We can find pH by taking the negative log of the hydronium ion concentration, using the expression pH = -log [H3O+]. I asked specifically for HCO3-: "Kb of bicarbonate is greater than Ka?". It's been a long time since I did my chemistry classes and I'm currently trying to analyze groundwater samples for hydrogeology purposes. Calculate \(K_b\) and \(pK_b\) of the butyrate ion (\(CH_3CH_2CH_2CO_2^\)). $$\ce{H2O + H2CO3 <=> H3O+ + HCO3-}$$ Bases accept protons or donate electron pairs. $$Cs = \ce{\frac{[HCO3-][H3O+]}{K1} + [HCO3-] + \frac{K2[HCO3-]}{[H3O+]}}$$ Let's start by writing out the dissociation equation and Ka expression for the acid. What is the Ka of a solution whose known values are given in the table: {eq}pH = -log[H^+]=-logx \rightarrow x = 10^-1.7 = 0.0199 {/eq}, {eq}K_a = (0.0199)^2/0.048 = 8.25*10^-3 {/eq}. The following example shows how to calculate Ka. Weak bases react with water to produce the hydroxide ion, as shown in the following general equation, where B is the parent base and BH+ is its conjugate acid: \[B_{(aq)}+H_2O_{(l)} \rightleftharpoons BH^+_{(aq)}+OH^_{(aq)} \label{16.5.4}\]. $$\frac{\ce{[HCO3-]}}{Cs} = \ce{\frac{K1[H3O+]}{[H3O+]^2 + K1[H3O+] + K1K2}} = \alpha1$$, So we got the expression for $\alpha1$, that has a curious structure: a fraction, where the denominator is a polynomial of degree 2, and the numerator its middle term. First, write the balanced chemical equation. The more A-^\text{-}-start superscript, start text, negative, end text, end superscript and HA molecules available, the less of an effect the addition of a strong acid or base will have on the pH of the solution. Because \(pK_b = \log K_b\), \(K_b\) is \(10^{9.17} = 6.8 \times 10^{10}\). At equilibrium the concentration of protons is equal to 0.00758M. Once again, water is not present. This assignment sounds intimidating at first, but we must remember that pH is really just a measurement of the hydronium ion concentration. As a member, you'll also get unlimited access to over 88,000 Many bicarbonates are soluble in water at standard temperature and pressure; in particular, sodium bicarbonate contributes to total dissolved solids, a common parameter for assessing water quality.[6]. Temperature is not fixed, but I will assume its close to room temperature; As other components are not mentioned, I will assume all carbonate comes from calcium carbonate. Relationship between \(pK_a\) and \(pK_b\) of a conjugate acidbase pair. Conversely, smaller values of \(pK_b\) correspond to larger base ionization constants and hence stronger bases. It is isoelectronic with nitric acid HNO 3. Get unlimited access to over 88,000 lessons. [10][11][12][13] The values of \(K_b\) for a number of common weak bases are given in Table \(\PageIndex{2}\). The same procedure can be repeated to find the expressions for the alphas of the other dissolved species. Can Martian regolith be easily melted with microwaves? Browse other questions tagged, Start here for a quick overview of the site, Detailed answers to any questions you might have, Discuss the workings and policies of this site. Bicarbonate is easily regulated by the kidney, which . The following questions will provide additional practice in calculating the acid (Ka) and base (Kb) dissociation constants. It is measured, along with carbon dioxide, chloride, potassium, and sodium, to assess electrolyte levels in an electrolyte panel test (which has Current Procedural Terminology, CPT, code 80051). This is the equation given by my textbook for hydrolysis of sodium carbonate: $$\ce {Na2CO3 + 2 H2O -> H2CO3 + 2 Na+ + 2 OH-}$$. flashcard sets. The Ka of NH 4+ is 5.6x10 -10 and the Kb of HCO 3- is 2.3x10 -8. It is about twice as effective in fire suppression as sodium bicarbonate. Hence this equilibrium also lies to the left: \[H_2O_{(l)} + NH_{3(aq)} \ce{ <<=>} NH^+_{4(aq)} + OH^-_{(aq)}\]. All chemical reactions proceed until they reach chemical equilibrium, the point at which the rates of the forward reaction and the reverse reaction are equal. Notice that water isn't present in this expression. Dawn has taught chemistry and forensic courses at the college level for 9 years. Does it change the "K" values? The equilibrium