1. CHEMICAL REACTIONS - Ch.8 Describing Chemical Change Reactants  Products Bonds are broken and new bonds are formed No atoms are created or destroyed Catalyst- a substance that speeds up the rate of a reaction but is not used up in the reaction. Since it is not a reactant or product, it is written above the arrow in the equation

2. Chemical Equations Skeleton equation does not show relative amounts of reactants and products Fe(s) + O 2 (g)  Fe 2 O 3 (s) H 2 O 2 (aq) H 2 O(l) + O 2 (g)

3. Common Symbols for EQ’s + separates 2 or more reactants or 2 or more products  yields (separates reactants and products)  a reversible reaction (s) solid state (l) liquid state (g) gaseous state (aq) aqueous solution (dissolved in water) Heat is supplied Catalyst is involved

4. Skeleton Equation Practice Solid sodium hydrogen carbonate reacts with hydrochloric acid to produce aqueous sodium chloride, water, and carbon dioxide gas NaHCO 3 (s) + HCl (aq)  NaCl (aq) + H 2 O (l) + CO 2 (g)

5. Balancing Chemical Equations Coefficients are placed in front of symbols to show relative amounts Since no atoms are lost, equation must be balanced (same # of each kind of atom on both sides of the equation) – The law of conservation of mass.

6. Rules for Balancing Equations Determine correct formulas for all reactants & products Place them on proper side of  Separate multiple reactants or products by using a + Count the number of atoms of each element on both sides (polyatomic ions are counted as a single unit if unchanged in the reaction)

7. Balance the elements 1 at a time by using coefficients. Begin with elements that appear only once on each side of equation. Do not ever change subscripts, you can’t change a correct formula. Balance the following particle types if they are present. –Balance the nonmetals –Balance the metals –Balance the hydrogen –Balance the oxygen last. –Check each atom or polyatomic ion to be sure that the equation is balanced

8. Balance These: AgNO 3 + H 2 S  Ag 2 S + HNO 3 MnO 2 + HCl  MnCl 2 + H 2 O + Cl 2 Zn(OH) 2 + H 3 PO 4  Zn 3 (PO 4 ) 2 + H 2 O Make sure all coefficients are in the lowest possible ratio that balances the full equation.

9. Types of Chemical Reactions Classifying Reactions Combination Decomposition Single-replacement Double-replacement Combustion

10. Combination or Synthesis Reactions 2 or more substances combine to form only 1 substance; product is always a compound Metal + Non-metal  Salt 2K(s) + Cl 2 (g)  2KCl(s) 2 nonmetals can form more than one product S(s) + O 2 (g)  SO 2 (g) sulfur dioxide and 2S(s) + 3O 2 (g)  2SO 3 (g) sulfur trioxide

11. Combination cont’d Transition metal and nonmetal can make more than 1 product Fe(s) + S(s)  FeS(s) iron (II) sulfide 2Fe(s) + 3 S(s)  Fe 2 S 3 (s) iron (III) sulfide Some nonmetallic oxides react with water to produce an acid (hydrogen ions in aqueous solution) SO 2 (g) + H 2 O(l)  H 2 SO 3 (aq) - sulfurous acid Metallic oxides react with water to give a base containing hydroxide ions CaO(s) + H 2 O(l)  Ca(OH) 2 (aq) - calcium hydroxide

12. Practice Al(s) + O 2 (g)  Cu(s) + S(s)  SO 3 (g) + H 2 O(l)  CO 2 (g) + H 2 O(l)  K 2 O(s) + H 2 O(l) 

13. Decomposition Reactions XY  X + Y A single compound is broken down into 2 or more products Can be difficult to predict products Usually require energy in the form of heat, light or electricity to happen Rapid decomposition reactions are often explosive

14. Decomposition Reactions If a binary compound, it will break into its constituent elements – don’t forget about the binary molecules. Metallic chlorates decompose into metallic chlorides and oxygen. Metallic carbonates typically decompose into metallic oxides and carbon dioxide. Metallic hydrogen carbonates decompose into metallic carbonates, carbon dioxide and water.

15. Practice- write the correct formula & don’t forget to balance Mg(ClO 3 ) 2  + PbO 2 (s)  HI   Na + Cl 2 CaCO 3 (s) 

16. Start reading and working problems for Chapter 8 page 202 - 235

17. Single-Replacement (displacement) Reactions XY + Z  ZY + X 1 element replaces a second element that forms a like charge in a compound Can determine if 1 metal will replace another by their relative reactivities Table 8.2 p. 217 lists metals in decreasing reactivity Nonmetals can replace another nonmetal usually the halogens, activity decreases as you go down the family, 7A

18. Single-Replacement (displacement) Reactions If one element is less reactive than the other, no reaction will take place. A reactive metal will replace any metal (or hydrogen) that is less reactive than it Metals that dissolve in water will form a metallic hydroxide (base) and release hydrogen gas. –Na(s) + H 2 O(l)  NaOH(aq) + H 2 (g)

