Acids and Bases Questions in English

(A) The conjugate acid of a species is formed by the addition of a proton $(H^+)$ to the given species.
For the species $H_2PO_4^-$,the conjugate acid is obtained by adding $H^+$:
$H_2PO_4^- + H^+ \rightarrow H_3PO_4$
Therefore,the conjugate acid of $H_2PO_4^-$ is $H_3PO_4$.

(D) The conjugate acid of a base is formed by adding a proton $(H^+)$ to the base.
For the base $S_2O_8^{2-}$,the conjugate acid is obtained by adding one $H^+$ ion:
$S_2O_8^{2-} + H^+ \rightarrow HS_2O_8^-$
Therefore,the correct option is $D$.

(B) According to the Brønsted-Lowry theory,there is an inverse relationship between the strength of an acid and its conjugate base.
$A$ strong acid has a very weak conjugate base,whereas a weak acid has a relatively strong conjugate base.
Therefore,the statement that the conjugate base of a weak acid is a strong base is correct.

(D) The conjugate base of a species is formed by removing a proton $(H^{+})$ from it.
For $OH^{-}$,the reaction is: $OH^{-} \to O^{2-} + H^{+}$.
Therefore,the conjugate base of $OH^{-}$ is $O^{2-}$.

(B) Lewis base is defined as a substance that can donate a lone pair of electrons.
In $C_2H_5OH$ (ethanol),the oxygen atom has two lone pairs of electrons,which it can donate to an electron-deficient species (Lewis acid).
Therefore,$C_2H_5OH$ acts as a Lewis base.

(C) $NH_4OH$ is a weak base because it undergoes partial dissociation in aqueous solution,producing a low concentration of $OH^{-}$ ions compared to strong bases like $NaOH$,$KOH$,and $Ca(OH)_2$,which dissociate completely. Therefore,$NH_4OH$ is the weakest base among the given options.

(A) In the given reaction,$HCl$ accepts a proton $(H^{+})$ from $HF$ to form $H_2Cl^{+}$.
Since $HCl$ acts as a proton acceptor,it behaves as a base.
Because $HF$ is a stronger acid than $HCl$,$HCl$ acts as a weak base in this medium.

(D) According to the $Br\text{ø}nsted-Lowry$ acid-base theory,the stronger an acid is,the weaker its conjugate base will be. $A$ strong acid dissociates completely in water,meaning its conjugate base has a very low affinity for protons. Therefore,the conjugate base of a strong acid is a weak base.

(C) In the auto-ionization reaction of water,$2H_2O \rightleftharpoons H_3O^{+} + OH^{-}$,one water molecule acts as a Bronsted-Lowry acid by donating a proton $(H^{+})$ to form $OH^{-}$.
Another water molecule acts as a Bronsted-Lowry base by accepting a proton $(H^{+})$ to form $H_3O^{+}$.
Since water performs both functions,it is amphoteric and acts as both a weak acid and a weak base.

(B) According to the Bronsted-Lowry theory,an acid-base conjugate pair differs by a single proton $(H^{+})$.
In the reaction $HCl + H_2O \rightleftharpoons H_3O^{+} + Cl^{-}$,$HCl$ acts as an acid by donating a proton to $H_2O$.
After donating the proton,$HCl$ becomes $Cl^{-}$,which is its conjugate base.
Therefore,$Cl^{-}$ is the conjugate base of $HCl$ acid.

(C) According to the Lewis theory of acid-base reactions,bases donate pairs of electrons and acids accept pairs of electrons.
$A$ Lewis acid is defined as any substance,such as the $H^+$ ion,that can accept a pair of nonbonding electrons.
In other words,a Lewis acid is an electron-pair acceptor.

(A) Water $(H_2O)$ can act as both an acid and a base,meaning it can donate a proton $(H^+)$ to form $OH^-$ or accept a proton to form $H_3O^+$.
Therefore,it is classified as an amphoteric substance.

(C) conjugate base is formed by the removal of a proton $(H^{+})$ from the acid.
For $NH_3$,the removal of one $H^{+}$ ion results in the formation of the amide ion $(NH_2^{-})$.
The reaction is: $NH_3 \rightleftharpoons NH_2^{-} + H^{+}$
Therefore,the conjugate base of $NH_3$ is $NH_2^{-}$.