arrow suggests that the concentration of the ions are equal to one another: {eq}K_a = \frac{[0.0006]^2}{[1.2]}=3*10^-7 mol/L {/eq}. Legal. The Ka formula and the Kb formula are very similar. There are no HCl molecules to be found because 100% of the HCl molecules have broken apart into hydrogen ions and chloride ions. Acid-Base Buffers: Calculating the pH of a Buffered Solution, Psychological Research & Experimental Design, All Teacher Certification Test Prep Courses, Maram Ghadban, Elizabeth (Nikki) Wyman, Dawn Mills, Using the Ka and Kb in Chemistry Problems, Experimental Chemistry and Introduction to Matter, LeChatelier's Principle: Disruption and Re-Establishment of Equilibrium, Equilibrium Constant (K) and Reaction Quotient (Q), Using a RICE Table in Equilibrium Calculations, Solubility Equilibrium: Using a Solubility Constant (Ksp) in Calculations, The Common Ion Effect and Selective Precipitation, Acid-Base Equilibrium: Calculating the Ka or Kb of a Solution, Titration of a Strong Acid or a Strong Base, NY Regents Exam - Physics: Help and Review, NY Regents Exam - Physics: Tutoring Solution, Middle School Earth Science: Help and Review, Middle School Earth Science: Tutoring Solution, Study.com ACT® Test Prep: Practice & Study Guide, ILTS Science - Environmental Science (112): Test Practice and Study Guide, Praxis Environmental Education (0831) Prep, ILTS Science - Earth and Space Science (108): Test Practice and Study Guide, Praxis Chemistry: Content Knowledge (5245) Prep, CSET Science Subtest II Life Sciences (217): Practice Test & Study Guide, How Acid & Base Structure Affect pH & pKa Values, How to Calculate the Acid Ionization Constant, Ionization Constants of Acids & Conjugate Bases, Wildlife Corridors: Definition & Explanation, Abiotic Factors in Freshwater vs. Potassium bicarbonate (IUPAC name: potassium hydrogencarbonate, also known as potassium acid carbonate) is the inorganic compound with the chemical formula KHCO3. Thus high HCO3 in water decreases the pH of water. Convert this to a ${K_a}$ value and we get about $5.0 \times 10^{-7}$. But at the same time it states that HCO3- will react as a base, because it's Kb >> Ka, True, $HCO_3^-$ will react as both an acid and a base. Following this lesson, you should be able to: To unlock this lesson you must be a Study.com Member. Find the pH. But how can I calculate $[\ce{HCO3-}]$ and $[\ce{CO3^2-}]$? Turns out we didn't need a pH probe after all. But at the same time it states that HCO3- will react as a base, because it's Kb >> Ka $\endgroup$ - then: +2 2 3 T [ HCO ][ ]H = CZ (13) - + 3 1 T [ HCO][ ] HK = CZ (14) 2312 [] T HCOKK CZ = (15) Figure 5.1. Acid ionization constant: \[K_a=\dfrac{[H_3O^+][A^]}{[HA]}\], Base ionization constant: \[K_b=\dfrac{[BH^+][OH^]}{[B]} \], Relationship between \(K_a\) and \(K_b\) of a conjugate acidbase pair: \[K_aK_b = K_w \], Definition of \(pK_a\): \[pKa = \log_{10}K_a \nonumber\] \[K_a=10^{pK_a}\], Definition of \(pK_b\): \[pK_b = \log_{10}K_b \nonumber\] \[K_b=10^{pK_b} \]. It can be assumed that the amount that's been dissociated is very small. We have an acetic acid (HC2H3O2) solution that is 0.9 M. Its hydronium ion concentration is 4 * 10^-3 M. What is the Ka for acetic acid? Ka = (4.0 * 10^-3 M) (4.0 * 10^-3 M) / 0.90 M. This Ka value is very small, so this is a weak acid. To solve it, we need at least one more independent equation, to match the number of unknows. HCO3 or more generally as: z = (H+) 2 + (H+) K 1 + K 1 K 2 where K 1 and K 2 are the first and second dissociation constants for the acid. This test measures the amount of bicarbonate, a form of carbon dioxide, in your blood. This order corresponds to decreasing strength of the conjugate base or increasing values of \(pK_b\). Titration Curves Graph & Function | How to Read a Titration Curve, R.I.C.E. It is a white solid. Potassium bicarbonate ( IUPAC name: potassium hydrogencarbonate, also known as potassium acid carbonate) is the inorganic compound with the chemical formula KHCO 3. 