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IONS

Section 7.1. IONS. Understanding the Key. Valence electrons hold the key to ion formation.

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IONS

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  1. Section 7.1 IONS

  2. Understanding the Key • Valence electrons hold the key to ion formation. • Gilbert Lewis formulated the Octet Rule: elements are more stable if they find a way to reach 8 electrons (by filling up to get 8 or emptying out the valence shell to get an inner shell to become the new outer shell that is full). • Helium is the exception to the Octet Rule.

  3. Showing Valence Electrons • Lewis Gilbert also created a notation for showing only the valence electrons of an atom: electron dot structures (a.k.a. Lewis dot diagram).

  4. Forming an Ion • The fewer the valence electrons, the more likely the atom is to give up electrons (empty out the valence shell), becoming positive ions known as cations.

  5. Forming an Ion • The closer to 8 valence electrons an atom is, the more likely the atom is to gain more electrons to try to reach a total of 8, becoming negative ions known as anions.

  6. Forming an Ion • The representative elements behave in a predictable manner when determining the type of ion they will form. • The group # (1A-8A) tells how many valence electrons each has (except for helium).

  7. Forming an Ion

  8. Naming Cations • For metallic elements, naming cations that are formed is easy. The name of the ion is the same as the name of the atom, just followed by the word “ion” or “cation.”

  9. Transition Metal Cations • Transition metals are not as easy to predict. Some of them will readily form more than one type of ion; these must be memorized. • Others to know are silver (Ag+), gold (Au+), zinc (Zn2+) and cadmium (Cd2+).

  10. Naming Transition Metal Cations • Notice that these cation names are followed by Roman numerals, providing a way to distinguish them from each other. (Make flash cards!)

  11. Forming Anions • Atoms of nonmetals and metalloids form anions by gaining enough electrons to fill their valence shell. • This oxygen atom has room for 2 more valence electrons to fill the 2nd energy level with 8.

  12. Naming Anions • How do we name anions? • The atom name endings (suffix) is changed to –ide, and it is followed by either “ion” or “anion.” • Fluorine atom becomes a fluoride ion, sulfur atom becomes sulfide ion, and phosphorus atom becomes phosphide ion.

  13. Getting Some Practice For the following elements, draw the electron dot structure. Then, predict the ion it would form. Finish by naming the ion properly. • Strontium • Oxygen • Cesium

  14. Getting Some Practice • Boron • Chlorine • Arsenic • Krypton

  15. Breaking the rules • Why does copper form more than one type of ion? Consider the following: This is a pseudo noble-gas configuration since the preceding noble gas, Argon, doesn’t have 3d-level electrons.

  16. Breaking the rules • If the 2 electrons remained in the 4s sublevel, then there are 2 valence electrons to lose: Many transition metals have multiple ions they can form, and their reasons may vary. This example demonstrates that it is possible since 3d and 4s are so close in energy. Copper atom Cu Copper(II) ion Cu2+

  17. Mono- versus Poly- • All of these ions we’ve discussed are called monatomic ions – ions formed when one atom forms an ion. • Ions also exist that are polyatomic ions – ions formed when a group of atoms chemically bonded together has more or less electrons than it should. • Most polyatomic ions are anions, although two you may encounter are cations: the ammonium ion (NH4+) and hydronium ion (H3O+).

  18. Common Polyatomic Ions • You should make flashcards of these polyatomic ions to learn their symbols, charges and names.

  19. Common Polyatomic Ions • You should make flashcards of these polyatomic ions to learn their symbols, charges and names. We will be having a couple of ion quizzes, similar to the element quizzes, during this chapter.

  20. References • http://www.calzim.com/online/online2_1/class_material/unit1/ion.gif • http://www.roymech.co.uk/images14/lewis_elements.gif • http://www.middleschoolchemistry.com/img/content/multimedia/chapter_4/lesson_6/lewis_dot_table.jpg • http://www.kanescience.com/_images/chem_ionic/sodium_ions.png • http://www.calzim.com/online/online2_1/class_material/unit1/unit1.htm • http://www.personal.kent.edu/~cearley/ChemWrld/compounds/ionictable.gif • http://www.chem.uwec.edu/Chem103_F08_F0F/pages/resources/media/transitions_metal_ions_Silberberg_table_2.4.jpg • http://wps.prenhall.com/wps/media/objects/1053/1078773/tools/HPMP.table.2.4.gif

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