A solution has initially 0.1 M-HCOOH and 0.2 M-HCN . Ka of HCOOH =2.56 ×10^-4, Ka of HCN=9.6 ×10

step 1 Hello everyone so in this question it is given that the the two solutions are isohydric having t

A solution has initially 0.1 M-HCOOH and 0.2 M-HCN . Ka of...

A solution has initially $0.1 \mathrm{M}-\mathrm{HCOOH}$ and $0.2 \mathrm{M}-\mathrm{HCN} . K_{\mathrm{a}}$ of $\mathrm{HCOOH}$ $=2.56 \times 10^{-4}, K_{\mathrm{a}}$ of $\mathrm{HCN}=9.6 \times 10^{-10}$ The only incorrect statement for the solution is $(\log 2=0.3)$ (a) $\left[\mathrm{H}^{+}\right]=1.6 \times 10^{-3} \mathrm{M}$ (b) $\left[\mathrm{HCOO}^{-}\right]=1.6 \times 10^{-3} \mathrm{M}$ (c) $\left[\mathrm{CN}^{-}\right]=1.2 \times 10^{-7} \mathrm{M}$ (d) $\mathrm{pOH}=2.8$

A solution has initially $0.1 \mathrm{M}-\mathrm{HCOOH}$ and $0.2 \mathrm{M}-\mathrm{HCN} . K_{\mathrm{a}}$ of $\mathrm{HCOOH}$
$=2.56 \times 10^{-4}, K_{\mathrm{a}}$ of $\mathrm{HCN}=9.6 \times 10^{-10}$
The only incorrect statement for the solution is $(\log 2=0.3)$
(a) $\left[\mathrm{H}^{+}\right]=1.6 \times 10^{-3} \mathrm{M}$
(b) $\left[\mathrm{HCOO}^{-}\right]=1.6 \times 10^{-3} \mathrm{M}$
(c) $\left[\mathrm{CN}^{-}\right]=1.2 \times 10^{-7} \mathrm{M}$
(d) $\mathrm{pOH}=2.8$

Key Concepts

Weak Acid Dissociation Equilibrium

This concept involves the reversible reaction of a weak acid with water to produce its conjugate base and hydrogen ions. The equilibrium expression, given by Ka = [H+][A–]?[HA], describes the extent of dissociation in solution. It is fundamental in calculating the concentrations of species in a solution of a weak acid and understanding how changes in concentration affect the equilibrium state.

Acid Dissociation Constant (Ka)

Ka is a quantitative measure of the strength of a weak acid in solution. A larger Ka value indicates a stronger acid that dissociates more completely, whereas a smaller Ka implies a weaker acid that dissociates to a lesser extent. Comparing Ka values is critical in solutions containing multiple acids, as it helps predict which acid will contribute more to the overall acidity.

pH and pOH Calculations

pH and pOH are logarithmic measures of the hydrogen ion and hydroxide ion concentrations, respectively, and are linked through the relation pH + pOH = 14 at 25°C. Understanding this relationship is essential for determining the acid?base character of a solution from its ionic concentrations and for evaluating the validity of given pH or pOH values in various scenarios.

Relative Acid Strength in Mixed Solutions

When multiple weak acids are present in a solution, their individual contributions to the pH depend on their respective acid strengths, as indicated by their Ka values. The stronger acid (with a larger Ka) tends to have a more significant effect on the hydrogen ion concentration, while the weaker acid may contribute negligibly to the overall acidity. This concept helps in simplifying the analysis by focusing on the dominant acid in the mixture.

Logarithmic Approximations

Logarithmic approximations, such as using log 2 ? 0.3, are commonly employed in acid–base equilibrium calculations to simplify the computation of pH, pKa, and related values. These approximations allow for quicker estimations and facilitate understanding the scale and sensitivity of pH changes with respect to small variations in logarithmic values.

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