{"id":7691,"date":"2024-12-05T16:18:04","date_gmt":"2024-12-05T21:18:04","guid":{"rendered":"https:\/\/pressbooks.bccampus.ca\/pathology\/chapter\/acid-base-balance-overview\/"},"modified":"2025-11-11T23:44:49","modified_gmt":"2025-11-12T04:44:49","slug":"acid-base-balance-overview","status":"publish","type":"chapter","link":"https:\/\/pressbooks.bccampus.ca\/pathology\/chapter\/acid-base-balance-overview\/","title":{"raw":"What is pH? How does the Body Regulate pH Changes?","rendered":"What is pH? How does the Body Regulate pH Changes?"},"content":{"raw":"<div class=\"textbox textbox--learning-objectives\"><header class=\"textbox__header\">\r\n<p class=\"textbox__title\">Learning Objectives<\/p>\r\n\r\n<\/header>\r\n<div class=\"textbox__content\">\r\n\r\nBy the end of this section, you will be able to:\r\n<ul>\r\n \t<li>Identify abnormal pH ranges.<\/li>\r\n \t<li>Understand the basic factors which control and drive pH in the body.<\/li>\r\n \t<li>Describe the Blood Buffer System and the role of each of its parts.<\/li>\r\n \t<li>Calculate pH.<\/li>\r\n<\/ul>\r\n<\/div>\r\n<\/div>\r\n<h2>The Biochemistry of Blood pH<\/h2>\r\nTo understand how blood pH can change, one needs to understand the basic biochemistry that most of our cells undergo.\u00a0 This is particularly true for glucose (sugar) metabolism as this is the primary nutrient for our brain and other tissues.\r\n\r\nThe body metabolizes glucose sugar (C<sub>6<\/sub>H<sub>12<\/sub>O<sub>6<\/sub>) in an oxygenated environment\u00a0 by the following reaction.\r\n<div class=\"textbox shaded\">\r\n\r\nC<sub>6<\/sub>H<sub>12<\/sub>O<sub>6<\/sub> + 6 O<sub>2<\/sub> -&gt; 6 CO<sub>2<\/sub> + 6 H<sub>2<\/sub>O + Energy (ATP)\r\n\r\n<\/div>\r\nThe major byproducts are ATP, water, and carbon dioxide (CO<sub>2<\/sub>). When CO<sub>2<\/sub> dissolves into water it forms carbonic acid, raising the acidity of the blood. As you may be familiar from current conversations about global climate change which is the reaction largely responsible for the acidification of the world\u2019s oceans.\r\n\r\nIn the body the enzyme carbonic anhydrase catalyzes this reaction.\r\n<div class=\"textbox shaded\">\r\n\r\nH<sub>2<\/sub>O + CO<sub>2<\/sub> -&gt; H<sub>2<\/sub>CO<sub>3<\/sub>\r\n\r\n<\/div>\r\nWhen the hydrogen ion (H<sup>+<\/sup>) dissociates from carbonic acid (H<sub>2<\/sub>CO<sub>3<\/sub>), carbonic acid dissociates into bicarbonate (HCO<sub>3<\/sub><sup>-<\/sup>)\r\n<div class=\"textbox shaded\">\r\n\r\nH<sub>2<\/sub>CO<sub>3<\/sub> -&gt; H<sup>+<\/sup> + HCO<sub>3<\/sub><sup>-<\/sup>\r\n\r\n<\/div>\r\nthe following buffer system is responsible for maintaining a pH of 7.35 - 7.45 in the human body <sup>[3]<\/sup> <sup>[4]<\/sup>\r\n<div class=\"textbox shaded\">\r\n\r\nCO<sub>2<\/sub> + H<sub>2<\/sub>O &lt;-&gt; H<sub>2<\/sub>CO<sub>3<\/sub> &lt;-&gt; H<sup>+<\/sup> + HCO3<sup>-<\/sup>\r\n\r\n<\/div>\r\nShifts in the number of reactants will shift the number of products towards the maintenance of equilibria. Acids typically exist in equilibria in any aqueous solution. The amount of protons (H<sup>+<\/sup>) that disassociate from a certain concentration of acid dictate the strength of that acid. We describe this relationship with the introduction of the acid disassociation constant K<sub>a<\/sub>.