Mix Examples - Acids, Bases and Salts Questions in English

(C) The $Bronsted-Lowry$ theory defines an acid as a proton $(H^+)$ donor and a base as a proton $(H^+)$ acceptor.
Unlike the $Arrhenius$ theory,which is limited to aqueous solutions where $H^+$ or $OH^-$ ions are produced,the $Bronsted-Lowry$ theory does not require water as a solvent.
Therefore,it can be applied to both aqueous and non-aqueous solutions,as well as gas-phase reactions.

(D) According to the Bronsted-Lowry theory,an acid is a proton $(H^{+})$ donor.
When an acid is dissolved in water,it releases a proton $(H^{+})$.
Since a proton $(H^{+})$ cannot exist independently in an aqueous solution,it immediately reacts with a water molecule $(H_{2}O)$ to form a hydronium ion $(H_{3}O^{+})$.
Therefore,the acidity of an aqueous solution is due to the presence of hydronium ions $(H_{3}O^{+})$.

(A) The hydronium ion $(H_{3}O^{+})$ is formed when a water molecule $(H_{2}O)$ acts as a Lewis base and donates a lone pair of electrons to a hydrogen ion $(H^{+})$,which acts as a Lewis acid.
In a water molecule,the oxygen atom has two lone pairs of electrons.
One of these lone pairs is donated to the $H^{+}$ ion to form a coordinate covalent bond (also known as a dative bond).
Therefore,the coordinate covalent bond exists between the oxygen atom of the $H_{2}O$ molecule and the $H^{+}$ ion.

(C) According to the Bronsted-Lowry theory,an acid is defined as a substance that acts as a proton $(H^+)$ donor.
Conversely,a base is defined as a substance that acts as a proton $(H^+)$ acceptor.
Therefore,the substance that donates a proton to another substance is classified as an acid.

(B) According to the Bronsted-Lowry theory:
$1$. An acid is defined as a proton $(H^+)$ donor.
$2$. $A$ base is defined as a proton $(H^+)$ acceptor.
Since the question asks for a substance that accepts a proton from another substance,it is classified as a base.

(C) According to the $Brønsted-Lowry$ theory, an acid is defined as a substance that can donate a proton ($H^+$ ion), while a base is defined as a substance that can accept a proton ($H^+$ ion).
Therefore, a base is a proton acceptor.

(A) When hydrogen chloride $(HCl)$ gas is dissolved in water $(H_2O)$,it undergoes ionization to form hydronium ions $(H_3O^+)$ and chloride ions $(Cl^-)$.
The reaction is: $HCl + H_2O \rightarrow H_3O^+ + Cl^-$.
In this reaction,$HCl$ acts as a Brønsted-Lowry acid by donating a proton $(H^+)$ to the water molecule.
Water acts as a Brønsted-Lowry base by accepting the proton from $HCl$.

(B) When sulfur trioxide $(SO_3)$ reacts with water $(H_2O)$,it undergoes a hydration reaction to form sulfuric acid $(H_2SO_4)$.
The chemical equation for this reaction is:
$SO_3(g) + H_2O(l) \rightarrow H_2SO_4(aq)$
Sulfur trioxide is an acidic oxide,and when it dissolves in water,it forms a strong mineral acid known as sulfuric acid.

(B) When carbon dioxide $(CO_2)$ reacts with water $(H_2O)$,it forms carbonic acid $(H_2CO_3)$.
The chemical equation for this reaction is: $CO_2 + H_2O \rightarrow H_2CO_3$.
Carbonic acid is a weak acid formed when carbon dioxide dissolves in water.

(C) Metal oxides are generally basic in nature. When they react with water,they form metal hydroxides,which are basic solutions.
For example,sodium oxide $(Na_2O)$ reacts with water $(H_2O)$ to form sodium hydroxide $(NaOH)$,which is a strong base.
$Na_2O(s) + H_2O(l) \rightarrow 2NaOH(aq)$.

(A) An oxide of a metal is formed when a metal reacts with oxygen.
Calcium $(Ca)$ is a metal,and its reaction with oxygen produces calcium oxide $(CaO)$,which is a metallic oxide.
Carbon $(C)$,sulfur $(S)$,and oxygen $(O)$ are non-metals,so their oxides ($CO_2$,$SO_3$,$SO_2$) are non-metallic oxides.

