{"id":835,"date":"2019-09-27T15:23:28","date_gmt":"2019-09-27T15:23:28","guid":{"rendered":"http:\/\/chemdictionary.org\/?p=835"},"modified":"2019-09-27T15:23:30","modified_gmt":"2019-09-27T15:23:30","slug":"first-ionization-energy","status":"publish","type":"post","link":"https:\/\/chemdictionary.org\/first-ionization-energy\/","title":{"rendered":"First Ionization Energy"},"content":{"rendered":"\n<p><\/p>\n\n\n\n<h2><strong>Definition <\/strong><\/h2>\n\n\n\n<p>The\nionization energy of an element is the minimum energy required to remove an\nelectron from the valence shell of an isolated gaseous atom to form an ion [1].<\/p>\n\n\n\n<p>For\nexample:&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; Mg \u2192 Mg + e<sup>&#8211;<\/sup>&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; \u0394H = 738kjmol<sup>-1<\/sup> <\/p>\n\n\n\n<h2><strong>Explanation <\/strong><\/h2>\n\n\n\n<p>In\nthe gaseous phase, the atoms and ions are isolated and are free from all\nexternal influences. Thus, the ionization energy is the qualitative measure of\nthe stability of an atom. The higher ionization energy cause difficult to\nremove an electron. It is important because it can be used to predict the\nstrength of chemical bonds.<\/p>\n\n\n\n<p><strong>Unit:<\/strong> In chemistry, the unit of ionization energy is expressed as kilojoules per mole (Kj\/mol) or kilocalories per mole (kcal\/mol). In physics, it is measured in term of electronvolts [2]. <\/p>\n\n\n\n<p><strong>Types of ionization energy<\/strong><\/p>\n\n\n\n<p><strong>First ionization energy<\/strong><\/p>\n\n\n\n<p>The\nfirst ionization energy is the energy required to remove 1 electron from the valence\nshell. The first ionization energy of sodium ions is +496 kjmol<sup>-1<\/sup>.<\/p>\n\n\n\n<p>&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; Na\n\u2192 Na<sup>+<\/sup> + e<sup>&#8211;<\/sup>&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;\n\u0394H = +496 kJmol<sup>-1<\/sup>.<\/p>\n\n\n\n<p>&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; H \u2192\nH<sup>+<\/sup> + e<sup>&#8211;<\/sup>&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;\n\u0394H = -1312.0 kJmol<sup>-1<\/sup>.<\/p>\n\n\n\n<p>The\nfirst ionization energy of hydrogen is about half in a chemical reaction. When\nnatural gas burnt about 800 kJ of energy is released.<\/p>\n\n\n\n<p>&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; CH<sub>4<\/sub> (g)\n+ 2O<sub>2<\/sub> (g)&nbsp;&nbsp; \u2192&nbsp;&nbsp; CO<sub>2<\/sub> (g) + 2 H<sub>2<\/sub> O\n(g)&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; \u0394H = -802.4 kJmol<sup>-1<\/sup>&nbsp;&nbsp;&nbsp; <\/p>\n\n\n\n<h2><strong>Second ionization energy<\/strong><\/h2>\n\n\n\n<p>When\nthere is more than one electron in the valence shell, so they can be removed\none by one with the help of providing more energy. Like elements of group 2 and\n3 they have more than one electron in their outermost shells. That\u2019s why they\nhave more than one ionization energy. Each succeeding energy is larger than the\npreceding one.<\/p>\n\n\n\n<p>Second\nionization energy is higher than the first ionization energy because removing the\nfirst electron gives the atom stable electron shell like an alkali metal but\nremoving the second electron involves a new electron shell that is closer and\ntightly bound to the atomic nucleus [3]. <\/p>\n\n\n\n<p>&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; Mg\n\u2192 Mg<sup>+<\/sup> + e<sup>&#8211;<\/sup>&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;\n\u0394H = 738 kjmol<sup>-1<\/sup>.