\r\n\r\nK<sub>a<\/sub> = [H<sup>+<\/sup>][Conjugate Base] \/ [Acid]\r\n\r\nIf we want to relate K<sub>a<\/sub> to pH we must first remember\r\n\r\npH = -log[H<sup>+<\/sup>]\r\n<div class=\"textbox textbox--examples\"><header class=\"textbox__header\">\r\n<p class=\"textbox__title\">Relationship between Ka and pH<\/p>\r\n\r\n<\/header>\r\n<div class=\"textbox__content\">\r\n\r\nKa = [H<sup>+<\/sup>]*[Conjugate Base] \/ [Acid]\r\n\r\n-log[Ka] = -log[H<sup>+<\/sup>] * -log ([Conjugate Base]\/ [Acid])\r\n\r\npKa = pH * -log ([Conjugate Base]\/ [Acid])\r\n\r\npH = pKa * -log ([Conjugate Base]\/ [Acid])\r\n\r\n<\/div>\r\n<\/div>\r\nThe final orientation of this equation is known as the Henderson Hasselbalch equation for pH. In our blood buffer system technically the conjugate base is HCO<sub>3<\/sub><sup>-<\/sup> and the acid is H<sub>2<\/sub>CO<sub>3<\/sub>, however because H<sub>2<\/sub>CO<sub>3<\/sub> levels are directly correlated to CO<sub>2<\/sub> levels and we measure CO<sub>2<\/sub> in the blood, we can simply use our equilbiria into the following equation.\r\n<div class=\"textbox shaded\">\r\n\r\nCO<sub>2<\/sub> &lt;-&gt; HCO<sub>3<\/sub><sup>-<\/sup>\r\n\r\n<\/div>\r\nTherefore the modified Henderson Hasselbalch equation to calculate blood pH is:\r\n<div class=\"textbox shaded\">\r\n\r\npH = pKa+ log ([HCO<sub>3<\/sub><sup>-<\/sup>]\/ [CO<sub>2<\/sub>])\r\n\r\n<\/div>\r\nThe negative log of the acid disassociation constant (pKa) for this system at internal body temperature is\r\n<div class=\"textbox shaded\">\r\n\r\npKa =6.1\r\n\r\nModified Henderson Hasselbalch\r\npH = 6.1 + log ([HCO<sub>3<\/sub><sup>-<\/sup>]\/ [CO<sub>2<\/sub>]) <sup>[2]<\/sup>\r\n\r\nThough pKa value for this system is 6.1, the pKa value for blood itself is closer to 7.4 which is of course, equivalent to normal pH. This is achieved mainly due to the action of carbonic anhydrase, the existing CO<sub>2<\/sub> stores in the body, and other factors. <sup>[1]<\/sup>\r\n\r\n<\/div>\r\n<div class=\"textbox textbox--key-takeaways\"><header class=\"textbox__header\">\r\n<p class=\"textbox__title\">Key Takeaways<\/p>\r\n\r\n<\/header>\r\n<ul>\r\n \t<li>It is easy to see how changes in concentration of bicarbonate HCO<sub>3<\/sub><sup>-<\/sup> and CO<sub>2<\/sub> will move the equilibria of this equation to express different concentrations of protons.<\/li>\r\n \t<li>Try playing around with the equation and see how variable concentrations affect blood pH.<\/li>\r\n<\/ul>\r\n<\/div>\r\n&nbsp;\r\n<div class=\"textbox textbox--examples\"><header class=\"textbox__header\">\r\n<p class=\"textbox__title\">Exercise: Calculating pH<\/p>\r\n\r\n<\/header>\r\n<div class=\"textbox__content\">\r\n\r\npH = 7.36\r\n\r\nPaO<sub>2<\/sub> = 60mmHg\r\nPaCO<sub>2<\/sub> = 35mmHg\r\n\r\nHCO<sub>3<\/sub> = 22mEq\/L\r\n\r\nSaO<sub>2<\/sub> = 90%\r\n\r\npH is the negative log of the concentrations of protons in the blood sample\r\n\r\nPaO<sub>2<\/sub> is the partial pressure of