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AP Chemistry

AP Chemistry. Chapter 4: Aqueous Reactions & Solution Stoichiometry. Solutions. Homogeneous mixtures Have 2 parts: Solute Smaller quantities Dissolved in the solvent Solvent Larger quantities Dissolves the solute(s) Often water. Electrolytic Properties. Electrolytes

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AP Chemistry

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  1. AP Chemistry Chapter 4: Aqueous Reactions & Solution Stoichiometry

  2. Solutions • Homogeneous mixtures • Have 2 parts: • Solute • Smaller quantities • Dissolved in the solvent • Solvent • Larger quantities • Dissolves the solute(s) • Often water

  3. Electrolytic Properties • Electrolytes • Aqueous solutions that contain ions • Have ionic compounds or strong acids or bases as solute • Will conduct electricity • Nonelectrolytes • Aqueous solutions that do not contain ions • Have molecular compounds as solute

  4. Ionic Compounds in Water Ions attracted to one another to form a crystal lattice Ions dissociate from each other in water

  5. Ionic Compounds in Water Water is polar Each ion will besurrounded bywater molecules: Process is called solvation Water prevents ionsfrom recombiningto form a crystal Negative end of wateris attracted to positive ion

  6. Ionic Compounds in Water

  7. Molecular Compounds in Water Ions not usually formed Molecules remain in tact Strong acids are exceptions

  8. Strong vs. Weak Electrolytes • Strong Electrolytes: • Exist in solution completely (or nearly completely) as ions • Essentially all soluble ionic compounds & some molecular (some acids) fall into this category • Weak Electrolytes: • Exist in solution almost entirely as molecules (a small fraction may be ions) • Molecular compounds fall into this category

  9. Strong vs. Weak Electrolytes Strong Electrolyte Example: HCl(aq) H+(aq) + Cl-(aq) have no tendency to recombine to form HCl • Weak Electrolyte Example: • HC2H3O2(aq) H+(aq) + C2H3O2-(aq) • reaction is significant in both directions • is said to be in chemical equilibrium

  10. Strong vs. Weak Electrolyte IS NOT the same thing as solubility….compounds can be very soluble and be a weak electrolyte and vice versa…. Points To remember Soluble Ionic Compounds are Strong Electrolytes(metal & nonmetal or contain ammonium ion)

  11. Sample Exercise 4.1 The diagram represents an aqueous solution of one of the following compounds: MgCl2, KCl, or K2SO4. Which solution does it best represent? + + 2- 2- + + + 2- 2- + + +

  12. Precipitation Reactions Lead nitrate + Potassium Iodide

  13. Precipitation Reactions • Result in the formation of an insoluble compound • A precipitate is an insoluble compound formed by a reaction in solution

  14. Solubility Rules HAS CAN Halides All Nitrates Except: Hg22+ Ag Pb All Ammonium compounds All C2H3O2-compounds All Alkalimetals Hydroxides and Sulfides are soluble if bonded with Ca2+, Ba2+, or Sr2+, alkali metals or ammonium Sulfates Except: Ba2+, Sr2+, Hg22+, Ag, and Pb

  15. Precipitation Example Will a precipitate form when solutions of Mg(NO3)2 and NaOH are mixed?

  16. Sample Exercise 4.2 Classify the following ionic compounds as soluble or insoluble in water (a) sodium carbonate (b) lead sulfate

  17. Metathesis Reactions • Cation and anion appear to “change partners” • Precipitation reactions conform to this pattern AY + BX  AX + BY

  18. Completing and Balancing Metathesis: • Use chemical formulas of reactants to determine the ions that are present • Write the chemical formulas for the products by combining the cation from one reactant with the anion form the other reactant • Balance the equation * If all products are soluble, we say NO REACTION has occurred.

  19. Sample Exercise 4.3 • Predict the identity of the precipitate that forms when solutions of BaCl2 and K2SO4 are mixed. • Write the balanced equation for this reaction.

  20. Ionic Equations • Molecular equations show the chemical formulas for the reactants and the products • Complete ionic equations show the ions in solution Pb(NO3)2(aq) + 2KI (aq)  PbI2(s) + 2KNO3(aq) Pb2+(aq) +2NO3-(aq) + 2K+(aq) + 2I-(aq) PbI2(s) +2K+(aq) +2NO3-(aq)

  21. Ionic Equations • Spectator ions • appear in both the reactants and the products in the equation • play no part in the reaction • can be omitted (cancelled out)

  22. Ionic Equations • Net ionic equations show only the ions and molecules that are directly involved in the reaction Pb2+(aq) +2NO3-(aq) + 2K+(aq) + 2I-(aq) PbI2(s) +2K+(aq) +2NO3-(aq) NET IONIC EQUATION: Pb2+(aq) + 2I-(aq) PbI2(s)

  23. Reminders If every ion is a spectator, no reaction occurs Charge is conserved: Sum of charges of reactants must equal sum of charges of products

  24. To write net ionic equations: • Write a balanced molecular equation for the reaction • Rewrite to show ions that form in solution when each soluble, strong electrolyte dissociates into its component ions (Only strong electrolytes written in ionic form) • Identify and cancel the spectator ions

  25. Sample Exercise 4.4 Write the net ionic equation for the precipitation reaction that occurs when solutions of calcium chloride and sodium carbonate are mixed.

