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Homework Problems

Homework Problems Chapter 4 Homework Problems: 2, 8, 11, 14, 20, 22, 25, 28, 29, 30, 32, 42, 44, 48, 52, 54, 64, 68, 82, 84, 94, 124, 132, 150. CHAPTER 4 Reactions in Aqueous Solution. Solutions

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Homework Problems

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  1. Homework Problems Chapter 4 Homework Problems: 2, 8, 11, 14, 20, 22, 25, 28, 29, 30, 32, 42, 44, 48, 52, 54, 64, 68, 82, 84, 94, 124, 132, 150

  2. CHAPTER 4 Reactions in Aqueous Solution

  3. Solutions A solution is a homogeneous mixture of two or more pure chemical compounds. Solutions may be solids, liquids, or gases. Solvent - The major component of a solution (if there is one) Solute(s) - The minor component(s) of a solution Example: A solution is prepared by adding a small amount of sodium chloride (NaCl) to water. solvent = water (major component) solute = sodium chloride Aqueous solution - Solution where water is the solvent.

  4. Electrolytes and Nonelectrolytes Many chemical substances dissolve in water. A substance that forms ions when dissolved in water is called an electrolyte, while a substance that does not form ions when dissolved in water is called a nonelectrolyte. Electrolytes can be divided into two categories: Strong electrolyte - Completely dissociates (falls apart) when added to solution, to form ions. Weak electrolyte - Partially dissociates (falls apart) when added to solution, to form ions.

  5. Classification of Compounds as Electrolytes (in Water) We previously divided compounds into two general categories: Ionic compounds - Collections of cations (positive ions) and anions (negative ions) in a crystal structure. Usually form from a metal (or cation group) and a nonmetal (or anion group). Molecular compounds - Exist as individual molecules. Usually form from two or more nonmetals. We may use this as a starting point in classifying compounds as strong electrolytes, weak electrolytes, and nonelectrolytes.

  6. Ionic Compounds Ionic compounds - All ionic compounds are strong electrolytes. Examples: NaCl(s)  Na+(aq) + Cl-(aq) KOH(s)  K+(aq) + OH-(aq) Ca(NO3)2(s)  Ca2+(aq) + 2 NO3-(aq) Note that an insoluble ionic compound is considered a strong electrolyte. This is because even though very little of the compound dissolves in water, whatever does dissolve completely ionizes.

  7. Molecular Compounds Strong acids - Molecular compounds that are strong acids are also strong electrolytes. Examples: HCl(aq)  H+(aq) + Cl-(aq) HNO3(aq)  H+(aq) + NO3-(aq) Weak acids and bases - Molecular compounds that are weak acids or weak bases are also weak electrolytes. Note that an acid forms H+ ion when added to water, and a base forms OH- ion. Examples: HF(aq)  H+(aq) + F-(aq) NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq) The double arrow “” indicates a reaction that proceeds in both the for-ward and reverse direction, establishing equilibrium.

  8. Molecular Compounds (Continued) Other molecular compounds - Molecular compounds that are neither acids nor bases are nonelectrolytes. Examples: C6H12O6(s)  C6H12O6(aq) CH3OH(l)  CH3OH(aq) Note that water (H2O) is classified as a nonelectrolyte, even though a small number of H+ and OH- ions are formed in pure water. H2O(l)  H+(aq) + OH-(aq) For pure water at T = 25 ° there are only ~ 2 x 10-7 M of ions present. This is small enough that water is classified as a nonelectrolyte.

  9. Flow Chart For Classifying Compounds

  10. Precipitation Reactons A precipitation reactions is a reaction that forms a solid (a precipitate) out of soluble reactants. Examples: Pb(NO3)2(aq) + 2 KCl(aq)  PbCl2(s) + 2 KNO3(aq) 3 NaOH(aq) + FeBr3(aq)  Fe(OH)3(s) + 3 NaBr(aq)

  11. Solubility (Water) Solid compounds are divided into two general categories. Soluble compound - a compound that will dissolve in water Insoluble compound - a compound that will not dissolve in water The above classification scheme is an oversimplification. Most “insoluble” compounds will dissolve to a slight extent. For “soluble” compounds there will be a limit as to the amount of compound that will dissolve in a given amount of water. We can also discuss solubility in other solvents.

