2. Chapter 5: Ionic Compounds Review: Ions are atoms that have a net electrical charge. They have a different number of e - than p +. They only way for an atom to form an ion is for the atom to gain or lose e -. The difference in the numbers of e - and p + tells you the amount of charge on the ion.

3. Chapter 5 Learning Objectives 1.Describe how and why atoms form ions. 2.Draw Lewis dot structures for atoms and ions. 3.Describe how ionic compounds form and the nature of ionic bonds. 4.Define, identify and write formula using monoatomic and polyatomic ions. 5.Given the formula, write the name for ionic compounds. 6.Given the name, write the formula for ionic compounds. 7.Describe the general properties of ionic compounds. 8.Explain the concepts of crystal lattice and unit cell. 9.Describe the structure of silicates and hydrates

4. A few term for you to know: (yeah, now we’re learning!!!) Cation – An ion with a positive charge. An atom loses e - to form a cation. (CATIONS ARE “PAWSITIVE”) Anion – An ion with a negative charge. An atom must gain e - to form an anion. Isoelectronic – To have the same electron configuration as another element. Na + is isoelectronic with Ne. Octet – A group of eight. More specifically, 8 valence electrons, completely filling valence s and p orbitals.

5. Ion formation – How atoms form ions. For Main Group elements, the most stable e - configuration (lowest energy) is to have a completely filled valence energy level. This means that the s and p orbitals in the valence level are filled with 8 electrons (an octet). Atoms will gain or Atoms will gain or lose e - in order to achieve an octet. They will become isoelectronic with the nearest noble gas.

6. 1s2s2p3s F 1s2s2p3sMg 1s2s2p3sNe 1s2s2p3sNa1s2s2p3sO

7. O gains 2 e - to fill the 2p orbital and become isoelectronic with Ne……O -2 F gains 1 e - to fill the 2p orbital and become isoelectronic with Ne…….F - Ne already has a filled valence level (2 nd ) and will not form an ion. Na loses 1 e - to empty the 3 rd energy level and become isoelectronic with Ne…….Na + Mg loses 2 e - to empty the 3 rd energy level and become isoelectronic with Ne…….Mg +2

8. Practice: What ions will the following elements be expected to form? CaLi AlBr NBa RbSe

9. In general, an atoms with 3 or less valence e - will lose the e - to form + ions. Atoms with 5 or more valence e - will gain e - to form anions. Non-metals with 4 valence e - will not generally form ions (C, Si)  Metals will always form + ions!!!!!!!!!

10. Prediction of what cation that transition metals form are more difficult. Since they have d and f orbitals that are close in energy to the s and p orbitals, e - can be switched back and forth. They will form ions to increase orbital stability but do not follow the octet rule. Additionally, many can form more than 1 cation depending upon conditions. Example: Fe +2 Fe +3 Ti +2 Ti +4 Ti +2 Ti +4

11. The valence e - can be shown in a Lewis Dot Structure. To draw a Lewis Dot, write the symbol for the element and then 1 dot around the symbol for each valence e -. OFNe NaMg

12. Ionic compounds are ions bonded together. The bond between the ions is called an Ionic Bond (surprisingly enough) The bond is the electrostatic attraction between opposite charges – opposites attract. Ionic compounds always form so that the total charge of the compound is zero. The compound must have the same total number of + and – charges … not necessarily the same number of + and – ions.

13. Monatomic ions are ions formed from a single element that has lost or gained e -. Ex: Na +, Ca +2, S -2, Br - Polyatomic ions are groups of 2 or more bonded atoms that together form an ion. These ions are treated exactly as monatomic ions, making sure to write their entire formula-with subscripts. Ex: NH 4 +, OH -, SO 4 -2, CO 3 -2, PO 4 -3, NO 3 - Structures!

14. Rules for writing ionic formulas 1.Write the symbol for the + ion first. Do not write the charges in the formula. 2.Use subscripts to show the number of each ion required to give a net charge of zero. 3.Subscripts give the smallest ratio between the ions. 4.Use parentheses only with polyatomic ions to show more than one of them in the formula. 5.Do not write “1” as a subscript. No subscript means 1.

15. A useful method to find subscripts: If you have 2 ions, Ct +p and An -n and you want to find the formula, Ct x An y then x(p) = y(n) So = reduce to lowest fraction Practice: Write the formulas for the compounds that form between the following pairs of ions. Sodium and FluorideSodium and Oxide Calcium and chlorideMagnesium and sulfide x ynp

17. More practice: use your ion sheet to find symbols and charges for ions Aluminum and bromideAluminum and oxide Iron(II) and fluorideChromium(III) and oxide Titanium(IV) and oxide Potassium and hydroxide Calcium and hydroxide Lithium and sulfate Silver and phosphateCobalt(II) and phosphate

18. Writing names for ionic compounds: Notice on the previous page that you were given the names of each ion. If you remove the “and” that is the name of the compound!!! To name an ionic compound just write the names of the individual ions with the cation first. The only trick to this is that if an element forms more than 1 ion, you must tell which ion it is by using roman numerals. This is only needed for some cations. (See your ion sheet) The roman numeral in parentheses tells you the charge on the cation. Iron (III) = Fe +3 Iron (II) = Fe +2

19. To find which cation you have, you can use x/y = n/p You will have x, y and n – just find p as the charge on the cation! Ex: Write the name for SnO 2 x=1, y=2, n=2 so we can write: 1 / 2 = 2 / p So p must be 4 and the formula is Tin(IV) oxide

20. Practice: Write the names for the following ionic compounds: K 2 SSrF 2 AlBr 3 NaNO 3 CaSO 4 NH 4 Cl TiO 2 FeCl 3 V 2 (SO 4 ) 3 CrPO 4 Cu 2 OPb(C 2 H 3 O 2 ) 2

21. The structure of ionic compounds Term: Salt – Any ionic compound other than oxides and hydroxides. Note that when we write the formula for ionic compounds we always use the lowest ratio of ions. We do not give the exact numbers of cations and anions in a sample of the compound. This type of formula is called an empirical formula.

22. This is because each ion in the structure is electrostatically attracted to each oppositely charged ion. The ions will arrange themselves in complex structures so that each ion will be as close as possible to opposite ions and as far as possible from same charged ions. As more ions are added a network of repeating geometrical patterns will form this is a crystal. Ionic compounds do not exist as individual molecules but as networks of ions in fixed ratios bound together by electrostatic forces.

23. Terms: Crystal Lattice – The geometric pattern that ions form in a crystal. How the ions are packed together. Unit cell – The simplest repeating pattern in a lattice. Properties of ionic compounds: Very strong attraction between ions High melting and boiling points Generally hard and brittle Conduct electricity when melted or in solution

29. http://www.webmineral.com/

30. Many common ionic compounds found in nature (minerals) are silicate compounds. Silicates are complex compounds of Si, O (polyatomic ion) and various cations. They also often contain H, OH -, and H 2 O. Examples: granite – Na, K, Al silicate asbestos – Fe, Mg silicate-hydroxide asbestos – Fe, Mg silicate-hydroxide Hydrates – Ionic compounds with water molecules incorporated into the crystal lattice in very specific amounts. Examples: sodium carbonate hydrate- Na 2 CO 3. 10H 2 O Calcium sulfate hydrate – CaSO 4. 5H 2 O