When 60 ml of 0.1 M $\mathrm{Ca}{\left({\mathrm{NO}}_3\right)}_2$ is mixed with 40 ml of 0.125 M ${\mathrm{Na}}_2{\mathrm{CO}}_3$, ${\mathrm{CaCO}}_3$ precipitates. If K_sp of ${\mathrm{CaCO}}_3$ is $5\times {10}^{-9}$, then $\left[{\mathrm{CO}}_3^{2-}\right]$ in the resulting solution is

100 ml of 0.3 M NH₄OH is mixed with 100 ml of 0.2 M NaOH. K_b NH₄OH is $1.8\times {10}^{-5}$. The degree of dissociation of NH₄OH is

Solid AgNO₃ is added slowly to a buffer solution of pH=10 to precipitate AgOH. The [Ag⁺] concentration in the solution is $\left[{\mathrm{K}}_{\mathrm{sp}}\right(\mathrm{AgOH})={10}^{-10}]$

If pK_b for ${\mathrm{CN}}^{-}$ at $25°\mathrm{C}$ is 4.7, the pH of 0.5 M aqueous NaCN solution is

A buffer solution is prepared by mixing 20 ml of 0.1 M CH₃COOH and 40 ml of 0.5 M CH₃COONa and then diluted by adding 100 ml of distilled water. The pH of resulting buffer solution is (Given pKa CH₃COOH=4.76)

The ionization constant  of $\left[{\mathrm{NH}}_4^{+}\right]$ in water is $5.6\times {10}^{-10}$ at $25°\mathrm{C}$. The rate constant for the reaction of $\left[{\mathrm{NH}}_4^{+}\right]$ and $\left[{\mathrm{OH}}^{-}\right]$ to form NH₃ and H₂O is $3.4\times {10}^{10} \mathrm{lit} {\mathrm{mol}}^{-1 }{\sec }^{-1}$ at $25°\mathrm{C}$. The rate constant for the proton transfer from water to NH₃ in $\mathrm{lit} {\mathrm{mol}}^{-1} {\sec }^{-1}$ is

Equilibrium constant of ${\mathrm{NH}}_4^{+}$ to NH₃ and H⁺ is 10^-10. The rate constant for ${\mathrm{NH}}_4^{+}+{\mathrm{OH}}^{-}\to {\mathrm{NH}}_3+{\mathrm{H}}_2\mathrm{O}$ is 10¹⁰. The rate constant for ${\mathrm{NH}}_3+{\mathrm{H}}_2\mathrm{O}\to {\mathrm{NH}}_4^{+}+{\mathrm{OH}}^{-}$ is

An acid base indicator has ${\mathrm{K}}_{\mathrm{a}}=3\times {10}^{-5}$. The acid form of the indicator is red and the basic form is blue. By how much must the pH change in order to change the indicator from 75% red to 75% blue (log 3=0.4770)

The dissociation constants for aniline, acetic acid and ionic product of water at $25°\mathrm{C}$ are $3.83\times {10}^{-10}, 1.75\times {10}^{-5} \mathrm{and} 1.008\times {10}^{-14}$ respectively. The degree of hydrolysis of aniline acetate in a decinormal solution is

A solution of benzoic acid (a weak monobasic acid) is titrated with NaOH.The pH of the solution is 4.2 when half of the acid is neutralized. The dissociation constant of the acid will be:

The mixture that shows the maximum buffer capacity is:

The pH of a solution containing 0.1 mol of CH₃COOH, 0.2 mol of CH₃COONa, and 0.05 mol of NaOH in 1 L  is-

(pK_a of CH₃COOH=4.74 and log 5=0.7)

18 ml of mixture of acetic acid and sodium acetate required 6 ml of 0.1 M NaOH for neutralization of the acid and 12 ml of 0.1 M HCl for reaction with salt separately. If pK_a of the acid is 4.75, what is the pH of the mixture?

What is the minimum pH required to prevent the precipitation of ZnS in a solution that is 0.01 M ZnCl₂ and saturated with 0.10 M H₂S?

