Henderson Hassleback Equation
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Q. Ka for HCN is 5×10−10 at 25 ∘C. For maintaining a constant pH of 9, the volume of 5 M KCN solution required to be added to 10 mL of 2 M HCN solution is (log2=0.3 and log 5=0.69)
- 4 mL
- 8 mL
- 2 mL
- 10 mL
Q. The [HCO−3]/[H2CO3] ratio in the blood (pH = 7.4) is approximately (pKa=6.1 for H2CO3)
log 20= 1.301
log 20= 1.301
- 20
- 16
- 24
- 10
Q. When rain is accompanied by a thunderstorm, the collected rain water will have pH value:
- Which depends upon the amount of dust in the air
- Slightly lower than that of rain water without the thunderstorm
- Slightly higher than that of rain water without the thunderstorm
- Uninfluenced by occurrence of thunderstorm
Q.
Calculate the degree of ionization of 0.05M acetic acid if its pKa value is 4.74.
How is the degree of dissociation affected when its solution also contains
0.01 M
0.1 M in HCl?
Q. Calculate the pH of buffer solution prepared by mixing 50 ml of 0.1 M sodium acetate and 25 ml of 0.2 M CH3COOH if pKa of CH3COOH is 4.8.
- 4.8
- 10.2
- 5.8
- 14
Q. A solution is prepared by mixing 0.01 mole each of H2CO3, NaHCO3, Na2CO3, and NaOH in 100 mL of water. pH of the resulting solution is
[Given : pKa1 and pKa2 of H2CO3 are 6.37 and 10.32, respectively; log2=0.30]
[Given : pKa1 and pKa2 of H2CO3 are 6.37 and 10.32, respectively; log2=0.30]
Q. 1 ml of 13.6 M HCl is diluted with water to give 1 litre of the solution. Calculate Ph of the resulting solution.
Q. The pH of rain water is approximately :
- 7.5
- 5.6
- 7.0
- 6.5
Q. When 100 mL of 0.4 M CH3COOH is mixed with 100 mL of 0.2 M NaOH, the approximate value of [H3O+] in the solution mixture is? [Ka(CH3COOH)=1.8×10−5]
- 1.8×10−6
- 1.8×10−5
- 9×10−6
- 9×10−5
Q. Calculate the pH of a 0.01 M NH4OH solution containing 1 mol/litre of NH4Cl ? (pKb=4.76)
- 6.8
- 7.24
- 8.4
- 7
Q. What volume of 0.1 M HCOONa solution should be added to 50 mL of 0.05 M formic acid to produce a buffer solution of pH=4.0, pKa of formic acid =3.7 ?
- 60 mL
- 40 mL
- 30 mL
- 50 mL
Q. The pKa of a weak acid (HA) is 4.5. The pOH of an aqueous buffered solution of HA in which 50% of the acid is ionized is
- 2.5
- 9.5
- 7
- 4.5
Q. A weak acid HA after treating with 12 mL of 0.1 M strong base BOH has a pH of 5. At the end point the volume of same base required is 26.6 mL. Calculate the Ka of an acid.
- 1.8×10−5
- 8.22×10−6
- 1.8×10−6
- 8.2×10−5
Q.
The concentration of HCN and NaCN in a solution is 0.01 M each. Calculate the concentration of hydroxyl ions if the dissociation constant of HCN is 7.2×10−10.
- 1.4×10−5 M
- 1.4×10−7 M
- 14×10−5 M
- 14×10−7 M
Q. 20 ml of 0.1 M acetic acid is mixed with 10 mL of sodium acetate. pKa of CH3COOH is 4.8, then the concentration of sodium acetate, if pH of the mixture is 4.8 is:
- 0.1 M
- 0.4 M
- 0.2 M
- 1 M
Q. The pKb of cyanide ion is 4.7 The pH of a solution prepared by mixing 2.5 moles of KCN and 2.5 moles of HCN in water and making total volume upto 500 ml is :
- 9.3
- 7.3
- 10.3
- 8.3
Q. A certain buffer solution contains equal concentration of X− and HX. Kb for X− is 10−10. The pH of the buffer will be:
Q. Calculate the ratio of pH of a solution containing 1 mol of CH3COONa + 1 mol of HCl per litre to another solution containing 1 mol of CH3COONa + 1 mol of acetic acid per litre.