  26. Acid – Base Reactions • Acids: • Substances that ionize in solutions to from H+ ions • H+ is simply a proton • Acids often called proton donors • Monoprotic acids • Have only one H • Examples: HCl, HNO3 • Diprotic acids • Ionized in 2 steps • Examples:H2SO4 H2SO4(aq)  H+(aq) + HSO4-(aq) HSO4-(aq)  H+(aq) + SO42-(aq)

  27. Acid-Base Reactions • Bases • Substances that accept protons (H+) • Produce hydroxide ions (OH-) when dissolved in water • Include ionic hydroxide compounds • NaOH, MgOH, etc. • NH3(ammonia) is also a base…it accepts H+ from water leaving OH- • NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq) Only a small fraction will form NH4+ ammonia is a weak electrolyte

  28. Acid-Base Reactions • Acids and bases that are strong electrolytes are called strong acids and bases • Reactivity can be determined by strength of acid or on the strength of the anion • Example: HF is a weak acid, but is very reactive because of the F- ion

  29. Strong Acids HClO4 H2SO4 HBr HNO3 HI HCl HClO3

  30. Strong bases Group I Metal Hydroxides Ca, Sr, & Ba Hydroxides

  31. Sample Exercise 4.5 The following diagrams represent aqueous solutions of three acids (HX, HY, HZ) with water molecules omitted for clarity. Rank them from strongest to weakest. + + + - + - + + - - - + - + + - - + - - + + - + - - HY HX HZ

  32. Electrolytic Behavior of Soluble Compounds

  33. Sample Exercise 4.6 Classify each of the following substances as strong electrolyte, weak electrolyte, or nonelectrolyte: CaCl2, HNO3, C2H5OH (ethanol), HCHO2 (formic acid), KOH

  34. Acids vs. Bases

  35. Neutralization Reaction between acidic and basic solutions If the base is a metal hydroxide Products are a salt and water Compound whose cation comes from a base and anion comesfrom an acid

  36. Neutralization Example Molecular Equation: HCl(aq) + NaOH(aq)  NaCl(aq) + H2O(l) H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq)  Na+(aq) + Cl-(aq) + H2O(l) H+(aq) + OH-(aq)  H2O(l) Complete Ionic Equation: Net Ionic Equation:

  37. Sample Exercise 4.7 (a) Write a balanced molecular equation for the reaction between aqueous solutions of acetic acid (HC2H3O2) and barium hydroxide [Ba(OH)2]. (b) Write the net ionic equation for this reaction

  38. Gas Forming Reactions • Acids will form gases when they react with • Sulfides…H2S (g) forms • Carbonates & bicarbonates…CO2 (g) forms • First to form is the unstable carbonic acid (H2CO3) which decomposes to form carbon dioxide gas and water vapor • Example: Alka Seltzer: NaHCO3 reacts with stomach acid • Sodium bicarbonate & sodium carbonate are used to clean up acid spills

  39. Oxidation-Reduction Reactions • Reactions in which electrons are transferred between substances • Oxidation • Loss of electrons • Substance becomes more positive • Reduction • Gain of electrons • Substance becomes more negative

  40. Oxidation – Reduction Reactions • Oxidation of one substance is ALWAYS accompanied by the reduction of another

  41. Oxidation Numbers • Assigned to each atom in a neutral substance or in a charged species • This is the charge of the monatomic ions or the hypothetical charge assigned to the atom assuming the electrons are completely held by one atom or another

  42. Oxidation Numbers • In redox reactions, the ox. #’s change from products to reactants • Oxidation = increase in oxidation number • Reduction = decrease in oxidation number

  43. Assigning Oxidation Numbers…the Rules • Oxidation number is 0 for atoms in elemental form H2 H = 0 S8  S = 0 P4  P = 0 • Oxidation number = the charge of monatomic ions To distinguish between oxidation number and actual charge: Ca2+ oxidation number = +2

  44. Assigning Oxidation Numbers…the Rules • Nonmetals (usually negative ox. nos.) • Oxygen is usually -2 (both ionic & molecular)Exception: in peroxide, each oxygen is -1 • Hydrogen is +1 when bonded with nonmetals and -1 when bonded with metals • Fluorine is ALWAYS -1 • Other halogens are usually -1Exception: When bonded with oxygen, they are + • Sum of oxidation numbers in a compound is always 0: in a polyatomic ion the sum = the charge of the ion

  45. Sample Exercise 4.8 Determine the oxidation number of sulfur in each of the following: (a) H2S (b) S8 (c) SCl2 (d) Na2SO3 (e) SO42-

  46. Oxidation of Metals by Acids & Salts • When a metal is oxidized, it appears to be “eaten away” as it reacts • Called displacement reactions because the ion in solution is displaced (replaced) by the oxidation of an element

  47. Oxidation of Metals by Acids and Salts • Many metals will react with acids to form hydrogen gas and a salt • For example: Mg(s) + 2 HCl(aq)  MgCl2(aq) + H2(g) ox. #’s 0 +1 -1 +2 -1 0 H is reduced Mg is oxidized Cl is a spectator

  48. Oxidation of Metals by Acids & Salts Fe (s) + Ni(NO3)2 (aq)  Fe(NO3)2 (aq) + Ni (s)

  49. One more time: If something is oxidized….Something else mustbe reduced!

  50. Sample Exercise 4.9 Write the balanced molecular and net ionic equations for the reaction of aluminum with hydrobromic acid.

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