{"id":2921,"date":"2023-10-03T16:27:43","date_gmt":"2023-10-03T16:27:43","guid":{"rendered":"http:\/\/localhost\/ecole9ja\/?p=2921"},"modified":"2023-10-03T16:31:25","modified_gmt":"2023-10-03T16:31:25","slug":"week-7-ss2-first-term-chemistry-notes","status":"publish","type":"post","link":"https:\/\/ecolebooks.com\/nigeria\/posts\/week-7-ss2-first-term-chemistry-notes\/","title":{"rendered":"Week 7 – SS2 First Term Chemistry Notes"},"content":{"rendered":"

WEEK 7\u00a0\u00a0\u00a0\u00a0
\n<\/strong>ELECTROLYTIC CELLS AND ELECTROLYSIS OF SPECIFIED ELECTROLYTES
\n<\/strong>The passage of electricity through the electrolytes in their molten or dissolved state can cause chemical changes under suitable conditions. For example, the passage of electricity through the acidified water results in the formation of hydrogen and oxygen gases. The process of chemical decomposition of the electrolyte by he passage of electricity through its molten or dissolved state is called electrolysis.
\n\t\t\t<\/strong>\u00a0
\n\u00a0ELECTROL YTIC CELL<\/strong>
\n\t\tThe device in which the process of electrolysis is carried 1ut is called electrolytic cell. It consists of:
\n(i) Electrolytic tank, which is made of some nonconducting materials like glass, wood or bakelite.
\n(ii) Electrolyte in its dissolved state or molten state.
\n(iii) Source of electricity; an electrochemical cell or battery.
\n\"\"
\n\t\t(iv) Two metallic rods, suspended in the electrolyte and connected to the battery through conducting wires. These rods are called electrodes. The electrode connected to the negative terminal of battery is called cathode while the other one which is connected to the positive terminal is called anode. The apparatus used to constitute electrolytic cell has been shown in Fig. 34.2.<\/p>\n

\u00a0Mechanism and Criteria of Product Formation in Electrolytic Cell<\/strong>
\n\t\t\tMechanism and Criteria of Product Formation in Electrolytic Cells
\n\t\t\t<\/strong>The process of electrolysis can be explained on the basis of the theory of ionisation. When an electrolyte is dissolved in water, it splits up into charged particles called ions. The positively charged ions are called cations while the negatively charged ions are called anions. The ions are free to move about in aqueous solution. When electric current is passed through the solution, the ions respond to the applied potential difference and their movement is directed towards the oppositely charged electrodes. The cations move towards the negatively charged electrode while anions move towards the positively charged electrode. The formation of products at the respective electrodes is due to oxidation (loss of electrons) at the anode and reduction (gain of electrons) at the cathode.
\nFor example, when electricity is passed through the molten sodium chloride [NaCl)], sodium is deposited at the cathode while C~ gas is liberated at the anode. The process can be represented as:
\n\"\"
\n\t\tSimilarly, electrolysis of molten lead bromide [PbBr2(l)] produces lead at the cathode and Br2 at the anode.
\nIn case there is a possibility of formation of more than one products at the electrodes, or there is a competition between the liberation of ions at the electrodes, then the products formed depends upon the following factors in general:
\n\u2022 Position of ion\/species in the electromotive series
\n\u2022 Concentration of ion\/species
\n\u2022 Nature of electrodes.
\nOne of the dominant factors which control the criterion of product formation at the electrodes is the values of electrode potentials of the species. Let us understand it as follows:
\n(a) At the Cathode. Cathode involves reduction process at its surface. Therefore, for the different competing reduction processes, the one with higher reduction potential, will preferably take place. For example, during the electrolysis of aqueous solution of sodium chloride there is possibility of following reactions at the cathode:
\n\"\"
\n\t\tThe reduction of water will preferably take place at the cathode because E red of water is higher. Hence, the product of electrolysis of aqueous solution of NaCl at the cathode will be H2 gas instead of Na(s).
\nSimilarly, during electrolysis of aqueous solution of Copper (II) tetraoxosulphate (VI) reduction of Cu2+<\/sup> ions will take place at the cathode in preference to the reduction of 0 molecules because E2<\/sub> is greater than E2<\/sub>
\n\t\t\"\"
\n\t\t(b) At the Anode<\/strong>. Anode involves oxidation process at its surface. Therefore, for different competing oxidation processes, the one with higher oxidation potential (or lower reduction potential) will preferably occur. For example, if we carry out electrolysis of aqueous solution of copper(II) tetraoxosulphate(VI), the competing oxidation processes at the anode are as follows:
\n\"\"
\n\t\tAs oxidation potential of water is higher, the product formed at the anode will be O2 gas instead of S2<\/sub>O8O 2<\/sub>– ion.
\nSimilarly, if the electrolysis of copper sulphate solution is carried out using copper electrodes, then the process occurring at the anode will be oxidation of copper atoms to copper ions instead of oxidation of water because oxidation potential of Cu is higher.
\n\"\"
\n\t\tThus, in such a case copper from anode will go on dissolving into solution as Cu2+<\/sup> ions while Cu2+<\/sup> ions from solution will go on depositing at the cathode as copper atoms.
\nThe above discussion leads us to a general conclusion that for different competing reactions at the electrodes:
\nIt may be noted that in some cases the unexpected results are obtained due to overvoltage*. For example, let us compare the oxidation potentials of CI- ion and water<\/p>\n

