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Acidic Buffer Action

When 100 mL o...

Question

When 100 $m L$ of 1.0 $M$ $H C l$ was mixed with 100 $m L$ of 1.0 $M$ $N a O H$ in an insulated beaker at constant pressure, a temperature increase of ${5.7}^{\circ} C$ was measured for the beaker and its contents (Expt 1). Because the enthalpy of neutralization of a strong acid with a strong base is a constant (-57.0 $k J m o l^{- 1}$ ), this experiment could be used to measure the calorimeter constant. In a second experiment (Expt 2), 100 $m L$ of 2.0 $M$ acetic acid $(K_{a} = 2.0 \times 10^{- 5})$ was mixed with 100 $m L$ of 1.0 $M$ $N a O H$ (under identical conditions to Expt 1) where a temperature rise of ${5.6}^{\circ} C$ was measured.
(Consider heat capacity of all solutions as $4.2 J g^{- 1} K^{- 1}$ and density of all solutions as $1.0 g m L^{- 1}$ )

The pH of the solution after Expt. 2 is:

2.8

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4.7

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Solution

The correct option is B 4.7
Final solution contain 0.1 mole of $C H_{3} C O O H$ and $C H_{3} C O O N a$ each.
Hence, it is a buffer solution.
$p_{H} = p K_{a} + l o g \frac{[C H_{3} C O O^{-}]}{[C H_{3} C O O H]}$
$= 5 - l o g 2 + l o g \frac{0.1}{0.1} = 4.7$

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Similar questions

Q. When 100

m L

of 1.0

M

H C l

was mixed with 100

m L

of 1.0

M

N a O H

in an insulated beaker at constant pressure, a temperature increase of

{5.7}^{\circ} C

was measured for the beaker and its contents (Expt 1). Because the enthalpy of neutralization of a strong acid with a strong base is a constant (-57.0

k J m o l^{- 1}

), this experiment could be used to measure the calorimeter constant. In a second experiment (Expt 2), 100

m L

of 2.0

M

acetic acid

(K_{a} = 2.0 \times 10^{- 5})

was mixed with 100

m L

of 1.0

M

N a O H

(under identical conditions to Expt 1) where a temperature rise of

{5.6}^{\circ} C

was measured.
(Consider heat capacity of all solutions as

4.2 J g^{- 1} K^{- 1}

and density of all solutions as

1.0 g m L^{- 1}

)

The pH of the solution after Expt. 2 is:

Q. When 100

m L

of 1.0

M

H C l

was mixed with 100

m L

of 1.0

M

N a O H

in an insulated beaker at constant pressure, a temperature increase of

{5.7}^{\circ} C

was measured for the beaker and its contents (Expt 1). Because the enthalpy of neutralization of a strong acid with a strong base is a constant (-57.0

k J m o l^{- 1}

), this experiment could be used to measure the calorimeter constant. In a second experiment (Expt 2), 100

m L

of 2.0

M

acetic acid

(K_{a} = 2.0 \times 10^{- 5})

was mixed with 100

m L

of 1.0

M

N a O H

(under identical conditions to Expt 1) where a temperature rise of

{5.6}^{\circ} C

was measured.
(Consider heat capacity of all solutions as

4.2 J g^{- 1} K^{- 1}

and density of all solutions as

1.0 g m L^{- 1}

)

Enthalpy of dissociation (in

k J

m o l^{- 1}

) of acetic acid obtained from the Expt. 2 is:

Q. When 100

m L

of 1.0

M

H C l

was mixed with 100

m L

of 1.0

M

N a O H

in an insulated beaker at constant pressure, a temperature increase of

{5.7}^{\circ} C

was measured for the beaker and its contents (Expt 1). Because the enthalpy of neutralization of a strong acid with a strong base is a constant (-57.0

k J m o l^{- 1}

), this experiment could be used to measure the calorimeter constant. In a second experiment (Expt 2), 100

m L

of 2.0

M

acetic acid

(K_{a} = 2.0 \times 10^{- 5})

was mixed with 100

m L

of 1.0

M

N a O H

(under identical conditions to Expt 1) where a temperature rise of

{5.6}^{\circ} C

was measured.
(Consider heat capacity of all solutions as

4.2 J g^{- 1} K^{- 1}

and density of all solutions as

1.0 g m L^{- 1}

)

The pH of the solution after Expt. 2 is:

Q. When 100

m L

of 1.0

M

H C l

was mixed with 100

m L

of 1.0

M

N a O H

in an insulated beaker at constant pressure, a temperature increase of

{5.7}^{\circ} C

was measured for the beaker and its contents (Expt 1). Because the enthalpy of neutralization of a strong acid with a strong base is a constant (-57.0

k J m o l^{- 1}

), this experiment could be used to measure the calorimeter constant. In a second experiment (Expt 2), 100

m L

of 2.0

M

acetic acid

(K_{a} = 2.0 \times 10^{- 5})

was mixed with 100

m L

of 1.0

M

N a O H

(under identical conditions to Expt 1) where a temperature rise of

{5.6}^{\circ} C

was measured.
(Consider heat capacity of all solutions as

4.2 J g^{- 1} K^{- 1}

and density of all solutions as

1.0 g m L^{- 1}

)

Enthalpy of dissociation (in

k J

m o l^{- 1}

) of acetic acid obtained from the Expt. 2 is:

200

ml of

1

H C l

is mixed with

400

ml of

0.5

N a O H

. The temperature rise in the calorimeter was found to be

{4.0}^{o} C .

Water equivalent of calorimeter is

25

g and the specific heat of the solution is

1

cal/mL/degree. If the theoritical heat of neutralization of a strong acid and strong base is

- 13.5

kcal, then the percentage error (as nearest integer) incurred in this experiment while calculating the heat of neutralization is

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The correct option is B 4.7Final solution contain 0.1 mole of CH3COOH and CH3COONa each. Hence, it is a buffer solution. pH=pKa+log[CH3COO−][CH3COOH] =5−log 2+log0.10.1=4.7

The correct option is B 4.7
Final solution contain 0.1 mole of $C H_{3} C O O H$ and $C H_{3} C O O N a$ each.
Hence, it is a buffer solution.
$p_{H} = p K_{a} + l o g \frac{[C H_{3} C O O^{-}]}{[C H_{3} C O O H]}$
$= 5 - l o g 2 + l o g \frac{0.1}{0.1} = 4.7$