{"id":1226,"date":"2025-11-21T14:47:41","date_gmt":"2025-11-21T19:47:41","guid":{"rendered":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/?post_type=chapter&p=1226"},"modified":"2026-03-20T17:21:25","modified_gmt":"2026-03-20T21:21:25","slug":"5-effect-of-temperature-on-reactions","status":"publish","type":"chapter","link":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/chapter\/5-effect-of-temperature-on-reactions\/","title":{"raw":"5. Effect of Temperature on Reactions","rendered":"5. Effect of Temperature on Reactions"},"content":{"raw":"The relationship between \u0394G\u00b0 for a reaction and K is given by:\r\n\r\n\\[\r\n\\Delta G^\\circ = -RT \\ln K\r\n\\tag{443}\r\n\\]\r\n\r\nThe temperature dependence of the equilibrium constant can be found by differentiating:\r\n\r\n\\[\r\n\\frac{d\\ln K}{dT}\r\n= -\\,\\frac{d}{dT}\\!\\left(\\frac{\\Delta G^\\circ}{RT}\\right)\r\n\\tag{444}\r\n\\]\r\n\r\nFrom thermodynamics we get the expression,\r\n\r\n\\[\r\n\\frac{d}{dT}\\left(\\frac{\\Delta G^\\circ}{T}\\right)\r\n= -\\frac{\\Delta H^\\circ}{T^2}\r\n\\tag{445}\r\n\\]\r\n\r\nTherefore,\r\n\r\n\\[\r\n\\frac{d\\ln K}{dT}\r\n= \\frac{\\Delta H^\\circ}{RT^2}\r\n\\tag{446}\r\n\\]\r\n\r\nStated another way, this gives:\r\n\r\n\\[\r\n\\frac{d\\ln K}{d(1\/T)}\r\n= -\\frac{\\Delta H^\\circ}{R}\r\n\\tag{447}\r\n\\]\r\n\r\nwhere, again, \u0394G\u00b0 and \u0394H\u00b0 refer to free energy and enthalpy changes, respectively, for a reaction, under standard conditions. The last expression can be integrated if the temperature dependence of \u0394H\u00b0 is known. For reactions involving neutral (uncharged) species \u0394H\u00b0 is almost constant over a modest temperature range. Then we obtain, after integration,\r\n\r\n\\[\r\n\\ln K_2\r\n= \\ln K_1\r\n- \\Delta H^\\circ\\left(\\frac{1}{T_2} - \\frac{1}{T_1}\\right)\r\n\\tag{448}\r\n\\]\r\n\r\nwhere K2<\/sub> is the equilibrium constant at temperature T2<\/sub>, and K1<\/sub> the equilibrium constant at T1<\/sub>. Temperature is on the Kelvin scale. If K and \u0394H\u00b0 are known at one temperature, then K can be estimate at another temperature.\r\n\r\nHowever, for reactions involving ionic species \u0394H\u00b0 for the reaction depends strongly on temperature and the simple equation above cannot be used. This is severely limiting for reactions of interest in hydrometallurgy. Certainly we still know the qualitative effects of \u0394H\u00b0, as per Le Chatelier\u2019s principle. There are methods available to determine \u0394G\u00b0 as a function of temperature and hence equilibrium constants, for reactions involving ions.\r\n\r\nThere are also empirical correlations available for \u0394G\u00b0 as a function of temperature over limited temperature ranges. These take the form:\r\n\r\n\\[\r\n\\Delta G^\\circ\r\n= a + bT\r\n= -RT \\ln K\r\n\\tag{449}\r\n\\]\r\n\r\nwhere a and b are empirical constants, determined by experiment. These are valid up to about 75\u00b0C. A few examples are listed in the Table below.\r\n\r\n[table id=91 \/]\r\n\r\nWhere reactions can be added to obtain a new one, the \u0394G\u00b0 of the new reaction is the sum of the \u0394G\u00b0 values for the individual reactions. For example at 75\u00b0C,\r\n\r\n\\[\r\n\\ce{CuO(s) + 2H^+ = Cu^{2+} + H2O(l)}\r\n\\qquad\r\n\\Delta G^\\circ = -38{,}269\\ \\text{J\u00b7mol}^{-1}\r\n\\tag{450}\r\n\\]\r\n\r\n\\[\r\n\\ce{2H2O(l) = 2H^+ + 2OH^-}\r\n\\qquad\r\n\\Delta G^\\circ = 2 \\times 83{,}936\\ \\text{J\u00b7mol}^{-1}\r\n\\tag{451}\r\n\\]\r\n\r\n\\[\r\n\\ce{CuO(s) + H2O(l) = Cu^{2+} + 2OH^-}\r\n\\tag{452}\r\n\\]\r\n\r\n\\[\r\n\\Delta G^\\circ\r\n= -38{,}269 + 2(83{,}936)\r\n= 129{,}603\\ \\text{J\u00b7mol}^{-1}\r\n\\tag{453}\r\n\\]\r\n\r\n\\[\r\nK_{sp}\r\n= \\exp\\!