  12. Why Ionic Compounds Dissolve in Water The forces holding an ionic compound together are the strong electrical attraction that exists between cations and anions. It is therefore somewhat surprising that ionic compounds will dissolve in water. The reason some ionic compounds will dissolve in water is because the water molecules have a partial negative charge on the oxygen atom (-) and partial positive charges on the hydrogen atoms (+), where “” indicates a small positive or negative charge. The reason these partial charges exist will be discussed later in the semester. Because cations and anions can strongly interact with the negative and positive partial charges on several water molecules at the same time, it is possible for the ions in an ionic compound to be pulled out of the solid and go into solution.

  13. Hydration and Solvation The process of surrounding a cation or anion with water molecules to stabilize them in aqueous solution is called hydration. Often these ions will be surrounded by a definite number of water molecules, called a hydration shell. The more general process of surrounding ions or molecules with solvent molecules in a solution is called solvation. Water is unusual because the interaction of ions with water molecules is much stronger than the interactions that occur in most other solvents. hydration of anion hydration of cation

  14. Solubility Rules a) All compounds with group IA cations, or NH4+ (ammonium ion), are soluble. b) All compounds containing NO3- (nitrate ion), ClO4- (perchlorate ion), ClO3- (chlorate ion), or C2H3O2- (acetate ion) are soluble. c) Most compounds containing Cl-, Br-, and I- are soluble. (EXCEPTIONS: Cations Ag+, Pb2+, Hg22+). d) Most compounds containing SO42- (sulfate ion) are soluble. (EXCEPTIONS: Cations Ag+, Pb2+, Hg22+, Ca2+, Sr2+, Ba2+). e) Most compounds containing OH- (hydroxide ion) are insoluble. (EXCEPTIONS: Cations from group 1A, Ca2+, Sr2+, Ba2+). f) Most other compounds are insoluble. (EXCEPTIONS: Compounds covered by rules a-e).

  15. Use of Solubility Rules KClO4 (NH4)2SO4 CuCl2 PbI2 NiSO4 Mg(OH)2 Zn(IO3)2

  16. Use of Solubility Rules KClO4 soluble Rules a and b (NH4)2SO4 soluble Rule a CuCl2 soluble Rule c PbI2 insoluble Rule c NiSO4 soluble Rule d Mg(OH)2 insoluble Rule e Zn(IO3)2 insoluble Rule f

  17. Determining Whether a Precipitation Reaction Will Occur When solutions of two soluble substances mix we will not always form a precipitate. We may use the solubility rules to decide if a precipitate will form. 1) Write the ions formed from the starting compounds. 2) Look at all possible combinations of cations and anions. a) If there is a combination of ions that forms an insoluble compound, a precipitate will form and a precipitation reaction occurs. b) If there is not a combination of ions that forms an insoluble compound, then there will be no precipitation reaction.

  18. Examples: For each of the following mixtures of two solutions will a precipitate form? Note that all of the solutes are themselves soluble. Solution A solution B precipitate(?) 1) NaCl(aq) AgNO3(aq) 2) NaCl(aq) CuBr2(aq)

  19. Solution A solution B precipitate(?) 1) NaCl(aq) AgNO3(aq) yes, AgCl(s) Na+ Cl- Ag+ NO3- NaCl(aq) + AgNO3(aq)  NaNO3(aq) + AgCl(s) 2) NaCl(aq) CuBr2(aq) no Na+ Cl- Cu2+ Br- No reaction.

  20. Net Ionic Equations To this point we have written chemical equations in terms of neutral species (molecules, or formula units). However, in aqueous solution it is often more appropriate to write equations in terms of ions. There are three ways we normally write chemical equations: 1) Molecular equation. All reactants and products are written in terms of their chemical formula. 2) Total ionic equation. Strong electrolytes (strong acids, strong soluble bases, and soluble ionic compounds) are written as ions. 3) Net ionic equation. The total ionic equation after removing spectator ions. A spectator ion is an ion that appears on both the reactant and product side of the equation, and which does not participate directly in the chemical reaction.

