To find the pH of a vitamin C solution at equilibrium, we first need to understand the dissociation of ascorbic acid (C6H8O6) in water. Ascorbic acid is a weak acid and dissociates according to the following equilibrium reaction:C6H8O6 (aq)\u21ccH+ (aq)+C6H7O6- (aq)\\text{C6H8O6 (aq)} \\rightleftharpoons \\text{H+ (aq)} + \\text{C6H7O6- (aq)}C6H8O6 (aq)\u21ccH+ (aq)+C6H7O6- (aq)<\/p>\n\n\n\n
The equilibrium constant expression for this dissociation is given by:Ka=[H+][C6H7O6-][C6H8O6]K_a = \\frac{[\\text{H+}][\\text{C6H7O6-}]}{[\\text{C6H8O6}]}Ka\u200b=[C6H8O6][H+][C6H7O6-]\u200b<\/p>\n\n\n\n
Step 1: Convert the given mass of vitamin C to moles<\/h3>\n\n\n\n
We are given 15 g of vitamin C dissolved in 1.00 L of water. To find the number of moles of vitamin C, we use its molar mass. The molar mass of ascorbic acid (C6H8O6) is:M=6(12.01)+8(1.008)+6(16.00)=176.12\u2009g\/molM = 6(12.01) + 8(1.008) + 6(16.00) = 176.12 \\, \\text{g\/mol}M=6(12.01)+8(1.008)+6(16.00)=176.12g\/mol<\/p>\n\n\n\n
Now, calculate the number of moles of vitamin C:Moles of C6H8O6=15\u2009g176.12\u2009g\/mol=0.0852\u2009mol\\text{Moles of C6H8O6} = \\frac{15 \\, \\text{g}}{176.12 \\, \\text{g\/mol}} = 0.0852 \\, \\text{mol}Moles of C6H8O6=176.12g\/mol15g\u200b=0.0852mol<\/p>\n\n\n\n
Step 2: Write the ICE table for the dissociation<\/h3>\n\n\n\n
Let’s assume that the acid dissociates to a small extent. Let xxx be the concentration of H+ ions that dissociate from the acid.C6H8O6H+C6H7O6-Initial concentration0.0852\u2009M00Change\u2212x+x+xEquilibrium concentration0.0852\u2212xxx\\begin{array}{|c|c|c|c|} \\hline & \\text{C6H8O6} & \\text{H+} & \\text{C6H7O6-} \\\\ \\hline \\text{Initial concentration} & 0.0852 \\, \\text{M} & 0 & 0 \\\\ \\text{Change} & -x & +x & +x \\\\ \\text{Equilibrium concentration} & 0.0852 – x & x & x \\\\ \\hline \\end{array}Initial concentrationChangeEquilibrium concentration\u200bC6H8O60.0852M\u2212x0.0852\u2212x\u200bH+0+xx\u200bC6H7O6-0+xx\u200b\u200b<\/p>\n\n\n\n
Step 3: Set up the equilibrium expression<\/h3>\n\n\n\n
Using the Ka value of 8.0 x 10^-5, we can substitute the equilibrium concentrations into the Ka expression:Ka=x\u22c5x0.0852\u2212x=8.0\u00d710\u22125K_a = \\frac{x \\cdot x}{0.0852 – x} = 8.0 \\times 10^{-5}Ka\u200b=0.0852\u2212xx\u22c5x\u200b=8.0\u00d710\u22125<\/p>\n\n\n\n
This simplifies to:x20.0852\u2212x=8.0\u00d710\u22125\\frac{x^2}{0.0852 – x} = 8.0 \\times 10^{-5}0.0852\u2212xx2\u200b=8.0\u00d710\u22125<\/p>\n\n\n\n
For weak acids, the value of xxx (H+ concentration) will be small compared to the initial concentration of the acid. So, we can approximate 0.0852\u2212x\u22480.08520.0852 – x \\approx 0.08520.0852\u2212x\u22480.0852, simplifying the equation to:x20.0852=8.0\u00d710\u22125\\frac{x^2}{0.0852} = 8.0 \\times 10^{-5}0.0852×2\u200b=8.0\u00d710\u22125<\/p>\n\n\n\n
Now, solve for xxx:x2=(8.0\u00d710\u22125)\u00d70.0852=6.816\u00d710\u22126x^2 = (8.0 \\times 10^{-5}) \\times 0.0852 = 6.816 \\times 10^{-6}x2=(8.0\u00d710\u22125)\u00d70.0852=6.816\u00d710\u22126x=6.816\u00d710\u22126=2.61\u00d710\u22123\u2009Mx = \\sqrt{6.816 \\times 10^{-6}} = 2.61 \\times 10^{-3} \\, \\text{M}x=6.816\u00d710\u22126\u200b=2.61\u00d710\u22123M<\/p>\n\n\n\n
Step 4: Calculate the pH<\/h3>\n\n\n\n
The concentration of H+ ions is 2.61\u00d710\u22123\u2009M2.61 \\times 10^{-3} \\, \\text{M}2.61\u00d710\u22123M. The pH is given by:pH=\u2212log\u2061[H+]\\text{pH} = -\\log[\\text{H+}]pH=\u2212log[H+]<\/p>\n\n\n\n
Substituting the value of xxx:pH=\u2212log\u2061(2.61\u00d710\u22123)=2.58\\text{pH} = -\\log(2.61 \\times 10^{-3}) = 2.58pH=\u2212log(2.61\u00d710\u22123)=2.58<\/p>\n\n\n\n
Final Answer<\/h3>\n\n\n\n
The pH of the vitamin C solution at equilibrium is 2.58<\/strong>.<\/p>\n\n\n\n The Ka for vitamin C (also known as ascorbic acid, C6H8O6) is 8.0 x 10^-5. If 15 g of vitamin C is dissolved in 1.00 L of water, what is the pH of the solution at equilibrium The Correct Answer and Explanation is: To find the pH of a vitamin C solution at equilibrium, we […]<\/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-241943","post","type-post","status-publish","format-standard","hentry","category-exams-certification"],"_links":{"self":[{"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/posts\/241943","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=241943"}],"version-history":[{"count":0,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/posts\/241943\/revisions"}],"wp:attachment":[{"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/media?parent=241943"}],"wp:term":[{"taxonomy":"category","embeddable":true,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/categories?post=241943"},{"taxonomy":"post_tag","embeddable":true,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/tags?post=241943"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}![\"\"](\"https:\/\/learnexams.com\/blog\/wp-content\/uploads\/2025\/07\/learnexams-banner5-193.jpeg\")
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