\\left(\r\n-\\frac{129{,}603\\ \\text{J\u00b7mol}^{-1}}\r\n{8.314\\ \\text{J\u00b7mol}^{-1}\\text{K}^{-1} \\cdot 348.15\\ \\text{K}}\r\n\\right)\r\n= 3.6 \\times 10^{-20}\r\n\\tag{454}\r\n\\]\r\n\r\nThe case of bisulfate is important. Sulfuric acid is the most important acid in hydrometallurgy. Bisulfate is a weak acid, and it becomes weaker as temperature increases. At 0\u00b0C pKa2<\/sub> = 1.64, and at 75\u00b0C it is 2.54.\r\n\r\nPrecipitation of oxides and hydroxides is an important consideration in hydrometallurgy. Sometimes this is pursued as a way to recover a metal in this form, or remove an unwanted one (e.g. Fe+3<\/sup>), and sometimes care must be taken to avoid it. Take the case of Cu+2<\/sup> and CuO again. For the reaction:\r\n\r\n\\[\r\n\\ce{CuO(s) + 2H^+ = Cu^{2+} + H2O}\r\n\\tag{455}\r\n\\]\r\n\r\n\\[\r\nK = \\frac{[\\ce{Cu^{2+}}]}{[\\ce{H^+}]^{2}}\r\n\\tag{456}\r\n\\]\r\n\r\n\\[\r\n\\log K\r\n= \\log[\\ce{Cu^{2+}}]\r\n- 2\\log[\\ce{H^+}]\r\n\\tag{457}\r\n\\]\r\n\r\n\\[\r\n\\log K\r\n= \\log[\\ce{Cu^{2+}}]\r\n+ 2\\,\\text{pH}\r\n\\tag{458}\r\n\\]\r\n\r\n\\[\r\n\\log[\\ce{Cu^{2+}}]\r\n= \\log K - 2\\,\\text{pH}\r\n\\tag{459}\r\n\\]\r\n\r\nThe last equation is similar in form to that for the pH of a buffer comprised of a metal oxide and soluble metal salt; the two should not be confused! The solubility of CuO as a function of pH and as a function of temperature can be determined.\r\n
(a) What is the solubility of CuO (expressed as [Cu+2<\/sup>] in M) at 25\u00b0C and pH 3.5?<\/h5>\r\n\\[\r\nK = \\exp\\!\\left[-\\frac{a + bT}{RT}\\right]\r\n\\tag{460}\r\n\\]\r\n\r\nAt 25\u00b0C,\r\n\r\n\\[\r\nK\r\n= \\exp\\!\\left[\r\n-\\frac{-63{,}760 + 73.22 \\times 298.15}\r\n{8.314 \\times 298.15}\r\n\\right]\r\n= 2.22 \\times 10^{7}\r\n\\tag{461}\r\n\\]\r\n\r\n\\[\r\n\\log[\\ce{Cu^{2+}}]\r\n= \\log(2.22 \\times 10^{7})\r\n- 2(3.5)\r\n= 0.346\r\n\\quad\\Rightarrow\\quad\r\n[\\ce{Cu^{2+}}] = 2.22\\ \\text{M}\r\n\\tag{462}\r\n\\]\r\n\r\n(Using sulfuric acid we would obtain <2.2 M due to the solubility limit for CuSO\u00b75H2<\/sub>O, but is not important here.)\r\n
(b) How much of this Cu+2 is precipitated as CuO by raising the pH to 4.5 (decreases the H+ concentration by a factor of only 10)?<\/h5>\r\nAt pH 4.5, log[Cu+2<\/sup>] = -1.654\u00a0 (0.0222 M). The solubility of Cu+2<\/sup> has dropped by:\r\n\r\n\\[\r\n\\frac{2.22 - 0.0222}{2.22} \\times 100\r\n= 99\\%\r\n\\tag{463}\r\n\\]\r\n
(c) If a solution containing 2.22 M Cu+2 at 25\u00b0C is raised to 100\u00b0C, how much of the Cu+2 will precipitate?<\/h5>\r\n(The \u0394H\u00b0 for the reaction is -63.76 J\/mol at 25\u00b0C. Since this is negative, raising the temperature should lower the equilibrium constant, as per Le Chatelier\u2019s principle.)\r\n\r\nK at this temperature can be found using the a + bT relationship; this is a bit of an extrapolation, since the temperature is a bit higher than the range over which the relationship is known to be valid.\r\n\r\n\\[\r\nK = 1.26 \\times 10^{5}\r\n\\ (< 2.22 \\times 10^{7})\r\n\\tag{464}\r\n\\]\r\n\r\n\\[\r\n\\log[\\ce{Cu^{2+}}]\\ \\text{at pH 3.5}\r\n= -1.90\r\n\\quad\\Rightarrow\\quad\r\n[\\ce{Cu^{2+}}] = 0.0126\\ \\text{M}\r\n\\tag{465}\r\n\\]\r\n\r\nThe change in concentration is -99.4%, i.e. 99.4% of the Cu+2<\/sup> precipitates. Thus the combination of temperature and pH effects can have a strong influence on solubility. Such effects become even more pronounced in autoclave leaching vessels where temperatures can be as high as 300\u00b0C.\r\n\r\nAs an illustration, iminodiacetic acid ion exchange resins contain the functional group -NH-(CH2<\/sub>CO2<\/sub>H)2<\/sub> bonded to a polystyrene polymer. The acetic groups can ionize (weak acid) to form a dianion group. This plus the trivalent nitrogen atom complexes to Cu+2<\/sup> quite strongly and is a viable means of recovering cupric ion from dilute aqueous solution streams. And while the resin operates best at pH \u2265 4, precipitation of Cu+2<\/sup> can limit its use to pH 4 or less.","rendered":"