oxygen in the sample\r\nPaCO<sub>2<\/sub> is the partial pressure of CO<sub>2<\/sub> in the sample\r\nHCO<sub>3 <\/sub>is the milliequivilants per litre of bicarbonate in the sample\r\n\r\nSaO<sub>2<\/sub> is the percentage of oxygen saturated haemoglobin relative to total haemoglobin\r\n\r\nAt normal conditions\r\n\r\n[CO<sub>2<\/sub>] = 35 mmHg * 0.03 = 1.2 mmol\/L\r\n\r\n*Venous blood gases (VBG) read CO<sub>2<\/sub> in mmHg, the solubility coefficient is 0.03 (Therefore 1 mmHg CO<sub>2<\/sub> will dissolve into 0.03 mmol\/L of CO<sub>2<\/sub> in the blood)\r\n\r\n[HCO<sub>3<\/sub>] = 22 mmol\/L\r\n\r\n&nbsp;\r\n\r\nLet's use our modified Henderson Hasselbalch equation to calculate the pH using this VBG.\r\n\r\nHenderson Hasselbalch\r\npH = pKa + log ([HCO<sub>3<\/sub>]\/ [CO<sub>2<\/sub>])\r\n\r\npKa = 6.1 (calculated value)\r\n\r\npH = 6.1 + log(22\/1.2)\r\n\r\npH = 7.36\r\n\r\n&nbsp;\r\n\r\nIf we wanted to check our work using the formula for the acid disassociation constant\r\n\r\npKa = -log10(Ka)\r\n\r\n6.1 = -log(ka)\r\n\r\n10<sup>-6.1<\/sup> = Ka\r\n\r\nKa = 7.94*10<sup>-7<\/sup>\r\n\r\nKa = [H<sup>+<\/sup>][HCO<sub>3<\/sub><sup>-<\/sup>] \/ [CO<sub>2<\/sub>]\r\n\r\n7.94 * 10<sup>-7<\/sup> = 22\/1.2*[H<sup>+<\/sup>]\r\n\r\n[H<sup>+<\/sup>] = 4.33 10<sup>-8<\/sup>\r\n\r\nPh = -log[H<sup>+<\/sup>]\r\n\r\nPh = -log(4.33*10<sup>-8<\/sup>)\r\n\r\nPh = 7.36\r\n\r\n<\/div>\r\n<\/div>\r\n&nbsp;\r\n\r\nThe body has two major compensatory mechanisms for managing this buffer system: ventilation, and renal regulation.\r\n<h2>Ventilatory Regulation<\/h2>\r\nRespiratory rate determines how much CO<sub>2<\/sub> is exhaled, thus removing it from the blood. As mentioned CO<sub>2 <\/sub>+ water is easily transformed into carbonic acid H<sub>2<\/sub>CO<sub>3<\/sub>. Thus, excess CO<sub>2<\/sub> in the blood means there is more 'acid' in the blood, leading to a lowering of pH. This occurs with a decreased respiratory rate (hypoventilation) which results in retention of CO<sub>2<\/sub>. Conversely, excessive ventilation (hyperventilation) will remove CO<sub>2<\/sub> form the body meaning a loss of H<sup>+<\/sup>. This will result in a raise in pH making it more basic. <sup>[4]<\/sup>\r\n<h2>Renal Regulation<\/h2>\r\nSerum levels of bicarbonate are maintained by kidney function. When serum levels are high, the kidneys excrete additional bicarbonate when serum levels are low the kidneys retain bicarbonate, or produce it.\r\n\r\nWe refer to an undesirable shift in blood pH due to the management of Bicarbonate as metabolic acidosis, because it is managed by internal metabolic regulatory mechanisms. The kidneys are also responsible for proton excretion, by increasing or decreasing the number of protons excreted the kidneys will shift the blood pH, dysfunction of these mechanisms can also lead to metabolic acidosis\/alkalosis. <sup>[4]<\/sup>\r\n<h2>Other Systems of pH Regulation<\/h2>\r\nHemoglobin can bind hydrogen ions which is an additional buffering mechanism that can cause some undesirable effects discussed in our respiratory alkalosis chapter. The phosphate buffer system contains bases which