  21. Net Ionic Equation (Example) We may show the various ways of writing chemical reactions as follows: Molecular AgNO3(aq) + NaCl(aq)  NaNO3(aq) + AgCl(s) Total Ionic Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq)  Na+(aq) + NO3-(aq) + AgCl(s) Net Ionic Ag+(aq) + Cl-(aq)  AgCl(s)

  22. Finding the Total and New Ionic Equations We find the total and net ionic equations as follows. 1) Begin with a balanced molecular equation. 2) Break apart strong electrolytes into ions to obtain the total ionic equation. 3) Cancel ions that appear on both sides of the equation to obtain the net ionic equation. Example: molecular Pb(NO3)2(aq) + 2 KI(aq)  PbI2(s) + 2 KNO3(aq)

  23. molecular Pb(NO3)2(aq) + 2 KI(aq)  PbI2(s) + 2 KNO3(aq) total ionic Pb2+(aq) + 2 NO3-(aq) + 2 K+(aq) + 2 I-(aq)  PbI2(s) + 2 K+(aq) + 2 NO3-(aq) net ionic Pb2+(aq) + 2 I-(aq)  PbI2(s)

  24. By writing the net ionic equation we are able to focus on the important parts of a chemical reaction and observe similarities in reactions. mol. AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq) total Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq)  AgCl(s) + Na+(aq) + NO3-(aq) net Ag+(aq) + Cl-(aq)  AgCl(s) mol. AgNO3(aq) + KCl(aq)  AgCl(s) + KNO3(aq) total Ag+(aq) + NO3-(aq) + K+(aq) + Cl-(aq)  AgCl(s) + K+(aq) + NO3-(aq) net Ag+(aq) + Cl-(aq)  AgCl(s)

  25. Net Ionic Equation When No Precipitate Forms Writing the net ionic equation makes it clear when precipitation reactions do or do not occur. mol. 2 KBr(aq) + Zn(NO3)2(aq)  2 KNO3(aq) + ZnBr2(aq) K+ Br- Zn2+ NO3- total 2 K+(aq) + 2 Br-(aq) + Zn2+(aq) + 2 NO3-(aq)  2 K+(aq) + 2 NO3-(aq) + Zn2+(aq) + 2 Br-(aq) All the ions are spectator ions, and so no reaction occurs.

  26. Acids and Bases (Arrhenius) Arrhenius definition acid - A substance which, when added to water, forms H+ ion base - A substance which, when added to water, forms OH- ion Examples: HCl(aq)  H+(aq) + Cl-(aq) HF(aq)  H+(aq) + F-(aq) NaOH(s)  Na+(aq) + OH-(aq) NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq) The advantage of the Arrhenius definition is that it is simple and easy to implement. The disadvantage is that it is tied in to a particular solvent (water) and is not a general definition.

  27. Strong and Weak Acids An acid is a substance that forms H+ ion when added to water (Arrhenius definition). A strong acid forms one H+ ion for each acid molecule. A weak acid forms much less than one H+ ion for each acid molecule. Examples: HClO4(aq)  H+(aq) + ClO4-(aq) strong HNO2(aq)  H+(aq) + NO2-(aq) weak There are 7 common strong acids binaryternary HCl HI HClO3 HNO3 HBr HClO4 H2SO4 (1st proton)* All other acids can be assumed to be weak acids. * H2SO4(aq)  H+(aq) + HSO4-(aq)

  28. Strong and Weak Bases A base is a substance that forms OH- ion when added to water (Arrhenius definition). A strong soluble base is a soluble hydroxide compound that completely dissociates when added to water. An insoluble base is an insoluble hydroxide compound. There are also a few substances that act as weak bases in solution. These substances form much less than one OH- ion per base molecule. Examples: KOH(s)  K+(aq) + OH-(aq) strong soluble base Fe(OH)3(s)  no reaction insoluble base NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq) weak base There are 8 common strong soluble bases Group 1A - LiOH, NaOH, KOH, RbOH, CsOH Group 2A - Ca(OH)2, Sr(OH)2, Ba(OH)2 All other hydroxide compounds are insoluble bases.

  29. Acids and Bases (Bronsted) Bronsted definition acid - a proton (H+) donor base - a proton (H+) acceptor HF(aq) + H2O(l)  H3O+(aq) + F-(aq) acid base In the Bronsted theory, in an acid-base reaction an acid donates a proton, while a base accepts a proton. In addition, in Bronsted theory acids form hydronium ion (H3O+ ion) instead of hydrogen ion (H+ ion) when added to water. The Bronsted definition of an acid and a base is more general than the Arrhenius definition.

  30. Examples of Acids (Bronsted) Strong acid HCl(aq) + H2O(l)  H3O+(aq) + Cl-(aq) acid: HCl base: H2O Weak acid HC2H3O2(aq) + H2O(l)  H3O+(aq) + C2H3O2-(aq) acid: HC2H3O2 base: H2O The second reaction goes in both directions.

