{"id":254,"date":"2017-06-20T20:38:28","date_gmt":"2017-06-20T20:38:28","guid":{"rendered":"https:\/\/courses.lumenlearning.com\/chemistry2labs\/?post_type=chapter&p=254"},"modified":"2017-06-20T21:26:09","modified_gmt":"2017-06-20T21:26:09","slug":"lab-11-introduction","status":"publish","type":"chapter","link":"https:\/\/courses.lumenlearning.com\/chemistry2labs\/chapter\/lab-11-introduction\/","title":{"raw":"Lab 11 Introduction","rendered":"Lab 11 Introduction"},"content":{"raw":"
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\r\n\r\nChesapeake Campus \u2013 Chemistry 112 Laboratory\r\n\r\n#11- Electrochemistry\r\n\r\n<\/div>\r\n<\/div>\r\n
\r\n
\r\n\r\nObjectives\r\n\r\n<\/div>\r\n<\/div>\r\n
\r\n
\r\n
    \r\n \t
  • Calculate the reduction potential for a galvanic cell.<\/li>\r\n \t
  • Establish an activity series of metal ions based on their relative reduction potentials.Introduction\r\n\r\nZn(s) + Cu2+\u2192Zn2+ + Cu(s)<\/li>\r\n<\/ul>\r\n<\/div>\r\n<\/div>\r\n
    \r\n
    \r\n
    \r\n
    \r\n\r\nIt is physically impossible to measure the potential difference between a piece of metal and the solution in which it is immersed. We can, however, measure the difference between the potentials of two electrodes that\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n
    \r\n
    \r\n
    \r\n\r\ndip into the same solution, or more usefully, are in two different solutions. In the latter case, each electrode-solution pair constitutes an oxidation-reduction half-cell, and we are measuring the sum of the two half-cell potentials.\r\n\r\nThis arrangement is called a galvanic cell. A typical cell might consist of two pieces of metal, one zinc and the other copper, each immersed each in a solution containing a dissolved salt of the corresponding metal. The two solutions are separated by a porous barrier that prevents them from rapidly mixing but allows ions to diffuse through. This prevents a build-up of charge on either side of the cell and prolongs the life of the cell.\r\n\r\nIf we connect the zinc and copper by means of a metallic conductor, the excess electrons that remain when Zn2+ ions emerge from the zinc in the left cell would be able to flow through the external circuit and\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n
    \r\n
    \r\n
    \r\n\r\ninto the right electrode, where they could be delivered to the Cu2+ ions which become \"discharged\", that is, converted into Cu atoms at the surface of the copper electrode. (The mass of the zinc electrode is reduced as zinc becomes oxidized while the mass of the copper electrode increases). The net reaction is the oxidation of zinc by copper(II) ions:\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n
    \r\n
    \r\n\r\nBut unlike other reactions which occur in a single reaction vessel, the oxidation and reduction steps (half reactions) take place in separate locations:\r\n\r\n<\/div>\r\n<\/div>\r\n
    \r\n
    \r\n
    \r\n\r\nleft electrode: Zn(s) \u2192 Zn2+ + 2e\u2013 oxidation\r\n\r\n<\/div>\r\n<\/div>\r\n
    \r\n
    \r\n\r\nright electrode: Cu2+ + 2e\u2013\u2192 Cu(s) reduction\r\n\r\n<\/div>\r\n<\/div>\r\n
    \r\n
    \r\n\r\nElectrochemical Cells Allow Measurement and Control of a Redox Reaction\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n
    \r\n
    \r\n
    \r\n
    \r\n
    \r\n\r\nThe reaction can be started and stopped by connecting or disconnecting the two electrodes. If we place a variable resistance in the circuit, we can even control the rate of the net cell reaction by simply turning a knob. By connecting a battery or other source of current to the two electrodes, we can force the reaction to proceed in its non-spontaneous, or reverse direction. By placing an ammeter in the external circuit, we can measure the amount of electric charge that passes through the electrodes, and thus the number of moles of reactants that get transformed into products in the cell reaction.