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Chapter 4 Chemical Reactions

Chapter 4 Chemical Reactions

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Chapter 4 Chemical Reactions

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  1. Chapter 4Chemical Reactions

  2. Contents and Concepts Ions in Aqueous Solution Explore how molecular and ionic substances behave when they dissolve in water to form solutions. • Ionic Theory of Solutions and Solubility Rules • Molecular and Ionic Equations

  3. Types of Chemical Reactions Investigate several important types of reactions that typically occur in aqueous solution: precipitation reactions, acid–base reactions, and oxidation–reduction reactions. • Precipitation Reactions • Acid–Base Reactions • Oxidation–Reduction Reactions • Balancing Simple Oxidation–Reduction Equations

  4. Working with Solutions • Now that we have looked at how substances behave in solution, it is time to quantitatively describe these solutions using concentration. • 7. Molar Concentration • 8. Diluting Solutions • Quantitative Analysis • Using chemical reactions in aqueous solution, determine the amount of substance or species present in materials. • 9. Gravimetric Analysis • 10. Volumetric Analysis

  5. Learning Objectives • Ions in Aqueous Solution • 1. Ionic Theory of Solutions and Solubility Rules • a. Describe how an ionic substance can form ions in aqueous solution. • b. Explain how an electrolyte makes a solution electrically conductive. • c. Give examples of substances that are electrolytes. • d. Define nonelectrolyte and provide an example of a molecular substance that is a nonelectrolyte.

  6. e. Compare the properties of solutions that contain strong electrolytes and weak electrolytes. • f. Learn the solubility rules for ionic compounds. • g. Use the solubility rules.

  7. 2. Molecular and Ionic Equations • a. Write the molecular equation of a chemical reaction. • b. From the molecular equations for both strong electrolytes and weak electrolytes, determine the complete ionic equation. • c. From the complete ionic equation, write the net ionic equation. • d. Write net ionic equations.

  8. Types of Chemical Reactions • 3. Precipitation Reactions • a. Recognize precipitation (exchange) reactions. • b. Write molecular, complete ionic, and net ionic equations for precipitation reactions. • c. Decide whether a precipitation reaction will occur. • d. Determine the product of a precipitation reaction.

  9. 4. Acid–Base Reactions • a. Understand how an acid–base indicator is used to determine whether a solution is acidic or basic. • b. Define Arrhenius acid and Arrhenius base. • c. Write the chemical equation of an Arrhenius base in aqueous solution. • d. Define Brønsted–Lowry acid and Brønsted–Lowry base. • e. Write the chemical equation of a Brønsted–Lowry base in aqueous solution • f. Write the chemical equation of an acid in aqueous solution using the hydronium ion. • g. Learn the common strong acids and strong bases.

  10. h. Distinguish between a strong acid and a weak acid and the solutions they form. • i. Distinguish between a strong base and a weak base and the solutions they form. • j. Classify acids and bases as strong or weak. • k. Recognize neutralization reactions. • l. Write an equation for a neutralization reaction. • m. Write the reaction equations for a polyprotic acid in aqueous solution • n. Recognize acid–base reactions that lead to gas formation. • o. Write an equation for a reaction with gas formation.

  11. 5. Oxidation–Reduction Reactions • a. Define oxidation–reduction reaction. • b. Learn the oxidation-number rules. • c. Assign oxidation numbers. • d. Write the half-reactions of an oxidation–reduction reaction. • e. Determine the species undergoing oxidation and reduction. • f. Recognize combination reactions, decomposition reactions, displacement reactions, and combustion reactions. • g. Use the activity series to predict when displacement reactions will occur.

  12. 6. Balancing Simple Oxidation–Reduction Equations • a. Balance simple oxidation–reduction reactions by the half-reaction method.

  13. Working with Solutions • 7. Molar Concentration • a. Define molarity or molar concentration of a solution. • b. Calculate the molarity from mass and volume. • c. Use molarity as a conversion factor. • 8. Diluting Solutions • a. Describe what happens to the concentration of a solution when it is diluted. • b. Perform calculations associated with dilution. • c. Describe the process for diluting a solution.

  14. Quantitative Analysis • 9. Gravimetric Analysis • a. Determine the amount of a species by gravimetric analysis. • 10. Volumetric Analysis • a. Calculate the volume of reactant solution needed to perform a titration. • b. Understand how to perform a titration. • c. Calculate the quantity of substance in a titrated solution.

  15. A strong electrolyte dissolves to produce ions. The ions, as moving charges, complete the circuit. When a light bulb is attached to the circuit, it shines.

  16. A strong electrolyte is an electrolyte that exists in solution almost entirely as ions.

  17. A weak electrolyte is an electrolyte that dissolves in water to give a relatively small percentage of ions. As a result, the light bulb shines weakly.