(D) An oxide of a metal is formed when a metal reacts with oxygen.
$CaO$ (Calcium oxide) is a metal oxide.
$MgO$ (Magnesium oxide) is a metal oxide.
$Na_2O$ (Sodium oxide) is a metal oxide.
$CO_2$ (Carbon dioxide) is formed by carbon,which is a non-metal. Therefore,$CO_2$ is a non-metallic oxide (acidic oxide),not a metal oxide.

(C) strong acid is an acid that completely dissociates into its ions in an aqueous solution.
$HCl$ (Hydrochloric acid),$H_{2}SO_{4}$ (Sulfuric acid),and $HNO_{3}$ (Nitric acid) are all examples of strong mineral acids.
$CH_{3}COOH$ (Acetic acid) is a weak organic acid because it only partially dissociates into ions in an aqueous solution.
Therefore,$CH_{3}COOH$ is not a strong acid.

(D) The reactivity of metals with acids depends on their position in the reactivity series.
Metals that are more reactive than hydrogen can displace hydrogen from acids to form salt and hydrogen gas.
$Ca$,$Mg$,and $Zn$ are more reactive than hydrogen and react with acids to release $H_2$ gas.
$Au$ (Gold) is a noble metal and is placed below hydrogen in the reactivity series.
Therefore,$Au$ does not react with dilute acids to evolve hydrogen gas.

(C) When an acid reacts with a base, a neutralization reaction occurs.
In this reaction, the acid and base neutralize each other to produce salt and water.
The general chemical equation is: $\text{Acid} + \text{Base} \rightarrow \text{Salt} + \text{Water}$.

(B) $NaOH$ stands for Sodium Hydroxide.
It is a chemical compound that dissociates completely in water to produce hydroxide ions $(OH^-)$.
Because it ionizes completely in an aqueous solution,it is classified as a strong base.

(D) When an acid reacts with a metal carbonate or a metal hydrogen carbonate,it produces a corresponding salt,water,and carbon dioxide gas.
The general chemical equation is:
Metal Carbonate/Hydrogen Carbonate + Acid $\rightarrow$ Salt + Water + Carbon Dioxide $(CO_2)$
For example,the reaction of sodium carbonate $(Na_2CO_3)$ with hydrochloric acid $(HCl)$:
$Na_2CO_3(s) + 2HCl(aq) \rightarrow 2NaCl(aq) + H_2O(l) + CO_2(g)$
Therefore,the gas released is carbon dioxide.

(B) Sodium hydroxide $(NaOH)$ is a strong base. Most metals do not react with bases,but amphoteric metals like Aluminum $(Al)$,Zinc $(Zn)$,and Lead $(Pb)$ react with strong bases to form complex salts and release hydrogen gas $(H_2)$.
For the reaction of Aluminum with sodium hydroxide:
$2Al + 2NaOH + 6H_2O \rightarrow 2Na[Al(OH)_4] + 3H_2 \uparrow$
Here,$Na[Al(OH)_4]$ is sodium tetrahydroxoaluminate (a salt) and $H_2$ is hydrogen gas. Therefore,$Al$ is the correct metal.

(B) Molarity $(M)$ is defined as the number of moles of solute dissolved in $1 \ L$ of solution.
It is expressed as: $M = \frac{\text{Number of moles of solute}}{\text{Volume of solution in liters}}$.
Molality is the number of moles of solute per kilogram of solvent.
Normality is the number of gram equivalents of solute per liter of solution.

(B) Molarity $(M)$ is defined as the number of moles of solute dissolved in one liter of solution.
The formula for molarity is: $M = \frac{\text{moles of solute}}{\text{volume of solution in liters}}$.
Therefore,the unit of molarity is expressed as $\text{mole liter}^{-1}$ or $M$.

(C) Molar mass is defined as the mass of one mole of a substance.
It is expressed as the mass in grams per mole of the substance.
Therefore, the $SI$ unit for molar mass is grams per mole, which is written as $g \ mol^{-1}$.

(B) Molarity $(M)$ is defined as the number of moles of solute per liter of solution.
The formula is: $M = \frac{\text{moles of solute}}{\text{volume of solution in liters}}$.
Given:
Moles of solute = $1 \; mol$.
Volume of solution = $500 \; mL = 0.5 \; L$.
Calculation:
$M = \frac{1 \; mol}{0.5 \; L} = 2 \; mol/L = 2 \; M$.
Therefore,the molarity of the solution is $2 \; M$.