\u00a0
\n\u00a0
\n\u00a0\u00a0 ELECTROLYSIS OF SPECIFIED ELECTROLYTES<\/strong>
\n\t\tIn the light of above discussion let us discuss some examples of electrolytic reactions. Before we take up the actual examples, it is quite important to know about inert electrodes and active electrodes.
\nInert electrodes are those electrode which do not gain or lose electrons during the process. Two common examples are carbon (graphite) electrode and platinum electrodes.
\nActive electrodes are those electrodes which themselves participate in the gain or loss of electrons.
\n1. Electrolysis of molten lead (IT) bromide
\nElectrodes: Inert electrodes
\nElectrolyte: Molten lead (II) bromide Molten lead (II) bromide ionises as
\n\"\"
\n\t\tReaction at anode (oxidation of Br-)
\n\"\"
\n\t\t2. Electrolysis of concentration solution sodium chloride (Brine solution)<\/strong>
\n\t\tElectrodes: Inert electrodes
\nElectrolyte: Sodium chloride dissolved in water solvent.
\nNaCl ionises in aqueous solution as
\nNaCl (l) \u00e0 Na +<\/sup> + Cl \u2013<\/sup>
\n\t\tCathode reaction: Reduction of H2<\/sub>O occurs in preference to Na+ ions as Na+ is much higher in electromotive series
\n2H2<\/sub>O (l) –<\/sup> \u00e0 H2<\/sub> (g) + 2OH (aq)
\nAnode reaction: Oxidation of CI- ions occur in preference to H2<\/sub>O because of higher concentration of CI- ions and overvoltage
\n2Cl \u2013<\/sup>(aq) \u00e0 Cl2 (g) + 2e \u2013<\/sup>
\n\t\tThus, Cl2<\/sub> gas is produced at anode and H2<\/sub> gas is liberated at cathode. The pH of solution increases due to increase in the concentration of OH ions in solution.<\/p>\n