2. The first was took for carbonates only and MO for carbonate + bicarbonate weighed sum. Short story taking place on a toroidal planet or moon involving flying. Nikki has a master's degree in teaching chemistry and has taught high school chemistry, biology and astronomy. Hydrochloric acid, on the other hand, dissociates completely to chloride ions and protons: {eq}HCl_(aq) \rightarrow H^+_(aq) + Cl^-_(aq) {/eq}. pH is an acidity scale with a range of 0 to 14. The dissociation constant can be sought if information about the solution's pH was given. Weak acids and bases do not dissociate well (much, much less than 100%) in aqueous solutions. When using Ka or Kb expressions to solve for an unknown, make sure to write out the dissociation equation, or the dissociation expression, first. From your question, I can make some assumptions: Carbonic acid, $\ce{H2CO3}$, has two ionizable hydrogens, so it may assume three forms: The free acid itself, bicarbonate ion, $\ce{HCO3-}$(first-stage ionized form) and carbonate ion $\ce{CO3^2+}$(second-stage ionized form). [1] A fire extinguisher containing potassium bicarbonate. First, write the balanced chemical equation. How do I quantify the carbonate system and its pH speciation? But it is always helpful to know how to seek its value using the Ka formula, which is: Note that the unit of Ka is mole per liter. We need a weak acid for a chemical reaction. Both the Ka and Kb expressions for dissociation can be used to determine an unknown, whether it's Ka or Kb itself, the concentration of a substance, or even the pH. Note that sources differ in their ${K_a}$ values, and especially for carbonic acid, since there are two kinds - a pseudo-carbonic acid/hydrated carbon dioxide and the real thing (which exists in equilibrium with hydrated carbon dioxide but in a small concentration - about 4% of what what appears to be carbonic acid is true carbonic acid, with the rest simply being $\ce{H2O*CO_2}$. The conjugate base of a strong acid is a weak base and vice versa. Taking the world-renowned weak acid, acetic acid ({eq}CH_3COOH {/eq}), as an example: {eq}CH_3COOH_(aq)\rightleftharpoons CH_3COO^-_(aq) + H^+_(aq) {/eq}. General Ka expressions take the form Ka = [H3O+][A-] / [HA]. O A) True B) False 2) Why does rainwater have a pH of 5 to 6? Your kidneys also help regulate bicarbonate. Normal pH = 7.4. Try refreshing the page, or contact customer support. How to calculate the pH value of a Carbonate solution? In fact, for all acids we can use a general expression for dissociation using the generic acid HA: HA + H2O --> H3O+ + A-. What is the value of Ka? Why do small African island nations perform better than African continental nations, considering democracy and human development? The Ka value is the dissociation constant of acids. Because of the use of negative logarithms, smaller values of \(pK_a\) correspond to larger acid ionization constants and hence stronger acids. Can Martian regolith be easily melted with microwaves? You'll get a detailed solution from a subject matter expert that helps you learn core concepts. The Ka and Kb values for a conjugated acidbase pairs are related through the K. The conjugate base of a strong acid is a very weak base, and the conjugate base of a very weak acid is a strong base. How does CO2 'dissolve' in water (or blood)? For acid and base dissociation, the same concepts apply, except that we use Ka or Kb instead of Kc. It can substitute for baking soda (sodium bicarbonate) for those with a low-sodium diet,[4] and it is an ingredient in low-sodium baking powders.[5][6]. Why doesn't hydroxide concentration equal concentration of carbonic acid and bicarbonate in a sodium bicarbonate solution? What if the temperature is lower than or higher than room temperature? Site design / logo 2023 Stack Exchange Inc; user contributions licensed under CC BY-SA. Does Magnesium metal react with carbonic acid? Learn more about Stack Overflow the company, and our products. For example, hydrochloric acid is a strong acid that ionizes essentially completely in dilute aqueous solution to produce \(H_3O^+\) and \(Cl^\); only negligible amounts of \(HCl\) molecules remain undissociated. "The rate constants at all temperatures and salinities are given in . Great! They must sum to 1(100%), as in chemical reactions matter is neither created or destroyed, only changing between forms. An error occurred trying to load this video. For a given pH, the concentration of each species can be computed multiplying the respective $\alpha$ by the concentration of total calcium carbonate originally present. Find the concentration of its ions at equilibrium. 