{"id":6807,"date":"2020-12-14T23:58:25","date_gmt":"2020-12-14T23:58:25","guid":{"rendered":"https:\/\/courses.lumenlearning.com\/chemistryformajors\/?post_type=chapter&p=6807"},"modified":"2021-01-14T19:23:34","modified_gmt":"2021-01-14T19:23:34","slug":"assignment-electrochemistry","status":"publish","type":"chapter","link":"https:\/\/courses.lumenlearning.com\/chemistryformajors\/chapter\/assignment-electrochemistry\/","title":{"raw":"Assignment: Electrochemistry","rendered":"Assignment: Electrochemistry"},"content":{"raw":"
    \r\n \t
  1. \r\n
      \r\n \t
    1. Why do we balance the charge in redox reactions, but not in other chemical equations?<\/li>\r\n \t
    2. Balance each reaction below, and write a cell schematic representing the reaction as it would occur in a galvanic cell.\r\n
        \r\n \t
      1. Al(s)<\/em> + Zr4+<\/sup>(aq)<\/em> \u27f6 Al3+<\/sup>(aq)<\/em> + Zr(s)<\/em><\/li>\r\n \t
      2. Ag+<\/sup>(aq)<\/em> + NO(g)<\/em> \u27f6 Ag(s)<\/em> + NO3<\/sub>\u2013<\/sup>(aq)<\/em>\u00a0 (acidic solution)<\/li>\r\n \t
      3. CO3<\/sub>2\u2212<\/sup>(aq)<\/em> + Mg(s)<\/em> \u27f6 C(s)<\/em> + Mg(OH)2<\/sub>(s)\u00a0<\/em> (basic solution)<\/li>\r\n \t
      4. BrO3<\/sub>\u2013<\/sup>(aq)<\/em> + MnO2<\/sub>(s)<\/em> \u27f6 Br\u2212<\/sup>(aq)<\/em> + MnO4<\/sub>\u2212<\/sup>(aq)<\/em>\u00a0 (basic solution)<\/li>\r\n<\/ol>\r\n<\/li>\r\n \t
      5. Explain why a salt bridge is needed in galvanic cell reactions.<\/li>\r\n \t
      6. If a metal electrode loses mass in a galvanic cell reaction, was it the cathode or the anode? Explain your reasoning.<\/li>\r\n \t
      7. Write the balanced cell reaction for the cell schematic below, calculate the standard cell potential, and note whether the reaction is spontaneous under standard state conditions.<\/li>\r\n<\/ol>\r\nCu(s)\u2502Cu2+(aq)\u2551Ag1+(aq)\u2502Ag(s)\r\n
          \r\n \t
        1. Determine the cell reaction and standard cell potential at 25 \u00b0C for a cell made from a cathode half-cell consisting of a silver electrode in 1 M silver nitrate solution and an anode half-cell consisting of a zinc electrode in 1 M zinc nitrate. Is the reaction spontaneous at standard conditions?<\/li>\r\n<\/ol>\r\n<\/li>\r\n<\/ol>\r\n \r\n\r\nCalculate the Gibbs free energy change (G) for the following chemical reaction (The reaction occurs at 72 \u00b0F, the change in heat (H) = 18,070 cal, and the change in entropy (S) = 90 cal\/K.):\r\n

          ATP \u2192 ADP + Pi<\/p>\r\nCalculate the Gibbs free energy change (G) for the following chemical reaction (The reaction occurs at 72 \u00b0F, the change in heat (H) = 4104 cal, and the change in entropy (S) = 2.4 cal\/K.):\r\n

          glutamate + NH3<\/sub>\u2192 glutamine + H2<\/sub>O<\/p>\r\n\r\n

            \r\n \t
          1. Will either of these reactions occur spontaneously?<\/li>\r\n \t
          2. How does the Gibbs free energy in each of the two reactions change if the temperature were raised to normal body temperature (98.6 \u00b0F)?<\/li>\r\n \t
          3. Does putting the reaction at normal body temperature change the spontaneity of either of the reactions? Explain.<\/li>\r\n \t
          4. In an electrolysis reaction, water is being decomposed into oxygen gas and hydrogen gas.<\/li>\r\n<\/ol>\r\n \r\n\r\nLitmus can be used as an indicator in the water. In the following reaction, the litmus turns red in test tube #1 and blue in test tube #2.\r\n

