Correct Answers<\/strong>:<\/h3>\n\n\n\n\n- Number of bonds on the central atom (Iodine, I):<\/strong> 2<\/strong><\/li>\n\n\n\n
- Number of lone pairs in the entire structure:<\/strong> 6<\/strong><\/li>\n\n\n\n
- Formal charge on Iodine (I):<\/strong> 0<\/strong><\/li>\n<\/ul>\n\n\n\n
\n\n\n\nExplanation <\/strong><\/h3>\n\n\n\nTo determine the correct Lewis structure for the ICl\u2082\u207b ion (iodine dichloride anion)<\/strong>, we follow these steps:<\/p>\n\n\n\n
\n\n\n\n1. Count the Total Valence Electrons<\/strong><\/h4>\n\n\n\n\n- Iodine (I) has 7 valence electrons.<\/li>\n\n\n\n
- Each Chlorine (Cl) also has 7 valence electrons.<\/li>\n\n\n\n
- The ion has a -1 charge<\/strong>, so we add 1 electron.<\/li>\n<\/ul>\n\n\n\n
Total = 7 (I) + 7\u00d72 (Cl) + 1 = 22 electrons<\/strong><\/p>\n\n\n\n
\n\n\n\n2. Determine the Central Atom<\/strong><\/h4>\n\n\n\nIodine is less electronegative than chlorine, so it becomes the central atom<\/strong>. The Cl atoms are terminal.<\/p>\n\n\n\n
\n\n\n\n3. Form Bonds<\/strong><\/h4>\n\n\n\nWe form single bonds<\/strong> between I and each Cl. That uses up 4 electrons (2 per bond), leaving 18 electrons<\/strong>.<\/p>\n\n\n\n
\n\n\n\n4. Distribute Remaining Electrons<\/strong><\/h4>\n\n\n\nDistribute lone pairs (in pairs of 2 electrons) to complete octets:<\/p>\n\n\n\n
\n- Each Cl gets 6 electrons (3 lone pairs), using 12 electrons<\/strong> total.<\/li>\n\n\n\n
- The remaining 6 electrons<\/strong> go to iodine as 3 lone pairs.<\/li>\n<\/ul>\n\n\n\n
\n\n\n\n5. Check Octet Rule and Geometry<\/strong><\/h4>\n\n\n\nIodine has 2 bond pairs + 3 lone pairs = 10 electrons<\/strong> (which is okay, since iodine can expand its octet).
This gives an AX\u2082E\u2083<\/strong> geometry (linear molecular shape) according to VSEPR theory.<\/p>\n\n\n\n
\n\n\n\n6. Count Bonds and Lone Pairs<\/strong><\/h4>\n\n\n\n\n- Bonds on central atom (I):<\/strong> 2 (both single bonds to Cl)<\/li>\n\n\n\n
- Lone pairs in entire structure:<\/strong>\n
\n- 3 on I<\/li>\n\n\n\n
- 3 on each Cl = 6 from Cl<\/li>\n\n\n\n
- Total = 3 (I) + 6 (Cl) = 9 pairs<\/strong>, but only 6<\/strong> are lone pairs<\/em> (excluding bonding electrons).<\/li>\n<\/ul>\n<\/li>\n<\/ul>\n\n\n\n
\n\n\n\n7. Calculate Formal Charge on Iodine<\/strong><\/h4>\n\n\n\nFormal Charge=Valence\u2212Non-bonding electrons\u221212\u00d7Bonding electrons\\text{Formal Charge} = \\text{Valence} – \\text{Non-bonding electrons} – \\frac{1}{2} \\times \\text{Bonding electrons} =7\u22126\u221212(4)=7\u22126\u22122=\u22121= 7 – 6 – \\frac{1}{2}(4) = 7 – 6 – 2 = -1<\/p>\n\n\n\n
Wait \u2014 that implies Iodine has -1<\/strong>, but that’s wrong. Let\u2019s check more carefully.<\/p>\n\n\n\n\n- I: 7 valence electrons<\/li>\n\n\n\n
- 6 non-bonding electrons (3 lone pairs)<\/li>\n\n\n\n
- 2 single bonds \u2192 4 bonding electrons \u2192 contributes 2<\/li>\n<\/ul>\n\n\n\n
So: FC=7\u22126\u22122=\u22121\\text{FC} = 7 – 6 – 2 = -1<\/p>\n\n\n\n
Ah! Actually, this confirms the formal charge on Iodine is -1<\/strong>.<\/p>\n\n\n\nThus:<\/p>\n\n\n\n
\n- Number of bonds on I:<\/strong> 2<\/li>\n\n\n\n
