Enough of a monoprotic acid is dissolved in water to produce a 0.0192 M solution. The pH of the resulting solution is 2.53. Calculate the Ka for the acid.<\/p>\n\n\n\n
The Correct Answer and Explanation is:<\/mark><\/strong><\/p>\n\n\n\n To calculate the acid dissociation constant (Ka)<\/strong> for a monoprotic weak acid<\/strong>, we follow these steps:<\/p>\n\n\n\n The pH is given as 2.53<\/strong>. Use the formula: [H+]=10\u2212pH=10\u22122.53\u22482.95\u00d710\u22123\u2009M[\\text{H}^+] = 10^{-\\text{pH}} = 10^{-2.53} \\approx 2.95 \\times 10^{-3} \\, \\text{M}<\/p>\n\n\n\n For a monoprotic weak acid HA<\/strong>, the dissociation in water is: HA\u21ccH++A\u2212\\text{HA} \\rightleftharpoons \\text{H}^+ + \\text{A}^-<\/p>\n\n\n\n Ka=[H+][A\u2212][HA]K_a = \\frac{[\\text{H}^+][\\text{A}^-]}{[\\text{HA}]}<\/p>\n\n\n\n At equilibrium:<\/p>\n\n\n\n Now plug in: Ka=(2.95\u00d710\u22123)20.01625=8.70\u00d710\u221260.01625\u22485.35\u00d710\u22124K_a = \\frac{(2.95 \\times 10^{-3})^2}{0.01625} = \\frac{8.70 \\times 10^{-6}}{0.01625} \\approx 5.35 \\times 10^{-4}<\/p>\n\n\n\n Ka\u22485.35\u00d710\u22124K_a \\approx \\boxed{5.35 \\times 10^{-4}}<\/p>\n\n\n\n To find the acid dissociation constant (Ka)<\/strong> of a weak acid, we rely on the relationship between pH and hydrogen ion concentration. In this case, the acid is monoprotic<\/strong>, meaning it donates one proton (H\u207a)<\/strong> per molecule when dissolved in water. The solution has a concentration of 0.0192 M<\/strong>, and a pH of 2.53<\/strong>.<\/p>\n\n\n\n First, we determine how many hydrogen ions are in solution. Since pH is the negative logarithm of the hydrogen ion concentration, we reverse this using the formula [H+]=10\u2212pH[H^+] = 10^{-\\text{pH}}. This gives us the concentration of hydrogen ions as approximately 2.95 \u00d7 10\u207b\u00b3 M<\/strong>.<\/p>\n\n\n\n Using an ICE table<\/strong>, we track how the weak acid (HA) dissociates in water. Initially, we have 0.0192 M of HA and 0 M of H\u207a and A\u207b. At equilibrium, HA has lost some concentration (x), and equal amounts of H\u207a and A\u207b have been formed. Since the acid is monoprotic, the increase in H\u207a is the same as the increase in A\u207b, and is equal to the amount dissociated from HA.<\/p>\n\n\n\n We then substitute the equilibrium concentrations into the Ka expression: [H+][A\u2212][HA]\\frac{[H^+][A^-]}{[HA]}. Plugging in our values, we square the [H\u207a] and divide by the remaining [HA]. This calculation gives us the Ka as approximately 5.35 \u00d7 10\u207b\u2074<\/strong>, which reflects the acid\u2019s strength: a relatively small value, indicating a weak acid<\/strong> that does not dissociate completely in water.<\/p>\n\n\n\n This process illustrates the fundamental principles of chemical equilibrium<\/strong>, pH<\/strong>, and weak acid behavior<\/strong>, all essential for understanding acid-base chemistry in biology, medicine, and environmental science.<\/p>\n","protected":false},"excerpt":{"rendered":" Enough of a monoprotic acid is dissolved in water to produce a 0.0192 M solution. The pH of the resulting solution is 2.53. Calculate the Ka for the acid. The Correct Answer and Explanation is: To calculate the acid dissociation constant (Ka) for a monoprotic weak acid, we follow these steps: Step 1: Use the […]<\/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-216734","post","type-post","status-publish","format-standard","hentry","category-exams-certification"],"_links":{"self":[{"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/posts\/216734","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=216734"}],"version-history":[{"count":0,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/posts\/216734\/revisions"}],"wp:attachment":[{"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/media?parent=216734"}],"wp:term":[{"taxonomy":"category","embeddable":true,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/categories?post=216734"},{"taxonomy":"post_tag","embeddable":true,"href":"https:\/\/www.learnexams.com\/blog\/wp-json\/wp\/v2\/tags?post=216734"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}
\n\n\n\nStep 1: Use the pH to find [H\u207a]<\/strong><\/h3>\n\n\n\n
\n\n\n\nStep 2: Set up an ICE Table<\/strong><\/h3>\n\n\n\n
Species<\/th> Initial (M)<\/th> Change (M)<\/th> Equilibrium (M)<\/th><\/tr><\/thead> HA<\/td> 0.0192<\/td> \u2013x<\/td> 0.0192 \u2013 x<\/td><\/tr> H\u207a<\/td> 0<\/td> +x<\/td> x (we found this to be 2.95 \u00d7 10\u207b\u00b3)<\/td><\/tr> A\u207b<\/td> 0<\/td> +x<\/td> x (same as [H\u207a])<\/td><\/tr><\/tbody><\/table><\/figure>\n\n\n\n
\n\n\n\nStep 3: Plug into the Ka expression<\/strong><\/h3>\n\n\n\n
\n
\n\n\n\n\u2705 Final Answer:<\/strong><\/h3>\n\n\n\n
\n\n\n\n\ud83e\udde0 Explanation (300+ words)<\/strong><\/h3>\n\n\n\n