\r\n\r\nElectric charge q is measured in coulombs. The amount of charge carried by one mole of electrons is known as the Faraday, which we denote by F. Careful experiments have determined that 1 F = 96467 C. For most purposes, you can simply use 96,500 Coulombs as the value of the faraday. When we measure electric current, we are measuring the rate at which electric charge is transported through the circuit. A current of one ampere corresponds to the flow of one coulomb per second.\r\n\r\nCharge Transport within the Cell\r\n\r\nFor the cell to operate, not only must there be an external electrical circuit between the two electrodes, but the two electrolytes (the solutions) must be in contact. The need for this can be understood by considering what\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n
    \r\n
    \r\n
    \r\n\r\nwould happen if the two solutions were physically separated. Positive charge (in the form of Zn2+) is added to the electrolyte in the left compartment, and removed (as Cu2+) from the right side, causing the solution in contact with the zinc to acquire a net positive charge, while a net negative charge would build up in the solution on the copper side of the cell. These violations of electroneutrality would make it more difficult (require more work) to introduce additional Zn2+ ions into the positively-charged electrolyte or for electrons to flow into right compartment where they are needed to reduce the Cu2+ ions, thus effectively stopping the reaction after only a chemically insignificant amount has taken place.\r\n\r\nIn order to sustain the cell reaction, the charge carried by the\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n
    \r\n
    \r\n
    \r\n\r\nelectrons through the external circuit must be accompanied by a compensating transport of ions between the\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n
    \r\n
    \r\n
    \r\n\r\ntwo cells. This means that we must provide a path for ions to move directly from one cell to the other. This ionic transport involves not only the electroactive species Cu2+ and Zn2+, but also the counter ions, which in this example are nitrate, NO3-. Thus an excess of Cu2+ in the left compartment could be alleviated by the drift of these ions into the right side, or equally well by diffusion of nitrate ions to the left. More detailed studies reveal that both processes occur, and that the relative amounts of charge carried through the solution by positive and negative ions depends on their relative mobilities, which express the velocity with which the ions are able to make their way through the solution. Since negative ions tend to be larger than positive ions, the latter tend to have higher mobilities and carry the larger fraction of\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n
    \r\n
    \r\n
    \r\n\r\ncharge.\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n
    <\/div>\r\n<\/div>\r\n
    \u00a0In the simplest cells, the barrier between the two solutions can be a porous membrane, but for precise measurements, a more complicated arrangement, known as a salt bridge, is used. The salt bridge consists of an intermediate compartment filled with a concentrated solution of KCl and fitted with porous barriers at each end. The purpose of the salt bridge is to minimize the natural potential difference, known as the junction potential, that develops (as mentioned in the previous section) when any two phases (such as the two solutions) are in contact. This potential difference would combine with the two half-cell potentials so as introduce a degree of uncertainty into any measurement of the cell potential. With the salt bridge, we have two liquid junction potentials instead of one, but they tend to cancel each other out.<\/div>\r\n<\/div>\r\n
    \r\n
    \r\n
    \r\n
    \r\n\r\nCell description conventions\r\n\r\nIn order to make it easier to describe a given electrochemical cell, a special symbolic notation has been adopted. In this notation the cell we described above would be\r\n\r\nZn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)\r\nThere are several other conventions relating to cell notation and nomenclature that you are expected to know:\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n