70%75% of CO2 in the body is converted into carbonic acid (H2CO3), which is the conjugate acid of HCO3 and can quickly turn into it. This is used as a leavening agent in baking. The equilibrium constant for this dissociation is as follows: \[K=\dfrac{[H_3O^+][A^]}{[HA]} \label{16.5.2}\]. Do new devs get fired if they can't solve a certain bug? Initially, the protons produced will be taken up by the conjugate base (A-^\text{-}-start . $$pH = pK1 + log(\frac{\ce{[H2CO3]}}{[HCO3-]})$$. How to calculate the pH value of a Carbonate solution? We are given the \(pK_a\) for butyric acid and asked to calculate the \(K_b\) and the \(pK_b\) for its conjugate base, the butyrate ion. For the oxoacid, see, "Hydrocarbonate" redirects here. Nonetheless, I believe that your ${K_a}$ for carbonic acid is wrong; that number looks suspiciously like the ${K_a}$ instead for hydrogen carbonate ion (or the bicarbonate ion). * Compiled from Appendix 5 Chem 1A, B, C Lab Manual and Zumdahl 6th Ed. Ocean Biomes, Working Scholars Bringing Tuition-Free College to the Community. Like all equilibrium constants, acidbase ionization constants are actually measured in terms of the activities of \(H^+\) or \(OH^\), thus making them unitless. Plus, get practice tests, quizzes, and personalized coaching to help you Let's go into our cartoon lab and do some science with acids! We need to consider what's in a solution of carbonic acid. To solve this problem, we will need a few things: the equation for acid dissociation, the Ka expression, and our algebra skills. How does carbonic acid cause acid rain when $K_b$ of bicarbonate is greater than $K_a$? Is this a strong or a weak acid? Use the relationships pK = log K and K = 10pK (Equation 16.5.11 and Equation 16.5.13) to convert between \(K_a\) and \(pK_a\) or \(K_b\) and \(pK_b\). A solution of this salt is acidic . [H ][CO ] K (9.20b) The definition also takes into account that in reality instead of [H+] the pH is being measured based on a series of buffer solutions. What is the pKa of a solution whose Ka is equal to {eq}2*10^-5 mol/L {/eq}? A) Get the answers you need, now! HCl is the parent acid, H3O+ is the conjugate acid, and Cl- is the conjugate base. The distribution of carbonate species as a fraction of total dissolved carbonate in relation to . The problem provided us with a few bits of information: that the acetic acid concentration is 0.9 M, and its hydronium ion concentration is 4 * 10^-3 M. Since the equation is in equilibrium, the H3O+ concentration is equal to the C2H3O2- concentration. For acids, these values are represented by Ka; for bases, Kb. Why is this sentence from The Great Gatsby grammatical? Kb in chemistry is a measure of how much a base dissociates. See Answer Question: For which of the following equilibria does Kc correspond to the base-ionization constant, Kb, of HCO3? The pKa values for organic acids can be found in Appendix II of Bruice 5th Ed. Strong bases dissociate completely into ions, whereas weak bases dissociate poorly, much like the acid dissociation concept. Kb in chemistry is a measure of how much a base dissociates. For example normal sea water has around 8.2 pH and HCO3 is . By clicking Accept all cookies, you agree Stack Exchange can store cookies on your device and disclose information in accordance with our Cookie Policy. For the bicarbonate, for example: Thank you so much! Browse other questions tagged, Start here for a quick overview of the site, Detailed answers to any questions you might have, Discuss the workings and policies of this site. 1. rev2023.3.3.43278. {eq}[BOH] {/eq} is the molar concentration of the base itself. When HCO3 increases , pH value decreases. Its Ka value is {eq}1.3*10^-8 mol/L {/eq}. Bronsted-Lowry defines acids as chemical substances that have the ability to donate protons to other substances. An acidic solution's pH is lower than 7, a basic solution's pH is higher than 7. These shift the pH upward until in certain