<\/p>\n\n\n\n<p>&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; Mg\n\u2192 Mg<sup>2+<\/sup> + e<sup>&#8211;<\/sup>&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;\n\u0394H = 1451 kjmol<sup>-1<\/sup>.<\/p>\n\n\n\n<p>To\ncarry out chemical reactions, the role of ionization energy is significant. As\nit determines the probability of the reactants to form covalent or ionic bonds.\nFor example the ionization energy of sodium is 496 kJ\/mol but chlorine first\nionization energy is 1251.1 kJ\/mol, so due to difference in the energy values\nthe chemically combine and form ionic bonds. Elements that lies very close in\nthe periodic table form covalent bonds due to less difference in ionization\nenergies value. For example carbon and chlorine form CCl<sub>4<\/sub>.<\/p>\n\n\n\n<h2><strong>Periodic Table Trend<\/strong><\/h2>\n\n\n\n<p>In\nperiodic table ionization energies increases from left to right in a period\nwith the increase in proton number. Each period begins with an element which\nhas one electron in its valence shell and ends with the completion of an\nelectronic shell. Increase in atomic number is associated with increase in\nnuclear charge which leads to a strong force of attraction.&nbsp; <\/p>\n\n\n\n<p>Elements\non the left side of the periodic table have less ionization values as compare\nto the elements on the right side of the periodic table.<\/p>\n\n\n\n<table class=\"wp-block-table\"><tbody><tr><td>\n  &nbsp;\n  2<sup>nd<\/sup>\n  period elements\n  <\/td><td>\n  &nbsp;\n  <sup>3<\/sup>Li\n  <\/td><td>\n  &nbsp;\n  <sup>4<\/sup>Be\n  <\/td><td>\n  &nbsp;\n  <sup>5<\/sup>B\n  <\/td><td>\n  &nbsp;\n  <sup>6<\/sup>C\n  <\/td><td>\n  &nbsp;\n  <sup>7<\/sup>N\n  <\/td><td>\n  &nbsp;\n  <sup>8<\/sup>O\n  <\/td><td>\n  &nbsp;\n  <sup>9<\/sup>F\n  <\/td><td>\n  &nbsp;\n  <sup>10<\/sup>Ne\n  <\/td><\/tr><tr><td>\n  Ionization\n  energy (kJ\/mol)\n  <\/td><td>\n  &nbsp;\n  520\n  <\/td><td>\n  &nbsp;\n  899\n  <\/td><td>\n  &nbsp;\n  801\n  <\/td><td>\n  &nbsp;\n  1086\n  <\/td><td>\n  &nbsp;\n  1402\n  <\/td><td>\n  &nbsp;\n  1314\n  <\/td><td>\n  &nbsp;\n  1681\n  <\/td><td>\n  &nbsp;\n  2081\n  <\/td><\/tr><\/tbody><\/table>\n\n\n\n<p>In\ngroups, the ionization energy decreases due to increase in nuclear charge. Due\nto addition of electronic shells increases the distance between the nucleus and\nouter electron. That\u2019s why electron are easily removed with less energy. Force\nof attraction is also decreasing due to increase in shielding effect of\nintervening electrons.<\/p>\n\n\n\n<table class=\"wp-block-table\"><tbody><tr><td>\n  &nbsp;\n  1<sup>st<\/sup> group elements\n  <\/td><td>\n  &nbsp;\n  Ionization energy (kJ\/mol)\n  <\/td><\/tr><tr><td>\n  &nbsp;\n  &nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; <sup>3<\/sup>Li\n  <\/td><td>\n  &nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;\n  \n  &nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; 520\n  <\/td><\/tr><tr><td>\n  &nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; \n  &nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; <sup>11<\/sup>Na\n  &nbsp;\n  <\/td><td>\n  &nbsp;&nbsp;&nbsp;&nbsp;&nbsp;\n  \n  &nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; 496\n  <\/td><\/tr><tr><td>\n  &nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; \n  &nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; <sup>19<\/sup>K\n  <\/td><td>\n  &nbsp;\n  &nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; 