The relationship between \u0394G\u00b0 for a reaction and K is given by:<\/p>\n

\\[
\n\\Delta G^\\circ = -RT \\ln K
\n\\tag{443}
\n\\]<\/p>\n

The temperature dependence of the equilibrium constant can be found by differentiating:<\/p>\n

\\[
\n\\frac{d\\ln K}{dT}
\n= -\\,\\frac{d}{dT}\\!\\left(\\frac{\\Delta G^\\circ}{RT}\\right)
\n\\tag{444}
\n\\]<\/p>\n

From thermodynamics we get the expression,<\/p>\n

\\[
\n\\frac{d}{dT}\\left(\\frac{\\Delta G^\\circ}{T}\\right)
\n= -\\frac{\\Delta H^\\circ}{T^2}
\n\\tag{445}
\n\\]<\/p>\n

Therefore,<\/p>\n

\\[
\n\\frac{d\\ln K}{dT}
\n= \\frac{\\Delta H^\\circ}{RT^2}
\n\\tag{446}
\n\\]<\/p>\n

Stated another way, this gives:<\/p>\n

\\[
\n\\frac{d\\ln K}{d(1\/T)}
\n= -\\frac{\\Delta H^\\circ}{R}
\n\\tag{447}
\n\\]<\/p>\n

where, again, \u0394G\u00b0 and \u0394H\u00b0 refer to free energy and enthalpy changes, respectively, for a reaction, under standard conditions. The last expression can be integrated if the temperature dependence of \u0394H\u00b0 is known. For reactions involving neutral (uncharged) species \u0394H\u00b0 is almost constant over a modest temperature range. Then we obtain, after integration,<\/p>\n