accept hydrogen ions H<sup>+<\/sup> and important for the regulation of urine pH. <sup>[4]<\/sup>\r\n<h1>Review Questions<\/h1>\r\n[h5p id=\"441\"]\r\n<h1>References<\/h1>\r\n1: https:\/\/www.ncbi.nlm.nih.gov\/books\/NBK526028\/\r\n2: <a href=\"https:\/\/doi.org\/10.1007\/978-3-031-25810-7\">Fluid, Electrolyte and Acid-Base Disorders: Clinical Evaluation and Management | SpringerLink<\/a>\r\n3: https:\/\/view.officeapps.live.com\/op\/view.aspx?src=https%3A%2F%2Fwww.asep.org%2Fasep%2Fasep%2FBloodAcid-BaseBuffering.doc&amp;wdOrigin=BROWSELINK\r\n4: :<a href=\"https:\/\/www.ncbi.nlm.nih.gov\/books\/NBK507807\/\">https:\/\/www.ncbi.nlm.nih.gov\/books\/NBK507807\/<\/a>","rendered":"<div class=\"textbox textbox--learning-objectives\">\n<header class=\"textbox__header\">\n<p class=\"textbox__title\">Learning Objectives<\/p>\n<\/header>\n<div class=\"textbox__content\">\n<p>By the end of this section, you will be able to:<\/p>\n<ul>\n<li>Identify abnormal pH ranges.<\/li>\n<li>Understand the basic factors which control and drive pH in the body.<\/li>\n<li>Describe the Blood Buffer System and the role of each of its parts.<\/li>\n<li>Calculate pH.<\/li>\n<\/ul>\n<\/div>\n<\/div>\n<h2>The Biochemistry of Blood pH<\/h2>\n<p>To understand how blood pH can change, one needs to understand the basic biochemistry that most of our cells undergo.\u00a0 This is particularly true for glucose (sugar) metabolism as this is the primary nutrient for our brain and other tissues.<\/p>\n<p>The body metabolizes glucose sugar (C<sub>6<\/sub>H<sub>12<\/sub>O<sub>6<\/sub>) in an oxygenated environment\u00a0 by the following reaction.<\/p>\n<div class=\"textbox shaded\">\n<p>C<sub>6<\/sub>H<sub>12<\/sub>O<sub>6<\/sub> + 6 O<sub>2<\/sub> -&gt; 6 CO<sub>2<\/sub> + 6 H<sub>2<\/sub>O + Energy (ATP)<\/p>\n<\/div>\n<p>The major byproducts are ATP, water, and carbon dioxide (CO<sub>2<\/sub>). When CO<sub>2<\/sub> dissolves into water it forms carbonic acid, raising the acidity of the blood. As you may be familiar from current conversations about global climate change which is the reaction largely responsible for the acidification of the world\u2019s oceans.<\/p>\n<p>In the body the enzyme carbonic anhydrase catalyzes this reaction.<\/p>\n<div class=\"textbox shaded\">\n<p>H<sub>2<\/sub>O + CO<sub>2<\/sub> -&gt; H<sub>2<\/sub>CO<sub>3<\/sub><\/p>\n<\/div>\n<p>When the hydrogen ion (H<sup>+<\/sup>) dissociates from carbonic acid (H<sub>2<\/sub>CO<sub>3<\/sub>), carbonic acid dissociates into bicarbonate (HCO<sub>3<\/sub><sup>&#8211;<\/sup>)<\/p>\n<div class=\"textbox shaded\">\n<p>H<sub>2<\/sub>CO<sub>3<\/sub> -&gt; H<sup>+<\/sup> + HCO<sub>3<\/sub><sup>&#8211;<\/sup><\/p>\n<\/div>\n<p>the following buffer system is responsible for maintaining a pH of 7.35 &#8211; 7.45 in the human body <sup>[3]<\/sup> <sup>[4]<\/sup><\/p>\n<div class=\"textbox shaded\">\n<p>CO<sub>2<\/sub> + H<sub>2<\/sub>O &lt;-&gt; H<sub>2<\/sub>CO<sub>3<\/sub> &lt;-&gt; H<sup>+<\/sup> + HCO3<sup>&#8211;<\/sup><\/p>\n<\/div>\n<p>Shifts in the number of reactants will shift the number of products towards the maintenance of equilibria. Acids typically exist in equilibria in any aqueous solution. The amount of protons (H<sup>+<\/sup>) that disassociate from a certain concentration of acid dictate the strength of that acid. We describe this relationship with the introduction of the acid disassociation constant K<sub>a<\/sub>.