MCQ 11 Mark
The addition of HCl will not suppress the ionisation of:
- AAcetic acid
- BSulphuric acid
- CH2S
- DBenzoic acid
Answer
View full question & answer→- Sulphuric acid
50 questions · timed · auto-graded
Explanation:
Common ion effect is observed when a solution of weak electrolyte is mixed with a solution of strong electrolyte which provides an ion common to that provided by a weak electrolyte.
Ammonium hydroxide is a weak electrolyte and ammonium chloride is a strong electrolyte.
Ammonium chloride provides ammonium ion which is common to that provided by ammonium hydroxide.
Thus, the pair NH4OH + NH4Cl shows common ion effect. Ammonium chloride suppresses the ionization of ammonium hydroxide.
Explanation:
Ionic product depends only on temperature.
Explanation:
For the reaction, N2(g) + O2(g) ⇔ 2NO(g), the value of Δ n is 2 − (1 + 1) = 0. When there is no change in the number of moles of reactants and products, there is no effect on the equilibrium when pressure is changed.
In all other option, the number of mole of reactant and products are different, Hence, if pressure is changed the equilibrium will be affected.
Explanation:
2NH3(g) ⇌ N2(g) + 3H2(g)
upon increase in pressure the equilibrium will shift to the side of lesser gas molecules i.e. left. and thereby a decrease in number of moles of gas and volume of the system.
Explanation:
Sodium chloride is a salt of strong base NaOH and strong acid HCl. In its aqueous solution, following equilibrium is observed.
NaCl + H2O ⇌ Na+ + Cl− + H2O
Explanation:
According to Lewis concept, a positively charged or an electron deficient species acts as Lewis acid. BF3 is an electron deficient compound with B having 6 electrons only.
Explanation:
In lime kilns, CO2 formed continues to escape into the atmosphere and equilibrium is never established.
Explanation:
Chemical equilibrium of reversible reaction is not influenced by the catalyst. Catalyst decreases the activation barrier for a reaction so the reaction proceeds fast. In the presence of the catalyst, the equilibrium reaches faster but it doesn't affect the thermodynamic properties.
Explanation:
Hand OH- both will increase, therefore pH will decrease due to increase in H+ as Kw = [H+] [OH-] will increase with increase in temperature.
Explanation:
It is related with neutralisation of charge on colloidal particlr.
Explanation:
When a system is at equilibrium under constant temperature and pressure, its free energy change is zero $(\triangle\text{G}=0)$
Explanation:
With the increase in temperature from 298K to 1000K, the value of the equilibrium constant decreases from 5 × 10−3 to 2 × 10−3.
Thus, as the temperature increases, the equilibrium shifts to the left direction. This is possible for an exothermic reaction.
Explanation:
The conjugate bases should have one proton less in each case and therefore, the corresponding conjugate bases are $\text{F}^-,\text{HSO}^-_4\text{ and }\text{CO}^{2-}_3$ respectively.
Explanation:
Degree of ionization$(\alpha)$ depends on-
(1) Concentration of solute.
(2) Temperature.
(3) Nature of electrolysis.
(4) Nature of solvent.
(5) Dilution.
"α" does not depend on molecular mass of electrolyte.
Explanation:
With the increase of temperature, k value decreases, so that forward reaction decreases with increase of temperature. This implies that reaction will proceed in forward direction with decrease of temperature, i.e., heat is liberated and hence forward reaction is exothermic.
Explanation:
The given reaction is endothermic, so on increasing the temperature, it will shift in forward direction.
Explanation:
In the gaseous reaction,
$\text{N}_2(\text{g})+3\text{H}_2(\text{g})\rightleftharpoons2\text{NH}_3(\text{g}),$
reactants and products are in the homogeneous phase.
Similarly, for the reactants,
$\text{CH}_3\text{COOC}_2\text{H}_5(\text{aq})+\text{H}_2\text{O}(\text{l})\\\rightleftharpoons\text{CH}_3\text{COOH}(\text{aq})+\text{C}_2\text{H}_5\text{OH(aq)}$
and $\text{Fe}^{2+}(\text{aq})+\text{SCN}^-(\text{aq})\rightleftharpoons\text{[Fe(SCN)}]^{2+}(\text{aq})$
all the reactants and products are in the homogeneous solution phase.