circumstances the degree of alkalinity can become toxic to some organisms or can make other chemical constituents such as ammonia toxic. Examples include as buffering agent in medications, an additive in winemaking. Study Ka chemistry and Kb chemistry. Low values of Ka mean that the acid does not dissociate well and that it is a weak acid. B is the parent base, BH+ is the conjugate acid, and OH- is the conjugate base. 0.1M of solution is dissociated. [9], Potassium bicarbonate is an effective fungicide against powdery mildew and apple scab, allowed for use in organic farming. But it is my memory for chemical high school, focused on analytical chemistry in 1980-84 and subsequest undergrad lectures and labs. \(K_a = 1.4 \times 10^{4}\) for lactic acid; \(pK_b\) = 10.14 and \(K_b = 7.2 \times 10^{11}\) for the lactate ion. The flow of bicarbonate ions from rocks weathered by the carbonic acid in rainwater is an important part of the carbon cycle. The bicarbonate ion carries a negative one formal charge and is an amphiprotic species which has both acidic and basic properties. Science Chemistry Calculate the Kb values for the CO32- and C2H3O2- ions using the Ka values for HCO3- (4.7 x 10-11) and HC2H3O2 (1.8 x 10-5), respectively. If the molar concentrations of the acid and the ions it dissociates into are known, then Ka can be simply calculated by dividing the molar concentration of ions by the molar concentration of the acid: 14 chapters | All rights reserved. Learn how to use the Ka equation and Kb equation. To subscribe to this RSS feed, copy and paste this URL into your RSS reader. Consider, for example, the ionization of hydrocyanic acid (\(HCN\)) in water to produce an acidic solution, and the reaction of \(CN^\) with water to produce a basic solution: \[HCN_{(aq)} \rightleftharpoons H^+_{(aq)}+CN^_{(aq)} \label{16.5.6}\], \[CN^_{(aq)}+H_2O_{(l)} \rightleftharpoons OH^_{(aq)}+HCN_{(aq)} \label{16.5.7}\]. Once again, the concentration does not appear in the equilibrium constant expression.. (Kb > 1, pKb < 1). $$K2 = \frac{\ce{[H3O+][CO3^2-]}}{\ce{[HCO3-]}} \approx 4.69*10^-11 $$, You can also write a equation for the overrall reaction, by sum of each stage (and multiplication of the respective equilibrium constants): Remember that Henderson-Hasselbalch provides the equilibrium ratio of concentrations at a given pH. Site design / logo 2023 Stack Exchange Inc; user contributions licensed under CC BY-SA. Did any DOS compatibility layers exist for any UNIX-like systems before DOS started to become outmoded? The equilibrium constant expression for the ionization of HCN is as follows: \[K_a=\dfrac{[H^+][CN^]}{[HCN]} \label{16.5.8}\]. Our Kb expression is Kb = [NH4+][OH-] / [NH3]. These are the values for $\ce{HCO3-}$. {eq}HA_(aq) + H_2O_(l) \rightleftharpoons A^-_(aq) + H^+_(aq) {/eq}. It's a scale ranging from 0 to 14. From the equilibrium, we have: But what does that mean? General acid dissociation in water is represented by the equation HA + H2O --> H3O+ + A-. The parameter standard bicarbonate concentration (SBCe) is the bicarbonate concentration in the blood at a PaCO2 of 40mmHg (5.33kPa), full oxygen saturation and 36C. B) Due to oxides of sulfur and nitrogen from industrial pollution. So bicarb ion is. How can we prove that the supernatural or paranormal doesn't exist? In this case, we are given \(K_b\) for a base (dimethylamine) and asked to calculate \(K_a\) and \(pK_a\) for its conjugate acid, the dimethylammonium ion. Hence the ionization equilibrium lies virtually all the way to the right, as represented by a single arrow: \[HCl_{(aq)} + H_2O_{(l)} \rightarrow \rightarrow H_3O^+_{(aq)}+Cl^_{(aq)} \label{16.5.17}\]. 