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Chapter 11

Chapter 11. Chemical Reactions. 11.1 Goals. In this section you will be able to… Determine how to write a word equation Describe how to write a skeleton equation Describe the steps for writing balanced equations. YOU CAN DO IT!!. Describing Chemical Reactions.

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Chapter 11

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  1. Chapter 11 Chemical Reactions

  2. 11.1 Goals In this section you will be able to… • Determine how to write a word equation • Describe how to write a skeleton equation • Describe the steps for writing balanced equations YOU CAN DO IT!!

  3. Describing Chemical Reactions • Chemical equations use symbols and formulas to represent the identities and relative molar amounts of reactants and products in a chemical reaction • Chemical equations • Represent known facts • Must contain the correct formulas for reactants and products • Must follow the Law of Conservation of Mass

  4. Word equations use the names of reactants and products to describe the reaction Example 1: “Elementary hydrogen, which consists of diatomic molecules, reacts with oxygen molecules in the air to form molecules of liquid water”

  5. Formula Equation= 2H2 + O2 2H2O • Advantage of word equation = reactants and products can easily be identified, includes the process of the reaction • Disadvantage of word equation = amounts of reactants and products are not known

  6. Example 2: “Aluminum atoms combine with oxygen in the air to give the compound aluminum oxide. The piece of shiny aluminum becomes dull.” • Advantage = reactants and products are identified, includes description • Disadvantage = no amounts given!! • Formula Equation = 4 Al (s) + 3 O2 (g)  2 Al2O3 (s)

  7. Example 3: Word Equation = “Zinc atoms react with Iodine atoms in the presence of water to form Zinc Iodide” • “Skeleton” Formula Equation = Zn + I2 ZnI2 • Balanced Formula Equation = 1 mole Zn (s) + 1 mole I2 (s)  1 mole ZnI2 (s) • Does the mass of Zinc and Iodine used make a difference?

  8. Chemical Equations (Formula Equations) use chemical formulas to describe a chemical reaction • Skeleton Equations are formula equation that do not indicate the amounts of reactants and products (just shows the proper formulas)

  9. Yield Sign = an arrow similar to the equal sign in a math equation, it means produces/forms Physical States = indicated by (s), (l), (g), (aq) Other symbols – as shown in table 11.1 on p. 323 represent catalysts and energy changes involved in the reaction (E.g. reversible reactions shown by double arrow) • Remember! A chemical equation MUST be accurate in EVERY detail to be useful

  10. Balancing Chemical Equations • Coefficients are small whole numbers that appear in front of a chemical formula in an equation • Coefficients are used to show relative amounts of substances and follow the laws of conservation of mass and atoms

  11. Rules for Balancing Equations • Identify the names of the reactants and products and write a word equation • Write a formula (skeleton) equation • Balance the formula equation according to the laws of conservation of mass and atoms by adding coefficients • Count the atoms (again!) to be sure the equation is balanced

  12. Example 1: “Bottled gas, which consists mostly of propane (C3H8) burns in pure oxygen or in air to form gaseous carbon dioxide and liquid water. Write the balanced equation. • Facts = C3H8 (g) O2 (g) CO2 (g) H2O (l) • Basic Equation C3H8 (g) + O2 (g)  CO2 (g) + H2O (l) • Adjust Coefficients = by checking back and forth C3H8 + 5 O2 3 CO2 + 4 H2O • Check = for the lowest whole number ratio of the coefficients

  13. Example 2: Zinc reacts with hydrochloric acid to form zinc chloride and hydrogen • Facts = Zn HCl ZnCl2 H2 • Basic Equation = Zn (s) + HCl (l)  ZnCl2 (aq) + H2 (g) • Balanced Equation = Zn (s) + 2HCl (l)  ZnCl2 (aq) + H2 (g) • Check = for the simplest whole-number ratio  

  14. Example 3: Oxygen can be prepared in the lab by heating mercury (II) oxide, with mercury as a byproduct • Facts = HgO Hg O2 • Basic Equation = HgO (s) Hg (l)+ O2 (g) • Balanced Equation = 2 HgO (s)  2 Hg (l) + O2 (g) Check = for the simplest whole-number ratio

  15. Example 4: Sodium Hydrogen Carbonate reacts to form sodium carbonate, water, and carbon dioxide • Facts = NaHCO3 Na2CO3 H2O CO2 • Basic Equation = NaHCO3 Na2CO3 + H2O + CO2 Balanced Equation = 2 NaHCO3 (s)  Na2CO3 (s) + H2O (l) + CO2 (g)

  16. Example 5: Aluminum Sulfate and Calcium Hydroxide are added to water to remove suspended matter. The two chemicals react to form Aluminum Hydroxide and Calcium Sulfate • Facts = Al2(SO4)3 Ca(OH)2 Al(OH)3 CaSO4 • Basic Equation = Al2(SO4)3 + Ca(OH)2 Al(OH)3 + CaSO4 • Balanced Equation = Al2(SO4)3 (aq) + 3 Ca(OH)2 (aq)  2 Al(OH)3 (s) + 3CaSO4 (s)

  17. Hints • Write all formulas correctly using crossover method • NEVER change subscripts once you have the right formula • Recheck chapter 9 for help on formulas

  18. 11.2 Types of Chemical Reactions • There are 5 types of chemical reactions: Combination (Synthesis, Composition) Reactions General Form = A + X  AX Reactions of Elements with Oxygen & Sulfur = Ca (s) + O2 (g)CaO (s) Fe (s) + S (s)FeS (s) (ratio depends on the oxidation state of the metal) Reactions of two nonmetals = 2N2 (g) + O2 (g)  2N2O (g) Reactions of metals with Halogens = 2Co (s) + 3F2 (g) 2 CoF3 (s) Synthesis reactions of oxides = CaO (s) + H2O (l)  Ca(OH)2 (s)