Explanation:
It will favour backward reaction because process is exothermic.
Explanation:
As Kw increases $[\text{H}^+][\text{OH}^-]>10^{-14}$
As $[\text{H}^+]=[\text{OH}^-]$
or $[\text{H}^+]^2=10^{-14}$
or $[\text{H}^+]>10^{-7}\text{M}$
pH < 7
Explanation:
A buffer solution either is a mixture of a weak acid and its salt with strong base or a mixture of a weak base and its salt with strong acid. Hence, clearly CH3COONH4 is not a buffer solution.
Explanation:
Greater the boiling point, less is the vapour pressure.
Hence, the correct order of vapour pressures will be:
water < acetone < ether.
Explanation:
When carbon dioxide reacts with water it forms carbonic acid. Increasing the pressure of carbon dioxide makes the reaction feasible in the forward direction and hence solubility of CO2 increases.
Explanation:
A compound such as water (H2O) has many interesting properties. Water molecules can accept a proton to act as Bronsted-Lowry bases in certain circumstances. The following is an example of HCl dissolving in water:
$\text{HCl}+\text{H}_2\text{O}(\ell)\rightarrow\text{H}_3\text{O}^+_\text{(aq)}+\text{C}1^-_\text{(aq)}$
The another possibility is water can act like a Bronsted-Lowry acid by donating a proton. Water donates a proton to a proton-accepting amide ion in the presence of ammonia, resulting in the following product:
$\text{H}_2\text{O}\ell{(ℓ)}+\text{NH}^-_2\text{(aq)}\rightarrow\text{OH}^-_{\text{(aq)}}+\text{NH}_3\text{(aq)}$
Explanation:
pKa = −logKa
Higher the Ka, higher is the strength of the acid.
For higher Ka, pKa value is smaller.
Explanation:
When QC < KC then the reaction proceeds in the forward direction.
Explanation:
(a) and (c) are acidic.
Because they are salts of strong acid H2SO4 and HCl and weak bases Cu(OH)2 and Al(OH)2 respectively.
Explanation:
$\text{Ag}_2\text{C}_2\text{O}_4(\text{aq})\rightleftharpoons2\text{Ag}^+\ \ \ +\ \ \ \text{C}_2\text{O}^{2-}_4\\ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ 2.2\times10^{-4}\ \ \ \ \ 1.1\times10^{-4}$
$\text{K}_{\text{sp}}=(\text{Ag}^+)^2(\text{C}_2\text{O}^{2-}_4)$
$=(2.2\times10^{-4})^2=(1.1\times10^{-4})$
$5.3\times10^{-12}$
Explanation:
It's pH changes very little when a small amount of strong acid or base is added to it.
Buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of chemical applications. In nature, there are many systems that use buffering for pH regulation. For example, the bicarbonate buffering system is used to regulate the pH of blood.
Explanation:
$\ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \text{CH}_3\text{COOH}+\text{H}_2\text{O}\rightleftharpoons\text{H}_3\text{O}^++\text{CH}_3\text{COOH}^-\\ ^\text{Initial conc.} \ \ \ \ \ \ \ \ \ \ \ \ 0.01 \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ 0 \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ 0\\^\text{At equilibrium} \ \ \ \ \ \ \ \ \ 0.01-\text{x} \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \text{x} \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \text{x}$
$\text{K}_\text{a}=\frac{[\text{H}_3\text{O}^+][\text{CH}_3\text{COO}^-]}{[\text{CH}_3\text{COOH]}}=\frac{\text{x}^2}{0.01-\text{x}}$
Since $\text{x}<<0.01,$
Therefore, $0.01-\text{x}\approx0.01$
$\frac{\text{x}^2}{0.01}=1.74\times10^{-5}$
$\text{x}^2=1.74\times10^{-7}\text{ or x}=4.2\times10^{-4}$
$\text{pH}=-\log(4.2\times10^{-4})=3.4$
Explanation:
The equilibrium reaction is H2O(l) ⇌ H2O(g), when pressure is applied, the equilibrium will shift to left as the value of Δ n is positive. Hence, for boiling to occur more temperature is required. So, it will increase the boiling point.