2018ApHpHHCO3-NaHCO3. Similarly, the equilibrium constant for the reaction of a weak base with water is the base ionization constant (Kb). For an aqueous solution of a weak acid, the dissociation constant is called the acid ionization constant (Ka). In inorganic chemistry, bicarbonate (IUPAC-recommended nomenclature: hydrogencarbonate[2]) is an intermediate form in the deprotonation of carbonic acid. 120ch2co3ka1=4.2107ka2=5.61011nh3h2okb=1.7105hco3nh4+ohh+ 2nh2oh1fe2+fe3+ . The higher the Kb, the the stronger the base. The Kb formula is quite similar to the Ka formula. When the calcium carbonate dissolves, a equilibrium is established between its three forms, expressed by the respective equilibrium equations: First stage: The negative log base ten of the acid dissociation value is the pKa. On this Wikipedia the language links are at the top of the page across from the article title. EDIT: I see that you have updated your numbers. The same logic applies to bases. The concentration of H3O+ and F- are the same, so I replace them with x. I put 6.8 * 10^-4 for Ka, and 0.010 M for HF, then I solve for x. x = 0.0026, so our hydronium ion concentration equals 0.0026 M. To find pH, I take the negative log of that. John Wiley & Sons, 1998. Oceanogr., 27 (5), 1982, 849-855 p.851 table 1. Full text of the 'Sri Mahalakshmi Dhyanam & Stotram'. The Ka of a 0.6M solution is equal to {eq}1.54*10^-4 mol/L {/eq}. A) Due to carbon dioxide in the air. The equation is NH3 + H2O <==> NH4+ + OH-. Decomposition of the bicarbonate occurs between 100 and 120C (212 and 248F): This reaction is employed to prepare high purity potassium carbonate. I would definitely recommend Study.com to my colleagues. potassium hydrogencarbonate, potassium acid carbonate, InChI=1S/CH2O3.K/c2-1(3)4;/h(H2,2,3,4);/q;+1/p-1, InChI=1/CH2O3.K/c2-1(3)4;/h(H2,2,3,4);/q;+1/p-1, Except where otherwise noted, data are given for materials in their, "You Have the (Baking) Power with Low-Sodium Baking Powders", "Why Your Bottled Water Contains Four Different Ingredients", "Powdery Mildew - Sustainable Gardening Australia", "Efficacy of Armicarb (potassium bicarbonate) against scab and sooty blotch on apples", Safety Data sheet - potassium bicarbonate, https://en.wikipedia.org/w/index.php?title=Potassium_bicarbonate&oldid=1107665193, Pages using collapsible list with both background and text-align in titlestyle, Articles containing unverified chemical infoboxes, Wikipedia articles incorporating a citation from the New International Encyclopedia, Creative Commons Attribution-ShareAlike License 3.0, This page was last edited on 31 August 2022, at 05:54. EDIT 2: I think you've realized your mistake; as you say, the values are for $\ce{HCO_3^-}$, which is the hydrogen carbonate ion. In an acidbase reaction, the proton always reacts with the stronger base. If we are given any one of these four quantities for an acid or a base (\(K_a\), \(pK_a\), \(K_b\), or \(pK_b\)), we can calculate the other three. This is in-line with the value I obtained from a copy of Daniel C. Harris' Qualitative Chemical Analysis. Was ist wichtig fr die vierte Kursarbeit? Its \(pK_a\) is 3.86 at 25C. The expressions for the remaining two species have the same structure, just changing the term that goes in the numerator. Because the \(pK_a\) value cited is for a temperature of 25C, we can use Equation 16.5.16: \(pK_a\) + \(pK_b\) = pKw = 14.00. The magnitude of the equilibrium constant for an ionization reaction can be used to determine the relative strengths of acids and bases. In a solution of carbonic acid, we have 1) water and 2) carbonic acid in the main. In another laboratory scenario, our chemical needs have changed. Accessibility StatementFor more information contact us atinfo@libretexts.orgor check out our status page at https://status.libretexts.org. The equilibrium constant for this reaction is the base ionization constant (Kb), also called the base dissociation constant: \[K_b=\dfrac{[BH^+][OH^]}{[B]} \label{16.5.5}\]. [8], Potassium bicarbonate has widespread use in crops, especially for neutralizing acidic soil. Should it not create an alkaline solution? How does the relationship between carbonate, pH, and dissolved carbon dioxide work in water? In aqueous solution carbonic acid behaves as a dibasic acid.The Bjerrum plot shows typical equilibrium concentrations, in solution, in seawater, of carbon dioxide and the various species derived from it, as a function of pH. The relative strengths of some common acids and their conjugate bases are shown graphically in Figure 16.5. Sort by: Ka is the dissociation constant for acids. We know that Kb = 1.8 * 10^-5 and [NH3] is 15 M. We can make the assumption that [NH4+] = [OH-] and let these both equal x.