L'acide malonique HOOC-CH2-COOH est un diacide dont les deux acidités n'ont pas la même force :

step 1 So here we have that they give us the pH of a solution is 1 .47, and then the pH of another one

L'acide malonique HOOC-CH2-COOH est un diacide dont les deux...

L'acide malonique $\mathrm{HOOC}-\mathrm{CH}_2-\mathrm{COOH}$ est un diacide dont les deux acidités n'ont pas la même force : 1) $\mathrm{HOOC}-\mathrm{CH}_2-\mathrm{COOH}+\mathrm{H}_2 \mathrm{O} \rightleftharpoons \mathrm{HOOC}-\mathrm{CH}_2-\mathrm{COO}^{-}+\mathrm{H}_3 \mathrm{O}^{+}$ $$ K_1=1,7 \cdot 10^{-3} $$ 2) $$ \begin{aligned} & \mathrm{HOOC}-\mathrm{CH}_2-\mathrm{COO}^{-}+\mathrm{H}_2 \mathrm{O} \rightleftharpoons-\mathrm{OOC}-\mathrm{CH}_2-\mathrm{COO}^{-}+\mathrm{H}_3 \mathrm{O}^{+} \\ & K_2=2,2 \cdot 10^{-6} \end{aligned} $$ Quelle masse d'acide malonique faut-il dissoudre dans 1 litre d'eau pour obtenir une solution dont le $\mathrm{pH}$ soit 2,5 ?

L'acide malonique $\mathrm{HOOC}-\mathrm{CH}_2-\mathrm{COOH}$ est un diacide dont les deux acidités n'ont pas la même force :
1) $\mathrm{HOOC}-\mathrm{CH}_2-\mathrm{COOH}+\mathrm{H}_2 \mathrm{O} \rightleftharpoons \mathrm{HOOC}-\mathrm{CH}_2-\mathrm{COO}^{-}+\mathrm{H}_3 \mathrm{O}^{+}$
$$
K_1=1,7 \cdot 10^{-3}
$$
2)
$$
\begin{aligned}
& \mathrm{HOOC}-\mathrm{CH}_2-\mathrm{COO}^{-}+\mathrm{H}_2 \mathrm{O} \rightleftharpoons-\mathrm{OOC}-\mathrm{CH}_2-\mathrm{COO}^{-}+\mathrm{H}_3 \mathrm{O}^{+} \\
& K_2=2,2 \cdot 10^{-6}
\end{aligned}
$$

Quelle masse d'acide malonique faut-il dissoudre dans 1 litre d'eau pour obtenir une solution dont le $\mathrm{pH}$ soit 2,5 ?

Key Concepts

Molar Concentration and Mass Calculations

Determining the necessary mass of an acid to achieve a specific solution concentration involves understanding molar relationships. One must relate the desired hydronium ion concentration, as implied by the pH, to the molar concentration of the acid, and then convert that concentration into mass via the acid's molar mass. This process highlights the link between chemical equilibria and practical laboratory preparations.

pH and Hydronium Ion Concentration

pH is a measure of the acidity of a solution, defined as the negative logarithm of the hydronium ion concentration. This relationship allows one to directly connect the equilibrium position of acid dissociation with the observed acidity, making it possible to calculate the necessary concentration of acid required to achieve a desired pH.

Equilibrium Constants (K?)

The acid dissociation constant (K?) quantifies the strength of an acid in solution, representing the ratio of the concentrations of the deprotonated (conjugate base) and protonated forms multiplied by the inverse of the hydronium ion concentration. In multi-step processes, distinct K? values (such as K? and K?) indicate the varying ease with which each proton is lost.

Acid Dissociation Equilibrium

Acid dissociation equilibrium describes the reversible reaction in which an acid donates a proton to water, forming its conjugate base and hydronium ions. In polyprotic systems, each step of deprotonation is governed by its own equilibrium, and understanding these equilibria is essential for predicting the concentration of various ionic species in solution.

Polyprotic Acids

Polyprotic acids are acids that can donate more than one proton per molecule in solution. Each proton donation has its own distinct equilibrium constant, indicating that the acid releases its protons in successive steps with different strengths. This concept is crucial for understanding how different ionization stages contribute to the overall acidity of the solution.

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