Strength of Acids and Bases (pH Scale) Questions in English

(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.

(B) The $pH$ scale is a logarithmic scale used to measure the acidity or alkalinity of a solution.
$pH = -\log[H^+]$.
As the $pH$ value decreases,the concentration of hydrogen ions $[H^+]$ increases,making the solution more acidic.
Since $2 < 4$,the solution with $pH = 2$ has a higher concentration of $H^+$ ions compared to the solution with $pH = 4$.
Therefore,a solution with $pH = 2$ is more acidic than a solution with $pH = 4$.

(C) The $pH$ scale ranges from $0$ to $14$.
Solutions with a $pH$ less than $7$ are acidic,while solutions with a $pH$ greater than $7$ are basic.
As the $pH$ value increases above $7$,the concentration of hydroxide ions $(OH^-)$ increases,making the solution more basic.
Since $12 > 9$,the solution with $pH$ $12$ has a higher concentration of $OH^-$ ions than the solution with $pH$ $9$.
Therefore,a solution with $pH$ $12$ is more basic than a solution with $pH$ $9$.

(D) Blue litmus paper turns red in an acidic solution.
According to the $pH$ scale,acidic solutions have a $pH$ value less than $7$.
Therefore,the $pH$ range for an acidic solution is between $0$ and $7$.
Option $D$ is the correct answer.

(C) The $pH$ scale ranges from $0$ to $14$.
$pH$ paper (universal indicator) changes color based on the acidity or alkalinity of a solution.
$A$ pink or red color on $pH$ paper indicates a strongly acidic solution.
Among the given options,$pH = 2$ represents a strongly acidic solution,which would turn the $pH$ paper pink.
$pH = 7$ is neutral,while $pH = 8$ and $pH = 10$ are basic.

(B) The $pH$ scale ranges from $0$ to $14$.
$A$ $pH$ value of $7$ indicates that the solution is neutral.
Universal indicator or $pH$ paper shows a green or light green colour at $pH$ $7$.
Acidic solutions (below $pH$ $7$) turn the paper towards red/orange/yellow,while basic solutions (above $pH$ $7$) turn it towards blue/purple.

(A) $pH$ paper changes color based on the acidity or alkalinity of a solution.
$pH$ paper turns blue or purple in basic (alkaline) solutions,where the $pH$ value is greater than $7$.
Among the given options,$10$ is the only value greater than $7$,indicating a basic solution.
Therefore,the correct answer is $10$.

(B) The human body is highly sensitive to changes in $pH$. Most of the physiological processes and biochemical reactions in the human body occur within a narrow $pH$ range of $7.0$ to $7.8$. If the $pH$ of the body fluids,such as blood,deviates significantly from this range,it can lead to serious health complications.

(B) The $pH$ of soil is a critical factor for plant growth and fertility.
Most plants grow best in soil that is slightly acidic to neutral.
The ideal $pH$ range for most fertile agricultural land is generally considered to be between $6.5$ and $7.3$.
If the soil is too acidic or too alkaline,essential nutrients become unavailable to the plants,hindering their growth.

(A) The $pH$ scale is used to measure the acidity or alkalinity of a substance.
- $A$ $pH$ value of $7$ is considered neutral.
- $A$ $pH$ value less than $7$ indicates an acidic nature.
- $A$ $pH$ value greater than $7$ indicates a basic (alkaline) nature.
- For soil,a $pH$ range of $6.5$ to $7.3$ is generally considered neutral or optimal for most plants.
- Therefore,soil with a $pH$ value less than $6.5$ is considered acidic.

(C) Tooth decay starts when the $pH$ of the mouth falls below $5.5$.
This happens because bacteria present in the mouth produce acids by degrading sugar and food particles remaining in the mouth after eating.
These acids corrode the tooth enamel,which is made of calcium hydroxyapatite (a crystalline form of calcium phosphate),the hardest substance in the body.

(A) Acids are classified as strong or weak based on the extent of their ionization in an aqueous solution.
Strong acids are those that completely dissociate into ions in water (e.g.,$HCl$,$H_2SO_4$).
Weak acids are those that only partially dissociate into ions in water (e.g.,$CH_3COOH$,$H_2CO_3$).
Therefore,the correct basis for this classification is the degree of ionization.

(C) $NaOH$ is a strong base that dissociates completely in water as follows:
$NaOH \rightarrow Na^+ + OH^-$
Given the concentration of $NaOH$ is $1 \times 10^{-5} \; M$,the concentration of $OH^-$ ions $[OH^-]$ is $1 \times 10^{-5} \; M$.
We know that $pOH = -\log[OH^-]$.
$pOH = -\log(1 \times 10^{-5}) = 5$.
Using the relation $pH + pOH = 14$ at $25^{\circ}C$:
$pH = 14 - pOH = 14 - 5 = 9$.
Therefore,the $pH$ of the solution is $9$.

(D) $HCl$ is a strong acid that dissociates completely in water as: $HCl \rightarrow H^+ + Cl^-$.
Given the concentration of $HCl$ is $10^{-4} \; M$,the concentration of $H^+$ ions is $[H^+] = 10^{-4} \; M$.
The $pH$ of the solution is calculated as: $pH = -\log[H^+] = -\log(10^{-4}) = 4$.
Using the relationship $pH + pOH = 14$ at $25^{\circ}C$,we find $pOH = 14 - pH$.
Therefore,$pOH = 14 - 4 = 10$.

(B) strong acid is one that completely dissociates into its ions in an aqueous solution. Examples include $HCl$,$H_2SO_4$,and $HNO_3$.
$A$ weak acid is one that only partially dissociates into its ions in an aqueous solution. Acetic acid $(CH_3COOH)$ is a classic example of a weak acid because it does not ionize completely in water.
Therefore,the correct option is $B$.