19. Practice Zn(s) + H 2 SO 4 (aq)  Ca(s) + H 2 O(l)  Sn(s) + NaNO 3 (aq)  Cl 2 (g) + NaBr(aq) 

20. Double-Replacement Rx AB + CD  CB + AD Positive ions in 2 reacting compounds exchange places; usually ionic compounds in aqueous solutions (often form precipitates) The following are usual situations for double- replacement Insoluble precipitates (Solubility Rules Table 8.3)- Na 2 S(aq) + Cd(NO 3 )(aq)  CdS(s) + 2 NaNO 3 (aq) yellow precipitate forms because cadmium sulfide is only slightly soluble in water Production of a Gas - –2NaCN(aq) + H 2 SO 4 (aq)  2HCN(g) +Na 2 SO 4 (aq) One product is a molecular compound such as water, as in an acid-base reaction –Ca(OH) 2 (aq) + 2HCl(aq)  CaCl 2 (aq) + 2 H 2 O(l)

21. Practice (& balance) a precipitate of barium carbonate is formed BaCl 2 (aq) + K 2 CO 3 (aq)  Hydrogen sulfide gas (H 2 S) is formed FeS(s) + HCl (aq)  Iron hydroxide is a precipitate NaOH + Fe(NO 3 ) 3 

22. Combustion Reaction An element or compound reacts with oxygen, often producing heat or light Commonly involve hydrocarbons which produce products of CO 2 and H 2 O and large amount of heat are produced If supply of O 2 is insufficient, combustion will be incomplete and may form elemental Carbon and CO

23. Practice and Balance Mg (s) + O 2 (g)  S (s) + O 2 (g)  Benzene C 6 H 6 (l) + O 2 (g)  Methanol CH 3 OH (l) + O 2 (g) 

24. Homework 38,39,43,44,46-51,54

25. Reactions in Aqueous Solution 70% of Earth’s surface is covered by water and 66% of human body is water Many important chemical reactions take place in water (aqueous solution)

26. Net Ionic Reactions AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq) This equation is written with ionic compounds written in their formula units. Since most ionic compounds dissociate or separate into their cations and anions when they dissolve in water then the equation can be written as a complete ionic equation Ag + (aq) + NO 3 - (aq) + Na + (aq) + Cl - (aq)  AgCl(s) + Na + (aq) NO 3 -(aq)

27. Net Ionic Equations This equation can be simplified by eliminating spectator ions (ions that do not participate in the reaction) from both sides of the equation Ag + (aq) + NO 3 - (aq) + Na + (aq) + Cl - (aq)  AgCl (s) + Na + (aq) + NO 3 - (aq) Net ionic reaction is: Ag + (aq) + Cl - (aq)  AgCl (s)

28. You must balance the ionic charges Pb (s) + AgNO 3 (aq)  Ag (s) + Pb(NO 3 ) 2 (aq) Nitrate ion is the spectator ion  net ionic equation- (unbalanced with respect to charges)  Pb (s) + Ag + (aq)  Ag (s) + Pb 2+ (aq) Pb (s) + 2Ag + (aq)  2Ag (s) + Pb 2+ (aq) (charges and particles are balanced) Single and double replacement reactions can generally be written as net ionic reactions

29. Identify Spectator Ions and Write Balanced Net Ionic Equations HCl (aq) + ZnS (aq)  H 2 S (g) + ZnCl 2 (aq) Cl 2 (g) + NaBr (aq)  Br 2 (l) + NaCl (aq)

30. More practice Aqueous solutions of iron III chloride and potassium hydroxide are mixed, forming a precipitate of iron III hydroxide. Write the full equation and then the net ionic equation. FeCl 3 (aq) + KOH(aq)  KCl(s) + Fe(OH) 3 (s) Fe 3+ (aq) + OH - (aq)  Fe(OH) 3 (s) Balance the final reactions.

31. Predicting Precipitate Formation 1.Salts of alkali metals and ammonia are soluble (except some Li compounds) 2.Nitrate salts and Chlorate salts are soluble 3.Sulfate salts are soluble (except Pb, Ag, Hg, Ba, Sr, and Ca) 4.Chloride salts are soluble ( except Ag and some Hg and Pb) 5.Most carbonates, phosphates, chromates, sulfides, and hydroxides are insoluble (except compounds of alkali metals and ammonia)

32. Using Solubility Rules Predict Precipitate Formed: H 2 SO 4 (aq) + BaCl 2 (aq)  Al 2 (SO 4 ) 3 (aq) + NH 4 OH(aq)  AgNO 3 (aq) + H 2 S(aq)  NaCl(aq) + KNO 3 (aq)  CaCl 2 (aq) + Pb(NO 3 ) 2 (aq)  Ca(NO 3 ) 2 (aq) + Na 2 CO 3 (aq) 

33. Finish Homework, #55-58, 61 & 66 The End