{"id":246928,"date":"2025-07-07T14:01:21","date_gmt":"2025-07-07T14:01:21","guid":{"rendered":"https:\/\/learnexams.com\/blog\/?p=246928"},"modified":"2025-07-07T14:01:23","modified_gmt":"2025-07-07T14:01:23","slug":"a-student-proposes-the-following-lewis-structure-for-the-nitronium-no2-ion","status":"publish","type":"post","link":"https:\/\/www.learnexams.com\/blog\/2025\/07\/07\/a-student-proposes-the-following-lewis-structure-for-the-nitronium-no2-ion\/","title":{"rendered":"A student proposes the following Lewis structure for the nitronium NO2+ ion."},"content":{"rendered":"\n

A student proposes the following Lewis structure for the nitronium NO2+ ion. =N= Assign formal charge to each atom in the student’s Lewis structure. atom formal charge left 0 N right 0<\/p>\n\n\n\n

\"\"<\/figure>\n\n\n\n

The Correct Answer and Explanation is:<\/mark><\/strong><\/p>\n\n\n\n

To determine the formal charges of the atoms in the nitronium ion (NO2+), we follow these steps:<\/p>\n\n\n\n

Step 1: Determine the structure of NO2+<\/h3>\n\n\n\n

The nitronium ion (NO2+) consists of a central nitrogen (N) atom bonded to two oxygen (O) atoms. One of the oxygen atoms is double-bonded to nitrogen, while the other is single-bonded to nitrogen and carries a formal positive charge (due to the positive charge on the entire ion). The N-O bond length is shorter in the double bond and longer in the single bond.<\/p>\n\n\n\n

Step 2: Calculate the formal charges<\/h3>\n\n\n\n

The formal charge<\/strong> on an atom is given by the formula:Formal Charge=Valence electrons\u2212(Non-bonding electrons+12\u00d7Bonding electrons)\\text{Formal Charge} = \\text{Valence electrons} – (\\text{Non-bonding electrons} + \\frac{1}{2} \\times \\text{Bonding electrons})Formal Charge=Valence electrons\u2212(Non-bonding electrons+21\u200b\u00d7Bonding electrons)<\/p>\n\n\n\n

For the Nitrogen (N) atom:<\/h4>\n\n\n\n
    \n
  • Valence electrons<\/strong>: Nitrogen is in Group 15, so it has 5 valence electrons.<\/li>\n\n\n\n
  • Non-bonding electrons<\/strong>: Nitrogen has no lone pairs in the NO2+ ion.<\/li>\n\n\n\n
  • Bonding electrons<\/strong>: Nitrogen is bonded to two oxygen atoms. One oxygen has a double bond (4 electrons), and the other has a single bond (2 electrons).<\/li>\n<\/ul>\n\n\n\n

    Formal Charge on N=5\u2212(0+12\u00d7(4+2))=5\u22123=+1\\text{Formal Charge on N} = 5 – (0 + \\frac{1}{2} \\times (4 + 2)) = 5 – 3 = +1Formal Charge on N=5\u2212(0+21\u200b\u00d7(4+2))=5\u22123=+1<\/p>\n\n\n\n

    For the oxygen (O) with a double bond:<\/h4>\n\n\n\n
      \n
    • Valence electrons<\/strong>: Oxygen is in Group 16, so it has 6 valence electrons.<\/li>\n\n\n\n
    • Non-bonding electrons<\/strong>: Oxygen has two lone pairs (4 electrons).<\/li>\n\n\n\n
    • Bonding electrons<\/strong>: Oxygen has a double bond with nitrogen, which contributes 4 electrons.<\/li>\n<\/ul>\n\n\n\n

      Formal Charge on O (double-bonded)=6\u2212(4+12\u00d74)=6\u22126=0\\text{Formal Charge on O (double-bonded)} = 6 – (4 + \\frac{1}{2} \\times 4) = 6 – 6 = 0Formal Charge on O (double-bonded)=6\u2212(4+21\u200b\u00d74)=6\u22126=0<\/p>\n\n\n\n

      For the oxygen (O) with a single bond:<\/h4>\n\n\n\n
        \n
      • Valence electrons<\/strong>: This oxygen also has 6 valence electrons.<\/li>\n\n\n\n
      • Non-bonding electrons<\/strong>: This oxygen has three lone pairs (6 electrons).<\/li>\n\n\n\n
      • Bonding electrons<\/strong>: Oxygen has a single bond with nitrogen, contributing 2 electrons.<\/li>\n<\/ul>\n\n\n\n

        Formal Charge on O (single-bonded)=6\u2212(6+12\u00d72)=6\u22127=\u22121\\text{Formal Charge on O (single-bonded)} = 6 – (6 + \\frac{1}{2} \\times 2) = 6 – 7 = -1Formal Charge on O (single-bonded)=6\u2212(6+21\u200b\u00d72)=6\u22127=\u22121<\/p>\n\n\n\n