  18. Compounds that dissolve readily are said to be soluble. • Compounds that dissolve very little are said to be insoluble.

  19. 1. Group IA and ammonium compounds are soluble. • 2. Acetates and nitrates are soluble. • 3. Most chlorides, bromides, and iodides are soluble. • Exceptions: AgCl, Hg2Cl2, PbCl2; AgBr, Hg2Br2, HgBr2, PbBr2; AgI, Hg2I2, HgI2, PbI2 • 4. Most sulfates are soluble. • Exceptions: CaSO4, SrSO4, BaSO4, Ag2SO4, Hg2SO4, PbSO4 Solubility Rules

  20. 5. Most carbonates are insoluble. Exceptions: Group IA carbonates and (NH4)2SO4 6. Most phosphates are insoluble. Exceptions: Group IA phosphates and (NH4)3PO4 7. Most sulfides are insoluble. Exceptions: Group IA sulfides and (NH4)2S 8. Most hydroxides are insoluble. Exceptions: Group IA hydroxides, Ca(OH)2, Sr(OH)2, Ba(OH)2

  21. Molecular Equation • A chemical equation in which the reactants and products are written as if they were molecular substances, even though they may actually exist in solution as ions. • State symbols are include: (s), (l), (g), (aq). • For example: • AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq) • Although AgNO3, NaCl, and NaNO3 exist as ions in aqueous solutions, they are written as compounds in the molecular equation.

  22. Complete Ionic Equation • A chemical equation in which strong electrolytes are written as separate ions in the solution. Other reactants and products are written in molecular form. State symbols are included: (s), (l), (g), (aq). • For example: • AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq) • In ionic form: • Ag+(aq) + NO3-(aq) + Na+(aq)Cl-(aq)  • AgCl(s) + Na+(aq) + NO3-(aq)

  23. Spectator Ion • An ion in an ionic equation that does not take part in the reaction. It appears as both a reactant and a product.

  24. Net Ionic Equation • A chemical equation in which spectator ions are omitted. It shows the reaction that actually occurs at the ionic level. • For example: • Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq)  • AgCl(s) + Na+(aq) + NO3-(aq) • In net ionic form: • Ag+(aq) + Cl-(aq)  AgCl(s)

  25. Decide whether the following reaction occurs. If it does, write the molecular, ionic, and net ionic equations. • KBr + MgSO4 • Determine the product formulas: • K+ and SO42- make K2SO4 • Mg2+and Br- make MgBr2 • Determine whether the products are soluble: • K2SO4 is soluble • MgBr2 is soluble • KBr + MgSO4 no reaction

  26. Decide whether the following reaction occurs. If it does, write the molecular, ionic, and net ionic equations. • NaOH + MgCl2 • Determine the product formulas: • Na+ and Cl- make NaCl • Mg2+and OH- make Mg(OH)2 • Determine whether the products are soluble: • NaCl is soluble • Mg(OH)2 is insoluble

  27. Molecular Equation • (Balance the reaction and include state symbols) • 2NaOH(aq) + MgCl2(aq)  • 2NaCl(aq) + Mg(OH)2(s) • Ionic Equation • 2Na+(aq) + 2OH-(aq) + Mg2+(aq) + 2Cl-(aq)  • 2Na+(aq) + 2Cl-(aq) + Mg(OH)2(s) • Net Ionic Equation • 2OH-(aq) + Mg2+(aq)  Mg(OH)2(s)

  28. Decide whether the following reaction occurs. If it does, write the molecular, ionic, and net ionic equations. • K3PO4 + CaCl2 • Determine the product formulas: • K+ and Cl- make KCl • Ca2+and PO43- make Ca3(PO4)2 • Determine whether the products are soluble: • KCl is soluble • Ca3(PO4)2 is insoluble

  29. Molecular Equation • (Balance the reaction and include state symbols) • 2K3PO4(aq) + 3CaCl2(aq)  • 6KCl(aq) + Ca3(PO4)2(s) • Ionic Equation • 6K+(aq) + 2PO43-(aq) + 3Ca2+(aq) + 6Cl-(aq)  • 6K+(aq) + 6Cl-(aq) + Ca3(PO4)2(s) • Net Ionic Equation • 2PO43-(aq) + 3Ca2+(aq)  Ca3(PO4)2(s)

  30. Types of Chemical Reactions • Precipitation reactions: a solid ionic substance forms from the mixture of two solutions of ionic substances. • Acid–base reactions: reactions that involve the transfer of a proton (H+) between reactants • Oxidation–reductionreactions: reactions that involve the transfer of electrons between reactants.