    \r\n
    \r\n
    \r\n
      \r\n \t
    • \uf0b7 \u00a0The anode is where oxidation occurs, and the cathode is the site of reduction. In an actual cell, the identity of the electrodes depends on the direction in which the net cell reaction is occurring.<\/li>\r\n \t
    • \uf0b7 \u00a0If electrons flow from the left electrode to the right electrode (as depicted in the above cell notation) when the cell operates in its spontaneous direction, the potential of the right electrode will be higher than that of the left, and the cell potential will be positive.<\/li>\r\n \t
    • \uf0b7 \u00a0\"Conventional current flow\" is from positive to negative, which is opposite to the direction of the electron flow. This means that if the electrons are flowing from the left electrode to the right, a galvanometer placed in the external circuit would indicate a current flow from right to left.<\/li>\r\n<\/ul>\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n
      \r\n
      \r\n\r\nCell Potential\r\n\r\nThe cell potential is the driving force for electrons to move from one electrode to the other and is expressed in units of volts. The cell potential, Ecell, can be measured by placing a voltmeter between the two electrodes of voltaic cell. The cell potential is a measure of how easily electrons flow from the anode to the cathode. Equation 1 shows how to calculate the cell potential.\r\n\r\n\ud835\udc38\ud835\udc50\ud835\udc52\ud835\udc59\ud835\udc59 = \ud835\udc38\ud835\udc50\ud835\udc4e\ud835\udc61h\ud835\udc5c\ud835\udc51\ud835\udc52 \u2212 \ud835\udc38\ud835\udc4e\ud835\udc5b\ud835\udc5c\ud835\udc51\ud835\udc52 Equation 1\r\n\r\nA positive value for Ecell indicates that the transfer of electrons occurs spontaneously and the cell is capable of doing work. A negative value for Ecell indicates that energy must be supplied in order for the cell to function.\r\n\r\nStandard Reduction Potentials\r\n\r\nElectrical potentials are measured relative to one another. The convention in chemistry is to measure all reduction potentials of standard solutions relative to that of the Standard Hydrogen Electrode (SHE). The standard reduction potential, E0red, of the SHE is set as zero.\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>\r\n
      \r\n
      \r\n
      \r\n\r\n2H+ + 2e- \u2192 H2 E0red = 0.0V\r\n\r\nWe can calculate the reduction potentials of any other chemical species in relation to this standard. Species that are more easily reduced than H+ are considered to be good oxidizing agents and have positive E0red values. Species that have more negative E0red are less easily reduced than H+ are good reducing agents. For example consider the reduction potentials of Zn2+ and Cu2+ listed below.\r\n\r\nCu2+ + 2e- \u2192 Cu E0red = 0.34 V\r\n\r\nZn2+ + 2e- \u2192 Zn E0red = -0.76 V\r\n\r\nCopper has a more positive reduction potential and is therefore more easily reduced than Zn2+. In a voltaic cell the cathode half of the cell will always contain the species that is more easily reduced. The overall cell potential is calculated from Equation 1 as:\r\n\r\nE0cell = E0cathode - E0anode\r\n\r\nSince the potentials are always listed as reduction potentials, it is necessary to subtract the anode. Since the anode is where oxidation happens (the opposite of reduction) the value must be subtracted in the equation to account for the reverse reaction occurring. Therefore:\r\n\r\nE0cell = E0redCu2+ - E0redZn2+\r\nE0cell = -0.34 V \u2013 (-0.76 V) = -1.10 V\r\n\r\nExperiment\r\n\r\nIn this experiment the reduction potentials for a series of metal cations, M2+, will be determined relative to one other M2+ cation. Voltaic cells will be constructed using spot plate wells to house M2+ solutions, pieces of string soaked in KNO3 as salt bridges and pieces of metal as electrodes. A logger pro serves as a voltmeter.\r\n\r\n<\/div>\r\n<\/div>\r\n<\/div>","rendered":"
      \n
      \n
      \n