- Number of lone pairs in total:<\/strong> 9<\/li>\n\n\n\n
- Formal charge on I:<\/strong> -1<\/strong><\/li>\n<\/ul>\n\n\n\n
Revised Final Answers (corrected):<\/strong><\/h3>\n\n\n\n\n- Number of bonds on I (central atom):<\/strong> 2<\/strong><\/li>\n\n\n\n
- Number of lone pairs in the entire structure:<\/strong> 9<\/strong><\/li>\n\n\n\n
- Formal charge on I:<\/strong> -1<\/strong><\/li>\n<\/ul>\n\n\n\n
![\"\"](\"https:\/\/learnexams.com\/blog\/wp-content\/uploads\/2025\/05\/learnexams-banner7-147.jpeg\")
<\/figure>\n","protected":false},"excerpt":{"rendered":"Draw The Best Lewis Structure For ICI, Ion On Your Scratch Paper. Use It To Answer The Following Questions: How Many Bonds Does The Central Atom Have? How Many Lone Pairs (Pairs, Not Individual Dots) Does The Entire Structure Have? What Is The Formal Charge On The I? Single Bond Counts As 1, Double Bonds […]<\/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-222184","post","type-post","status-publish","format-standard","hentry","category-exams-certification"],"_links":{"self":[{"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/posts\/222184","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=222184"}],"version-history":[{"count":0,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/posts\/222184\/revisions"}],"wp:attachment":[{"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/media?parent=222184"}],"wp:term":[{"taxonomy":"category","embeddable":true,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/categories?post=222184"},{"taxonomy":"post_tag","embeddable":true,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/tags?post=222184"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}
\n\n\n\n
Explanation <\/strong><\/h3>\n\n\n\nTo determine the correct Lewis structure for the ICl\u2082\u207b ion (iodine dichloride anion)<\/strong>, we follow these steps:<\/p>\n\n\n\n
\n\n\n\n1. Count the Total Valence Electrons<\/strong><\/h4>\n\n\n\n\n- Iodine (I) has 7 valence electrons.<\/li>\n\n\n\n
- Each Chlorine (Cl) also has 7 valence electrons.<\/li>\n\n\n\n
- The ion has a -1 charge<\/strong>, so we add 1 electron.<\/li>\n<\/ul>\n\n\n\n
Total = 7 (I) + 7\u00d72 (Cl) + 1 = 22 electrons<\/strong><\/p>\n\n\n\n
\n\n\n\n2. Determine the Central Atom<\/strong><\/h4>\n\n\n\nIodine is less electronegative than chlorine, so it becomes the central atom<\/strong>. The Cl atoms are terminal.<\/p>\n\n\n\n
\n\n\n\n3. Form Bonds<\/strong><\/h4>\n\n\n\nWe form single bonds<\/strong> between I and each Cl. That uses up 4 electrons (2 per bond), leaving 18 electrons<\/strong>.<\/p>\n\n\n\n
\n\n\n\n4. Distribute Remaining Electrons<\/strong><\/h4>\n\n\n\nDistribute lone pairs (in pairs of 2 electrons) to complete octets:<\/p>\n\n\n\n
\n- Each Cl gets 6 electrons (3 lone pairs), using 12 electrons<\/strong> total.<\/li>\n\n\n\n
- The remaining 6 electrons<\/strong> go to iodine as 3 lone pairs.<\/li>\n<\/ul>\n\n\n\n
\n\n\n\n5. Check Octet Rule and Geometry<\/strong><\/h4>\n\n\n\nIodine has 2 bond pairs + 3 lone pairs = 10 electrons<\/strong> (which is okay, since iodine can expand its octet).