            Test tube 1: 2 H2<\/sub>O(l) \u2192 O2<\/sub>(g) + 4H+<\/sup>(aq) + 4e-<\/p>\r\n

            Test tube 2: 4 H2<\/sub>O(l) + 4e- \u2192 2H2<\/sub>(g) + 4OH\u2013(aq)<\/p>\r\n\r\n

              \r\n \t
            1. \r\n
                \r\n \t
              1. Find the oxidation numbers for the species in each of the two test tubes, then identify which species is being oxidized and which is being reduced?<\/li>\r\n \t
              2. Which test tube contains the anode?<\/li>\r\n \t
              3. Explain why the litmus turns blue in test tube #2<\/li>\r\n<\/ol>\r\n<\/li>\r\n<\/ol>","rendered":"
                  \n
                1. \n
                    \n
                  1. Why do we balance the charge in redox reactions, but not in other chemical equations?<\/li>\n
                  2. Balance each reaction below, and write a cell schematic representing the reaction as it would occur in a galvanic cell.\n
                      \n
                    1. Al(s)<\/em> + Zr4+<\/sup>(aq)<\/em> \u27f6 Al3+<\/sup>(aq)<\/em> + Zr(s)<\/em><\/li>\n
                    2. Ag+<\/sup>(aq)<\/em> + NO(g)<\/em> \u27f6 Ag(s)<\/em> + NO3<\/sub>\u2013<\/sup>(aq)<\/em>\u00a0 (acidic solution)<\/li>\n
                    3. CO3<\/sub>2\u2212<\/sup>(aq)<\/em> + Mg(s)<\/em> \u27f6 C(s)<\/em> + Mg(OH)2<\/sub>(s)\u00a0<\/em> (basic solution)<\/li>\n
                    4. BrO3<\/sub>\u2013<\/sup>(aq)<\/em> + MnO2<\/sub>(s)<\/em> \u27f6 Br\u2212<\/sup>(aq)<\/em> + MnO4<\/sub>\u2212<\/sup>(aq)<\/em>\u00a0 (basic solution)<\/li>\n<\/ol>\n<\/li>\n
                    5. Explain why a salt bridge is needed in galvanic cell reactions.<\/li>\n
                    6. If a metal electrode loses mass in a galvanic cell reaction, was it the cathode or the anode? Explain your reasoning.<\/li>\n
                    7. Write the balanced cell reaction for the cell schematic below, calculate the standard cell potential, and note whether the reaction is spontaneous under standard state conditions.<\/li>\n<\/ol>\n

                      Cu(s)\u2502Cu2+(aq)\u2551Ag1+(aq)\u2502Ag(s)<\/p>\n

                        \n
                      1. Determine the cell reaction and standard cell potential at 25 \u00b0C for a cell made from a cathode half-cell consisting of a silver electrode in 1 M silver nitrate solution and an anode half-cell consisting of a zinc electrode in 1 M zinc nitrate. Is the reaction spontaneous at standard conditions?<\/li>\n<\/ol>\n<\/li>\n<\/ol>\n

                         <\/p>\n

                        Calculate the Gibbs free energy change (G) for the following chemical reaction (The reaction occurs at 72 \u00b0F, the change in heat (H) = 18,070 cal, and the change in entropy (S) = 90 cal\/K.):<\/p>\n

                        ATP \u2192 ADP + Pi<\/p>\n

                        Calculate the Gibbs free energy change (G) for the following chemical reaction (The reaction occurs at 72 \u00b0F, the change in heat (H) = 4104 cal, and the change in entropy (S) = 2.4 cal\/K.):<\/p>\n