\\[
\n\\ln K_2
\n= \\ln K_1
\n– \\Delta H^\\circ\\left(\\frac{1}{T_2} – \\frac{1}{T_1}\\right)
\n\\tag{448}
\n\\]<\/p>\n

where K2<\/sub> is the equilibrium constant at temperature T2<\/sub>, and K1<\/sub> the equilibrium constant at T1<\/sub>. Temperature is on the Kelvin scale. If K and \u0394H\u00b0 are known at one temperature, then K can be estimate at another temperature.<\/p>\n

However, for reactions involving ionic species \u0394H\u00b0 for the reaction depends strongly on temperature and the simple equation above cannot be used. This is severely limiting for reactions of interest in hydrometallurgy. Certainly we still know the qualitative effects of \u0394H\u00b0, as per Le Chatelier\u2019s principle. There are methods available to determine \u0394G\u00b0 as a function of temperature and hence equilibrium constants, for reactions involving ions.<\/p>\n

There are also empirical correlations available for \u0394G\u00b0 as a function of temperature over limited temperature ranges. These take the form:<\/p>\n

\\[
\n\\Delta G^\\circ
\n= a + bT
\n= -RT \\ln K
\n\\tag{449}
\n\\]<\/p>\n

where a and b are empirical constants, determined by experiment. These are valid up to about 75\u00b0C. A few examples are listed in the Table below.<\/p>\n\n\n\n\n\n\n\n\n\n\n\n\n\n\n\n
Table 5.1 - Variation of \u0394G\u00b0 with temperature for some reactions of hydrometallurgical interest. T is in Kelvin.<\/th>\n<\/tr>\n
Reaction<\/th>\nConstants for \u0394G\u00b0 = a + bT<\/th>\n<\/tr>\n
a (J\/mol)<\/th>\nb (J\/mol K)<\/th>\n<\/tr>\n<\/thead>\n
H2<\/sub>O = H+<\/sup> + OH-<\/sup><\/td>\n55840<\/td>\n80.70<\/td>\n<\/tr>\n
0.5Al2<\/sub>O3 (s)<\/sub> + 2.5H2<\/sub>O = [Al(OH)4<\/sub>]-<\/sup> + H+<\/sup><\/td>\n62420<\/td>\n-69.96<\/td>\n<\/tr>\n
Al(OH)3 (s)<\/sub> + OH-<\/sup> = [Al(OH)4<\/sub>]-<\/sup><\/td>\n72830<\/td>\n1441.2<\/td>\n<\/tr>\n
PbCl2 (s)<\/sub> = Pb+2<\/sup> + 2Cl-<\/sup><\/td>\n22980<\/td>\n13.85<\/td>\n<\/tr>\n
CuO(s)<\/sub> + 2H+<\/sup> = Cu+2<\/sup> + H2<\/sub>O(l)<\/sub><\/td>\n-63760<\/td>\n73.22<\/td>\n<\/tr>\n
CuS(s)<\/sub> + 2H+<\/sup> = Cu+2<\/sup> + H2<\/sub>S(aq)<\/sub><\/td>\n74070<\/td>\n42.25<\/td>\n<\/tr>\n
NiS(s)<\/sub> + 2H+<\/sup> = Ni+2<\/sup> + H2<\/sub>S(aq)<\/sub><\/td>\n-10700<\/td>\n90.17<\/td>\n<\/tr>\n
Ni(OH)2<\/sub>(s)<\/sub> = Ni+<\/sup> + 2OH-<\/sup><\/td>\n14000<\/td>\n265<\/td>\n<\/tr>\n
HSO4<\/sub>-<\/sup> = H+<\/sup> + SO4<\/sub>2-<\/sup><\/td>\n-21930<\/td>\n111.7<\/td>\n<\/tr>\n<\/tbody>\n<\/table>\n

<\/p>\n

Where reactions can be added to obtain a new one, the \u0394G\u00b0 of the new reaction is the sum of the \u0394G\u00b0 values for the individual reactions. For example at 75\u00b0C,<\/p>\n

