More on Reactions

1. More on Reactions Predicting Products Solubility Rules

2. Synthesis • In a simple synthesis reaction, two elements can combine to form a binary ionic compound • Remember to correctly write the ionic compound’s formula, you must balance charges • Balance the equation when you finish.

3. Decomposition • In a simple decomposition reaction, we can predict that a binary compound will break down into its elements • Remember that some elements are diatomic in their elemental form • Balance the equation when you finish.

4. Combustion • If a hydrocarbon (any compound with C and H) is combusted, it will react with oxygen gas (O2) to burn. • The products will be carbon dioxide (CO2)and water, usually in the form of water vapor or steam (H2O). • Then balance the equation with coefficients to have correct #’s of C,H, and O on both sides.

5. Single Replacement One element replaces another in a compound • Cationic: a metal replaces the cation (another metal) in a compound • H can act as the cation in these reactions • Anionic: a nonmetal replaces the anion (another nonmetal) in a compound

6. Single Replacement • For this class, you can assume a single replacement reaction will occur if you are asked to predict its product. • Don’t forget that some elements are diatomic when they are alone (not in a compound). You also must balance charges to write correct formulas for the compounds.

7. Double Replacement • The cation of one compound switches places with the cation of the other compound • Don’t forget that you will need to balance charges to write the correct formula for each compound. • Balance the overall equation after you have written correct formulas.

8. Types of Double Replacement • Production of a gas • Two compounds combine to form a gas and another product • Production of a precipitate • Two aqueous solutions combine to form a solid precipitate and another product • Acid/Base Neutralization* • An acid combines with a base to form a salt (any ionic compound) and water

9. Acid/Base Neutralization Terms* • Remember, acids begin with H (not water) • Examples: HCl, HNO3 • Bases are ionic compounds with hydroxide (OH) as the anion • Examples: NaOH, Ba(OH)2 • The word “salt” refers to any ionic compound formed from reacting an acid with a base • Examples: CaBr2, NaCl

10. Solubility Rules • Solubility Rules are used to give us a general idea of whether or not an ionic compound will dissolve in water to form an aqueous solution • These are very useful in determining states of matter for your compounds in reactions, especially double replacement precipitation reactions.

11. What Solubility Rules Mean • Solubility Rules give you general guidelines based on cations and anions present in an ionic compound. • For each compound you have in the presence of water, look for a rule that will describe it.

12. How to use the Solubility Rules • If the compound is insoluble in water, it will not dissolve and will fall out as a solid precipitate. • If the compound is soluble in water, it will dissolve and form an aqueous solution.

13. Solubility Rules • Most nitrate salts are soluble. • Most salts of Na+, K+, and NH4+ are soluble. • Most chloride salts are soluble; notable exceptions are Ag+, Pb2+, and Hg2 2+. • Most sulfate salts are soluble; notable exceptions are Ba2+, Pb2+, and Ca2+. • Most hydroxide compounds are mostly insoluble; the important exceptions are Na+ and K+ and Ba2+ and Ca2+ are mostly soluble. • Most sulfide, carbonate, and phosphate salts are mostly insoluble except Na+, K+, and NH4+. • Most dichromate salts are soluble except Ag+ and Pb+2.  * insoluble compounds will fall out of solutions as precipitates