1. Chapter 10 Chemical Reactions

2. Understanding Chemical Reactions A chemical reaction occurs when: A change in energy occurs Exothermic –gives off energy Endothermic – absorbs energy The color changes A new gas is created A precipitate is formed from 2 solutions

3. Parts of Reaction Reactants- starting “ingredients” for the reaction (written to LEFT of arrow) Products- final result(s) of reaction (written to RIGHT of arrow) Arrow- signifies a change is taking place (read “react to form” or “yield”) Symbols- used to specify properties of reactants and products or conditions of the reaction (s), (l), (g), (aq)

4. Word Equations CuSO 4(aq) + Ag (s)  Ag 2 SO 4(aq) + Cu (s) “aqueous copper (II) sulfate and solid silver react to produce aqueous silver sulfate and solid copper.”

5. Balancing Equations Why? Law of Conservation of Mass (LCM)- mass is neither created nor destroyed in chemical reactions (Antione Lavosier, 1700’s) Coefficients- added so reactions obey the LCM (DON’T touch subscripts!!)

6. Procedure for Balancing Write formulas for reactants and product on respective sides of the equation Perform initial “atom inventory” on each side of the equation Add coefficients to make inventories match Update and repeat as needed Verify answer

7. Types of Reactions 1.Combustion 2.Synthesis 3.Decomposition 4.Single replacement 5.Double replacement

8. Combustion Reaction with oxygen that releases energy (heat or light) Hydrocarbons (contain only C and H) complete combustion always produces CO 2 (g) and H 2 O(g) note: incomplete combustion may result in C(s) or CO(g)

9. Combustion Examples Complete combustion of propane (C 3 H 8 ) Combustion of aluminum

10. Synthesis (or Combination) Two or more substances combine to form a “larger” compound. Group 1A & 2A metals- use common charges to predict ionic compound formed Transition metals or 2 nonmetals- many products are possible

11. Synthesis Examples Synthesis of barium chloride Combination of copper and sulfur

12. Decomposition More complex molecule breaks down into simpler substances Binary compounds- decompose into constituent elements Ternary or more complex compounds- decompose into either constituent elements or into a combination of compounds SUMMARY OF REACTIONS in PACKET

13. Decomposition Examples Decomposition of sulfur trioxide gas Decomposition of ammonium nitrate to produce dinitrogen oxide and water

14. Decomposition vs Dissociation Decomposition- atoms with NEW properties are created Dissociation- ions without new properties are created

15. Single Replacement Possible reaction between an element and an aqueous compound Metals ONLY replace metals and nonmetals ONLY replace nonmetals Activity series (front of packet) used to predict whether replacement will occur

16. S.R. Examples Lithium is combined with calcium nitrate Silver is placed into a solution of magnesium carbonate Chlorine gas is bubbled through a solution of aluminum bromide

17. Double Replacement Possible exchange of partners by two compounds Opposites attract- two positive ions WILL NOT combine (two negative ions won’t either) Solubility Rules (handout)- used to predict if double replacement will occur

18. D.R. Examples Solutions of magnesium nitrate and aluminum phosphate are mixed Solutions of sodium chloride and aluminum nitrate are mixed

19. Net Ionic Equations Eliminate spectator ions Solutions of sodium chloride and silver nitrate are mixed

20. Double Replacement to Form Water Acids and bases react in a double replacement reaction and produce water Acid- produces H + ions, pH < 7 Base- produces OH - ions, pH > 7 Neutralization- producing neutral water (pH = 7)

21. Neutralization Examples Hydrochloric acid and sodium hydroxide are mixed Sulfuric acid and calcium hydroxide are mixed

22. TEST