\\[
\n\\ce{CuO(s) + 2H^+ = Cu^{2+} + H2O(l)}
\n\\qquad
\n\\Delta G^\\circ = -38{,}269\\ \\text{J\u00b7mol}^{-1}
\n\\tag{450}
\n\\]<\/p>\n

\\[
\n\\ce{2H2O(l) = 2H^+ + 2OH^-}
\n\\qquad
\n\\Delta G^\\circ = 2 \\times 83{,}936\\ \\text{J\u00b7mol}^{-1}
\n\\tag{451}
\n\\]<\/p>\n

\\[
\n\\ce{CuO(s) + H2O(l) = Cu^{2+} + 2OH^-}
\n\\tag{452}
\n\\]<\/p>\n

\\[
\n\\Delta G^\\circ
\n= -38{,}269 + 2(83{,}936)
\n= 129{,}603\\ \\text{J\u00b7mol}^{-1}
\n\\tag{453}
\n\\]<\/p>\n

\\[
\nK_{sp}
\n= \\exp\\!\\left(
\n-\\frac{129{,}603\\ \\text{J\u00b7mol}^{-1}}
\n{8.314\\ \\text{J\u00b7mol}^{-1}\\text{K}^{-1} \\cdot 348.15\\ \\text{K}}
\n\\right)
\n= 3.6 \\times 10^{-20}
\n\\tag{454}
\n\\]<\/p>\n

The case of bisulfate is important. Sulfuric acid is the most important acid in hydrometallurgy. Bisulfate is a weak acid, and it becomes weaker as temperature increases. At 0\u00b0C pKa2<\/sub> = 1.64, and at 75\u00b0C it is 2.54.<\/p>\n

Precipitation of oxides and hydroxides is an important consideration in hydrometallurgy. Sometimes this is pursued as a way to recover a metal in this form, or remove an unwanted one (e.g. Fe+3<\/sup>), and sometimes care must be taken to avoid it. Take the case of Cu+2<\/sup> and CuO again. For the reaction:<\/p>\n

\\[
\n\\ce{CuO(s) + 2H^+ = Cu^{2+} + H2O}
\n\\tag{455}
\n\\]<\/p>\n

\\[
\nK = \\frac{[\\ce{Cu^{2+}}]}{[\\ce{H^+}]^{2}}
\n\\tag{456}
\n\\]<\/p>\n

\\[
\n\\log K
\n= \\log[\\ce{Cu^{2+}}]
\n– 2\\log[\\ce{H^+}]
\n\\tag{457}
\n\\]<\/p>\n

\\[
\n\\log K
\n= \\log[\\ce{Cu^{2+}}]
\n+ 2\\,\\text{pH}
\n\\tag{458}
\n\\]<\/p>\n

\\[
\n\\log[\\ce{Cu^{2+}}]
\n= \\log K – 2\\,\\text{pH}
\n\\tag{459}
\n\\]<\/p>\n

The last equation is similar in form to that for the pH of a buffer comprised of a metal oxide and soluble metal salt; the two should not be confused! The solubility of CuO as a function of pH and as a function of temperature can be determined.<\/p>\n

(a) What is the solubility of CuO (expressed as [Cu+2<\/sup>] in M) at 25\u00b0C and pH 3.5?<\/h5>\n

\\[
\nK = \\exp\\!\\left[-\\frac{a + bT}{RT}\\right]
\n\\tag{460}
\n\\]<\/p>\n

At 25\u00b0C,<\/p>\n

\\[
\nK
\n= \\exp\\!\\left[
\n-\\frac{-63{,}760 + 73.22 \\times 298.15}
\n{8.314 \\times 298.15}
\n\\right]
\n= 2.22 \\times 10^{7}
\n\\tag{461}
\n\\]<\/p>\n

\\[
\n\\log[\\ce{Cu^{2+}}]
\n= \\log(2.22 \\times 10^{7})
\n– 2(3.5)
\n= 0.346
\n\\quad\\Rightarrow\\quad
\n[\\ce{Cu^{2+}}] = 2.22\\ \\text{M}
\n\\tag{462}
\n\\]<\/p>\n