(B) The molecular mass of $NaOH$ is calculated by adding the atomic masses of its constituent elements:
Atomic mass of $Na = 23 \text{ u}$
Atomic mass of $O = 16 \text{ u}$
Atomic mass of $H = 1 \text{ u}$
Molecular mass of $NaOH = 23 + 16 + 1 = 40 \text{ g mol}^{-1}$.

(D) To calculate the molarity $(M)$,we use the formula: $M = \frac{\text{moles of solute}}{\text{volume of solution in liters}}$.
First,calculate the molar mass of $NaOH$: $Na = 23, O = 16, H = 1$. So,$23 + 16 + 1 = 40 \ g/mol$.
Next,calculate the number of moles of $NaOH$: $\text{moles} = \frac{\text{given mass}}{\text{molar mass}} = \frac{8 \ g}{40 \ g/mol} = 0.2 \ mol$.
Convert the volume from $mL$ to $L$: $100 \ mL = 0.1 \ L$.
Finally,calculate molarity: $M = \frac{0.2 \ mol}{0.1 \ L} = 2 \ M$.
Therefore,the correct option is $D$.

(A) The concentration of a solution in Molarity $(M)$ is defined as the number of moles of solute dissolved per liter of solution.
Formula: $M = \frac{\text{Number of moles of solute}}{\text{Volume of solution in liters}}$
Given:
Number of moles of solute = $1\;\text{mole}$
Volume of solution = $2\;\text{liters}$
Calculation:
$M = \frac{1\;\text{mole}}{2\;\text{liters}} = 0.5\;M$
Therefore,the concentration of the solution is $0.5\;M$.

(C) The $pH$ scale was introduced by the Danish chemist $S.P.L. Sorensen$ in $1909$.
It is a logarithmic scale used to specify the acidity or basicity of an aqueous solution.

(C) Bases are substances that dissociate in aqueous solutions to produce hydroxyl ions $(OH^{-})$.
These $OH^{-}$ ions are responsible for the characteristic chemical properties of bases,such as their bitter taste,soapy feel,and ability to react with acids to form salt and water (neutralization reaction).

(A) Bases are substances that dissociate in water to produce hydroxyl ions $(OH^-)$.
These $OH^-$ ions are responsible for the characteristic chemical properties of bases,such as their bitter taste,soapy feel,and ability to react with acids to form salt and water (neutralization reaction).
In contrast,acids are responsible for producing hydronium ions $(H_3O^+)$ in aqueous solutions.

(B) In an aqueous solution at $25^{\circ}C$,the product of the concentrations of hydronium ions and hydroxide ions is constant,given by $[H_3O^+][OH^-] = 10^{-14}$.
For a neutral solution,$[H_3O^+] = [OH^-] = 10^{-7} \; M$.
For a basic solution,the concentration of hydroxide ions is greater than the concentration of hydronium ions,meaning $[OH^-] > [H_3O^+]$.
Since $[H_3O^+][OH^-] = 10^{-14}$,if $[OH^-] > 10^{-7} \; M$,then $[H_3O^+] < 10^{-7} \; M$.
Therefore,for a basic solution,the correct relation is $[OH^-] > 10^{-7} \; M$.

(A) In pure water or a neutral aqueous solution at $25^{\circ}C$,the concentration of hydronium ions $[H_3O^+]$ is equal to $10^{-7} \ M$.
An acidic solution is defined as a solution where the concentration of hydrogen ions (or hydronium ions) is higher than that of a neutral solution.
Therefore,for an acidic solution,$[H_3O^+] > 10^{-7} \ M$.
Since $pH = -\log[H_3O^+]$,a higher concentration of $[H_3O^+]$ results in a $pH$ value of less than $7$.

(B) The $pH$ of a solution is defined as the negative logarithm of the molar concentration of hydrogen ions $[H^+]$ or hydronium ions $[H_3O^+]$ present in the solution.
Mathematically,it is expressed as: $pH = -\log_{10}[H_3O^+]$.
Therefore,the correct definition refers to the negative logarithm to the base $10$.

(B) The $pH$ of a solution is defined by the formula: $pH = -\log_{10}[H_3O^+]$.
Given that $[H_3O^+] = 10^{-6} \ M$.
Substituting the value into the formula:
$pH = -\log_{10}(10^{-6})$.
Using the logarithmic property $\log(a^b) = b \cdot \log(a)$:
$pH = -(-6) \cdot \log_{10}(10)$.
Since $\log_{10}(10) = 1$,we get:
$pH = 6$.
Therefore,the $pH$ of the solution is $6$.