\u00a03. Electrolysis of very dilute solution of sodium chloride<\/strong>
\n\t\tElectrode: Inert electrodes.
\nElectrolyte: Very dilute brine (NaCl) solution.
\nNaCl ionises as
\n\"\"
\n\t\tCathode reaction: Reduction of water occurs in preference to Na+ ions
\n4H2<\/sub>O (l) + 4e \u2013 \u00e0 2H2<\/sub> (g) + 4 OH \u2013 (aq)
\nHere, H2 is liberated at cathode whereas O2 is liberated at the anode. The pH of solution does not change due to equal consumption of H+ and OH ions. However, concentration of NaCl gradually increase due to decomposition of water.
\n\u00a0
\n\u00a04. Electrolysis of aqueous solution of copper (II) tetraoxosulphate (VI) solution: using platinum electrodes<\/strong>.
\nElectrodes: Inert electrodes (Pt).
\nElectrolyte: CuSO 4<\/sub> dissolved in water CuSO 4 <\/sub>ionises in aqueous solution as
\n\"\"
\n\t\tCathode reaction: eu2+<\/sup> undergoes reduction in preference to H2<\/sub>O because H2<\/sub>O lies above Cu2+<\/sup> in electrochemical series
\nCu2<\/sub>+ (aq) + 2e \u2013 \u00e0 Cu (s)
\nThus, copper is deposited at cathode and 0 2 is liberated at anode. The pH of the solution decreases gradually due to increase in concentration of H+ ions. Blue colour of the solution gradually fades due to decrease in concentration of eu2+<\/sup> ions.
\n\u00a0
\n\u00a05. Electrolysis of aqueous solution of copper(II) tetraoxosulphate(VI) solution using copper electrodes<\/strong>
\n\t\tElectrodes: Active electrodes.
\nElectrolyte: CuSO4 dissolved in water.
\nCuSO4 ionises in aqueous solution as:
\n\"\"
\n\t\tCathode reaction: Cu2+ ions undergo reduction in preference ~0 molecules
\neu2+<\/sup> (aq) + 2e– <\/sup>\u00e07 Cu (s)
\nAnode reaction: Cu atoms from the copper electrode get oxidised in preference to ~0 because Cu lies below Hp in activity series.
\nCu (s) \u00e0 Cu2+<\/sup> (aq)
\nThus, copper dissolves from anode and get deposited at cathode. Mass of cathode increases and that of anode decreases. Blue colour of the solution does not fade because concentration of Cu2+<\/sup> ions in solution remains unaltered. The pH of the solution also does not change as the electrolysis proceeds.
\nApplications of Electrolysis
\n<\/strong>Some of the important industrial applications of electrolysis are as FOLLOWS.
\n\u00a0
\n\u00a0ELECTROMETALLURGY<\/strong>
\n\t\tMany of the highly reactive metals such as metal of group 1 or group 2 are extracted from their ores by electrolysis of their molten ores. This is because chemical reduction of their compounds .is either chemically not viable or highly expensive. Metals like sodium and magnesium, are manufactured by the electrolysis of their molten chlorides. Pure aluminium is obtained at Valco Industry from a solution of its oxide in molten cryolite. Details of this extraction and that of some other metals is discussed in later units.
\n\u00a0
\n\u00a0ELECTROREFINING OF METALS<\/strong>
\n\t\tRefining of many of the crude metals such as copper, silver; aluminium, etc. is carried out by the process of electrolysis. In this. process, the block of crude metal is made anode while a thin sheet of pure metal is made cathode and the electrolyte is aqueous solution of some salt of the metal. As the electrolysis proceeds metal from anode dissolves whereas it gets deposited in pure form at cathode. The impurities settle down just below anode in the form of anode sludge or anode mud.
\n\u00a0
\n\u00a0MANUFACTURE OF CHEMICAL SUBSTANCES<\/strong>
\n\t\tElectrolysis is used in the manufacture of some important substances such as hydrogen, chlorine, sodium hydroxide, sodium oxochlorate (I) and oxygen.<\/p>\n

\u00a0Electroplating
\n<\/strong>It is another important industrial application of electrolysis. Electroplating is an art of coating a layer of costlier metal like gold, silver, etc. over the cheaper metal like iron. The purpose of electroplating is
\n\u2022 protection of cheaper metal like iron from corrosion
\n\u2022 beautification of articles like, earrings, bangles, tings, parts of wrist watches, etc.
\n\u2022 repair of the broken parts of delicate machinary where welding is not possible.
\nThe principle of electroplating is similar to that of purification of metals by electrolysis. The article to be plated, is thoroughly cleaned with aq H2<\/sub>SO 4<\/sub> and washed with distilled water. It is then made the cathode of the electrolytic cell. The anode is pure sheet of metal to be coated or plated. The electrolyte is a solution of a salt of the metal to be plated. During electrolysis, the metal to be electroplated is transferred from the anode to the cathode.
\nAPPLICATION OF ELECTROLYSIS
\nElectrolysis has several uses in industry. Its main application has been in the fields of manufacture of chemicals and in the purification of metals for which other purification methods prove either too difficult or highly expensive to apply. Some applications of electrolysis are as discussed below:
\n\u00a0
\n\u00a0\u00a0
\n\u00a0<\/p>\n

