbital.

  20. ions of transition metals • Transition metals lose all of their s electrons first, then they will lose d electrons to satisfy one of the following: • Completely filled d subshell • half filled d subshell • completely empty d subshell • What will Ti form? Fe?

  21. ion size • anions = big • cations = small • If species are isoelectronic, more positive charge is smallest.

More Related