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Chapter 11 Chemical Reactions

Chapter 11 Chemical Reactions. Chemistry 2. Describing Chemical Reactions 11.1. Writing Chemical Equations 11.1. Reactants  products  = yields, gives, or reacts to produce Iron + Oxygen  iron(III) oxide Hydrogen peroxide  water + oxygen

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Chapter 11 Chemical Reactions

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  1. Chapter 11Chemical 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) + O2(g)  Fe2O3(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

  4. 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

  5. Here are some practice problems. • __NaCl + __BeF2 __NaF + __BeCl2 • 2. __FeCl3 + __Be3(PO4)2 __BeCl2 + __FePO4 • 3. __AgNO3 + __LiOH __AgOH + __LiNO3 • 4. __CH4 + __O2  __CO2 + __H2O • 5. __Mg + __Mn2O3  __MgO + __Mn • Practice Problems: Page 327 # 3 – 4, Page 328 #5-6

  6. Types of Chemical Reactions 11.2

  7. Classifying Reactions 11.2 • 5 types of reactions • Combination or Synthesis – 2 or mores substances react to form a single new substance • Group A metal + nonmetal = 2K + Cl2 2KCl • 2 nonmetals can have more than 1 product • S + O2  SO2 • 2S + 3O2  2SO3 • Transition metal and nonmetal can have more than 1 product • Fe + S  FeS • 2Fe + 3S  Fe2S3 • Practice Problems page 331 # 13 - 14

  8. 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 • 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 • Br2 + NaI  NaBr + I2 • Br2 + NaCl  no reaction

  9. 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

  10. 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 • Na2S(aq) + CD(NO3)(aq)  CdS(s) + 2NaNO3(aq) • CdS precipitated out • One of products is a gas • One product is a molecular compound • Practice Problems Page 335 # 18 - 19

  11. 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

  12. 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

  13. Reactions in Aqueous Solution 11.3

  14. 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 • 2Na1+(aq) + SO42-(aq) + Ba2+(aq) + 2Cl1-(aq) 2Na1+(aq) + 2Cl1-(aq) + BaSO4(s) • Cross out ions that appear unchanged on both sides = spectator ions • Na1+(aq) + SO42-(aq) + Ba2+(aq) + 2Cl1-(aq) 2Na1+(aq) + 2Cl1-(aq) + BaSO4(s) • Write the net ionic equation • Ba2+(aq) + SO42-(aq) BaSO4(s) • Then balance

  15. 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 (NH4)+= Soluble • Nitrate (NO3)- salts and chlorate (ClO3)- salts = Soluble • Sulfate (SO4)2- slats except compounds with Pb2+, Ag+, Hg22+, Ba2+, Sr2+, and Ca2+ = Soluble • Chloride salts, except compound with Ag+, Pb2+, and Hg22+ = Soluble • Carbonates (CO3)2-, Phosphates (PO4)3-, Chromates(CrO4)2-, Sulfides, and Hydroxides (OH)-

  16. Precipitate example: • Na2Co3(aq) + Ba(NO3)2(aq) ?? Precipitate formed • Separate = 2Na+(aq)+ CO32- (aq) + Ba2+ (aq) + 2NO3-(aq) • Would form NaNO3 and BaCO3 • Na = Alkali = soluble • Nitrate salts = soluble Carbonates generally insoluble = BaCO3 will precipitate out • Ba2+(aq) + CO32- (aq) BaCO3(s) • Practice Problems Page 343 # 28 - 29

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