(C) The $pH$ of a solution is defined by the formula: $pH = -\log[H_3O^+]$.
Given that the concentration of hydronium ions is $[H_3O^+] = 1 \; M$.
Substituting the value into the formula: $pH = -\log(1)$.
Since $\log(1) = 0$,the $pH$ of the solution is $0$.

(C) Given: $[H_3O^+] = 10^{-5} \; M$.
First,calculate the $pH$ of the solution using the formula: $pH = -\log[H_3O^+]$.
$pH = -\log(10^{-5}) = 5$.
We know that at $25^{\circ}C$,the relationship between $pH$ and $pOH$ is $pH + pOH = 14$.
Substituting the value of $pH$: $5 + pOH = 14$.
Therefore,$pOH = 14 - 5 = 9$.

(C) The relationship between $pH$ and $pOH$ in an aqueous solution at $25^{\circ}C$ is given by the equation: $pH + pOH = 14$.
Given that $pH = 6$,we can substitute this value into the equation:
$6 + pOH = 14$
$pOH = 14 - 6$
$pOH = 8$.
Therefore,the correct answer is $8$.

(C) In pure distilled water at $298 \; K$,the water undergoes auto-ionization according to the equation: $2H_2O(l) \rightleftharpoons H_3O^+(aq) + OH^-(aq)$.
For pure water,the concentration of hydronium ions $[H_3O^+]$ is equal to the concentration of hydroxide ions $[OH^-]$.
The ionic product of water $(K_w)$ at $298 \; K$ is $1 \times 10^{-14}$.
Since $K_w = [H_3O^+][OH^-]$,and $[H_3O^+] = [OH^-]$,we can write $K_w = [H_3O^+]^2$.
Therefore,$[H_3O^+] = \sqrt{K_w} = \sqrt{1 \times 10^{-14}} = 1 \times 10^{-7} \; M$.

(C) At $298 \; K$,the ionic product of water $(K_w)$ is $1 \times 10^{-14}$.
In pure distilled water,the concentration of hydrogen ions $[H^+]$ is equal to the concentration of hydroxide ions $[OH^-]$.
Therefore,$[H^+] = [OH^-]$.
Since $K_w = [H^+][OH^-]$,we can write $K_w = [OH^-]^2$.
Substituting the value of $K_w$: $1 \times 10^{-14} = [OH^-]^2$.
Taking the square root on both sides: $[OH^-] = \sqrt{1 \times 10^{-14}} = 1 \times 10^{-7} \; M$.

(A) The $pH$ scale is a measure of the hydrogen ion concentration in a solution.
For an acidic solution,the concentration of hydrogen ions $(H^+)$ is greater than the concentration of hydroxide ions $(OH^-)$.
According to the $pH$ scale,a solution with a $pH$ value less than $7$ is considered acidic.
$A$ $pH$ value of $7$ indicates a neutral solution (like pure water),and a $pH$ value greater than $7$ indicates a basic (alkaline) solution.
Therefore,for an acidic aqueous solution,the $pH$ is always less than $7$ $(pH < 7)$.

(B) The $pH$ scale is a measure of the hydrogen ion concentration in a solution.
For a neutral aqueous solution at $25^{\circ}C$,the concentration of hydrogen ions $[H^+]$ is equal to the concentration of hydroxide ions $[OH^-]$,which is $10^{-7} \ M$.
Therefore,the $pH$ is calculated as $pH = -\log[H^+] = -\log(10^{-7}) = 7$.
Thus,a $pH$ value of $7$ indicates a neutral solution.

(B) The $pH$ scale is used to measure the acidity or alkalinity of an aqueous solution.
- $A$ solution with a $pH$ value of $7$ is considered neutral (e.g.,pure water).
- $A$ solution with a $pH$ value less than $7$ $(pH < 7)$ is acidic.
- $A$ solution with a $pH$ value greater than $7$ $(pH > 7)$ is basic or alkaline.
Therefore,for a basic aqueous solution,the $pH$ value is always greater than $7$.

(B) The $pH$ scale is used to measure the acidity or alkalinity of an aqueous solution.
- $A$ solution with a $pH$ value of $7$ is considered neutral (e.g.,pure water).
- $A$ solution with a $pH$ value less than $7$ $(pH < 7)$ is acidic.
- $A$ solution with a $pH$ value greater than $7$ $(pH > 7)$ is basic (alkaline).
Therefore,for a basic aqueous solution,the $pH$ value is always greater than $7$.