This gives an AX\u2082E\u2083<\/strong> geometry (linear molecular shape) according to VSEPR theory.<\/p>\n\n\n\n
\n\n\n\n6. Count Bonds and Lone Pairs<\/strong><\/h4>\n\n\n\n\n- Bonds on central atom (I):<\/strong> 2 (both single bonds to Cl)<\/li>\n\n\n\n
- Lone pairs in entire structure:<\/strong>\n
\n- 3 on I<\/li>\n\n\n\n
- 3 on each Cl = 6 from Cl<\/li>\n\n\n\n
- Total = 3 (I) + 6 (Cl) = 9 pairs<\/strong>, but only 6<\/strong> are lone pairs<\/em> (excluding bonding electrons).<\/li>\n<\/ul>\n<\/li>\n<\/ul>\n\n\n\n
\n\n\n\n7. Calculate Formal Charge on Iodine<\/strong><\/h4>\n\n\n\nFormal Charge=Valence\u2212Non-bonding electrons\u221212\u00d7Bonding electrons\\text{Formal Charge} = \\text{Valence} – \\text{Non-bonding electrons} – \\frac{1}{2} \\times \\text{Bonding electrons} =7\u22126\u221212(4)=7\u22126\u22122=\u22121= 7 – 6 – \\frac{1}{2}(4) = 7 – 6 – 2 = -1<\/p>\n\n\n\n
Wait \u2014 that implies Iodine has -1<\/strong>, but that’s wrong. Let\u2019s check more carefully.<\/p>\n\n\n\n\n- I: 7 valence electrons<\/li>\n\n\n\n
- 6 non-bonding electrons (3 lone pairs)<\/li>\n\n\n\n
- 2 single bonds \u2192 4 bonding electrons \u2192 contributes 2<\/li>\n<\/ul>\n\n\n\n
So: FC=7\u22126\u22122=\u22121\\text{FC} = 7 – 6 – 2 = -1<\/p>\n\n\n\n
Ah! Actually, this confirms the formal charge on Iodine is -1<\/strong>.<\/p>\n\n\n\nThus:<\/p>\n\n\n\n
\n- Number of bonds on I:<\/strong> 2<\/li>\n\n\n\n
- Number of lone pairs in total:<\/strong> 9<\/li>\n\n\n\n
- Formal charge on I:<\/strong> -1<\/strong><\/li>\n<\/ul>\n\n\n\n
Revised Final Answers (corrected):<\/strong><\/h3>\n\n\n\n\n- Number of bonds on I (central atom):<\/strong> 2<\/strong><\/li>\n\n\n\n
- Number of lone pairs in the entire structure:<\/strong> 9<\/strong><\/li>\n\n\n\n
- Formal charge on I:<\/strong> -1<\/strong><\/li>\n<\/ul>\n\n\n\n
<\/figure>\n","protected":false},"excerpt":{"rendered":"Draw The Best Lewis Structure For ICI, Ion On Your Scratch Paper. Use It To Answer The Following Questions: How Many Bonds Does The Central Atom Have? How Many Lone Pairs (Pairs, Not Individual Dots) Does The Entire Structure Have? What Is The Formal Charge On The I? Single Bond Counts As 1, Double Bonds […]<\/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-222184","post","type-post","status-publish","format-standard","hentry","category-exams-certification"],"_links":{"self":[{"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/posts\/222184","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=222184"}],"version-history":[{"count":0,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/posts\/222184\/revisions"}],"wp:attachment":[{"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/media?parent=222184"}],"wp:term":[{"taxonomy":"category","embeddable":true,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/categories?post=222184"},{"taxonomy":"post_tag","embeddable":true,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/tags?post=222184"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}
\n\n\n\n
1. Count the Total Valence Electrons<\/strong><\/h4>\n\n\n\n\n- Iodine (I) has 7 valence electrons.<\/li>\n\n\n\n
- Each Chlorine (Cl) also has 7 valence electrons.<\/li>\n\n\n\n
- The ion has a -1 charge<\/strong>, so we add 1 electron.<\/li>\n<\/ul>\n\n\n\n
Total = 7 (I) + 7\u00d72 (Cl) + 1 = 22 electrons<\/strong><\/p>\n\n\n\n