419\n  <\/td><\/tr><tr><td>\n  &nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; \n  &nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; 37Rb\n  <\/td><td>\n  &nbsp;\n  &nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; 403\n  <\/td><\/tr><tr><td>\n  &nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; \n  &nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; <sup>55<\/sup>Cs\n  <\/td><td>\n  &nbsp;&nbsp;&nbsp;&nbsp;\n  \n  &nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; 377\n  <\/td><\/tr><\/tbody><\/table>\n\n\n\n<h2><strong>Factors Affecting Ionizing Trend<\/strong><\/h2>\n\n\n\n<p>Ionization\npotential is an older term used for ionization energy because in history\nionization energy was measured through ionizing a sample and accelerating the\nelectron. The electrons are removed through electrostatic potential.<\/p>\n\n\n\n<p>Here\nare some factors which are affecting ionization energy:<\/p>\n\n\n\n<ol><li><strong>Nuclear\ncharge<\/strong><\/li><\/ol>\n\n\n\n<p>Higher the magnitude of nuclear\ncharge the more closely electrons are held by the nucleus. That\u2019s why more will\nbe the ionization energy.<\/p>\n\n\n\n<ul><li><strong>Number\nof electron shells<\/strong><\/li><\/ul>\n\n\n\n<p>On the size of atom\nionization energy will be measured. If the size of atom is greater less tightly\nthe electrons are held so the ionization energy will be less.<\/p>\n\n\n\n<ul><li><strong>Effective\nnuclear charge (Z <sub>eff<\/sub>)<\/strong><\/li><\/ul>\n\n\n\n<p>Higher the magnitude of\nelectrons and shielding power. The electrons are less tightly present by the\nnucleus and that\u2019s why there will be less ionization energy.<\/p>\n\n\n\n<ul><li><strong>Type\nof orbital ionized<\/strong><\/li><\/ul>\n\n\n\n<p>If the atom has stable\nelectronic configuration then it will be very difficult to remove electron from\noutermost shell which results in higher ionization energy.<\/p>\n\n\n\n<ul><li><strong>Occupancy\nof orbital matters<\/strong><\/li><\/ul>\n\n\n\n<p>If the orbit is completely filled or\nhalf-filled then it is very difficult to remove the electrons.<\/p>\n\n\n\n<p>The\nionization energy of (n+1)th is higher than the nth ionization energy because\nremoval of next electron from same shell increases the ionization energy due to\nincreased net charge of the ion. Element also has greater electrostatic force\nof attraction because of less distance between the nucleus and the electron. These\ntwo factors also increase the ionization energy of atom.<\/p>\n\n\n\n<p>However\nthere are certain elements which show irregular trends in period 3. For example\nionization energy of Mg and P are higher than those of Al and S. In case of Mg\n(1s<sup>2<\/sup>, 2s<sup>2<\/sup>, 2p<sup>6<\/sup>, 3s<sup>2<\/sup>) it is more\ndifficult to remove electron from completely filled 3s orbital while in case of\nAl (1s<sup>2<\/sup>, 2s<sup>2<\/sup>, 2p<sup>6<\/sup>, 3s<sup>2<\/sup>, 3p<sup>1<\/sup>)\nit is easier to remove the same from partially filled 3p orbital. Since to\nremove an electron from a 3s-orbital of Mg atom requires more energy than to\nremove the same from a 3p- orbital of Al atom that is why ionization energy of\nMg is higher than that of Al.