    \n
  1. Extraction of metals<\/strong>\n\t\t<\/li>\n<\/ol>\n

    Reactive metals exist as compounds, so it is difficult to extract them from their compounds. Extraction of metals such as sodium, potassium, magnesium and calcium necessarily requires an electrolytic process. For example, electrolytic method is used to extract sodium from molten sodium chloride:
    \n2NaCl(l)<\/sub> \u2192 2Na(l)<\/sub> + Cl2(g)<\/sub>
    \n\t\u00a0
    \n\u00a0<\/p>\n

      \n
    1. Purification of metals<\/strong>\n\t\t<\/li>\n<\/ol>\n

      Some metals can be purified by means of electrolysis. This process is used in industry to purify copper, which must be very pure 99.9% for electrical wiring. Copper made by roasting the sulphide ore is about 99.5% pure (so it has an impurity level of 0.5%). This level of impurity cuts down electrical conductivity significantly.
      \nThis is how the electrolytic purification (refining) process is carried out:
      \nThe anode is made of a large block of impure copper. The cathode is a thin sheet of pure copper. The electrolyte is copper (II) sulphate solution.
      \nDuring the refining process, the copper atoms of the impure block become ions (the anode dissolves).
      \n\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0 Cu \u2192 Cu2+<\/sup> + 2e–<\/sup> \u00a0
      \nThe ions from the solution become atoms.
      \n\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0 Cu2+<\/sup> + 2e–<\/sup> \u2192 Cu(s)<\/sub>
      \n\tThey stick onto the cathode. A layer of pure copper builds up on the cathode. As electrolysis takes place, the cathode gains mass as copper is deposited on it. As a result, the cathode gets smaller while the cathode gets bigger as electrolysis proceeds. Eventually the whole cathode dissolves.
      \n\u00a0\u00a0\u00a0 \"\"\/
      \n\tOnly pure copper sticks to the cathode. Most impurities fall to the bottom of the electrolytic cell. They form a solid material (anode sludge<\/strong> or slime<\/strong>) which contains small quantities of precious metals such as silver, gold and platinum. The precious metals recovered from the slime are purified and sold.<\/p>\n

        \n
      1. Industrial manufacture of chemicals from brine<\/strong>\n\t\t<\/li>\n<\/ol>\n

        The electrolysis of a concentrated sodium chloride solution (brine) is an important industrial process. The electrolysis of brine produces hydrogen, chlorine and sodium hydroxide. The electrolysis is carried out in an apparatus called Down’s cell.
        \nThe diagram below shows one of several types of cell that are used. The cell is sometimes referred to as the membrane cell<\/em>.
        \n\"\"\/
        \n\tThe anode is made of titanium and the cathode of nickel. Anion-exchange up the middle lets sodium ions through but keeps the gases apart. The sodium ions move freely to the cathode.
        \nAt the cathode<\/em>: hydrogen bubbles off:
        \n\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0 2H+\u00a0 <\/sup>+ 2e– <\/sup>\u2192 H2(g)<\/sub>
        \n\tAt the anode<\/em>: chlorine bubbles off:
        \n\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0 2Cl–<\/sup> \u2192 Cl2<\/sub> + 2e–<\/sup>
        \n\tNa+<\/sup> and OH–<\/sup> ions are left behind, which means a solution of sodium hydroxide forms. Some is evaporated to a more concentrated solution, and some evaporated completely to give solid sodium hydroxide.
        \nThe products of this electrolysis have a wide range of uses.
        \nHydrogen is used to make ammonia and in the manufacture of margarine.
        \nChlorine is used to kill bacteria in swimming pools and domestic water supplies. It is also used to make plastics.
        \nSodium hydroxide is used to make soap, paper, bleach, and rayon fibres.
        \nThe products also have several other uses not mentioned here.<\/p>\n