  31. Examples of Bases (Bronsted) Strong base NaOH(s)  Na+(aq) + OH-(aq) Considered an ionization reaction in Bronsted theory, not an acid-base reaction. Weak base NH3(aq) + H2O(l)  NH4+(aq) + OH-(aq) base: NH3 acid: H2O The second reaction goes in both directions.

  32. Example Chlorous acid (HClO2) is a weak acid. Pyridine (C5H5N) is a weak base. Indicate the behavior of these two substances when added to water, according to Bronsted theory.

  33. Example Chlorous acid (HClO2) is a weak acid. Pyridine (C5H5N) is a weak base. Indicate the behavior of these two substances when added to water, according to Bronsted theory. HClO2(aq) + H2O(l)  H3O+(aq) + ClO2-(aq) C5H5N(aq) + H2O(l)  C5H5NH+(aq) + OH-(aq) Notice that, depending on the reaction, water can act as either a Bronsted base or a Bronsted acid. A substance that can act as either a Bronsted acid or a Bronsted base is called amphoteric.

  34. Polyprotic Acid A polyprotic acid has two or more ionizable protons that can be donated in an acid-base reaction. Monoprotic (one ionizable proton) HCl, HNO2, HC2H3O2 Diprotic (two ionizable protons) H2CO3, H2SO4 Triprotic (three ionizable protons) H3PO3 Example: H2SO4 (sulfuric acid) 1st proton H2SO4(aq) + H2O(l)  H3O+(aq) + HSO4-(aq) 2nd proton HSO4-(aq) + H2O(l)  H3O+(aq) + SO42-(aq) Example: H2CO3 (carbonic acid) 1st proton H2CO3(aq) + H2O(l)  H3O+(aq) + HCO3-(aq) 2nd proton HCO3-(aq) + H2O(l)  H3O+(aq) + CO32-(aq)

  35. Other Amphoteric Species We previously pointed out that depending on the reaction taking place water can act as either a Bronsted acid or a Bronsted base. Other species have this same property. Example: HCO3- ion as acid HCO3-(aq) + OH-(aq)  CO32- (aq) + H2O(l) as base HCO3-(aq) + H3O+(aq)  H2CO3(aq) + H2O(l) Like water, HCO3- ion is classified as amphoteric. Note that the above property of the hydrogen carbonate ion (and other amphoteric species) can be useful in acid-base systems.

  36. Dibasic Base A dibasic base is a strong base that releases two OH- ions per formula unit of base. Example: Ba(OH)2 (barium hydroxide) Ba(OH)2(s)  Ba2+(aq) +2 OH-(aq)

  37. Acid-Base (Neutralization) Reactions An acid-base reactions (or neutralization reaction) is the reaction of an acid with a base. These reactions usually form a salt and water. salt - An ionic compound that can be formed by the reaction of an acid with a base. Note that if either the acid or the base is strong the reaction will go completely from left to right. Examples: 2 HClO4(aq) + Ca(OH)2(aq)  Ca(ClO4)2(aq) + 2 H2O(l) HNO2(aq) + NaOH(aq)  NaNO2(aq) + H2O(l) 2 HBr(aq) + Cu(OH)2(s)  CuBr2(aq) + 2 H2O(l) The last reaction above illustrates why insoluble hydroxide compounds are considered insoluble bases. Note that ionic compounds other than hydroxides (OH-) or oxides (O2-) are classified as salts, as they can be formed from acid-base reactions.

  38. Strong Acid + Strong Base All strong acid + strong base reactions have the same net ionic equation. mol. 2 HClO4(aq) + Sr(OH)2(aq)  Sr(ClO4)2(aq) + 2 H2O() total 2 H+(aq) + 2 ClO4-(aq) + Sr2+(aq) + 2 OH-(aq)  Sr2+(aq) + 2 ClO4-(aq) + 2 H2O() net H+(aq) + OH-(aq)  H2O() mol. HCl(aq) + NaOH(aq)  NaCl(aq) + H2O() total H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq)  Na+(aq) + Cl-(aq) + H2O() net H+(aq) + OH-(aq)  H2O()

  39. Net Ionic Equation for Weak Acid or Base Reaction When the acid or base in an acid base reaction is weak, the net ionic equation will be slightly more complicated. mol. HNO2(aq) + KOH(aq)  KNO2(aq) + H2O() mol. HBr(aq) + NH3(aq)  NH4Br(aq)

  40. mol. HNO2(aq) + KOH(aq)  KNO2(aq) + H2O() total HNO2(aq) + K+(aq) + OH-(aq)  K+(aq) + NO2-(aq) + H2O() net HNO2(aq) + OH-(aq)  NO2-(aq) + H2O() mol. HBr(aq) + NH3(aq)  NH4Br(aq) total H+(aq) + Br-(aq) + NH3(aq)  NH4+(aq) + Br-(aq) net H+(aq) + NH3(aq)  NH4+(aq) Notice that we do not break apart the weak acid or weak base, as they are weak electrolytes, and therefore exist mainly as molecules rather than ions.

