(B) For a weak monobasic acid,the dissociation constant $K_a$ is given by the formula $K_a = C\alpha^2,$ where $C$ is the concentration and $\alpha$ i

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(B) For a weak monobasic acid,the dissociation constant $K_a$ is given by the formula $K_a = C\alpha^2,$ where $C$ is the concentration and $\alpha$ is the degree of dissociation.Given: Concentration $C = 0.10 \ M$ and degree of dissociation $\alpha = 3.7 \% = 0.037.$Substituting the values:$K_a = 0.10 \times (0.037)^2$$K_a = 0.10 \times 0.001369$$K_a = 1.369 \times 10^{-4}$Rounding to two significant figures,we get $K_a \approx 1.4 \times 10^{-4}.$

Class 11 Chemistry 6-2.Equilibrium-II (Ionic Equilibrium)pH of weak Acids and weak Bases

Difficult

Accumulation of lactic acid $(HC_3H_5O_3),$ a monobasic acid in tissues,leads to pain and a feeling of fatigue. In a $0.10 \ M$ aqueous solution,lactic acid is $3.7 \%$ dissociated. The value of the dissociation constant,$K_a,$ for this acid will be:

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A
$1.4 \times 10^{-5}$
B
$1.4 \times 10^{-4}$
C
$3.7 \times 10^{-4}$
D
$2.8 \times 10^{-4}$

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Class 11

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6-2.Equilibrium-II (Ionic Equilibrium)

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pH of weak Acids and weak Bases

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What is the dissociation constant for $NH_4OH$ if at a given temperature its $0.1 \ N$ solution has $pH = 11.27$ and the ionic product of water is $7.1 \times 10^{-15}$ (antilog $0.73 = 5.37$)?

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$K_b$ for $NH_4OH$ is $1.8 \times 10^{-5}$. The $[OH^{-}]$ of $0.1 \ M \ NH_4OH$ is

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The $K_a$ of a weak acid $HA$ is $1.00 \times 10^{-5}$. If $0.100 \ mol$ of this acid is dissolved in $1 \ L$ of water,what is the percentage dissociation of the acid at equilibrium?

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$1.0 \ mol$ of $HCl$ and $1.0 \ mol$ of $CH_3COONa$ are dissolved in water to make a $1 \ L$ solution. What is the concentration of $H^+$ in the solution? ($K_a$ for $CH_3COOH = 1.6 \times 10^{-5}$)

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What is the $pH$ of $0.001 \,M$ aniline solution? The ionization constant of aniline can be taken from the table. Calculate the degree of ionization of aniline in the solution. Also,calculate the ionization constant of the conjugate acid of aniline.

Base $K_{b}$

Dimethylamine,$(CH_{3})_{2}NH$ $5.4 \times 10^{-4}$

Triethylamine,$(C_{2}H_{5})_{3}N$ $6.45 \times 10^{-5}$

Ammonia,$NH_{3}$ $1.77 \times 10^{-5}$

Quinine $1.10 \times 10^{-6}$

Pyridine,$C_{5}H_{5}N$ $1.77 \times 10^{-9}$

Aniline,$C_{6}H_{5}NH_{2}$ $4.27 \times 10^{-10}$

Urea,$CO(NH_{2})_{2}$ $1.3 \times 10^{-14}$

Medium

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Base	$K_{b}$
Dimethylamine,$(CH_{3})_{2}NH$	$5.4 \times 10^{-4}$
Triethylamine,$(C_{2}H_{5})_{3}N$	$6.45 \times 10^{-5}$
Ammonia,$NH_{3}$	$1.77 \times 10^{-5}$
Quinine	$1.10 \times 10^{-6}$
Pyridine,$C_{5}H_{5}N$	$1.77 \times 10^{-9}$
Aniline,$C_{6}H_{5}NH_{2}$	$4.27 \times 10^{-10}$
Urea,$CO(NH_{2})_{2}$	$1.3 \times 10^{-14}$

Accumulation of lactic acid $(HC_3H_5O_3),$ a monobasic acid in tissues,leads to pain and a feeling of fatigue. In a $0.10 \ M$ aqueous solution,lactic acid is $3.7 \%$ dissociated. The value of the dissociation constant,$K_a,$ for this acid will be:

Similar Questions

What is the dissociation constant for $NH_4OH$ if at a given temperature its $0.1 \ N$ solution has $pH = 11.27$ and the ionic product of water is $7.1 \times 10^{-15}$ (antilog $0.73 = 5.37$)?

$K_b$ for $NH_4OH$ is $1.8 \times 10^{-5}$. The $[OH^{-}]$ of $0.1 \ M \ NH_4OH$ is

The $K_a$ of a weak acid $HA$ is $1.00 \times 10^{-5}$. If $0.100 \ mol$ of this acid is dissolved in $1 \ L$ of water,what is the percentage dissociation of the acid at equilibrium?

$1.0 \ mol$ of $HCl$ and $1.0 \ mol$ of $CH_3COONa$ are dissolved in water to make a $1 \ L$ solution. What is the concentration of $H^+$ in the solution? ($K_a$ for $CH_3COOH = 1.6 \times 10^{-5}$)

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