          \n
        1. Electroplating <\/strong>\n\t\t<\/li>\n<\/ol>\n

          Electroplating is the coating of a metal with a layer of another metal by means of electrolysis. Electrolysis can be used to coat a thin layer of a less reactive metal onto a more reactive metal. The thin layer of less reactive metal will provide protection from corrosion for the more reactive metal underneath. It may also make the product more attractive.
          \nThe object to be coated should be made the cathode and the coating material should be the electrolyte. The most commonly used metals for electroplating are copper, chromium, silver and tin.
          \nSteel can be electroplated with chromium or tin. This prevents the steel from rusting and gives it a shiny, silver finish. This is also the idea behind chromium-plating articles such as car bumpers, kettles, bath taps, etc. Chromium does not corrode, it is a hard metal that resists scratching and wear, and can also be polished to give an attractive finish.
          \nNickel can be electroplated with silver. This will make nickel more attractive.
          \nThe diagram below shows how a steel jug is electroplated with silver. The jug becomes the cathode of an electrolytic cell. The anode is made of silver. The electrolyte is a solution of a silver compound, for example silver nitrate.
          \n\"\"\/
          \n\tAt the anode<\/em>: The silver dissolves, forming ions in solution:
          \n\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0 \u00a0\u00a0\u00a0\u00a0Ag \u2192 Ag+<\/sup> + e–<\/sup>
          \n\tAt the cathode<\/em>: The silver ions receive electrons, forming a coat of silver on the jug:
          \n\u00a0\u00a0\u00a0\u00a0 \u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0Ag+<\/sup> + e–<\/sup> \u2192 Ag (s)<\/sub>
          \n\tWhen the layer of silver is thick enough, the jug is removed.
          \n\u00a0
          \n\u00a0In general, to electroplate any object with metal M, the set up is:
          \nCathode \u2013 object to be electroplated
          \nAnode \u2013 metal M
          \nElectrolyte \u2013 solution of a soluble compound of M
          \n\u00a0
          \n\u00a0<\/p>\n

            \n
          1. Anodizing aluminium<\/strong>\n\t\t<\/li>\n<\/ol>\n

            Although aluminium is quite reactive, corrosion in air is not a problem. This is because a thin layer of aluminium oxide quickly forms and acts as a barrier to oxygen. The layer can be made thicker, to give a more protection by a process called anodizing.
            \nThe aluminium is used as the anode of a cell. Dilute sulphuric acid is the electrolyte. When this acid is electrolysed, as you learned early, oxygen forms at the anode.
            \n\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0 4OH–<\/sup>\u21922H2<\/sub>O(l) <\/sub>+ O2(g) <\/sub>+ 4e–<\/sup>
            \n\tThe oxygen liberated at the aluminium anode reacts with aluminium to form a protective layer of the oxide on the metal:
            \n\u00a0\u00a0 \u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a04Al(s)<\/sub> + 3O2(g)<\/sub> \u2192 2Al2<\/sub>O3(g)<\/sub>
            \n\tThe anodized aluminum absorbs dyes and pigments easily.It is used for a whole range of things including masts for yachts and windsurfing boards, facings for buildings, and window frames.
            \n\u00a0
            \n\u00a0<\/p>\n

              \n
            1. Producing hydrogen from water<\/strong>\n\t\t<\/li>\n<\/ol>\n

              Hydrogen may become a much more important fuel in the future. It can be produced by the electrolysis of acidified water. Pure water is a very poor conductor of electricity. However, it can be made to decompose if some sulphuric acid is added.
              \nSulphuric acid, when dissolved in water, forms ions:
              \n\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0\u00a0 H2<\/sub>SO4(aq)<\/sub> \u2192 2H+<\/sup>(aq)<\/sub> + SO4<\/sub>2-<\/sup>(aq)<\/sub>
              \n\tWater ionizes thus:\u00a0\u00a0 H2<\/sub>O\u21d4H+<\/sup> + OH–<\/sup>
              \n\tTherefore, a dilute solution of the acid contains H+<\/sup> ions from water and acid, OH–<\/sup> ions from the water and SO4<\/sub>2- <\/sup>ions from the acid.<\/p>\n