<\/p>\n\n\n\n<p><strong>Electron Affinity<\/strong><\/p>\n\n\n\n<h3><strong>Definition<\/strong><\/h3>\n\n\n\n<p>The\nelectron affinity of an atom is the energy released when an electron adds to an\nempty or partially filled orbital of an isolated gaseous atom in its valence\nenergy level to form an anion having a unit negative charge. <\/p>\n\n\n\n<p>&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; Cl (g) + e<sup>&#8211;<\/sup>&nbsp;&nbsp; \u2192 Cl<sup>&#8211;<\/sup> (g)&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; \u0394H=-349kJ\/mol<\/p>\n\n\n\n<p>Since\nenergy is released, so electron affinity is given the negative sign. Electron\naffinity is the measure of the attraction of the nucleus of an atom for the\nextra electrons. <\/p>\n\n\n\n<p><strong>Factor Influencing Electron Affinity<\/strong><\/p>\n\n\n\n<p>Like\nionization energy, electron affinity is influenced by atomic radius, the\nnuclear charge and the shielding effects of inner electrons. As the force of\nattraction between the valence electrons and the nucleus decreases with the\nincrease in the atomic radius, the electron affinities usually decreases.<\/p>\n\n\n\n<p><strong>Variation in the Periodic Table<\/strong><\/p>\n\n\n\n<p>In\na period the atomic radius decreases due to increase in the nuclear charge.\nThus the electron affinities of element increases from left to right in the\nperiodic table. That is why the alkali metals have the lowest and the halogens\nhave the highest electron affinities. <\/p>\n\n\n\n<p>In\ngroups atomic radii increase with the increase in the proton number due to\nsuccessive increase of electronic shells. This also exerts a shielding effect\non the force of attraction between nucleus and the valence electron. Thus\nelectron affinities usually decreases from top to bottom.<\/p>\n\n\n\n<p><strong>Prediction of Ionic and covalent\nbonds<\/strong><\/p>\n\n\n\n<p>There\nare many types of bonding and forces are present that bind molecules but two\nbasic types are ionic and covalent bond [4].<\/p>\n\n\n\n<p><strong>Ionic bond &nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; <\/strong><\/p>\n\n\n\n<p>It\nis defined as complete transfer of valence electron between atoms.&nbsp; In ionic bonding it require one electron\ndonor (metal) and one electron acceptor (non-metal). The metal loses electrons\nto become a positively charged cation and non-metal accepts that electron to\nbecome a negatively charged anion.<\/p>\n\n\n\n<p>For\nexample: Sodium (Na) is donating 1 electron to the chlorine (Cl) atom. The net\ncharge of the resulting compound is zero.<\/p>\n\n\n\n<p>&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; &nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;Na&nbsp;&nbsp; + Cl \u2192 Na<sup>+<\/sup> Cl<sup>&#8211;<\/sup><\/p>\n\n\n\n<p>Overall\npredicted that ionization energy of metal and electron affinity of non-metal\ncompound is positive. The attraction between these particles release energy to\ncarry out the reaction. Due to opposite charge of both compounds they are\nmostly dissolve in polar solvents.<\/p>\n\n\n\n<p><strong>Covalent bond<\/strong><\/p>\n\n\n\n<p>It\nis defined as complete sharing of electrons between atoms. The bonding occurs\nbetween atoms having same electronegativity because both have no tendency to\ndonate electrons but they share electrons to complete their outermost shell and\nbecome stable.<\/p>\n\n\n\n<p>For\nexample: Carbon share its 4 electrons through single, double and triple bonds.<\/p>\n\n\n\n<p><strong>Electron binding energy<\/strong><\/p>\n\n\n\n<p>It\nis also known as ionization potential or energy, it is that amount of energy\nwhich is required to free electrons from a molecule or an atom or an ion.<\/p>\n\n\n\n<p>Binding\nenergy of a single proton or neutron in nucleus is million times greater than\nelectron in an atom [5].