1. 6.3 – NOTES Formulas of Ionic Compounds

2. III. Names and formulas for Ionic compounds
A. Formulas for ionic compounds
1. Determining charge using the periodic table elements are stable with a noble gas configuration; atoms can gain/lose/share electrons to get an octet;

3. Group 1 has one valence electron so it is easier to lose one valence and have a full shell than it is to gain seven electrons to have a full shell.
Result: all group I metals will lose one electron (forming a cation with a charge of +1)
Na – [Ne] 3s1 To get octet, could lose 1 e- to become Na+ with a configuration of [Ne] (octet!)
OR could gain 7 e- to become Na-7 with a configuration of [Ar] (octet!)
*gaining 7 e- would yield an octet, but not realistic

4. Group 16 has six valence electrons and can lose 6 electrons or gain 2 electrons to obtain a full s and p (octet!).
It is easier to gain 2 e- so nonmetals in group 16 will gain 2 electrons (forming an anion with a charge of -2)

5. Transition metals vary and most of the time you should consult your ion list. However, some you can predict.
Sc and Y have a configuration of s2d1 and will lose all three electrons resulting in a charge of +3
Ti and Zr have a configuration of s2d2 and will lose four electrons resulting in a charge of +4

6. Fill in charges!

7. 2. Use the charges to write a formula
2. Use the charges to write a formula. The cation is first, then the anion. The total charge must be zero, so we use subscripts to indicate the number of ions. The ions must have the smallest ratio possible (empirical formula). To obtain the smallest ratio, reduce charges BEFORE criss-crossing if possible
Examples:
sodium and bromine calcium and fluorine CaF2
Na1Br1 * the number 1 is assumed so it is written as NaBr

8. potassium and oxygen barium and sulfur
K2O Ba S-2 *can reduce charges to 1
= BaS

9. aluminum and chlorine magnesium and nitrogen
AlCl3
Mg3N2
lithium and sulfur scandium and oxygen
Li2S
Sc2O3

10. Polyatomic ions in formulas
Definition: ions made up of one or more than one atom; there is a list on the back of your periodic table;
Can be anions or cations
Use parentheses if more than 1 polyatomic ion is needed. Again use the lowest ratio of ions (empirical formula).

11. Examples:
ammonium and sulfur potassium and carbonate
*ammonium on ion list, sulfur on PT potassium on PT and carbonate on list
NH S K + CO3 -2
(NH4)2S K2CO3
*must put ( ) around NH4 since the 2 carries down *do not need ( ) b/c only have 1 ion

12. aluminum sulfate magnesium cyanide
Al2(SO4)3 Mg+2 CN - *CN has to be in ( ) b/c more than 1 is present and a polyatomic
Mg(CN)2

13. Exceptional ions: Peroxide, mercury (I), and oxalate
Exceptional ions: Peroxide, mercury (I), and oxalate. Often these do not obey the simplest formula rule. There are 3 exceptions: peroxide is O2-2, mercury I is Hg2+2 and oxalate is C2O4-2. These ions do not get reduced. It is best to reduce charges BEFORE criss-cross to ensure the correct answer.

14. Peroxide – O2-2 **Must keep subscript “2” w/ group I: Na+ O2-2  Na2O2
charges cannot be reduced so final formula cannot be reduced;
Mercury (I) – Hg2+2 ** Must keep superscript “2” w/ group 7A: Hg2+2 F-  Hg2F2
Oxalate – C2H4-2 **Must keep subscripts w/ group I: Na+ C2O4 -2  Na2C2O4