  19. Decomposition Reactions General Form = AX  A + X Decomposition of binary compounds = 2 H2O  2 H2 + O2 Decomposition of Metal Carbonates = when heated, form metallic oxides and carbon dioxide Metallic Carbonate  metallic oxide + carbon dioxide(g)

  20. Demo – Heating CuCO3 and flame-testing the gas CuCO3CuO + CO2(g) Decomposition of Metal Hydroxides = when heated, decompose into metallic oxides and water Metallic Hydroxide  Metallic Oxide + water MOH MO + H2O Demo – Ca(OH)2CaO + H2O Decomposition of Metal Chlorates = when heated, decompose into metallic chlorides and oxygen Metallic Chlorate  Metallic Chloride + oxygen MClO3MCl + O2 (g) Lab from earlier – 2 KClO3 2 KCl + 3 O2 (g)

  21. Decomposition of acids = when heated, decompose into nonmetallic oxides and water Acid  Nonmetallic Oxide + water Acid NmO + H2O Nm = nonmetal Demo – H2CO3 CO2 + H2O Carbonic Acid Decomposition of oxides = a very few will decompose when heated Oxide  Metal/Metallic oxide/Nonmetal + Oxygen Oxide MO + O2 (g) 2HgO  2 Hg + O2 (g) Disc/Demo: 2 H2O2 2 H2O + O2 (g) Decomposition by electric current (electrolysis) substance  element/compound + element AB  A + B 2 H20  2 H2 (g) + O2 (g) Some decomposition reactions are very exothermic – such as NH4NO3 (ammonium nitrate) used as an explosive and a fertilizer

  22. Single-Replacement (Displacement) Reactions – Note Activity Series on p. 333! General Form = A + BX  AX + B or Y + BX  BY + X Replacement of a metal in a compound by a more reactive metal Ionic compound + metal  metal 2 + ionic compound 2 AX + B  BX + A Demo: Zn + CuSO4 ZnSO4 + Cu (s) Replacement of Hydrogen in water by a metal Metal + H2O  Metallic Hydroxide + H2 A + H2O  AOH + H2 Ca + H2O  Ca(OH)2 + H2 (g) Demo?: Na + H2O NaOH + H2 (g) Replacement of Hydrogen in an Acid by a metal Metal + Acid  Ionic Compound + H2 (g) A + Acid X  AX + H2 (g) Demo: Zn + H2SO4 ZnSO4 + H2 (g)

  23. Replacement of Halogens based on Activity Halogen1 + Ionic Compound2 Ionic Compound1 + Halogen2 A + XB  XA + B Cl2 + 2KBr  2 KCl + Br2 MORE active elements replace less active ones

  24. Rules for completing Replacement Reaction equations • Locate the element on the reactant side and decide whether it is a metal or a nonmetal • Decide what element in the compound it will replace • Check the activity series to see if the element in question is above the one it is replacing – If so, it will replace it – If not, write NO REACTION • Write the correct formula of the compound formed by checking oxidation numbers • Balance the equation!

  25. Double-Replacement (Displacement) Reactions General Form = AX + BY  AY + BX or A(aq) + Y (aq)  AY (s) Formation of a Precipitate = AgNO3 (aq) + NaCl (aq) AgCl (s) + NaNO3 (aq) Formation of a Gas = FeS (s) + 2 HCl (aq)  H2S (g) + FeCl2 (aq) Formation of water (acid-base rxn.) = HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l)

  26. Combustion Reactions – General form = Hydrocarbon + Oxygen  Carbon Dioxide + Water Example (propane) = C3H8 (g) + 5 O2 (g)  3 CO2 (g) + 4 H2O (g)

  27. Reversible Reactions Example 1 = 4 H2 + Fe3O4 3 Fe + 4 H2O (g) AND 3 Fe + 4 H2O  Fe3O4 + 4 H2 (g) Conditions: Must be in a CLOSED system and set up an equilibrium Example 2 = 2 HgO↔2 Hg + O2 closed system (we hope!) Example 3 = Haber process – forms NH3 when a N2 and H2 gas mixture is sparked N2 + 3 H2↔2 NH3 (g) + 22 kcal Predicting the Products of a Chemical Reaction See Road Map!!!

  28. Net Ionic Equations Complete Ionic Equation – shows the formulas for dissolved ionic compounds as dissociated free ions Ag+ (aq) + NO3- (aq) + Na+ (aq) + Cl- (aq) AgCl (s) + Na+ (aq) + NO3– (aq) Spectator Ions – ions that appear on both sides of the equation (often nitrates) Net Ionic Equation – shows only those particles that are directly involved in the chemical change (no spectators!) Must be balanced by mass and charge Example – Fe (s) + CuSO4 (aq)  FeSO4 (aq) + Cu (s) becomes Fe (s) + Cu2+ (aq)  Fe2+ (aq) + Cu (s)

  29. Predicting the Formation of a Precipitate Definition – an insoluble solid that forms as the result of a double-displacement (ionic) reaction The formation of a precipitate depends upon the solubility of the products in a reaction and can be predicted as follows (Table 11.3 on p. 344) Salts of alkali metals and ammonia = soluble Nitrate salts and chlorate salts = soluble Sulfate salts (except compounds of Pb2+, Ag+, Hg22+, Ba2+, Sr2+, and Ca2+) = soluble Chloride salts (except compounds of Ag+, Pb2+, and Hg22+) = soluble Carbonates, phosphates, chromates, sulfides, and hydroxides = MOST are Insoluble  

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