      Chesapeake Campus \u2013 Chemistry 112 Laboratory<\/p>\n

      #11- Electrochemistry<\/p>\n<\/div>\n<\/div>\n

      \n
      \n

      Objectives<\/p>\n<\/div>\n<\/div>\n

      \n
      \n
        \n
      • Calculate the reduction potential for a galvanic cell.<\/li>\n
      • Establish an activity series of metal ions based on their relative reduction potentials.Introduction\n

        Zn(s) + Cu2+\u2192Zn2+ + Cu(s)<\/li>\n<\/ul>\n<\/div>\n<\/div>\n

        \n
        \n
        \n
        \n

        It is physically impossible to measure the potential difference between a piece of metal and the solution in which it is immersed. We can, however, measure the difference between the potentials of two electrodes that<\/p>\n<\/div>\n<\/div>\n<\/div>\n

        \n
        \n
        \n

        dip into the same solution, or more usefully, are in two different solutions. In the latter case, each electrode-solution pair constitutes an oxidation-reduction half-cell, and we are measuring the sum of the two half-cell potentials.<\/p>\n

        This arrangement is called a galvanic cell. A typical cell might consist of two pieces of metal, one zinc and the other copper, each immersed each in a solution containing a dissolved salt of the corresponding metal. The two solutions are separated by a porous barrier that prevents them from rapidly mixing but allows ions to diffuse through. This prevents a build-up of charge on either side of the cell and prolongs the life of the cell.<\/p>\n

        If we connect the zinc and copper by means of a metallic conductor, the excess electrons that remain when Zn2+ ions emerge from the zinc in the left cell would be able to flow through the external circuit and<\/p>\n<\/div>\n<\/div>\n<\/div>\n

        \n
        \n
        \n

        into the right electrode, where they could be delivered to the Cu2+ ions which become “discharged”, that is, converted into Cu atoms at the surface of the copper electrode. (The mass of the zinc electrode is reduced as zinc becomes oxidized while the mass of the copper electrode increases). The net reaction is the oxidation of zinc by copper(II) ions:<\/p>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n

        \n
        \n

        But unlike other reactions which occur in a single reaction vessel, the oxidation and reduction steps (half reactions) take place in separate locations:<\/p>\n<\/div>\n<\/div>\n

        \n
        \n
        \n

        left electrode: Zn(s) \u2192 Zn2+ + 2e\u2013 oxidation<\/p>\n<\/div>\n<\/div>\n

        \n
        \n

        right electrode: Cu2+ + 2e\u2013\u2192 Cu(s) reduction<\/p>\n<\/div>\n<\/div>\n

        \n
        \n

        Electrochemical Cells Allow Measurement and Control of a Redox Reaction<\/p>\n<\/div>\n<\/div>\n<\/div>\n<\/div>\n

        \n
        \n
        \n
        \n
        \n

        The reaction can be started and stopped by connecting or disconnecting the two electrodes. If we place a variable resistance in the circuit, we can even control the rate of the net cell reaction by simply turning a knob. By connecting a battery or other source of current to the two electrodes, we can force the reaction to proceed in its non-spontaneous, or reverse direction. By placing an ammeter in the external circuit, we can measure the amount of electric charge that passes through the electrodes, and thus the number of moles of reactants that get transformed into products in the cell reaction.<\/p>\n

        Electric charge q is measured in coulombs. The amount of charge carried by one mole of electrons is known as the Faraday, which we denote by F. Careful experiments have determined that 1 F = 96467 C. For most purposes, you can simply use 96,500 Coulombs as the value of the faraday. When we measure electric current, we are measuring the rate at which electric charge is transported through the circuit. A current of one ampere corresponds to the flow of one coulomb per second.<\/p>\n

        Charge Transport within the Cell<\/p>\n

        For the cell to operate, not only must there be an external electrical circuit between the two electrodes, but the two electrolytes (the solutions) must be in contact. The need for this can be understood by considering what<\/p>\n<\/div>\n<\/div>\n<\/div>\n

        \n
        \n
        \n

        would happen if the two solutions were physically separated. Positive charge (in the form of Zn2+) is added to the electrolyte in the left compartment, and removed (as Cu2+) from the right side, causing the solution in contact with the zinc to acquire a net positive charge, while a net negative charge would build up in the solution on the copper side of the cell. These violations of electroneutrality would make it more difficult (require more work) to introduce additional Zn2+ ions into the positively-charged electrolyte or for electrons to flow into right compartment where they are needed to reduce the Cu2+ ions, thus effectively stopping the reaction after only a chemically insignificant amount has taken place.<\/p>\n

        In order to sustain the cell reaction, the charge carried by the<\/p>\n<\/div>\n<\/div>\n<\/div>\n

        \n
        \n
        \n

        electrons through the external circuit must be accompanied by a compensating transport of ions between the<\/p>\n<\/div>\n<\/div>\n<\/div>\n

