Isocyanic acid (HNCO) can be prepared by heating sodium...

Isocyanic acid (HNCO) can be prepared by heating sodium cyanate in the presence of solid oxalic acid according to the equation $$2 \mathrm{NaOCN}(s)+\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}(s) \longrightarrow 2 \mathrm{HNCO}(l)+\mathrm{Na}_{2} \mathrm{C}_{2} \mathrm{O}_{4}(s)$$ Upon isolating pure HNCO(I), an aqueous solution of HNCO can be prepared by dissolving the liquid HNCO in water. What is the $\mathrm{pH}$ of a $100 .$ -mL solution of HNCO prepared from the reaction of $10.0 \mathrm{g}$ each of $\mathrm{NaOCN}$ and $\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4},$ assuming all of the HNCO produced is dissolved in solution? $(K_{\mathrm{a}}$ of HNCO $\left.=1.2 \times 10^{-4} .\right)$

Isocyanic acid (HNCO) can be prepared by heating sodium cyanate in the presence of solid oxalic acid according to the equation
$$2 \mathrm{NaOCN}(s)+\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}(s) \longrightarrow 2 \mathrm{HNCO}(l)+\mathrm{Na}_{2} \mathrm{C}_{2} \mathrm{O}_{4}(s)$$
Upon isolating pure HNCO(I), an aqueous solution of HNCO can be prepared by dissolving the liquid HNCO in water. What is the $\mathrm{pH}$ of a $100 .$ -mL solution of HNCO prepared from the reaction of $10.0 \mathrm{g}$ each of $\mathrm{NaOCN}$ and $\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4},$ assuming all of the HNCO produced is dissolved in solution? $(K_{\mathrm{a}}$ of HNCO $\left.=1.2 \times 10^{-4} .\right)$

Key Concepts

Stoichiometry

This concept involves using the balanced chemical equation to convert masses of reactants to moles, determine the amounts of products formed, and identify the limiting reactant. By accurately calculating the number of moles, one can predict the total concentration of a substance produced in a reaction, which is crucial for further equilibrium calculations.

Limiting Reactant Determination

Determining the limiting reactant is key when reactants are not present in stoichiometric amounts. It identifies which reactant will be completely consumed first, thereby limiting the maximum amount of product that can be formed, and sets the basis for calculating the molarity of the resultant acid in solution.

Acid-Base Equilibrium

Acid-base equilibrium concerns the partial dissociation of weak acids in water. The acid dissociation constant (Ka) quantifies the extent of this dissociation and is used to set up an equilibrium expression that relates the concentrations of the undissociated acid and its ionization products, such as hydrogen ions and conjugate bases.

pH Calculation

Calculating pH involves determining the concentration of hydrogen ions (H?) present in solution and applying the formula pH = -log[H?]. For weak acids, this requires setting up the equilibrium expression using the Ka value, solving for [H?], and then taking the negative logarithm of the resulting concentration.

Equilibrium Approximations

When dealing with weak acid equilibrium, approximations can simplify calculations. For instance, if the degree of ionization is small compared to the initial concentration, the change in concentration during ionization may be neglected, which allows one to avoid solving more complex equations and still achieve an accurate estimation of the pH.

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