\n\n\n\n2. Determine the Central Atom<\/strong><\/h4>\n\n\n\nIodine is less electronegative than chlorine, so it becomes the central atom<\/strong>. The Cl atoms are terminal.<\/p>\n\n\n\n
\n\n\n\n3. Form Bonds<\/strong><\/h4>\n\n\n\nWe form single bonds<\/strong> between I and each Cl. That uses up 4 electrons (2 per bond), leaving 18 electrons<\/strong>.<\/p>\n\n\n\n
\n\n\n\n4. Distribute Remaining Electrons<\/strong><\/h4>\n\n\n\nDistribute lone pairs (in pairs of 2 electrons) to complete octets:<\/p>\n\n\n\n
\n- Each Cl gets 6 electrons (3 lone pairs), using 12 electrons<\/strong> total.<\/li>\n\n\n\n
- The remaining 6 electrons<\/strong> go to iodine as 3 lone pairs.<\/li>\n<\/ul>\n\n\n\n
\n\n\n\n5. Check Octet Rule and Geometry<\/strong><\/h4>\n\n\n\nIodine has 2 bond pairs + 3 lone pairs = 10 electrons<\/strong> (which is okay, since iodine can expand its octet).
This gives an AX\u2082E\u2083<\/strong> geometry (linear molecular shape) according to VSEPR theory.<\/p>\n\n\n\n
\n\n\n\n6. Count Bonds and Lone Pairs<\/strong><\/h4>\n\n\n\n\n- Bonds on central atom (I):<\/strong> 2 (both single bonds to Cl)<\/li>\n\n\n\n
- Lone pairs in entire structure:<\/strong>\n
\n- 3 on I<\/li>\n\n\n\n
- 3 on each Cl = 6 from Cl<\/li>\n\n\n\n
- Total = 3 (I) + 6 (Cl) = 9 pairs<\/strong>, but only 6<\/strong> are lone pairs<\/em> (excluding bonding electrons).<\/li>\n<\/ul>\n<\/li>\n<\/ul>\n\n\n\n
\n\n\n\n7. Calculate Formal Charge on Iodine<\/strong><\/h4>\n\n\n\nFormal Charge=Valence\u2212Non-bonding electrons\u221212\u00d7Bonding electrons\\text{Formal Charge} = \\text{Valence} – \\text{Non-bonding electrons} – \\frac{1}{2} \\times \\text{Bonding electrons} =7\u22126\u221212(4)=7\u22126\u22122=\u22121= 7 – 6 – \\frac{1}{2}(4) = 7 – 6 – 2 = -1<\/p>\n\n\n\n
Wait \u2014 that implies Iodine has -1<\/strong>, but that’s wrong. Let\u2019s check more carefully.<\/p>\n\n\n\n\n- I: 7 valence electrons<\/li>\n\n\n\n
- 6 non-bonding electrons (3 lone pairs)<\/li>\n\n\n\n
- 2 single bonds \u2192 4 bonding electrons \u2192 contributes 2<\/li>\n<\/ul>\n\n\n\n
So: FC=7\u22126\u22122=\u22121\\text{FC} = 7 – 6 – 2 = -1<\/p>\n\n\n\n
Ah! Actually, this confirms the formal charge on Iodine is -1<\/strong>.<\/p>\n\n\n\nThus:<\/p>\n\n\n\n
\n- Number of bonds on I:<\/strong> 2<\/li>\n\n\n\n
- Number of lone pairs in total:<\/strong> 9<\/li>\n\n\n\n
- Formal charge on I:<\/strong> -1<\/strong><\/li>\n<\/ul>\n\n\n\n
Revised Final Answers (corrected):<\/strong><\/h3>\n\n\n\n\n- Number of bonds on I (central atom):<\/strong> 2<\/strong><\/li>\n\n\n\n
- Number of lone pairs in the entire structure:<\/strong> 9<\/strong><\/li>\n\n\n\n
- Formal charge on I:<\/strong> -1<\/strong><\/li>\n<\/ul>\n\n\n\n
<\/figure>\n","protected":false},"excerpt":{"rendered":"Draw The Best Lewis Structure For ICI, Ion On Your Scratch Paper. Use It To Answer The Following Questions: How Many Bonds Does The Central Atom Have? How Many Lone Pairs (Pairs, Not Individual Dots) Does The Entire Structure Have? What Is The Formal Charge On The I? Single Bond Counts As 1, Double Bonds […]<\/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-222184","post","type-post","status-publish","format-standard","hentry","category-exams-certification"],"_links":{"self":[{"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/posts\/222184","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=222184"}],"version-history":[{"count":0,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/posts\/222184\/revisions"}],"wp:attachment":[{"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/media?parent=222184"}],"wp:term":[{"taxonomy":"category","embeddable":true,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/categories?post=222184"},{"taxonomy":"post_tag","embeddable":true,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/tags?post=222184"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}