$[\mathrm{Given} : {\mathrm{K}}_{\mathrm{sp}} \mathrm{of} \mathrm{ZnS}={10}^{-21}, {\mathrm{K}}_{{\mathrm{a}}_1}\times {\mathrm{K}}_{{\mathrm{a}}_2} \mathrm{of} {\mathrm{H}}_2\mathrm{S}={10}^{-20}]$

Which of the following is an amphiprotic (can accept and give protons) ion?

Calculate the ratio of F^- and HCOO^- in a mixture of 0.1 M HF $({\mathrm{K}}_{\mathrm{a}}=6.6\times {10}^{-4})$ and 0.2 M HCOOH $({\mathrm{K}}_{\mathrm{a}}=2\times {10}^{-4}).$

A buffer solution is prepared by mixing 10 mL of 1.0 M acetic acid with 20 mL of 0.5 M sodium acetate which is then diluted
to 100 mL with distilled water. If the pK_a of ${\mathrm{CH}}_3\mathrm{COOH}$ is 4.76, the pH of the buffer solution prepared is -

The pH of a solution obtained by mixing equal volume of solutions having pH=3 and pH=4. [log 5.5=0.7404]

An acid-base indicator has ${\mathrm{K}}_{\mathrm{a}}=3.0\times {10}^{-5}$. The acid form of the indicator is red and the basic form is blue. The [H⁺] required to change the indicator from 75% red to 75% blue is

The self-ionization constant for pure formic acid, K = [$HCOO{H}_2^{+}\left]\right[HCO{O}^{-}]$ ] has been estimated as ${10}^{-4}$ at room temperature.The percentage of formic acid molecules in pure formic acid that are converted to formate ions is

$(\mathrm{Given}:\hspace{0.17em}{\mathrm{d}}_{\mathrm{HCOOH}}=1.22 \mathrm{g}/\mathrm{cc})$ 

AgOH is added to NaCl solution to form AgCl precipitate. After the precipitation, the pH of the solution is 8. The [Cl^-] is $({\mathrm{K}}_{\mathrm{sp}} \mathrm{of} \mathrm{AgCl}={10}^{-12}, {\mathrm{K}}_{\mathrm{sp}} \mathrm{of} \mathrm{AgOH}={10}^{-10})$

For the reaction ${\left[\mathrm{Ag}{\left(\mathrm{CN}\right)}_2\right]}^{-}\left(\mathrm{aq}\right)\rightleftharpoons {\mathrm{Ag}}^{+}\left(\mathrm{aq}\right)+2{\mathrm{CN}}^{-}\left(\mathrm{aq}\right),$ the equilibrium constant at $25°\mathrm{C}$ is $4.0\times {10}^{-19}.$ Calculate the silver ion concentration in a solution which was originally 0.10 M in KCN and 0.03 M in AgNO₃.

Separate solutions of four sodium salts NaW, NaX, NaY and NaZ had pH 7.0, 9.0, 10.0 and 11.0 respectively. When each solution is 0.1 M, the strongest acid is

Silver acetate is a slightly soluble salt of a weak acid $({\mathrm{K}}_{\mathrm{a}}=1.75\times {10}^{-5})$. At $25°\mathrm{C}$, 100 g of water dissolves 1.04 g of crystalline silver acetate. The density of saturated solution of silver acetate at $20°\mathrm{C}$ is 1.01 g/cc. The solubility product constant for silver acetate at $20°\mathrm{C}$

0.1 M solution of three different sodium salts NaX, NaY and NaZ have pH values 7.0, 9.0 and 11.0 respectively. The correct order of dissociation constant values of these acids is

When equal volume of the following solutions are mixed, which of the following gives maximum precipitate? $({\mathrm{K}}_{\mathrm{sp}} \mathrm{of} \mathrm{AgCl}={10}^{-12})$

The percentage hydrolysis of NaCN in $\left(\frac{\mathrm{N}}{80}\right)$ aqueous solution [Dissociation constant of HCN is $1.3\times {10}^{-9}$ and ${\mathrm{K}}_{\mathrm{w}}=1.0\times {10}^{-14}$] is

Hydrolysis constant K_A and K_B of two salts of weak acids HA and HB are 10^-8 and 10^-6 respectively. If the dissociation constant of third acid HC is 10^-2. The order of acidic strengths of three acids will be

A solution containing NH₄Cl and NH₄OH has a hydroxide ion concentration of ${10}^{-6} \mathrm{mol} {\mathrm{litre}}^{-1}$, which of the following hydroxides could be precipitated when this solution is added in equal volume to a solution containing 0.1 M of metal ions?