        Step 3: Assign formal charges<\/h3>\n\n\n\n
          \n
        • Nitrogen (N)<\/strong>: +1<\/li>\n\n\n\n
        • Double-bonded Oxygen (O)<\/strong>: 0<\/li>\n\n\n\n
        • Single-bonded Oxygen (O)<\/strong>: -1<\/li>\n<\/ul>\n\n\n\n

          Step 4: Overall charge of the molecule<\/h3>\n\n\n\n

          The formal charges sum up to give the overall charge of the ion:+1+0+(\u22121)=+1+1 + 0 + (-1) = +1+1+0+(\u22121)=+1<\/p>\n\n\n\n

          Conclusion:<\/h3>\n\n\n\n

          In the correct Lewis structure for the nitronium ion (NO2+), the nitrogen atom carries a formal charge of +1<\/strong>, the double-bonded oxygen carries a formal charge of 0<\/strong>, and the single-bonded oxygen carries a formal charge of -1<\/strong>. These charges are consistent with the overall positive charge on the NO2+ ion.<\/p>\n\n\n\n

          \"\"<\/figure>\n","protected":false},"excerpt":{"rendered":"

          A student proposes the following Lewis structure for the nitronium NO2+ ion. =N= Assign formal charge to each atom in the student’s Lewis structure. atom formal charge left 0 N right 0 The Correct Answer and Explanation is: To determine the formal charges of the atoms in the nitronium ion (NO2+), we follow these steps: […]<\/p>\n","protected":false},"author":1,"featured_media":0,"comment_status":"closed","ping_status":"closed","sticky":false,"template":"","format":"standard","meta":{"site-sidebar-layout":"default","site-content-layout":"","ast-site-content-layout":"default","site-content-style":"default","site-sidebar-style":"default","ast-global-header-display":"","ast-banner-title-visibility":"","ast-main-header-display":"","ast-hfb-above-header-display":"","ast-hfb-below-header-display":"","ast-hfb-mobile-header-display":"","site-post-title":"","ast-breadcrumbs-content":"","ast-featured-img":"","footer-sml-layout":"","ast-disable-related-posts":"","theme-transparent-header-meta":"","adv-header-id-meta":"","stick-header-meta":"","header-above-stick-meta":"","header-main-stick-meta":"","header-below-stick-meta":"","astra-migrate-meta-layouts":"default","ast-page-background-enabled":"default","ast-page-background-meta":{"desktop":{"background-color":"","background-image":"","background-repeat":"repeat","background-position":"center center","background-size":"auto","background-attachment":"scroll","background-type":"","background-media":"","overlay-type":"","overlay-color":"","overlay-opacity":"","overlay-gradient":""},"tablet":{"background-color":"","background-image":"","background-repeat":"repeat","background-position":"center center","background-size":"auto","background-attachment":"scroll","background-type":"","background-media":"","overlay-type":"","overlay-color":"","overlay-opacity":"","overlay-gradient":""},"mobile":{"background-color":"","background-image":"","background-repeat":"repeat","background-position":"center center","background-size":"auto","background-attachment":"scroll","background-type":"","background-media":"","overlay-type":"","overlay-color":"","overlay-opacity":"","overlay-gradient":""}},"ast-content-background-meta":{"desktop":{"background-color":"var(--ast-global-color-5)","background-image":"","background-repeat":"repeat","background-position":"center center","background-size":"auto","background-attachment":"scroll","background-type":"","background-media":"","overlay-type":"","overlay-color":"","overlay-opacity":"","overlay-gradient":""},"tablet":{"background-color":"var(--ast-global-color-5)","background-image":"","background-repeat":"repeat","background-position":"center center","background-size":"auto","background-attachment":"scroll","background-type":"","background-media":"","overlay-type":"","overlay-color":"","overlay-opacity":"","overlay-gradient":""},"mobile":{"background-color":"var(--ast-global-color-5)","background-image":"","background-repeat":"repeat","background-position":"center center","background-size":"auto","background-attachment":"scroll","background-type":"","background-media":"","overlay-type":"","overlay-color":"","overlay-opacity":"","overlay-gradient":""}},"footnotes":""},"categories":[25],"tags":[],"class_list":["post-246928","post","type-post","status-publish","format-standard","hentry","category-exams-certification"],"_links":{"self":[{"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/posts\/246928","targetHints":{"allow":["GET"]}}],"collection":[{"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/posts"}],"about":[{"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/types\/post"}],"author":[{"embeddable":true,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/users\/1"}],"replies":[{"embeddable":true,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/comments?post=246928"}],"version-history":[{"count":0,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/posts\/246928\/revisions"}],"wp:attachment":[{"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/media?parent=246928"}],"wp:term":[{"taxonomy":"category","embeddable":true,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/categories?post=246928"},{"taxonomy":"post_tag","embeddable":true,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/tags?post=246928"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}