(A) species that can act as both an acid and a conjugate base is called an amphoteric species.
$H_2SO_4$ acts as an acid and dissociates as follows: $H_2SO_4 \rightleftharpoons H^{+} + HSO_4^-$.
Here,$HSO_4^-$ is the conjugate base of $H_2SO_4$.
However,$HSO_4^-$ can further lose a proton to act as an acid: $HSO_4^- \rightleftharpoons H^{+} + SO_4^{2-}$.
Therefore,$HSO_4^-$ is both a conjugate base (of $H_2SO_4$) and an acid (which forms $SO_4^{2-}$).

(C) According to the Bronsted-Lowry concept,an acid is defined as a proton $(H^+)$ donor,while a base is defined as a substance that acts as a proton $(H^+)$ acceptor.

(C) The hydronium ion is represented by the formula $H_3O^{+}$.
It is formed when a proton $(H^{+})$ attaches to a water molecule $(H_2O)$.
In aqueous solutions,it is often represented as $H_3O^{+}$,although it exists in more complex hydrated forms such as $H_9O_4^{+}$.

(D) The conjugate base of an acid is formed by the removal of one proton $(H^+)$ from the acid molecule.
For the acid $H_2SO_4$,the dissociation reaction is:
$H_2SO_4 \rightleftharpoons H^+ + HSO_4^-$
Thus,$HSO_4^-$ is the conjugate base of $H_2SO_4$.

(B) According to the $Bronsted-Lowry$ theory of acids and bases,an acid is defined as a substance that acts as a proton donor.
Therefore,an acid is a compound which furnishes a proton ($H^+$ ion).

(C) The conjugate base of an acid is formed by removing one proton $(H^+)$ from the acid molecule.
For sulphuric acid $(H_2SO_4)$,the reaction is:
$H_2SO_4 \rightarrow H^+ + HSO_4^-$.
Therefore,the conjugate base of $H_2SO_4$ is the bisulphate ion $(HSO_4^-)$.

(B) According to the Brønsted-Lowry theory,an acid is a substance that donates a proton $(H^+)$.
$HNO_3$ acts as an acid in an aqueous solution and dissociates as follows:
$HNO_3(aq) + H_2O(l) \rightarrow H_3O^+(aq) + NO_3^-(aq)$
Thus,the solution contains the conjugate base $NO_3^-$.

(C) In an aqueous solution,an acid releases $H^{+}$ ions.
Since $H^{+}$ ions are highly reactive,they immediately combine with water molecules $(H_2O)$ to form hydronium ions $(H_3O^{+})$.
The reaction is: $H^{+} + H_2O \rightarrow H_3O^{+}$.
Therefore,the acidic nature in an aqueous medium is characterized by the presence of $H_3O^{+}$ ions.

(C) According to the $Br\o nsted-Lowry$ theory,a conjugate acid is formed when a base accepts a proton $(H^+)$.
$NH_3$ (ammonia) acts as a base and accepts a proton to form the ammonium ion $(NH_4^+)$.
Therefore,$NH_4^+$ is the conjugate acid of the base $NH_3$.
The dissociation equilibrium is: $NH_4^+ ⇌ NH_3 + H^+$.

(C) $Bronsted-Lowry$ acid is a chemical species that donates one or more hydrogen ions $(H^+)$ in a reaction.
$A$ $Lewis$ acid is a chemical species that accepts an electron pair.
$NH_2^{-}$,$O^{2-}$,and $OH^{-}$ all possess lone pairs of electrons,making them $Lewis$ bases.
$BF_3$ acts as a $Lewis$ acid because the boron atom is electron-deficient and possesses an empty $p$-orbital,allowing it to accept a lone pair of electrons.
However,$BF_3$ does not contain any hydrogen atoms to donate as protons; therefore,it cannot act as a $Bronsted$ acid.

(D) The strength of a conjugate base is inversely proportional to the strength of its corresponding acid.
$HCl$ is a strong acid,therefore its conjugate base $Cl^{-}$ is a very weak base.
$H_{2}$,$CH_{4}$,and $CH_{3}OH$ are much weaker acids than $HCl$,making their conjugate bases ($H^{-}$,$CH_{3}^{-}$,and $CH_{3}O^{-}$) much stronger bases than $Cl^{-}$.
Thus,$Cl^{-}$ is the weakest base among the given options.

(D) Lewis base is a substance that can donate a lone pair of electrons.
$NH_3$,$PH_3$,and $(CH_3)_3N$ all have a lone pair of electrons on the central atom (nitrogen or phosphorus) and act as Lewis bases.
$HN_3$ (Hydrazoic acid) does not have a lone pair available for donation in the same manner as the others and is generally considered to act as a Lewis acid in specific contexts due to its electron-deficient nature.
Therefore,the correct option is $D$.