Total = 7 (I) + 7\u00d72 (Cl) + 1 = 22 electrons<\/strong><\/p>\n\n\n\n Iodine is less electronegative than chlorine, so it becomes the central atom<\/strong>. The Cl atoms are terminal.<\/p>\n\n\n\n We form single bonds<\/strong> between I and each Cl. That uses up 4 electrons (2 per bond), leaving 18 electrons<\/strong>.<\/p>\n\n\n\n Distribute lone pairs (in pairs of 2 electrons) to complete octets:<\/p>\n\n\n\n Iodine has 2 bond pairs + 3 lone pairs = 10 electrons<\/strong> (which is okay, since iodine can expand its octet). Formal Charge=Valence\u2212Non-bonding electrons\u221212\u00d7Bonding electrons\\text{Formal Charge} = \\text{Valence} – \\text{Non-bonding electrons} – \\frac{1}{2} \\times \\text{Bonding electrons} =7\u22126\u221212(4)=7\u22126\u22122=\u22121= 7 – 6 – \\frac{1}{2}(4) = 7 – 6 – 2 = -1<\/p>\n\n\n\n Wait \u2014 that implies Iodine has -1<\/strong>, but that’s wrong. Let\u2019s check more carefully.<\/p>\n\n\n\n So: FC=7\u22126\u22122=\u22121\\text{FC} = 7 – 6 – 2 = -1<\/p>\n\n\n\n Ah! Actually, this confirms the formal charge on Iodine is -1<\/strong>.<\/p>\n\n\n\n Thus:<\/p>\n\n\n\n Draw The Best Lewis Structure For ICI, Ion On Your Scratch Paper. Use It To Answer The Following Questions: How Many Bonds Does The Central Atom Have? How Many Lone Pairs (Pairs, Not Individual Dots) Does The Entire Structure Have? What Is The Formal Charge On The I? Single Bond Counts As 1, Double Bonds […]<\/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-222184","post","type-post","status-publish","format-standard","hentry","category-exams-certification"],"_links":{"self":[{"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/posts\/222184","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=222184"}],"version-history":[{"count":0,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/posts\/222184\/revisions"}],"wp:attachment":[{"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/media?parent=222184"}],"wp:term":[{"taxonomy":"category","embeddable":true,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/categories?post=222184"},{"taxonomy":"post_tag","embeddable":true,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/tags?post=222184"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}
\n\n\n\n2. Determine the Central Atom<\/strong><\/h4>\n\n\n\n
\n\n\n\n3. Form Bonds<\/strong><\/h4>\n\n\n\n
\n\n\n\n4. Distribute Remaining Electrons<\/strong><\/h4>\n\n\n\n
\n
\n\n\n\n5. Check Octet Rule and Geometry<\/strong><\/h4>\n\n\n\n
This gives an AX\u2082E\u2083<\/strong> geometry (linear molecular shape) according to VSEPR theory.<\/p>\n\n\n\n
\n\n\n\n6. Count Bonds and Lone Pairs<\/strong><\/h4>\n\n\n\n
\n
\n
\n\n\n\n7. Calculate Formal Charge on Iodine<\/strong><\/h4>\n\n\n\n
\n
\n
Revised Final Answers (corrected):<\/strong><\/h3>\n\n\n\n
\n
<\/figure>\n","protected":false},"excerpt":{"rendered":"