<\/p>\n<p>K<sub>a<\/sub> = [H<sup>+<\/sup>][Conjugate Base] \/ [Acid]<\/p>\n<p>If we want to relate K<sub>a<\/sub> to pH we must first remember<\/p>\n<p>pH = -log[H<sup>+<\/sup>]<\/p>\n<div class=\"textbox textbox--examples\">\n<header class=\"textbox__header\">\n<p class=\"textbox__title\">Relationship between Ka and pH<\/p>\n<\/header>\n<div class=\"textbox__content\">\n<p>Ka = [H<sup>+<\/sup>]*[Conjugate Base] \/ [Acid]<\/p>\n<p>-log[Ka] = -log[H<sup>+<\/sup>] * -log ([Conjugate Base]\/ [Acid])<\/p>\n<p>pKa = pH * -log ([Conjugate Base]\/ [Acid])<\/p>\n<p>pH = pKa * -log ([Conjugate Base]\/ [Acid])<\/p>\n<\/div>\n<\/div>\n<p>The final orientation of this equation is known as the Henderson Hasselbalch equation for pH. In our blood buffer system technically the conjugate base is HCO<sub>3<\/sub><sup>&#8211;<\/sup> and the acid is H<sub>2<\/sub>CO<sub>3<\/sub>, however because H<sub>2<\/sub>CO<sub>3<\/sub> levels are directly correlated to CO<sub>2<\/sub> levels and we measure CO<sub>2<\/sub> in the blood, we can simply use our equilbiria into the following equation.<\/p>\n<div class=\"textbox shaded\">\n<p>CO<sub>2<\/sub> &lt;-&gt; HCO<sub>3<\/sub><sup>&#8211;<\/sup><\/p>\n<\/div>\n<p>Therefore the modified Henderson Hasselbalch equation to calculate blood pH is:<\/p>\n<div class=\"textbox shaded\">\n<p>pH = pKa+ log ([HCO<sub>3<\/sub><sup>&#8211;<\/sup>]\/ [CO<sub>2<\/sub>])<\/p>\n<\/div>\n<p>The negative log of the acid disassociation constant (pKa) for this system at internal body temperature is<\/p>\n<div class=\"textbox shaded\">\n<p>pKa =6.1<\/p>\n<p>Modified Henderson Hasselbalch<br \/>\npH = 6.1 + log ([HCO<sub>3<\/sub><sup>&#8211;<\/sup>]\/ [CO<sub>2<\/sub>]) <sup>[2]<\/sup><\/p>\n<p>Though pKa value for this system is 6.1, the pKa value for blood itself is closer to 7.4 which is of course, equivalent to normal pH. This is achieved mainly due to the action of carbonic anhydrase, the existing CO<sub>2<\/sub> stores in the body, and other factors. <sup>[1]<\/sup><\/p>\n<\/div>\n<div class=\"textbox textbox--key-takeaways\">\n<header class=\"textbox__header\">\n<p class=\"textbox__title\">Key Takeaways<\/p>\n<\/header>\n<ul>\n<li>It is easy to see how changes in concentration of bicarbonate HCO<sub>3<\/sub><sup>&#8211;<\/sup> and CO<sub>2<\/sub> will move the equilibria of this equation to express different concentrations of protons.<\/li>\n<li>Try playing around with the equation and see how variable concentrations affect blood pH.