  41. Oxidation Number An oxidation number is a number assigned to an atom in a molecule or ion indicating whether that atom is electron rich (negative oxidation number), electron poor (positive oxidation number), or neutral (oxidation number of zero). Oxidation numbers do not indicate charges of atoms, but changes in oxidation number indicate whether atoms have gained or lost electron density. A reaction where some of the oxidation numbers change in going from reactants to products is called an oxidation-reduction (redox) reaction.

  42. Rules For Assigning Oxidation Numbers 1) The sum of the oxidation numbers of a species is equal to the charge of the species. 2) Atoms in elemental forms have an oxidation number of 0. 3) Hydrogen has an oxidation number of +1 when bonded with nonmetals (molecular compounds) and -1 when combined with metals (ionic compounds). 4) Oxygen usually has an oxidation number of -2. Exceptions: Elemental forms, peroxides (H2O2), OF2. H2O2 H +1, O -1 OF2 F -1, O +2 5) Halogens. F in compounds is always -1. Other halogens are often -1, but in compounds with oxygen or other halogens the oxidation number may take on different values. 6) For assigning oxidation numbers in ionic compounds, it is useful to break the compounds up into ions.

  43. Examples: Assign oxidation numbers for each atom in the following substances. H2O Fe(NO3)2 SO42- O3 HClO2 MnO4-

  44. H2O H +1; O -2 Fe(NO3)2 = Fe2+ and NO3- so Fe +2, O -2, and so N +5 SO42- O -2, and so S +6 O3 elemental form, so O = 0 HClO2 = H+ and ClO2- (even though molecular) so H +1, O -2, and so Cl +3 MnO4- O -2, and so Mn +7

  45. Common Oxidation Numbers

  46. Oxidation-Reduction Reactions In an oxidation-reduction (redox) reaction the oxidation numbers of some of the atoms involved in the reaction change when going from reactants to products. 2 Fe2O3(s) + 3 C(s)  4 Fe(s) + 3 CO2(g) +3 -2 0 0 +4 -2 We use the following terms in reference to redox reactions. Oxidation - An increase in the value for the oxidation number. Reduction - A decrease in the value for the oxidation number. There will always be one oxidation and one reduction process per reaction. C 0 to +4, and so is oxidized. Fe +3 to 0, and so is reduced.

  47. Oxidizing and Reducing Agents We call the species that is oxidized in a redox reaction the reducing agent, and the species that is reduced in a redox reaction an oxidizing agent. This seems confusing, but is based on the idea that there must be both an oxidation and a reduction taking place in a redox reaction. C(s) + O2(g)  CO2(g) 0 0 +4 -2 In this reaction carbon is oxidized (from 0 to +4) and so C is a reducing agent, and oxygen is reduced (from 0 to -2) and so O2 is an oxidizing agent. SO2(g) + 2 Fe3+(aq) + 2 H2O(l) 2 Fe2+(aq) + SO42-(aq) + 4 H+(aq) +4 -2 +3 +1 -2 +2 +6 -2 +1 In this reaction S is oxidized (from +4 to +6) and so SO2 is a reducing agent, and Fe is reduced (from +3 to +2) and so Fe3+ is an oxidizing agent.

  48. Activity Series The higher the element in the activity series, the easier it is to oxidize. The lower the anion in the activity series, the easier it is to reduce.

  49. Active and Noble Metals Some metals in the activity series (those at or near the top of the series) are so reactive that they are never found in the metallic state in nature. Such metals are called active metals. The alkali metals and alkaline earth metals are active metals. Metals at or near the bottom of the activity series are not very reactive, and so are often found in the metallic state in nature. Such metals are called noble metals. Examples are gold (Au), platinum (Pt), and mercury (Hg).

  50. Use of the Activity Series The activity series can be used to predict which metals will reduce a particular ion, or which ions will oxidize a particular metal. Example: A solution contains cobalt II ions (Co2+). Which of the following metals would be expected to react if placed in the solution? Cu, Fe, Ni, Zn

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