        \n
        \n
        \n

        two cells. This means that we must provide a path for ions to move directly from one cell to the other. This ionic transport involves not only the electroactive species Cu2+ and Zn2+, but also the counter ions, which in this example are nitrate, NO3-. Thus an excess of Cu2+ in the left compartment could be alleviated by the drift of these ions into the right side, or equally well by diffusion of nitrate ions to the left. More detailed studies reveal that both processes occur, and that the relative amounts of charge carried through the solution by positive and negative ions depends on their relative mobilities, which express the velocity with which the ions are able to make their way through the solution. Since negative ions tend to be larger than positive ions, the latter tend to have higher mobilities and carry the larger fraction of<\/p>\n<\/div>\n<\/div>\n<\/div>\n

        \n
        \n
        \n

        charge.<\/p>\n<\/div>\n<\/div>\n<\/div>\n

        <\/div>\n<\/div>\n
        \u00a0In the simplest cells, the barrier between the two solutions can be a porous membrane, but for precise measurements, a more complicated arrangement, known as a salt bridge, is used. The salt bridge consists of an intermediate compartment filled with a concentrated solution of KCl and fitted with porous barriers at each end. The purpose of the salt bridge is to minimize the natural potential difference, known as the junction potential, that develops (as mentioned in the previous section) when any two phases (such as the two solutions) are in contact. This potential difference would combine with the two half-cell potentials so as introduce a degree of uncertainty into any measurement of the cell potential. With the salt bridge, we have two liquid junction potentials instead of one, but they tend to cancel each other out.<\/div>\n<\/div>\n
        \n
        \n
        \n
        \n

        Cell description conventions<\/p>\n

        In order to make it easier to describe a given electrochemical cell, a special symbolic notation has been adopted. In this notation the cell we described above would be<\/p>\n

        Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s)
        \nThere are several other conventions relating to cell notation and nomenclature that you are expected to know:<\/p>\n<\/div>\n<\/div>\n<\/div>\n

        \n
        \n
        \n
          \n
        • \uf0b7 \u00a0The anode is where oxidation occurs, and the cathode is the site of reduction. In an actual cell, the identity of the electrodes depends on the direction in which the net cell reaction is occurring.<\/li>\n
        • \uf0b7 \u00a0If electrons flow from the left electrode to the right electrode (as depicted in the above cell notation) when the cell operates in its spontaneous direction, the potential of the right electrode will be higher than that of the left, and the cell potential will be positive.<\/li>\n
        • \uf0b7 \u00a0“Conventional current flow” is from positive to negative, which is opposite to the direction of the electron flow. This means that if the electrons are flowing from the left electrode to the right, a galvanometer placed in the external circuit would indicate a current flow from right to left.<\/li>\n<\/ul>\n<\/div>\n<\/div>\n<\/div>\n
          \n
          \n

          Cell Potential<\/p>\n

          The cell potential is the driving force for electrons to move from one electrode to the other and is expressed in units of volts. The cell potential, Ecell, can be measured by placing a voltmeter between the two electrodes of voltaic cell. The cell potential is a measure of how easily electrons flow from the anode to the cathode. Equation 1 shows how to calculate the cell potential.<\/p>\n

          \ud835\udc38\ud835\udc50\ud835\udc52\ud835\udc59\ud835\udc59 = \ud835\udc38\ud835\udc50\ud835\udc4e\ud835\udc61h\ud835\udc5c\ud835\udc51\ud835\udc52 \u2212 \ud835\udc38\ud835\udc4e\ud835\udc5b\ud835\udc5c\ud835\udc51\ud835\udc52 Equation 1<\/p>\n

          A positive value for Ecell indicates that the transfer of electrons occurs spontaneously and the cell is capable of doing work. A negative value for Ecell indicates that energy must be supplied in order for the cell to function.<\/p>\n