<\/li>\n<\/ul>\n<\/div>\n<p>&nbsp;<\/p>\n<div class=\"textbox textbox--examples\">\n<header class=\"textbox__header\">\n<p class=\"textbox__title\">Exercise: Calculating pH<\/p>\n<\/header>\n<div class=\"textbox__content\">\n<p>pH = 7.36<\/p>\n<p>PaO<sub>2<\/sub> = 60mmHg<br \/>\nPaCO<sub>2<\/sub> = 35mmHg<\/p>\n<p>HCO<sub>3<\/sub> = 22mEq\/L<\/p>\n<p>SaO<sub>2<\/sub> = 90%<\/p>\n<p>pH is the negative log of the concentrations of protons in the blood sample<\/p>\n<p>PaO<sub>2<\/sub> is the partial pressure of oxygen in the sample<br \/>\nPaCO<sub>2<\/sub> is the partial pressure of CO<sub>2<\/sub> in the sample<br \/>\nHCO<sub>3 <\/sub>is the milliequivilants per litre of bicarbonate in the sample<\/p>\n<p>SaO<sub>2<\/sub> is the percentage of oxygen saturated haemoglobin relative to total haemoglobin<\/p>\n<p>At normal conditions<\/p>\n<p>[CO<sub>2<\/sub>] = 35 mmHg * 0.03 = 1.2 mmol\/L<\/p>\n<p>*Venous blood gases (VBG) read CO<sub>2<\/sub> in mmHg, the solubility coefficient is 0.03 (Therefore 1 mmHg CO<sub>2<\/sub> will dissolve into 0.03 mmol\/L of CO<sub>2<\/sub> in the blood)<\/p>\n<p>[HCO<sub>3<\/sub>] = 22 mmol\/L<\/p>\n<p>&nbsp;<\/p>\n<p>Let&#8217;s use our modified Henderson Hasselbalch equation to calculate the pH using this VBG.<\/p>\n<p>Henderson Hasselbalch<br \/>\npH = pKa + log ([HCO<sub>3<\/sub>]\/ [CO<sub>2<\/sub>])<\/p>\n<p>pKa = 6.1 (calculated value)<\/p>\n<p>pH = 6.1 + log(22\/1.2)<\/p>\n<p>pH = 7.36<\/p>\n<p>&nbsp;<\/p>\n<p>If we wanted to check our work using the formula for the acid disassociation constant<\/p>\n<p>pKa = -log10(Ka)<\/p>\n<p>6.1 = -log(ka)<\/p>\n<p>10<sup>-6.1<\/sup> = Ka<\/p>\n<p>Ka = 7.94*10<sup>-7<\/sup><\/p>\n<p>Ka = [H<sup>+<\/sup>][HCO<sub>3<\/sub><sup>&#8211;<\/sup>] \/ [CO<sub>2<\/sub>]<\/p>\n<p>7.94 * 10<sup>-7<\/sup> = 22\/1.2*[H<sup>+<\/sup>]<\/p>\n<p>[H<sup>+<\/sup>] = 4.33 10<sup>-8<\/sup><\/p>\n<p>Ph = -log[H<sup>+<\/sup>]<\/p>\n<p>Ph = -log(4.33*10<sup>-8<\/sup>)<\/p>\n<p>Ph = 7.36<\/p>\n<\/div>\n<\/div>\n<p>&nbsp;<\/p>\n<p>The body has two major compensatory mechanisms for managing this buffer system: ventilation, and renal regulation.<\/p>\n<h2>Ventilatory Regulation<\/h2>\n<p>Respiratory rate determines how much CO<sub>2<\/sub> is exhaled, thus removing it from the blood. As mentioned CO<sub>2 <\/sub>+ water is easily transformed into carbonic acid H<sub>2<\/sub>CO<sub>3<\/sub>. Thus, excess CO<sub>2<\/sub> in the blood means there is more &#8216;acid&#8217; in the blood, leading to a lowering of pH. This occurs with a decreased respiratory rate (hypoventilation) which results in retention of CO<sub>2<\/sub>. Conversely, excessive ventilation (hyperventilation) will remove CO<sub>2<\/sub> form the body meaning a loss of H<sup>+<\/sup>. This will result in a raise in pH making it more basic. <sup>[4]<\/sup><\/p>\n<h2>Renal Regulation<\/h2>\n<p>Serum levels of bicarbonate are maintained by kidney function. When serum levels are high, the kidneys excrete additional bicarbonate when serum levels are low the kidneys retain bicarbonate, or produce it.