In the precipitation titration of KCl against AgNO₃, K₂CrO₄ is used as an indicator since, AgCl is white coloured. End point is detected by appearance of deep yellow coloured precipitate of Ag₂CrO₄. The minimum concentration of chromate ion required for detection of end point is $[{\mathrm{K}}_{\mathrm{sp}} \mathrm{of} \mathrm{AgCl}=2.5\times {10}^{-10 }\mathrm{and} {\mathrm{K}}_{\mathrm{sp} }\mathrm{of} {\mathrm{Ag}}_2{\mathrm{CrO}}_4=1.8\times {10}^{-12}]$

The concentration of hydroxyl ion in a solution left after mixing 100 mL of 0.1 M MgCl₂ and 100 mL of 0.2 M NaOH $[{\mathrm{K}}_{\mathrm{sp}} \mathrm{of} \mathrm{Mg}{\left(\mathrm{OH}\right)}_2=1.2\times {10}^{-11}]$ is

When H₂S is passed through an aqueous solution of an equimolar mixture of Zn²⁺ and Pb²⁺ acidified with dilute acetic acid, ZnS is not precipitated, because

The type of bond formed between the ${\mathrm{BF}}_3\&{\mathrm{NH}}_3$ is-

The decreasing order of basic strength of the following conjugate bases is-

${\mathrm{OH}}^{-}, {\mathrm{RO}}^{-}{\mathrm{CH}}_3{\mathrm{COO}}^{-}, {\mathrm{Cl}}^{-}$

If the solubility of BaSO₄ in water is  $8\times {10}^{-4} \mathrm{mol} {\mathrm{dm}}^{-3}$. The solubility BaSO₄ in 0.01 mol dm^-3 of H₂SO₄ is :  

At a certain temperature and pressure of 10⁵ Pa, iodine vapour contains 40% by volume of I atoms. The K_p for the equilibrium of the reaction would be-

 $I_2 \left(g\right) \leftrightharpoons 2I \left(g\right)$

Given the reaction 2HI _(g) $\rightleftharpoons$ H_2 (g) + I_2 (g)

A sample of HI_(g) is placed in a flask at a pressure of 0.2 atm. At equilibrium, the partial pressure of HI_(g) is 0.04 atm.  The ${\mathrm{K}}_{\mathrm{P}}$ for the given equilibrium would be:

A mixture of 1.57 mol of N₂, 1.92 mol of H_2, and 8.13 mol of NH₃ is introduced into a 20 L vessel at 500 K. At this temperature, the equilibrium constant, K_c for the reaction N_2 (g) + 3H_2 (g) $\rightleftharpoons$ 2NH_3 (g) is 1.7 × 10². The direction of the net reaction is

Given the reaction:

2BrCl _(g) $\leftrightharpoons$ Br_2 (g) + Cl_2 (g); K_c= 32 at 500 K. If the initial concentration of BrCl is 3.3 × 10^-3 mol L^–1, the molar concentration of BrCl in the mixture at equilibrium would be:

For the reaction, NO (g) + 1/2O₂ (g) $\rightleftharpoons$ NO₂ (g)

 ∆f_G° (NO₂) = 52.0 kJ/mol , ∆f_G° (NO) = 87.0 kJ/mol and ∆f_G° (O₂) = 0 kJ/mol.

The equilibrium constant for the formation of NO₂ from NO and O₂ at 298K would be

The equilibrium constant for the following reaction is 1.6 ×10⁵ at 1024K

H_2(g) + Br_2(g)  $\leftrightharpoons$2HBr_(g)

If 10.0 bar of HBr is introduced into a sealed container at 1024 K, the equilibrium pressure of HBr is -

Which of the following is/are the Lewis acids?