  31. A precipitate is an insoluble solid compound formed during a chemical reaction in solution. • Predicting Precipitation Reactions • 1. Predict the products (exchange of parts). • 2. Determine the state of each product: (s), (l), (g), (aq). • 3. If all products are aqueous (aq), no net reaction occurred.

  32. Arrhenius Acid • A substance that produces hydrogen ions, H+, when it dissolves in water. • Arrhenius Base • A substance that produces hydroxide ions, OH-, when it dissolves in water.

  33. Brønsted–Lowry Acid • The species (molecule or ion) that donates a proton to another species in a proton-transfer reaction • Brønsted–Lowry Base • The species (molecule or ion) that accepts a proton from another species in a proton-transfer reaction

  34. Household Acids and Bases

  35. Acid-Base Indicator • A dye used to distinguish between an acidic and basic solution by means of the color changes it undergoes in these solutions. The sample beakers show a red cabbage indicator in beakers varying in acidity from highly acidic (left) to highly basic (right).

  36. Strong Acid • An acid that ionizes completely in water. It is present entirely as ions; it is a strong electrolyte. • Common strong acids: • HNO3H2SO4HClO4 • HCl HBr HI

  37. Weak Acid • An acid that only partly ionizes in water. It is present primarily as molecules and partly as ions; it is a weak electrolyte. • If an acid is not strong, it is weak.

  38. In Figure A, a solution of HCl (a strong acid) illustrated on a molecular/ionic level, shows the acid as all ions. In Figure B, a solution of HF (a weak acid) also illustrated on a molecular/ionic level, shows mostly molecules with very few ions.

  39. Strong Base • A base that ionizes completely in water. It is present entirely as ions; it is a strong electrolyte. • Common strong bases: • LiOH NaOH KOH • Ca(OH)2 Sr(OH)2 Ba(OH)2

  40. Weak Base • A base that is only partly ionized in water. It is present primarily as molecules and partly as ions; it is a weak electrolyte. These are often nitrogen bases such as NH3: • NH3(aq) + H2O(l) NH4+(aq) + OH-(aq) • If a base is not strong, it is weak.

  41. Classify the following as strong or weak acids or bases: • a. KOH • b. H2S • c. CH3NH2 • d. HClO4 a. KOH is a strong base. b. H2S is a weak acid. c. CH3NH2 is a weak base. d. HClO4 is a strong acid

  42. Polyprotic Acid • An acid that results in two or more acidic hydrogens per molecule • For example: H2SO4, sulfuric acid

  43. Neutralization Reaction • A reaction of an acid and a base that results in an ionic compound (a salt) and possibly water.

  44. Write the molecular, ionic, and net ionic equations for the neutralization of sulfurous acid, H2SO3, by potassium hydroxide, KOH.

  45. Molecular Equation • (Balance the reaction and include state symbols) • H2SO3(aq) + 2KOH(aq)  2H2O(l) + K2SO3(aq) • Ionic Equation • H2SO3(aq) + 2K+(aq) + 2OH-(aq)  • 2H2O(l) + 2K+(aq) + SO32-(aq) • Net Ionic Equation • H2SO3(aq) + 2OH-(aq)  2H2O(l) + SO32-(aq)

  46. Acid-Base Reaction with Gas Formation • Some salts, when treated with an acid, produce a gas. Typically sulfides, sulfites, and carbonates behave in this way producing hydrogen sulfide, sulfur trioxide, and carbon dioxide, respectively. The photo to the right shows baking soda (sodium hydrogen carbonate) reacting with acetic acid in vinegar to give bubbles of carbon dioxide.

  47. Gas-forming acid–base reactions: • Na2S(aq) + 2HCl(aq)  2NaCl(aq) + H2S(g) • Na2CO3(aq) + 2HCl(aq)  • 2NaCl(aq) + H2O(l) + CO2(g) • Na2SO3(aq) + 2HCl(aq)  • 2NaCl(aq) + H2O(l) + SO2(g)

  48. Write the molecular, ionic, and net ionic equations for the reaction of copper(II) carbonate with hydrochloric acid.

  49. Molecular Equation • (Balance the reaction and include state symbols) • CuCO3(s) + 2HCl(aq)  • CuCl2(aq) + H2O(l) + CO2(g) • Ionic Equation • CuCO3(s) + 2H+(aq) + 2Cl-(aq)  • Cu2+(aq) + 2Cl-(aq) + H2O(l) + CO2(g) • Net Ionic Equation • CuCO3(s) + 2H+(aq)  Cu2+(aq) + H2O(l) + CO2(g)

  50. Oxidation Number • For a monatomic ion, the actual charge of the atom or a hypothetical charge assigned to the atom in the substance using simple rules.