                        glutamate + NH3<\/sub>\u2192 glutamine + H2<\/sub>O<\/p>\n

                          \n
                        1. Will either of these reactions occur spontaneously?<\/li>\n
                        2. How does the Gibbs free energy in each of the two reactions change if the temperature were raised to normal body temperature (98.6 \u00b0F)?<\/li>\n
                        3. Does putting the reaction at normal body temperature change the spontaneity of either of the reactions? Explain.<\/li>\n
                        4. In an electrolysis reaction, water is being decomposed into oxygen gas and hydrogen gas.<\/li>\n<\/ol>\n

                           <\/p>\n

                          Litmus can be used as an indicator in the water. In the following reaction, the litmus turns red in test tube #1 and blue in test tube #2.<\/p>\n

                          Test tube 1: 2 H2<\/sub>O(l) \u2192 O2<\/sub>(g) + 4H+<\/sup>(aq) + 4e-<\/p>\n

                          Test tube 2: 4 H2<\/sub>O(l) + 4e- \u2192 2H2<\/sub>(g) + 4OH\u2013(aq)<\/p>\n

                            \n
                          1. \n
                              \n
                            1. Find the oxidation numbers for the species in each of the two test tubes, then identify which species is being oxidized and which is being reduced?<\/li>\n
                            2. Which test tube contains the anode?<\/li>\n
                            3. Explain why the litmus turns blue in test tube #2<\/li>\n<\/ol>\n<\/li>\n<\/ol>\n","protected":false},"author":17533,"menu_order":11,"template":"","meta":{"_candela_citation":"[]","CANDELA_OUTCOMES_GUID":"","pb_show_title":"on","pb_short_title":"","pb_subtitle":"","pb_authors":[],"pb_section_license":""},"chapter-type":[],"contributor":[],"license":[],"class_list":["post-6807","chapter","type-chapter","status-publish","hentry"],"part":2970,"_links":{"self":[{"href":"https:\/\/courses.lumenlearning.com\/chemistryformajors\/wp-json\/pressbooks\/v2\/chapters\/6807","targetHints":{"allow":["GET"]}}],"collection":[{"href":"https:\/\/courses.lumenlearning.com\/chemistryformajors\/wp-json\/pressbooks\/v2\/chapters"}],"about":[{"href":"https:\/\/courses.lumenlearning.com\/chemistryformajors\/wp-json\/wp\/v2\/types\/chapter"}],"author":[{"embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/chemistryformajors\/wp-json\/wp\/v2\/users\/17533"}],"version-history":[{"count":4,"href":"https:\/\/courses.lumenlearning.com\/chemistryformajors\/wp-json\/pressbooks\/v2\/chapters\/6807\/revisions"}],"predecessor-version":[{"id":7991,"href":"https:\/\/courses.lumenlearning.com\/chemistryformajors\/wp-json\/pressbooks\/v2\/chapters\/6807\/revisions\/7991"}],"part":[{"href":"https:\/\/courses.lumenlearning.com\/chemistryformajors\/wp-json\/pressbooks\/v2\/parts\/2970"}],"metadata":[{"href":"https:\/\/courses.lumenlearning.com\/chemistryformajors\/wp-json\/pressbooks\/v2\/chapters\/6807\/metadata\/"}],"wp:attachment":[{"href":"https:\/\/courses.lumenlearning.com\/chemistryformajors\/wp-json\/wp\/v2\/media?parent=6807"}],"wp:term":[{"taxonomy":"chapter-type","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/chemistryformajors\/wp-json\/pressbooks\/v2\/chapter-type?post=6807"},{"taxonomy":"contributor","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/chemistryformajors\/wp-json\/wp\/v2\/contributor?post=6807"},{"taxonomy":"license","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/chemistryformajors\/wp-json\/wp\/v2\/license?post=6807"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}