- 1:1
- 2:1
- 1:2
- 2:3
Q. HCl gas is passed through an aqueous solution of 0.1 M 1-aminopropane (PrNH2) till the pH reaches 9.71. Calculate the ratio of [PrNH2]/[PrNH+3] in this solution. [Ka(PrNH+3)=1.96×10−11]
log (1.96)=0.292, log (0.1)=−1
log (1.96)=0.292, log (0.1)=−1
- 10:1
- 8:1
- 4:1
- 1:10
Q.
pH of 0.5 M Ba(CN)2 solution (pKb of CN−=9.30) is :
8.35
3.35
9.35
9.50
Q. pKa of CH3COOH is 4.74. Thus, the buffer range for CH3COOH and CH3COONa buffer is
- 3.74 to 4.74
- 4.74 to 5.74
- 3.74 to 5.74
- 1.0 to -1.0
Q. Two buffers, (X) and (Y) of pH 4.0 and 6.0 respectively are prepared from acid HA and the salt NaA. Both the buffers are 0.50 M of HA. What would be the pH of the solution obtained by mixing equal volumes of the buffers ? KHA=1.0×10−5
- 8.4
- 4.3
- 5.7
- 7.5
Q. In an acidic buffer the ratio [acid]/[salt] is 0.5. If we make the concentration of acid, twice of its initial value then what is value of 10 ΔpH. [log2=0.3]
Q. What is the resultant pH of the solution if equal volumes of two solutions with pH=5 and pH=9 are mixed at 25 oC?
- 7
- 6.5
- 5.5
- 8.5
Q. Which of the following would have the highest pH?
- HN3
- C2H5OH
- MgO
- SO2
- CO2
Q. In an acidic indicator HIn has ionisation constant =10−8. The acid form of indicator is yellow and alkaline form is red. Which is correct statement: (Given : log2=0.3, log3=0.48)
- The pH range of indicator is 7 to 9
- Change in pH is 0.96 when 75% acidic form of the indicator changes to 75% alkaline form
- This indicator is suitable for the titration of strong acid Vs strong base
- pH of indicator is 8.3 when ratio of acid form to alkaline form is 2
Q. The [HCO−3]/[H2CO3] ratio in the blood (pH = 7.4) is approximately (pKa=6.1 for H2CO3)
log 20= 1.301
log 20= 1.301
- 20
- 24
- 10
- 16
Q. The pH of a buffer solution containing 2.0 mol per litre CH3COONa and 1.5 mol per litre CH3COOH is: (Ka for acetic acid is 1.8×10−5)
- 4.87
- 9.2
- 5.8
- 2.4
Q.
To 1L solution containing 0.1 mol each of NH3 and NH4Cl, 0.05 mol of NaOH is added. The change in pH will be (pKb for NH3=4.74)
-0.48
0.48
0.30
-0.30
Q. At 25∘C, consider 1 L of two solutions as following:
Solution A : Contains 1 mol CH3COONa+0.5 mol HCl
Solution B : Contains 1 mol CH3COONa+0.5 mol CH3COOH
If Ka(CH3COOH)=10−5 then ratio of [H+ (aq.)] of solution A to that of solution B is,
(Given: log 1=0 and log 2=0.3, log(5×10−6=−5.3)
Solution A : Contains 1 mol CH3COONa+0.5 mol HCl
Solution B : Contains 1 mol CH3COONa+0.5 mol CH3COOH
If Ka(CH3COOH)=10−5 then ratio of [H+ (aq.)] of solution A to that of solution B is,
(Given: log 1=0 and log 2=0.3, log(5×10−6=−5.3)
- 1:2
- 2:1
- 1:4
- 4:1