          Standard Reduction Potentials<\/p>\n

          Electrical potentials are measured relative to one another. The convention in chemistry is to measure all reduction potentials of standard solutions relative to that of the Standard Hydrogen Electrode (SHE). The standard reduction potential, E0red, of the SHE is set as zero.<\/p>\n<\/div>\n<\/div>\n<\/div>\n

          \n
          \n
          \n

          2H+ + 2e- \u2192 H2 E0red = 0.0V<\/p>\n

          We can calculate the reduction potentials of any other chemical species in relation to this standard. Species that are more easily reduced than H+ are considered to be good oxidizing agents and have positive E0red values. Species that have more negative E0red are less easily reduced than H+ are good reducing agents. For example consider the reduction potentials of Zn2+ and Cu2+ listed below.<\/p>\n

          Cu2+ + 2e- \u2192 Cu E0red = 0.34 V<\/p>\n

          Zn2+ + 2e- \u2192 Zn E0red = -0.76 V<\/p>\n

          Copper has a more positive reduction potential and is therefore more easily reduced than Zn2+. In a voltaic cell the cathode half of the cell will always contain the species that is more easily reduced. The overall cell potential is calculated from Equation 1 as:<\/p>\n

          E0cell = E0cathode – E0anode<\/p>\n

          Since the potentials are always listed as reduction potentials, it is necessary to subtract the anode. Since the anode is where oxidation happens (the opposite of reduction) the value must be subtracted in the equation to account for the reverse reaction occurring. Therefore:<\/p>\n

          E0cell = E0redCu2+ – E0redZn2+
          \nE0cell = -0.34 V \u2013 (-0.76 V) = -1.10 V<\/p>\n

          Experiment<\/p>\n

          In this experiment the reduction potentials for a series of metal cations, M2+, will be determined relative to one other M2+ cation. Voltaic cells will be constructed using spot plate wells to house M2+ solutions, pieces of string soaked in KNO3 as salt bridges and pieces of metal as electrodes. A logger pro serves as a voltmeter.<\/p>\n<\/div>\n<\/div>\n<\/div>\n","protected":false},"author":23588,"menu_order":2,"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-254","chapter","type-chapter","status-publish","hentry"],"part":29,"_links":{"self":[{"href":"https:\/\/courses.lumenlearning.com\/chemistry2labs\/wp-json\/pressbooks\/v2\/chapters\/254","targetHints":{"allow":["GET"]}}],"collection":[{"href":"https:\/\/courses.lumenlearning.com\/chemistry2labs\/wp-json\/pressbooks\/v2\/chapters"}],"about":[{"href":"https:\/\/courses.lumenlearning.com\/chemistry2labs\/wp-json\/wp\/v2\/types\/chapter"}],"author":[{"embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/chemistry2labs\/wp-json\/wp\/v2\/users\/23588"}],"version-history":[{"count":2,"href":"https:\/\/courses.lumenlearning.com\/chemistry2labs\/wp-json\/pressbooks\/v2\/chapters\/254\/revisions"}],"predecessor-version":[{"id":299,"href":"https:\/\/courses.lumenlearning.com\/chemistry2labs\/wp-json\/pressbooks\/v2\/chapters\/254\/revisions\/299"}],"part":[{"href":"https:\/\/courses.lumenlearning.com\/chemistry2labs\/wp-json\/pressbooks\/v2\/parts\/29"}],"metadata":[{"href":"https:\/\/courses.lumenlearning.com\/chemistry2labs\/wp-json\/pressbooks\/v2\/chapters\/254\/metadata\/"}],"wp:attachment":[{"href":"https:\/\/courses.lumenlearning.com\/chemistry2labs\/wp-json\/wp\/v2\/media?parent=254"}],"wp:term":[{"taxonomy":"chapter-type","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/chemistry2labs\/wp-json\/pressbooks\/v2\/chapter-type?post=254"},{"taxonomy":"contributor","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/chemistry2labs\/wp-json\/wp\/v2\/contributor?post=254"},{"taxonomy":"license","embeddable":true,"href":"https:\/\/courses.lumenlearning.com\/chemistry2labs\/wp-json\/wp\/v2\/license?post=254"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}