Phosphoric acid (H3PO4) is a weak acid naturally found in the body that supports kidney function and also involved in fat storage. It is a weak acid with a Ka value of 6.9×10-3 at STP. (i) Determine the H+ concentration of a 0.25 M solution of phosphoric acid at STP (assume that the acid is monoacidic). (ii) Calculate the pH of the 0.25 M H3PO4 solution at STP. (iii) What effect on the pH of the solution would you expect to observe if the temperature of the weak solution was increased. Give a brief explanation for your answer.
Phosphoric acid (H3PO4) is a weak acid naturally found in the body that supports kidney function and also involved in fat storage. It is a weak acid with a Ka value of 6.9×10-3 at STP.
(i) Determine the H+ concentration of a 0.25 M solution of phosphoric acid at STP (assume that the acid is monoacidic).
(ii) Calculate the pH of the 0.25 M H3PO4 solution at STP.
(iii) What effect on the pH of the solution would you expect to observe if the temperature of the weak solution was increased. Give a brief explanation for your answer.
The correct answer and explanation is :
Phosphoric acid (H₃PO₄) is a triprotic acid, meaning it can donate up to three protons (H⁺) in solution. However, for simplification, we’ll consider it as a monoprotic acid, focusing only on its first dissociation:
[ \text{H}_3\text{PO}_4 \rightleftharpoons \text{H}^+ + \text{H}_2\text{PO}_4^- ]
The acid dissociation constant (Kₐ) for this equilibrium is approximately 7.5 × 10⁻³ at 25°C (standard temperature and pressure, STP). citeturn0search2
(i) Determining the H⁺ Concentration of a 0.25 M H₃PO₄ Solution at STP:
For a weak monoprotic acid, the dissociation constant Kₐ is given by:
[ K_a = \frac{[\text{H}^+][\text{A}^-]}{[\text{HA}]} ]
Assuming the initial concentration of H₃PO₄ is 0.25 M and the degree of dissociation is small, we let the concentration of dissociated H⁺ be ‘x’. Thus:
- [H⁺] = x
- [H₂PO₄⁻] = x
- [H₃PO₄] ≈ 0.25 – x ≈ 0.25 M (since x is small)
Substituting into the Kₐ expression:
[ 7.5 \times 10^{-3} = \frac{x \times x}{0.25} ]
Simplifying:
[ x^2 = 7.5 \times 10^{-3} \times 0.25 ]
[ x^2 = 1.875 \times 10^{-3} ]
[ x = \sqrt{1.875 \times 10^{-3}} ]
[ x \approx 0.0433 \, \text{M} ]
Therefore, the H⁺ concentration is approximately 0.0433 M.
(ii) Calculating the pH of the 0.25 M H₃PO₄ Solution at STP:
pH is defined as:
[ \text{pH} = -\log[\text{H}^+] ]
Substituting the H⁺ concentration:
[ \text{pH} = -\log(0.0433) ]
[ \text{pH} \approx 1.36 ]
Thus, the pH of the 0.25 M H₃PO₄ solution is approximately 1.36.
(iii) Effect of Temperature Increase on the pH of the Weak Acid Solution:
The dissociation of weak acids is temperature-dependent. Generally, increasing the temperature favors the endothermic dissociation process, leading to greater ionization of the acid. This results in a higher concentration of H⁺ ions, thereby lowering the pH of the solution. However, it’s important to note that the exact effect of temperature on pH can vary depending on the specific acid and its dissociation characteristics. citeturn0search9
In summary, for a 0.25 M phosphoric acid solution at standard temperature and pressure, the H⁺ concentration is approximately 0.0433 M, corresponding to a pH of about 1.36. Increasing the temperature is expected to increase the ionization of the acid, leading to a decrease in pH.