(Using sulfuric acid we would obtain <2.2 M due to the solubility limit for CuSO\u00b75H2<\/sub>O, but is not important here.)<\/p>\n

(b) How much of this Cu+2 is precipitated as CuO by raising the pH to 4.5 (decreases the H+ concentration by a factor of only 10)?<\/h5>\n

At pH 4.5, log[Cu+2<\/sup>] = -1.654\u00a0 (0.0222 M). The solubility of Cu+2<\/sup> has dropped by:<\/p>\n

\\[
\n\\frac{2.22 – 0.0222}{2.22} \\times 100
\n= 99\\%
\n\\tag{463}
\n\\]<\/p>\n

(c) If a solution containing 2.22 M Cu+2 at 25\u00b0C is raised to 100\u00b0C, how much of the Cu+2 will precipitate?<\/h5>\n

(The \u0394H\u00b0 for the reaction is -63.76 J\/mol at 25\u00b0C. Since this is negative, raising the temperature should lower the equilibrium constant, as per Le Chatelier\u2019s principle.)<\/p>\n

K at this temperature can be found using the a + bT relationship; this is a bit of an extrapolation, since the temperature is a bit higher than the range over which the relationship is known to be valid.<\/p>\n

\\[
\nK = 1.26 \\times 10^{5}
\n\\ (< 2.22 \\times 10^{7})
\n\\tag{464}
\n\\]<\/p>\n

\\[
\n\\log[\\ce{Cu^{2+}}]\\ \\text{at pH 3.5}
\n= -1.90
\n\\quad\\Rightarrow\\quad
\n[\\ce{Cu^{2+}}] = 0.0126\\ \\text{M}
\n\\tag{465}
\n\\]<\/p>\n

The change in concentration is -99.4%, i.e. 99.4% of the Cu+2<\/sup> precipitates. Thus the combination of temperature and pH effects can have a strong influence on solubility. Such effects become even more pronounced in autoclave leaching vessels where temperatures can be as high as 300\u00b0C.<\/p>\n

As an illustration, iminodiacetic acid ion exchange resins contain the functional group -NH-(CH2<\/sub>CO2<\/sub>H)2<\/sub> bonded to a polystyrene polymer. The acetic groups can ionize (weak acid) to form a dianion group. This plus the trivalent nitrogen atom complexes to Cu+2<\/sup> quite strongly and is a viable means of recovering cupric ion from dilute aqueous solution streams. And while the resin operates best at pH \u2265 4, precipitation of Cu+2<\/sup> can limit its use to pH 4 or less.<\/p>\n","protected":false},"author":2529,"menu_order":5,"template":"","meta":{"pb_show_title":"on","pb_short_title":"","pb_subtitle":"","pb_authors":[],"pb_section_license":""},"chapter-type":[],"contributor":[],"license":[],"class_list":["post-1226","chapter","type-chapter","status-publish","hentry"],"part":1188,"_links":{"self":[{"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/pressbooks\/v2\/chapters\/1226","targetHints":{"allow":["GET"]}}],"collection":[{"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/pressbooks\/v2\/chapters"}],"about":[{"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/wp\/v2\/types\/chapter"}],"author":[{"embeddable":true,"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/wp\/v2\/users\/2529"}],"version-history":[{"count":10,"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/pressbooks\/v2\/chapters\/1226\/revisions"}],"predecessor-version":[{"id":3791,"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/pressbooks\/v2\/chapters\/1226\/revisions\/3791"}],"part":[{"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/pressbooks\/v2\/parts\/1188"}],"metadata":[{"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/pressbooks\/v2\/chapters\/1226\/metadata\/"}],"wp:attachment":[{"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/wp\/v2\/media?parent=1226"}],"wp:term":[{"taxonomy":"chapter-type","embeddable":true,"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/pressbooks\/v2\/chapter-type?post=1226"},{"taxonomy":"contributor","embeddable":true,"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/wp\/v2\/contributor?post=1226"},{"taxonomy":"license","embeddable":true,"href":"https:\/\/pressbooks.bccampus.ca\/hydrometallurgy\/wp-json\/wp\/v2\/license?post=1226"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}