1. Classifying Chemical Reactions There are five (5) basic types of chemical reactions. You must know how to identify each and balance each type.

2. Synthesis A reaction in which two or more substances react to produce a SINGLE product In generic terms, A + B → AB Ex: 2Na (s) + Cl 2(g) → 2NaCl (s) C (s) + O 2 (g) → CO 2 (g)

3. Combustion Reactions In a combustion reaction, oxygen combines with a substance and releases energy in the form. of heat and light. C (s) + O 2 (g) → CO 2 (g) + Energy 2H 2(g) + O 2(g) → 2H 2 O (g) + Energy Note that the combustion reactions above are also synthesis reactions. However, not all combustion reactions are synthesis reactions.

4. Combustion For example, the reaction involving methane gas, CH 4(g), and oxygen gas, O 2, illustrates a combustion reaction in which one substance replaces another in the formation of products. CH 4(g) + 2O 2(g) → CO 2(g) + 2H 2 O (g) Methane oxygen carbon dioxide water Methane is a hydrocarbon. All hydrocarbons burn in oxygen to yield the same products as methane does– carbon dioxide and water.

5. Practice Problems page 285 Write chemical equations for the following reactions and classify each reaction as either synthesis and/or combustion. 1.The solids aluminum and sulfur react to produce aluminum sulfide. 2.Water and dinitrogen pentoxide gas react to produce aqueous nitric acid. 3.The gases nitrogen dioxide and oxygen react to produce dinitrogen pentoxide gas. 4.Ethane gas (C 2 H 6 ) burns in air, producing carbon dioxide gas and water vapor.

6. Decomposition Reactions A decomposition reaction is one in which a single compound, reactant, breaks down into two or more elements or new compounds. In generic terms AB → A + B Decomposition reactions are the OPPOSITE of Synthesis reactions.

7. Examples 1. 2H 2 O (g) → 2H 2(g) + O 2(g) 2. NH 4 NO 3(s) → N 2 O (g) + 2H 2 O (g) 3. 2NaN 3(s) → 2Na (s) + 3N 2(g)

8. Replacement Reactions There are two types of replacement reactions. Single replacement and Double Replacement

9. Single Replacement Reactions A reaction in which the atoms of one element replace the atoms of another element in a compound is called a single replacement reaction. A + BX → AX + B There is always one single element on each side of the arrow. Usually, but not always, a metal is one of the single elements.

10. Single Replacement Examples Cu (s) + 2AgNO 3(aq) → 2Ag (s) + Cu(NO 3 ) 2(aq) single single 2Li (s) + 2H 2 O (l) → 2LiOH (aq) + H 2(g) single single

11. Reactivity Series Refer to the figure 10-10 on page 288 of text. A metal will not always replace another metal in a compound dissolved in water. This is because metals differ in their reactivities. A metal’s reactivity is its ability to react with another substance. Figure 10-10 orders metals by their reactivity with other metals. The most active metals, which are those that do replace the metal in a compound are at the TOP of the list. The least active are at the bottom.

12. Refer to Activity Series Ag (s) + Cu(NO 3 ) 2 (aq) → NR no reaction F 2(g) + 2NaBr (aq) → 2NaF (aq) + Br 2(l)

13. Practice Problems Page 289 21. K (s) + ZnCl 2(aq) → 22. Cl 2(g) + HF (aq) → 23. Fe (s) + Na 3 PO 4(aq) →

14. Double Replacement In DOUBLE REPLACEMENT ions between two compounds replace each other. AX + BY → AY + BX

15. 2NaOH (aq) + CuCl 2 (aq) → 2NaCl (aq) + Cu(OH) 2 (s) In this reaction, the anions (OH - and Cl - ) changed places and are now associated with the other cations (Na + and Cu 2+ ). The product, copper II hydroxide, is a solid known as a precipitate. The precipitate is shown as a (s) just after the subscript.

16. Practice problems Do the practice problems on page 291 of text.