<\/p>\n<p>We refer to an undesirable shift in blood pH due to the management of Bicarbonate as metabolic acidosis, because it is managed by internal metabolic regulatory mechanisms. The kidneys are also responsible for proton excretion, by increasing or decreasing the number of protons excreted the kidneys will shift the blood pH, dysfunction of these mechanisms can also lead to metabolic acidosis\/alkalosis. <sup>[4]<\/sup><\/p>\n<h2>Other Systems of pH Regulation<\/h2>\n<p>Hemoglobin can bind hydrogen ions which is an additional buffering mechanism that can cause some undesirable effects discussed in our respiratory alkalosis chapter. The phosphate buffer system contains bases which accept hydrogen ions H<sup>+<\/sup> and important for the regulation of urine pH. <sup>[4]<\/sup><\/p>\n<h1>Review Questions<\/h1>\n<div id=\"h5p-441\">\n<div class=\"h5p-iframe-wrapper\"><iframe id=\"h5p-iframe-441\" class=\"h5p-iframe\" data-content-id=\"441\" style=\"height:1px\" src=\"about:blank\" frameBorder=\"0\" scrolling=\"no\" title=\"Biochemistry of pH\"><\/iframe><\/div>\n<\/div>\n<h1>References<\/h1>\n<p>1: https:\/\/www.ncbi.nlm.nih.gov\/books\/NBK526028\/<br \/>\n2: <a href=\"https:\/\/doi.org\/10.1007\/978-3-031-25810-7\">Fluid, Electrolyte and Acid-Base Disorders: Clinical Evaluation and Management | SpringerLink<\/a><br \/>\n3: https:\/\/view.officeapps.live.com\/op\/view.aspx?src=https%3A%2F%2Fwww.asep.org%2Fasep%2Fasep%2FBloodAcid-BaseBuffering.doc&amp;wdOrigin=BROWSELINK<br \/>\n4: :<a href=\"https:\/\/www.ncbi.nlm.nih.gov\/books\/NBK507807\/\">https:\/\/www.ncbi.nlm.nih.gov\/books\/NBK507807\/<\/a><\/p>\n","protected":false},"author":1076,"menu_order":3,"template":"","meta":{"pb_show_title":"on","pb_short_title":"","pb_subtitle":"","pb_authors":["callen05-y6watyomw8"],"pb_section_license":""},"chapter-type":[],"contributor":[264],"license":[],"class_list":["post-7691","chapter","type-chapter","status-publish","hentry","contributor-callen05-y6watyomw8"],"part":7690,"_links":{"self":[{"href":"https:\/\/pressbooks.bccampus.ca\/pathology\/wp-json\/pressbooks\/v2\/chapters\/7691","targetHints":{"allow":["GET"]}}],"collection":[{"href":"https:\/\/pressbooks.bccampus.ca\/pathology\/wp-json\/pressbooks\/v2\/chapters"}],"about":[{"href":"https:\/\/pressbooks.bccampus.ca\/pathology\/wp-json\/wp\/v2\/types\/chapter"}],"author":[{"embeddable":true,"href":"https:\/\/pressbooks.bccampus.ca\/pathology\/wp-json\/wp\/v2\/users\/1076"}],"version-history":[{"count":25,"href":"https:\/\/pressbooks.bccampus.ca\/pathology\/wp-json\/pressbooks\/v2\/chapters\/7691\/revisions"}],"predecessor-version":[{"id":9804,"href":"https:\/\/pressbooks.bccampus.ca\/pathology\/wp-json\/pressbooks\/v2\/chapters\/7691\/revisions\/9804"}],"part":[{"href":"https:\/\/pressbooks.bccampus.ca\/pathology\/wp-json\/pressbooks\/v2\/parts\/7690"}],"metadata":[{"href":"https:\/\/pressbooks.bccampus.ca\/pathology\/wp-json\/pressbooks\/v2\/chapters\/7691\/metadata\/"}],"wp:attachment":[{"href":"https:\/\/pressbooks.bccampus.ca\/pathology\/wp-json\/wp\/v2\/media?parent=7691"}],"wp:term":[{"taxonomy":"chapter-type","embeddable":true,"href":"https:\/\/pressbooks.bccampus.ca\/pathology\/wp-json\/pressbooks\/v2\/chapter-type?post=7691"},{"taxonomy":"contributor","embeddable":true,"href":"https:\/\/pressbooks.bccampus.ca\/pathology\/wp-json\/wp\/v2\/contributor?post=7691"},{"taxonomy":"license","embeddable":true,"href":"https:\/\/pressbooks.bccampus.ca\/pathology\/wp-json\/wp\/v2\/license?post=7691"}],"curies":[{"name":"wp","href":"https:\/\/api.w.org\/{rel}","templated":true}]}}