1 / 29

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.

jiro
Télécharger la présentation

Chapter 11

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  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  

More Related