3. The sensitivity of a spectrophotometer allows detection of I indicator (for e

Question

3. The sensitivity of a spectrophotometer allows detection of I indicator (for example HIn) in the presence of 99 percent of the other colored form (In). Based on this fact, over what pH range relative to its pK can an indicator be used if a spectrophotometer is employed? 4. A major buffering system in the blood is H,CO,/HCO,. At first glance this would appear to be a poor choice. At the normal blood pH (7.4) a solution of H,CO, and HCO, will change pH abruptly if small amounts of acid or base are added since the pK, of H,CO, is 6.4. Offer an explaination for this apparent contradiction. Support this argumen with appropriate chemical equations

Explanation / Answer

Answer:

3.) When the sensitivity of a spectrophoometer allows detection of 1 percent of one colored forms of an indicator, in presence of 99 percent of the other colored form of the same indicator, the pH range applicable for the indicator relative to the pK of the indicator would be in the range of +/-2.0 units.

4)

The pH of blood is given by the Henderson-Hasselbalch equation:

pH = pKa + log([HCO3-]/[H2CO3]) = 7.4

When small amounts of acid are added, it reacts with HCO3-:

H+(aq) + HCO3-(aq) => H2CO3(aq)

When small amounts of base are added, it reacts with H2CO3:

OH-(aq) + H2CO3(aq) => HCO3-(aq) + H2O(l)

In each case, the ratio [HCO3-]/[H2CO3] only changes slightly

=> log([HCO3-]/[H2CO3]) remains essentially constant

=> pH = pKa + log([HCO3-]/[H2CO3]) remains constant at 7.4 and does not change abruptly

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