(A) The concentration of hydronium ions $[H_{3}O^{+}]$ is directly related to the acidity of a solution.
According to the $pH$ scale,$pH = -log[H_{3}O^{+}]$.
As the concentration of $[H_{3}O^{+}]$ ions increases,the $pH$ value decreases.
$A$ lower $pH$ value indicates a higher degree of acidity.
Therefore,an increase in $[H_{3}O^{+}]$ concentration makes the solution more acidic.

(C) The acidity of an aqueous solution is directly proportional to the concentration of hydronium ions $([H_{3}O^{+}])$.
As the concentration of $[H_{3}O^{+}]$ decreases,the solution becomes less acidic,meaning its acidity decreases.
This is also reflected in the $pH$ value,where a decrease in $[H_{3}O^{+}]$ leads to an increase in $pH$,indicating a shift towards a more neutral or basic state.

(B) The basicity of an aqueous solution is directly proportional to the concentration of hydroxide ions $[OH^-]$.
As the concentration of $[OH^-]$ increases,the solution becomes more basic,meaning its $pH$ value increases (moving further away from $7$ towards $14$).
Therefore,the basicity of the solution increases.

(A) The $pOH$ of a solution is defined by the formula: $pOH = -\log_{10}[OH^-]$.
Since $pOH$ is inversely proportional to the logarithm of the hydroxide ion concentration,as the concentration of $[OH^-]$ increases,the value of $\log_{10}[OH^-]$ also increases.
Because of the negative sign in the formula,an increase in the value of $\log_{10}[OH^-]$ leads to a decrease in the $pOH$ value.
Therefore,if the concentration of $[OH^-]$ increases,the $pOH$ decreases.

(B) The $pH$ scale is defined as $pH = -\log[H^+]$.
In an aqueous solution,the product of the concentrations of hydrogen ions and hydroxide ions is constant at a given temperature,represented by the ionic product of water: $[H^+][OH^-] = K_w = 10^{-14}$.
If the concentration of $[OH^-]$ increases,the concentration of $[H^+]$ must decrease to maintain the constant value of $K_w$.
Since $pH$ is inversely proportional to the concentration of $[H^+]$ (specifically,$pH = 14 + \log[OH^-]$),a decrease in $[H^+]$ concentration leads to an increase in the $pH$ value.
Therefore,as the basicity of the solution increases due to higher $[OH^-]$ concentration,the $pH$ value increases.

(A) The basicity of an aqueous solution is directly proportional to the concentration of hydroxide ions $[OH^{-}]$.
As the concentration of $[OH^{-}]$ ions decreases,the solution becomes less basic.
Therefore,the basicity of the solution decreases.

(B) The $pOH$ of a solution is defined by the formula: $pOH = -\log_{10}[OH^-]$.
Since the $pOH$ is inversely proportional to the logarithm of the hydroxide ion concentration,as the concentration of $[OH^-]$ decreases,the value of $\log_{10}[OH^-]$ becomes more negative.
Consequently,the negative sign in the formula makes the $pOH$ value increase.
Therefore,if the concentration of $[OH^-]$ decreases,the $pOH$ increases.

(A) The $pH$ of a solution is related to the concentration of $[H^+]$ ions by the formula $pH = -\log[H^+]$.
In an aqueous solution,the product of $[H^+]$ and $[OH^-]$ is constant at a given temperature,represented by the ion product of water,$K_w = [H^+][OH^-] = 10^{-14}$.
If the concentration of $[OH^-]$ decreases,the concentration of $[H^+]$ must increase to maintain the constant $K_w$.
As the concentration of $[H^+]$ increases,the $pH$ value decreases because $pH$ is inversely proportional to the logarithm of $[H^+]$ concentration.
Therefore,a decrease in $[OH^-]$ concentration leads to a decrease in $pH$.

(A) The $pH$ scale is defined as the negative logarithm of the hydrogen ion concentration,expressed as $pH = -\log[H^+]$.
This definition specifically requires the presence of hydrogen ions in a solvent,which is characteristic of aqueous solutions.
Therefore,the $pH$ scale is primarily applicable to aqueous solutions where water acts as the solvent to facilitate the dissociation of acids or bases.