              \n\n\n\n\n\n\n
              \u00a0\u00a0\u00a0\u00a0 \u00a0At cathode <\/td>\nAt anode<\/td>\n<\/tr>\n
              Ions<\/em>:\u00a0\u00a0\u00a0\u00a0 \u00a0\u00a0\u00a0\u00a0\u00a0H+<\/sup><\/td>\n\u00a0OH–<\/sup>\u00a0 and \u00a0SO4<\/sub>2-<\/sup><\/td>\n<\/tr>\n
              Reaction<\/em>:\u00a0 2H+ <\/sup>+2e– <\/sup>\u2192H2(g)<\/sub><\/td>\n4OH–<\/sup>\u21922H2<\/sub>O(l)<\/sub>+O2(g)<\/sub>+4e–<\/sup><\/td>\n<\/tr>\n
              Product<\/em>:\u00a0\u00a0 Hydrogen is liberated\u00a0<\/td>\nOxygen is liberated.<\/td>\n<\/tr>\n<\/tbody>\n<\/table>\n<\/div>\n

              A Hofmann voltammeter shown below can be used to keep the gases produced separate. The gas collected above the cathode is hydrogen. Oxygen collects at the anode. The ratio of the volume is approximately 2:1. Effectively this is the electrolysis of water.
              \n\"\"\/
              \n\tIn the future, hydrogen from the electrolysis of water may be piped into homes and used as a fuel for cooking and heating. It will also be used in batteries called fuel cells to power homes and electric cars.
              \nEVALUATION
              \n<\/strong>1. With the aid of a well labelled diagram, describe the electrolysis of dilute sodium tetraoxosulphate(vi) using suitable electrodes
              \n2. Name the electrode<\/p>\n","protected":false},"excerpt":{"rendered":"

              WEEK 7\u00a0\u00a0\u00a0\u00a0 ELECTROLYTIC CELLS AND ELECTROLYSIS OF SPECIFIED ELECTROLYTES The passage of electricity through the…<\/p>\n","protected":false},"author":1,"featured_media":0,"comment_status":"open","ping_status":"open","sticky":false,"template":"","format":"standard","meta":{"footnotes":""},"categories":[1,235],"tags":[],"class_list":["post-2921","post","type-post","status-publish","format-standard","hentry","category-posts","category-first-term-ss2-chemistry"],"aioseo_notices":[],"_links":{"self":[{"href":"https:\/\/ecolebooks.com\/nigeria\/wp-json\/wp\/v2\/posts\/2921","targetHints":{"allow":["GET"]}}],"collection":[{"href":"https:\/\/ecolebooks.com\/nigeria\/wp-json\/wp\/v2\/posts"}],"about":[{"href":"https:\/\/ecolebooks.com\/nigeria\/wp-json\/wp\/v2\/types\/post"}],"author":[{"embeddable":true,"href":"https:\/\/ecolebooks.com\/nigeria\/wp-json\/wp\/v2\/users\/1"}],"replies":[{"embeddable":true,"href":"https:\/\/ecolebooks.com\/nigeria\/wp-json\/wp\/v2\/comments?post=2921"}],"version-history":[{"count":1,"href":"https:\/\/ecolebooks.com\/nigeria\/wp-json\/wp\/v2\/posts\/2921\/revisions"}],"predecessor-version":[{"id":2922,"href":"https:\/\/ecolebooks.com\/nigeria\/wp-json\/wp\/v2\/posts\/2921\/revisions\/2922"}],"wp:attachment":[{"href":"https:\/\/ecolebooks.com\/nigeria\/wp-json\/wp\/v2\/media?parent=2921"}],"wp:term":[{"taxonomy":"category","embeddable":true,"href":"https:\/\/ecolebooks.com\/nigeria\/wp-json\/wp\/v2\/categories?post=2921"},{"taxonomy":"post_tag","embeddable":true,"href":"https:\/\/ecolebooks.com\/nigeria\/wp-json\/wp\/v2\/tags?post=2921"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}