(C) The strength of an acid is inversely proportional to its $pK_a$ value.
Therefore,the smaller the $pK_a$ value,the stronger the acid.
Among the given options,$1$ is the smallest value,so it represents the strongest acid.

(D) is the correct answer.
$NH_3$ (ammonia) is a weak base and undergoes partial ionisation in water to form $NH_4^+$ and $OH^-$ ions.
In contrast,$H_2SO_4$,$NaCl$,and $HNO_3$ are strong electrolytes that undergo complete ionisation in water.

(A) The $Br\text{ø}nsted-Lowry$ theory defines acids as proton $(H^+)$ donors and bases as proton acceptors.
$H_2SO_4$ and $HNO_3$ act as proton donors,fitting the $Br\text{ø}nsted-Lowry$ definition.
$AlCl_3$ and $SO_2$ do not contain ionizable protons and act as Lewis acids (electron pair acceptors),which is explained by the Lewis theory,not the $Br\text{ø}nsted-Lowry$ theory.
Since both $AlCl_3$ and $SO_2$ are listed as options,and the question asks for an example not covered by the theory,both are technically correct; however,$AlCl_3$ is the classic example of a Lewis acid.

(D) According to the $Br\o nsted-Lowry$ theory,an acid is a proton donor and a base is a proton acceptor.
When a base accepts a proton $(H^+)$,it forms its conjugate acid.
For example,$NH_3 + H^+ \rightarrow NH_4^+$,where $NH_4^+$ is the conjugate acid of the base $NH_3$.

(C) The reaction is represented as: $NH_3 + H_2O ⇌ NH_4^+ + OH^-$.
In this reaction,$H_2O$ donates a proton $(H^+)$ to $NH_3$ to form $NH_4^+$.
According to the $Br\o nsted-Lowry$ theory,a substance that donates a proton is an acid.
Therefore,$H_2O$ acts as an acid.

(B) According to the $Br\o nsted-Lowry$ theory,a conjugate base is formed when an acid donates a proton $(H^{+})$.
In the given reaction: $H_2SO_4 + H_2O \rightleftharpoons H_3O^{+} + HSO_4^-$
$H_2SO_4$ acts as an acid and donates a proton to $H_2O$.
After losing a proton,$H_2SO_4$ becomes $HSO_4^-$.
Therefore,$HSO_4^-$ is the conjugate base of the acid $H_2SO_4$.

(C) The neutralization reaction between an acid and a base is defined as the reaction of $H^{+}$ ions from the acid with $OH^{-}$ ions from the base to form water.
$NaOH + HCl \rightarrow NaCl + H_2O$
In this process,the essential product formed is water $(H_2O)$.

(A) The conjugate acid of a species is formed by the addition of a proton $(H^+)$ to the given base.
For the base $HPO_4^{2-}$,the conjugate acid is obtained by adding one $H^+$ ion:
$HPO_4^{2-} + H^+ \rightarrow H_2PO_4^-$
Therefore,the correct option is $A$.

(C) According to the Lewis concept,a base is a substance that can donate an electron pair.
$OH^{-}$,$H_2O$,and $NH_3$ all possess lone pairs of electrons that they can donate.
$Ag^{+}$ is a Lewis acid because it is an electron-deficient species (cation) that can accept an electron pair.
Therefore,$Ag^{+}$ is not a base.

(B) The conjugate acid of a base is formed by adding a proton $(H^+)$ to the base.
For the base $NH_3$,the conjugate acid is formed as follows:
$NH_3 + H^+ \rightarrow NH_4^+$
Therefore,the conjugate acid of $NH_3$ is $NH_4^+$.

(A) According to the Lewis acid-base theory,a Lewis acid is defined as a substance that can accept a lone pair of electrons to form a coordinate covalent bond. Therefore,the correct option is $(A)$.

(B) The conjugate base is formed by the removal of a proton $(H^+)$ from the given species.
For the bicarbonate ion $(HCO_3^-)$,the removal of one $H^+$ ion results in the carbonate ion $(CO_3^{2-})$.
$HCO_3^- \rightarrow H^+ + CO_3^{2-}$
Therefore,the conjugate base of $HCO_3^-$ is $CO_3^{2-}$.

