1. Chapter 11 Chemical Reactions Chemistry 2

2. Describing Chemical Reactions 11.1

3. Writing Chemical Equations 11.1 Reactants  products ▫  = yields, gives, or reacts to produce Iron + Oxygen  iron(III) oxide Hydrogen peroxide  water + oxygen ▫Hydrogen peroxide decomposes to form water and oxygen gas Skeleton equation – chemical equation that does not indicate the relative amount of the reactant and products Fe(s) + O 2 (g)  Fe 2 O 3 (s) ▫(s) = solid ▫(l) = liquid ▫(g) = gas ▫(aq) = aqueous (dissolved in water) Catalyst – substance that speeds up the reaction but is not used up in the reaction ▫Formula is written above the arrow Practice Problems Page 324 # 1-2

5. Balancing Chemical Equations 11.1 Coefficients – small whole #s placed in front of formulas to balance Each side of equation should have same # of each type of atom Law of Conservatio nof mass: mass is neither created or destroyed in chemical reaction ▫Atoms rearranged ▫Bonds broken and bond formed

7. Here are some practice problems. 1.__NaCl + __BeF 2  __NaF + __BeCl 2 2. __FeCl 3 + __Be 3 (PO 4 ) 2  __BeCl 2 + __FePO 4 3. __AgNO 3 + __LiOH  __AgOH + __LiNO 3 4. __CH 4 + __O 2  __CO 2 + __H 2 O 5. __Mg + __Mn 2 O 3  __MgO + __Mn Practice Problems: Page 327 # 3 – 4, Page 328 #5-6

8. Types of Chemical Reactions 11.2

9. Classifying Reactions 11.2 5 types of reactions 1.Combination or Synthesis – 2 or mores substances react to form a single new substance  Group A metal + nonmetal = 2K + Cl 2  2KCl  2 nonmetals can have more than 1 product  S + O 2  SO 2  2S + 3O 2  2SO 3  Transition metal and nonmetal can have more than 1 product  Fe + S  FeS  2Fe + 3S  Fe 2 S 3  Practice Problems page 331 # 13 - 14

10. 2.Decomposition Reactions – chemical change in which a single compound breaks down into two or more simpler products ▫Opposite of synthesis ▫2HgO  2Hg + O2 ▫Difficult to predict product ▫Most required energy input (endothermic) ▫Practice Problems page 332 # 15 – 16 3.Single-Replacement Reaction – chemical change in which one element replaces a second element in a compound ▫2K + 2H2O  2KOH + H2 ▫Both reactant and product contain of an element and a compound ▫Activity series – list metals in order of decreasing reactivity = table 11.2 page 333  Halogens = reactivity decreases down column  Br 2 + NaI  NaBr + I 2  Br 2 + NaCl  no reaction

11. Metals from Li to Na will replace H from acids and water Metals from Mg to Pb will replace H from acids only Practice Problems Page 334 # 17

12. 4.Double Replacement Reactions – a chemical change involving an exchange of positive ions bvetween 2 compounds ▫To occur, one of the following is usually true:  1 of produces is slightly soluble and precipitates from solution  Na 2 S(aq) + CD(NO 3 )(aq)  CdS(s) + 2NaNO 3 (aq) ▫CdS precipitated out  One of products is a gas  One product is a molecular compound ▫Practice Problems Page 335 # 18 - 19

13. 5.Combustion Reaction – chemical change in which an element or a compound reacts with oxygen often producing energy in the form of heat and light ▫ALWAYS involves oxygen ▫Often uses hydrocarbons  Complete combustion form CO2 and water (and energy)  Incomplete combustion forms CO  Supply of oxygen is limited ▫Practice Problems Page 337 # 20 - 21

14. Predicting the Products of a Chemical Reaction 11.2 # of elements/compounds reacting is a good indicator of possible reaction type ▫Combination = 2 or more reactants  single product ▫Decomposition = single compound  2 or more substances ▫Single-replacement = element + compound  element + compound ▫Double-Replacement = 2 ionic compounds  2 new compounds ▫Combustion = Oxygen + hydrocarbon (usually)  water and carbon dioxide

18. Reactions in Aqueous Solution 11.3

19. Net Ionic Equations 11.3 Many of chemical reactions take place in water (aqueous solution) ▫70% Earth surface covered by water ▫66% of human body is water Complete Ionic equation – an equation that shows dissolved ionic compounds as dissociated free ions ▫2Na 1+ (aq) + SO 4 2- (aq) + Ba 2+ (aq) + 2Cl 1- (aq)  2Na 1+ (aq) + 2Cl 1- (aq) + BaSO 4(s)  Cross out ions that appear unchanged on both sides = spectator ions ▫Na 1+ (aq) + SO 4 2- (aq) + Ba 2+ (aq) + 2Cl 1- (aq)  2Na 1+ (aq) + 2Cl 1- (aq) + BaSO 4(s)  Write the net ionic equation ▫Ba 2+ (aq) + SO4 2- (aq)  BaSO 4(s)  Then balance

20. Predicting the Formation of a Precipitate 11.3 Mixing ionic compounds can sometimes form a precipitate (insoluble salt) Solubility Rules for Ionic Compounds ▫Salts of alkali metals (1A) and ammonia (NH 4 ) + = Soluble ▫Nitrate (NO 3 ) - salts and chlorate (ClO 3 ) - salts = Soluble ▫Sulfate (SO 4 ) 2- slats except compounds with Pb 2+, Ag +, Hg 2 2+, Ba 2+, Sr 2+, and Ca 2+ = Soluble ▫Chloride salts, except compound with Ag +, Pb 2+, and Hg 2 2+ = Soluble ▫Carbonates (CO 3 ) 2-, Phosphates (PO 4 ) 3-, Chromates(CrO 4 ) 2-, Sulfides, and Hydroxides (OH) -

21. Precipitate example: ▫Na 2 Co 3 (aq) + Ba(NO 3 ) 2 (aq)  ?? Precipitate formed  Separate = 2Na + (aq) + CO 3 2- (aq) + Ba 2+ (aq) + 2NO 3 - (aq)  Would form NaNO 3 and BaCO 3 ▫Na = Alkali = soluble ▫Nitrate salts = soluble Carbonates generally insoluble = BaCO3 will precipitate out ▫Ba 2+ (aq) + CO 3 2- (aq)  BaCO 3 (s) Practice Problems Page 343 # 28 - 29