(i) BF₃     (ii) H₂O    (iii) H⁺       $\left(\mathrm{iv}\right) {\mathrm{AlF}}_3$

A 0.02 M solution of pyridinium hydrochloride has a pH of 3.44. The ionization constant of pyridine will be:

The solubility product for a salt of type AB is $4\times {10}^{-8}$.  The molarity of its standard solution will be:

 

The K_sp of Ag₂CrO₄ and AgBr is 1.1 × 10^–12 and 5.0 × 10^–13 respectively.

The molarity ratio of saturated solutions of Ag₂CrO₄ and AgBr will be:

When equal volumes of 0.002 M solutions of sodium iodate and cupric chlorate are mixed together (K_{sp  (cupric iodate} = 7.4 × 10^–8 ), the correct observation would be 

The minimum volume of water required to dissolve 1g of calcium sulphate at 298 K is

(For CaSO₄ , K_sp is 9.1 × 10^–6)

For the reaction, 2NOCl (g) $\rightleftharpoons$ 2NO (g) + Cl₂ (g); K_p= 1.8 × 10^–2 atm at 500 K.

The value of K_c for above mentioned reaction would be 

$$\begin{gathered} 1. 4.33 \times {10}^{-4} \mathrm{mol} {\mathrm{L}}^{-1} \\ 2. 4.33 \times {10}^4 \mathrm{mol} {\mathrm{L}}^{-1} \\ 3. 1.65 \times {10}^{-5} \mathrm{mol} {\mathrm{L}}^{-1} \\ 4. 2.39 \times {10}^{-3} \mathrm{mol} {\mathrm{L}}^{-1} \end{gathered}$$

For the following equilibrium, K_c= 6.3 × 10¹⁴ at 1000 K

$\mathrm{NO}\left(\mathrm{g}\right) + {\mathrm{O}}_3 \left(\mathrm{g}\right) \to {\mathrm{NO}}_2\left(\mathrm{g}\right) + {\mathrm{O}}_2 \left(\mathrm{g}\right)$

What would be the k_c for the reverse reaction?

$$\begin{gathered} 1. 2.33 \times {10}^{-16} \\ 2. 1.59 \times {10}^{-15} \\ 3. 2.67 \times {10}^{-13} \\ 4. 4.47 \times {10}^{14} \end{gathered}$$

${\mathrm{K}}_{\mathrm{c}} = \frac{{[{\mathrm{NH}}_3 ]}^4 {[ {\mathrm{O}}_2 ]}^5}{{\left[\mathrm{NO}\right]}^4 {\left[{\mathrm{H}}_2\mathrm{O}\right]}^6}$

The balanced chemical equation corresponding to above mentioned expression is 

$$\begin{gathered} 1. 4{\mathrm{NO}}_{\left(\mathrm{g}\right)} + 6{\mathrm{H}}_2{\mathrm{O}}_{\left(\mathrm{g}\right)} \rightleftharpoons 4{{\mathrm{NH}}_3}_{\left(\mathrm{g}\right)} + 5{{\mathrm{O}}_2}_{\left(\mathrm{g}\right)} \\ 2. 4{{\mathrm{NH}}_3}_{\left(\mathrm{g}\right)} +5{{\mathrm{O}}_2}_{\left(\mathrm{g}\right)} \rightleftharpoons 4{\mathrm{NO}}_{\left(\mathrm{g}\right)} +\hspace{0.17em}6{\mathrm{H}}_2{\mathrm{O}}_{\left(\mathrm{g}\right)} \\ 3. 2{\mathrm{NO}}_{\left(\mathrm{g}\right)} +\hspace{0.17em}3{\mathrm{H}}_2{\mathrm{O}}_{\left(\mathrm{g}\right)} \rightleftharpoons 4{{\mathrm{NH}}_3}_{\left(\mathrm{g}\right)} +\hspace{0.17em}3{{\mathrm{O}}_2}_{\left(\mathrm{g}\right)} \\ 4. {{\mathrm{NH}}_3}_{\left(\mathrm{g}\right)} +\hspace{0.17em}3{\mathrm{H}}_2{\mathrm{O}}_{\left(\mathrm{g}\right)} \rightleftharpoons 2{\mathrm{NO}}_{\left(\mathrm{g}\right)}\hspace{0.17em}+\hspace{0.17em}3{{\mathrm{O}}_2}_{\left(\mathrm{g}\right)} \end{gathered}$$