<\/p>\n\n\n\n<p>For example: Hydrogen atom has one electron in its outermost shell so it is very harder to remove electron because there is less distance between the nucleus and the electron greater electrostatic force of repulsion is required to remove an electron. <\/p>\n\n\n\n<h4><strong>References<\/strong><\/h4>\n\n\n\n<ol><li><a href=\"https:\/\/en.wikipedia.org\/wiki\/Ionization_energy\">https:\/\/en.wikipedia.org\/wiki\/Ionization_energy<\/a><\/li><li><a href=\"http:\/\/chemwiki.ucdavis.edu\/Inorganic_Chemistry\/Descriptive_Chemistry\/Periodic_Table_of_the_Elements\/Ionization_Energy\">&#8220;Ionization\nEnergy&#8221;<\/a>.&nbsp;<em>ChemWiki<\/em>. University of California, Davis.\n2013-10-02.<\/li><li><a href=\"https:\/\/www.thoughtco.com\/ionization-energy-and-trend-604538\">https:\/\/www.thoughtco.com\/ionization-energy-and-trend-604538<\/a><\/li><li><a href=\"https:\/\/chem.libretexts.org\/Bookshelves\/Organic_Chemistry\/Supplemental_Modules_(Organic_Chemistry)\/Fundamentals\/Ionic_and_Covalent_Bonds\">https:\/\/chem.libretexts.org\/Bookshelves\/Organic_Chemistry\/Supplemental_Modules_(Organic_Chemistry)\/Fundamentals\/Ionic_and_Covalent_Bonds<\/a>.<\/li><li>https:\/\/www.britannica.com\/science\/binding-energy<\/li><\/ol>\n","protected":false},"excerpt":{"rendered":"<p>Definition The ionization energy of an element is the minimum energy required to remove an electron from the valence shell of an isolated gaseous atom to form an ion [1]. For example:&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; Mg \u2192 Mg + e&#8211;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; \u0394H = 738kjmol-1 Explanation In the gaseous phase, the atoms and ions are isolated and are free from &#8230; <a title=\"First Ionization Energy\" class=\"read-more\" href=\"https:\/\/chemdictionary.org\/first-ionization-energy\/\" aria-label=\"More on First Ionization Energy\">Read more<\/a><\/p>\n","protected":false},"author":6,"featured_media":0,"comment_status":"closed","ping_status":"open","sticky":false,"template":"","format":"standard","meta":{"rating_form_position":"","rating_results_position":"","mr_structured_data_type":""},"categories":[10],"tags":[],"multi-rating":{"mr_rating_results":[{"adjusted_star_result":1,"star_result":1,"total_max_option_value":5,"adjusted_score_result":1,"score_result":1,"percentage_result":20,"adjusted_percentage_result":20,"count":3,"post_id":835}]},"_links":{"self":[{"href":"https:\/\/chemdictionary.org\/wp-json\/wp\/v2\/posts\/835"}],"collection":[{"href":"https:\/\/chemdictionary.org\/wp-json\/wp\/v2\/posts"}],"about":[{"href":"https:\/\/chemdictionary.org\/wp-json\/wp\/v2\/types\/post"}],"author":[{"embeddable":true,"href":"https:\/\/chemdictionary.org\/wp-json\/wp\/v2\/users\/6"}],"replies":[{"embeddable":true,"href":"https:\/\/chemdictionary.org\/wp-json\/wp\/v2\/comments?post=835"}],"version-history":[{"count":2,"href":"https:\/\/chemdictionary.org\/wp-json\/wp\/v2\/posts\/835\/revisions"}],"predecessor-version":[{"id":869,"href":"https:\/\/chemdictionary.org\/wp-json\/wp\/v2\/posts\/835\/revisions\/869"}],"wp:attachment":[{"href":"https:\/\/chemdictionary.org\/wp-json\/wp\/v2\/media?parent=835"}],"wp:term":[{"taxonomy":"category","embeddable":true,"href":"https:\/\/chemdictionary.org\/wp-json\/wp\/v2\/categories?post=835"},{"taxonomy":"post_tag","embeddable":true,"href":"https:\/\/chemdictionary.org\/wp-json\/wp\/v2\/tags?post=835"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}