Write examples of weak acids and weak bases and give ionic equilibrium in its aqueous solution.
Write examples of weak acids and weak bases and give ionic equilibrium in its aqueous solution.
| $(A)$ Weak Acids | Ionic Equilibriums |
| $(1)$ Acetic acid $\left(\mathrm{CH}_{3} \mathrm{COOH}\right)$ | $\mathrm{CH}_{3} \mathrm{COOH}_{(\text {aq })}+\mathrm{H}_{2} \mathrm{O}_{(l)} \square \mathrm{H}_{3} \mathrm{O}_{\text {(aq) }}^{+}+\mathrm{CH}_{3} \mathrm{COO}_{\text {(aq) }}^{-}$ |
| $(2)$ Benzoic acid $\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COOH}\right)$ | $\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COOH}_{(\text {aq })}+\mathrm{H}_{2} \mathrm{O}_{(l)} \square \mathrm{H}_{3} \mathrm{O}_{(\text {aq })}^{+}+\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{COO}_{(\text {aq })}^{-}$ |
| $(3)$ Hydrocyanic acid $(HCN)$ | $\mathrm{HCN}_{(\text {aq })}+\mathrm{H}_{2} \mathrm{O}_{(l)} \square \mathrm{H}_{3} \mathrm{O}_{\text {(aq) }}^{+}+\mathrm{CN}_{(\text {aq })}^{-}$ |
| $(4)$ Formic acid ($HCOOH$) | $\mathrm{HCOOH}_{(\text {aq })}+\mathrm{H}_{2} \mathrm{O}_{(l)} \square \mathrm{H}_{3} \mathrm{O}_{\text {(aq) }}^{+}+\mathrm{HCOO}_{\text {(aq) }}^{-}$ |
| $(5)$ Hypochlorous acid $(HOCl)$ | $\mathrm{HOCl}_{\text {(aq) }}+\mathrm{H}_{2} \mathrm{O}_{(l)} \square \mathrm{OCl}_{\text {(aq) }}^{-}+\mathrm{H}_{3} \mathrm{O}_{\text {(aq) }}^{+} \text {etc. }$ |
| $(B)$ Weak Bases | Ionic Equilibriums |
| $(1)$ Ammonia $\left(\mathrm{NH}_{3}\right)$ | $\mathrm{NH}_{3(\mathrm{aq})}+\mathrm{H}_{2} \mathrm{O}_{(l)} \square \mathrm{NH}_{4(\mathrm{aq})}^{+}+\mathrm{OH}_{(\mathrm{aq})^{-}}^{-}$ |
| $(2)$ Aniline $\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{2}\right)$ | $\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{2(\mathrm{aq})}+\mathrm{H}_{2} \mathrm{O}_{(l)} \square \mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{3(\mathrm{aq})}^{+}+\mathrm{OH}_{(\mathrm{aq})}^{-}$ |
| $(3)$ Hydrazine $\left(\mathrm{NH}_{2} \mathrm{NH}_{2}\right)$ | $\mathrm{NH}_{2} \mathrm{NH}_{2(\mathrm{aq})}+\mathrm{H}_{2} \mathrm{O}_{(l)} \square \mathrm{NH}_{2} \mathrm{NH}_{3(\mathrm{aq})}^{+}+\mathrm{OH}_{(\mathrm{aq})}$ |
| $(4)$ Methyl Amine $\left(\mathrm{CH}_{3} \mathrm{NH}_{2}\right)$ | $\mathrm{CH}_{3} \mathrm{NH}_{2(\mathrm{aq})}+\mathrm{H}_{2} \mathrm{O}_{(l)} \square \mathrm{CH}_{3} \mathrm{NH}_{3(\mathrm{aq})}^{+}+\mathrm{OH}_{(\mathrm{aq})}^{-}$ |
| $(5)$ Dimethyl Amine $\left(\mathrm{CH}_{3}\right)_{2} \mathrm{NH}$ | $\left.\left(\mathrm{CH}_{3}\right)_{2} \mathrm{NH}_{(\mathrm{aq})}+\mathrm{H}_{2} \mathrm{O}_{(l)} \square\left(\mathrm{CH}_{3}\right)_{2} \mathrm{NH}_{3(\mathrm{aq})}^{+}+\mathrm{OH}_{(\mathrm{aq})}^{-}\right)$etc. |
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When $100 \ mL$ of $1.0 \ M \ HCl$ was mixed with $100 \ mL$ of $1.0 \ M \ NaOH$ 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 $\left(-57.0 \ kJ \ mol ^{-1}\right)$, this experiment could be used to measure the calorimeter constant. In a second experiment (Expt. $2$), $100 \ mL$ of $2.0 \ M$ acetic acid $\left(K_a=2.0 \times 10^{-5}\right)$ was mixed with $100 \ mL$ of $1.0 M \ NaOH$ (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 $1.0 \ g mL ^{-1}$ )
$1.$ Enthalpy of dissociation (in $kJ mol ^{-1}$ ) of acetic acid obtained from the Expt. $2$ is
$(A)$ $1.0$ $(B)$ $10.0$ $(C)$ $24.5$ $(D)$ $51.4$
$2.$ The $pH$ of the solution after Expt. $2$ is
$(A)$ $2.8$ $(B)$ $4.7$ $(C)$ $5.0$ $(D)$ $7.0$
Give the answer question $1$ and $2.$
When $100 \ mL$ of $1.0 \ M \ HCl$ was mixed with $100 \ mL$ of $1.0 \ M \ NaOH$ 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 $\left(-57.0 \ kJ \ mol ^{-1}\right)$, this experiment could be used to measure the calorimeter constant. In a second experiment (Expt. $2$), $100 \ mL$ of $2.0 \ M$ acetic acid $\left(K_a=2.0 \times 10^{-5}\right)$ was mixed with $100 \ mL$ of $1.0 M \ NaOH$ (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 $1.0 \ g mL ^{-1}$ )
$1.$ Enthalpy of dissociation (in $kJ mol ^{-1}$ ) of acetic acid obtained from the Expt. $2$ is
$(A)$ $1.0$ $(B)$ $10.0$ $(C)$ $24.5$ $(D)$ $51.4$
$2.$ The $pH$ of the solution after Expt. $2$ is
$(A)$ $2.8$ $(B)$ $4.7$ $(C)$ $5.0$ $(D)$ $7.0$
Give the answer question $1$ and $2.$
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In aqueous solution the ionization constants for carbonic acid are
$K_1 = 4.2 \times 10^{-7}$ and $K_2 = 4.8 \times 10^{-11}$
Select the correct statement for a saturated $0.034\, M$ solution of the carbonic acid.
In aqueous solution the ionization constants for carbonic acid are
$K_1 = 4.2 \times 10^{-7}$ and $K_2 = 4.8 \times 10^{-11}$
Select the correct statement for a saturated $0.034\, M$ solution of the carbonic acid.
- [AIEEE 2010]