Explanation:
Adding a common ion prevents the weak acid or weak base from ionizing as much as it would without the added common ion.The common ion effect suppresses the ionization of a weak acid by adding more of an ion that is a product of this equilibrium.Due to this common ion effect, when we add sodium acetate dissociation of acetic acid decreases and solution will have less number of hydrogen ion and so, pH increases. (as pH = −log[H+])
Explanation:
The solution of buffer resists changes in pH.
A buffer solution is defined as a solution which resists drastic changes in pH upon the addition of a small amount of either an acid or a base.
Explanation:
Buffer solutions have the capacity to react with small amounts of added acid or base without affecting the hydrogen ion concentration of the solution.
Explanation:
Adding a common ion prevents the weak acid or weak base from ionizing as much as it would without the added common ion. The common ion effect suppresses the ionization of a weak acid by adding more of an ion that is a product of this equilibrium.
Explanation:
The ph of boiling water is 6.4. This implies that boiling water is neutral. When water is boiled, both hydrogen ion and hydroxide ion concentration increases to same extent. Hence it is neutral. With increase in the hydrogen ion concentration, pH decreases from 7 to 6.4. Also, the value of Kw also increases as the degree of dissociation of water increases with increase in temperature.
Explanation:
The equilibrium constant for basic ionization called basic ionization constant and is represented by Kb.
$\text{K}_{\text{b}}=\frac{[\text{M}^+][\text{OH}^-]}{[\text{MOH}]}$
Explanation:
As we are combining the two equations.
$\therefore$ Equilibrium constant for combined reaction, i.e. K = K1 × K2
Explanation:
For the reaction PCl5 ⇔ PCl3 + Cl2, the forward reaction occurs with increase in the number of moles from 1 to 2. Also the reverse reaction occurs with decrease in the number of moles from 2 to 1. When the pressure of the system is increased, the backward reaction will be favoured as the reverse reaction occurs with decrease in the number of moles. Thus the dissociation of PCl5 is suppressed. Hence, the degree of dissociation decreases. Pressure will be more in 5 L vessel than in 10 L vessel. Thus the extent of dissociation of PCl5 will be more in 10L vessel.
Explanation:
In case of a conjugate acid-base pair,
$\text{K}_{\text{a}}\times\text{K}_{\text{b}}=\text{K}_{\text{w}}.$
Explanation:
Since, NaC1 is soluble to a very significant extent, when AgC1 is added to NaCl solution, the common ion [Cl−] increases in the solution. To have the solubility product or Ksp of AgCl constant, [Ag+] will decrease or AgCl will percipitate out from the solution. This is common ion effect.
Explanation:
An acidic buffer contains equimolar quantities of weak acid and its salt with strong base. A basic buffer contains equinolar quantities of weak base and its salt with strong acid. Sodium acetate is a salt with strong base but sodium propionate is not weak acid, it is also a salt.
Explanation:
Buffer solutions have the capacity to react with small amounts of added acid or base without affecting the hydrogen ion concentration of the solution. Thus, the buffer solutions help to keep the pH value constant in a chemical reaction.
Explanation:
When NH4Cl is added to NH4OH solution, concentration of NH4+ ions increases so the equilibrium shift towards left.So the dissociation of ammonium hydroxide is reduced.
Explanation:
For the reaction, $\text{H}_2\text{S}\rightleftharpoons\text{H}^++\text{HS}^-$
$\text{K}_{\text{a}_1}=\frac{[\text{H}^+][\text{HS}^-]}{[\text{H}_2\text{S}]}$
For the reaction, $\text{H}\text{S}^-\rightleftharpoons\text{H}^++\text{S}^{2-}$
$\text{K}_{\text{a}_2}=\frac{[\text{H}^+][\text{S}^{2-}]}{[\text{H}\text{S}^-]}$
When the above two reaction are added, their equilibrium constants are multiplied. Thus
$\text{K}_{\text{a}_3}=\frac{[\text{H}^+]^2[\text{S}^{2-}]}{[\text{H}_2\text{S}]}=\text{K}_{\text{a}_1}\times\text{K}_{\text{a}_2}$
Hence, $\text{K}_{\text{a}_3}=\text{K}_{\text{a}_1}\times\text{K}_{\text{a}_2}$