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## weak acids and bases

Hence this equilibrium also lies to the left: $H_2O_{(l)} + NH_{3(aq)} \ce{ <<=>} NH^+_{4(aq)} + OH^-_{(aq)}$. Acid ionization constant: $K_a=K[H_2O]=\dfrac{[H_3O^+][A^−]}{[HA]}$, Base ionization constant: $K_b=K[H_2O]=[BH^+][OH^−][B]$, Relationship between $$K_a$$ and $$K_b$$ of a conjugate acid–base pair: $K_aK_b = K_w$, Definition of $$pK_a$$: $pKa = −\log_{10}K_a \ref{16.5.11}$ $K_a=10^{−pK_a}$, Definition of $$pK_b$$: $pK_b = −\log_{10}K_b \ref{16.5.13}$ $K_b=10^{−pK_b}$. Answer: $$K_a = 1.4 \times 10^{−4}$$ for lactic acid; $$pK_b$$ = 10.14 and $$K_b = 7.2 \times 10^{−11}$$ for the lactate ion. Calculate $$K_a$$ for lactic acid and $$pK_b$$ and $$K_b$$ for the lactate ion. Weak Polyprotic Acids and Bases. Consequently, aqueous solutions of acetic acid contain mostly acetic acid molecules in equilibrium with a small concentration of $$H_3O^+$$ and acetate ions, and the ionization equilibrium lies far to the left, as represented by these arrows: $\ce{ CH_3CO_2H_{(aq)} + H_2O_{(l)} <<=> H_3O^+_{(aq)} + CH_3CO_{2(aq)}^- }$. The leveling effect applies to solutions of strong bases as well: In aqueous solution, any base stronger than OH− is leveled to the strength of OH− because OH− is the strongest base that can exist in equilibrium with water. AP® is a registered trademark of the College Board, which has not reviewed this resource. Once again, the concentration of water is constant, so it does not appear in the equilibrium constant expression; instead, it is included in the $$K_b$$. Thus 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. The conjugate acid–base pairs are listed in order (from top to bottom) of increasing acid strength, which corresponds to decreasing values of $$pK_a$$. In this case, the sum of the reactions described by $$K_a$$ and $$K_b$$ is the equation for the autoionization of water, and the product of the two equilibrium constants is $$K_w$$: Thus if we know either $$K_a$$ for an acid or $$K_b$$ for its conjugate base, we can calculate the other equilibrium constant for any conjugate acid–base pair. Relating Ka and Kb to pH, and calculating percent dissociation. Thus nitric acid should properly be written as $$HONO_2$$. The base ionization constant $$K_b$$ of dimethylamine ($$(CH_3)_2NH$$) is $$5.4 \times 10^{−4}$$ at 25°C. The conjugate base of a strong acid is a weak base and vice versa. Similarly, the equilibrium constant for the reaction of a weak base with water is the base ionization constant (Kb). Donate or volunteer today! The LibreTexts libraries are Powered by MindTouch® and are supported by the Department of Education Open Textbook Pilot Project, the UC Davis Office of the Provost, the UC Davis Library, the California State University Affordable Learning Solutions Program, and Merlot. Calculate $$K_b$$ and $$pK_b$$ of the butyrate ion ($$CH_3CH_2CH_2CO_2^−$$). For example. So, therefore, in an acid-base equilibrium where an acid reacts with a base, you have the proton (or H + ion) being transferred from the acid to the base. CH3COOH (acetic acid) HCOOH (formic acid) HF (hydrofluoric acid) HCN (hydrocyanic acid) HNO2 (nitrous acid) HSO4- (hydrogen sulfate ion) Bases Strong. We also acknowledge previous National Science Foundation support under grant numbers 1246120, 1525057, and 1413739. Two species that differ by only a proton constitute a conjugate acid–base pair. Strong acid add all their H+ to will weak acid only add some H+ to solution. The corresponding expression for the reaction of cyanide with water is as follows: $K_b=\dfrac{[OH^−][HCN]}{[CN^−]} \label{16.5.9}$. Some common weak acids and bases are given here. Weak acid and base ionization reactions and the related equilibrium constants, Ka and Kb. Relating Ka and Kb to pH, and calculating percent dissociation. The Brønsted-Lowry theory of acids and bases is that: acids are proton donators and bases are proton acceptors. Substituting the values of $$K_b$$ and $$K_w$$ at 25°C and solving for $$K_a$$, $K_a(5.4 \times 10^{−4})=1.01 \times 10^{−14}$. Because the $$pK_a$$ value cited is for a temperature of 25°C, we can use Equation 16.5.16: $$pK_a$$ + $$pK_b$$ = pKw = 14.00. The values of $$K_b$$ for a number of common weak bases are given in Table $$\PageIndex{2}$$. Example 1 - Finding the K a of a weak acid from the pH of its solution. 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. Khan Academy is a 501(c)(3) nonprofit organization. There are many more weak acids than strong acids. Salts such as $$K_2O$$, $$NaOCH_3$$ (sodium methoxide), and $$NaNH_2$$ (sodamide, or sodium amide), whose anions are the conjugate bases of species that would lie below water in Table $$\PageIndex{2}$$, are all strong bases that react essentially completely (and often violently) with water, accepting a proton to give a solution of $$OH^−$$ and the corresponding cation: $K_2O_{(s)}+H_2O_{(l)} \rightarrow 2OH^−_{(aq)}+2K^+_{(aq)} \label{16.5.18}$, $NaOCH_{3(s)}+H_2O_{(l)} \rightarrow OH^−_{(aq)}+Na^+_{(aq)}+CH_3OH_{(aq)} \label{16.5.19}$, $NaNH_{2(s)}+H_2O_{(l)} \rightarrow OH^−_{(aq)}+Na^+_{(aq)}+NH_{3(aq)} \label{16.5.20}$. Acids have pH values from 1 to 7. The ionization of weak acids and bases is a chemical … This only time this becomes important is at very low (< 10-6 M) concentrations of acids or bases, when water will be the main source of H + and OH-. Here is a partial list, ordered from strongest to weakest.

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