(C) According to the $Br\o nsted-Lowry$ theory,an acid is a proton donor and a base is a proton acceptor.
Water $(H_2O)$ can act as both an acid and a base (amphoteric nature).
It can donate a proton to form $OH^-$ and accept a proton to form $H_3O^+$.
The auto-ionization reaction is: $H_2O + H_2O \rightleftharpoons H_3O^+ + OH^-$.

(A) . $H_2O$ undergoes self-ionization to produce $OH^{-}$ ions as shown in the following equilibrium reaction:
$H_2O(l) + H_2O(l) \rightleftharpoons H_3O^{+}(aq) + OH^{-}(aq)$

(D) Lewis acid is an electron pair acceptor.
$BF_3$ and $AlCl_3$ are electron-deficient compounds and act as Lewis acids.
$HCl$ can act as a Lewis acid due to the presence of an empty orbital on the hydrogen atom or by accepting electron pairs.
$LiAlH_4$ contains the $AlH_4^-$ ion,which can donate a hydride ion $(H^-)$,acting as a nucleophile or Lewis base.
Therefore,$LiAlH_4$ is not a Lewis acid.

(C) solvent that does not contain acidic hydrogen atoms is unable to donate a proton,and if it lacks lone pairs or basic sites,it cannot accept a proton. Such solvents are classified as $Aprotic$ solvents. Examples include $Benzene$ and $Carbon \ tetrachloride$ $(CCl_4)$.

(B) In an aqueous solution,an acid releases $H^{+}$ ions. These $H^{+}$ ions are highly reactive and immediately combine with water molecules $(H_2O)$ to form hydronium ions $(H_3O^{+})$. Therefore,the characteristic species of an acid in an aqueous medium is the $H_3O^{+}$ ion.

(B) According to the Arrhenius theory,an acid is a substance that dissociates in water to produce $H^{+}$ ions.
In the gaseous state,$HCl$ exists as a covalent molecule and does not dissociate into ions.
Therefore,gaseous $HCl$ does not behave as an Arrhenius acid because it cannot provide $H^{+}$ ions in the absence of an aqueous medium.

(C) The strength of a conjugate base is inversely proportional to the strength of its corresponding acid.
The order of acidic strength of the hydrohalic acids is $HI > HBr > HCl > HF$.
Therefore,the order of basic strength of their conjugate bases is $F^{-} > Cl^{-} > Br^{-} > I^{-}$.
$F^{-}$ is the strongest conjugate base because $HF$ is the weakest acid among the given hydrohalic acids due to the high electronegativity and small size of the fluorine atom.
The reaction is: $H^{+} + F^{-} \to HF$.

(D) . The conjugate base of an acid is formed by the removal of one proton $(H^+)$ from the acid molecule.
For the acid $H_2PO_4^-$,the reaction is:
$H_2PO_4^- \to H^+ + HPO_4^{2-}$
Thus,the conjugate base of $H_2PO_4^-$ is $HPO_4^{2-}$.

(A) The conjugate base of an acid is formed by the removal of a proton $(H^+)$ from the acid.
For the acid $HSO_4^-$,the removal of a proton results in the formation of the sulfate ion $(SO_4^{2-})$.
The reaction is: $HSO_4^- \to H^+ + SO_4^{2-}$.
Therefore,the conjugate base of $HSO_4^-$ is $SO_4^{2-}$.

(B) Baking soda is $NaHCO_3$ (sodium bicarbonate).
It is formed by the partial neutralization of a strong base $(NaOH)$ and a weak acid $(H_2CO_3)$.
Since it contains one replaceable hydrogen atom,it is classified as an acidic salt.

(C) $NH_3$ acts as a Lewis base and donates an electron pair to the $H_3O^{+}$ ion.
$H_3O^{+}$ acts as a Lewis acid and accepts the electron pair from $NH_3$.
The reaction is: $NH_3 + H_3O^{+} \rightarrow NH_4^{+} + H_2O$.
As a result,the concentration of $H_3O^{+}$ ions decreases.

(C) $HCl$ is a strong acid that undergoes complete dissociation in water: $HCl(aq) + H_2O(l) \rightarrow H_3O^+(aq) + Cl^-(aq)$.
Since it is a strong acid,it provides the highest concentration of hydronium ions $(H_3O^+)$ compared to the other options,which are either weak acids,salts,or bases.

(B) $NaHSO_3$ is an acidic salt because it contains a replaceable $H^+$ ion.
When dissolved in water,it undergoes ionization: $NaHSO_3 \rightleftharpoons Na^+ + H^+ + SO_3^{2-}$.
The presence of the $H^+$ ion makes the solution acidic.