Acids, Bases, and Salts

The governing concept is that the pH scale is logarithmic: pH = −log[H⁺]. The difference between pH 2 and pH 5 is 3 units. Therefore the hydrogen-ion concentration ratio is 10³ = 1000. The pH 2 solution has one thousand times more H⁺ ions and is consequently one thousand times more acidic, so option C is correct. A difference of three is not a threefold acidity ratio.

The governing concepts are the pH scale and partial neutralisation. Both pH 3 and pH 4 are below 7, so the solution remains acidic. However, increasing pH by one unit means the hydrogen-ion concentration has fallen by a factor of ten: [H⁺] at pH 4 is one-tenth of that at pH 3. Therefore acidity decreased but did not disappear, making option A correct. It did not become a strong base.

The governing concept is the logarithmic pH scale: pH = −log₁₀[H⁺]. A rise from pH 2 to pH 3 is an increase of one pH unit, so the hydrogen-ion concentration becomes 10 times smaller. For example, it changes from about 10⁻² mol/L to 10⁻³ mol/L. The solution is still acidic because its pH is below 7, but its acidity, measured through hydrogen-ion concentration, decreases tenfold. Therefore A is correct.

Salt hydrolysis explains the observation. A salt produced from a strong acid and a weak base commonly gives an acidic aqueous solution because its positive ion hydrolyses water and increases hydrogen-ion concentration. Thus option A is the best conclusion, although the wording “may be” is important because solution behaviour also depends on relative strengths and concentration. Strong acid–strong base salts are generally neutral.

Blue litmus turning red establishes that the unknown solution is acidic. Acids react with zinc, an active metal, to form a salt and hydrogen gas: acid + Zn → salt + H2. The same acid reacts with sodium carbonate to form a salt, water, and carbon dioxide: acid + carbonate → salt + H2O + CO2. Therefore option A correctly identifies the two gases.

Use the definition pH = −log[H+]. If the hydrogen-ion concentration decreases by a factor of 100, that is a decrease of 10², so the pH increases by 2 units. Starting from pH 3, the new value is 3 + 2 = 5. Therefore option B is correct. A decrease in hydrogen-ion concentration cannot lower the pH to 1, while the other listed values imply a much larger change.

The pH scale is logarithmic, and pOH = 14 − pH at 25°C. At pH 13, pOH is 1; at pH 11, pOH is 3. Thus the hydroxide-ion concentration changes from about 10⁻¹ M to 10⁻³ M, a decrease by 10², or 100 times. The solution remains basic at pH 11, but it is much less basic. Therefore option A is correct.

Neutralisation depends on the stoichiometric amounts of hydrogen and hydroxide ions, not merely on the volumes mixed. Equal volumes can have different concentrations or different strengths, so one reactant may remain in excess. For example, equal volumes of 0.1 M acid and 0.01 M base cannot neutralise completely. The final pH therefore need not be 7. Hence option A is the scientific reason.

The governing relationship is pH = −log[H⁺]. A tenfold decrease in hydrogen-ion concentration raises pH by one unit. Ten thousand equals 10⁴, so making the solution 10,000 times less acidic means [H⁺] decreases by four powers of ten and pH rises by four units. Starting from pH 2: 2 + 4 = 6. Thus option B is correct. This assumes the stated change refers to hydrogen-ion concentration and ignores minor water-ion effects.

A pH of 6 is below 7, and the slight red colour of blue litmus confirms that the salt solution is mildly acidic. In school-level salt hydrolysis, a salt formed from a strong acid and a weak base generally gives an acidic solution because the conjugate acid of the weak base hydrolyses in water. A strong base with a weak acid usually gives a basic solution, while strong acid plus strong base is approximately neutral. Hence option B is correct.

In the chlor-alkali process, electrolysis of concentrated brine produces chlorine at the anode, hydrogen at the cathode and sodium hydroxide in solution. These products must be separated because chlorine can react with sodium hydroxide, and chlorine and hydrogen can form a hazardous mixture under suitable conditions. A membrane or another separation arrangement limits such contact. Thus unwanted reactions and safety risks, not insolubility or freezing, explain option A.

Gypsum is calcium sulphate dihydrate, CaSO₄·2H₂O. Controlled heating removes part of its water of crystallisation and produces plaster of Paris, calcium sulphate hemihydrate, CaSO₄·½H₂O. If heating is excessive, more water can be removed and anhydrous calcium sulphate may form. That material does not hydrate and set as effectively, so the plaster quality falls. Hence option A is correct.

The bleaching action of bleaching powder depends mainly on its available or active chlorine, which produces oxidising species in the presence of moisture. In a moist, open environment, the substance can slowly react with atmospheric moisture and carbon dioxide and lose available chlorine. Its oxidising capacity therefore decreases during storage. It does not turn into sodium metal, gypsum, or acid rain, so option A is correct.

The chlor-alkali process electrolyses concentrated aqueous sodium chloride, or brine. At the cathode, water gains electrons and produces hydrogen gas and hydroxide ions. At the anode, chloride ions produce chlorine gas. The sodium ions remaining in solution combine with hydroxide ions to form sodium hydroxide. Therefore option B is correct; calcium hydroxide, oxygen, carbon dioxide, and ammonium hydroxide are not the products in this process.

Dilution of concentrated acid with water is highly exothermic, so a considerable amount of heat can be released. Adding acid slowly to a larger quantity of water allows the heat to disperse through the water and prevents a sudden local temperature rise. This reduces violent boiling and dangerous splashing. Option A is correct. The safe laboratory rule is always ‘add acid to water’, never water to concentrated acid, with stirring and suitable protection.

The governing idea is salt hydrolysis and the relative strength of the parent acid and base. Ammonium chloride comes from a strong acid and weak base; the ammonium ion hydrolyses and makes the solution acidic. Sodium carbonate comes from a weak acid and strong base; carbonate hydrolysis produces hydroxide ions, making it basic. Thus A is correct.

Acidic behaviour in this test requires the gas to dissolve in water and ionise, producing hydrogen ions or hydronium ions that interact with the indicator. Dry litmus provides no water film, so the necessary ions are not formed at the indicator surface and no colour change is observed. Therefore A is correct; the gas still has chemical properties.

Using $\text{pH}=-\log_{10}[H^+]$, we have $[H^+] \propto 10^{-\text{pH}}$.
So
\[
\frac{[H^+]_{\text{pH }1}}{[H^+]_{\text{pH }4}}=10^{(4-1)}=10^3=1000.
\]
Hence the first solution has $1000$ times greater hydrogen ion concentration.

The governing concept is salt hydrolysis and the strength of the parent acid and base. A pH of 9 and the change of red litmus to blue show that the salt solution is basic. Salts formed from a weak acid and a strong base commonly give basic solutions because the conjugate base of the weak acid hydrolyses water and produces hydroxide ions. Therefore option B is correct. A strong acid-weak base salt is generally acidic, while a strong acid-strong base salt is approximately neutral.

Litmus changes colour because of acidic or basic species present in an aqueous medium, not simply because a dry gas touches the paper. Turning red litmus blue demonstrates a basic effect, while the absence of a change on dry litmus shows that moisture is required for the gas to dissolve, ionise, or otherwise produce the relevant basic species. The safest conclusion is therefore that the gas forms a basic solution in moisture, making option A correct. The observation does not prove that the gas is a metal or always neutral.

The pH scale is logarithmic, so a three-unit decrease corresponds to a factor of 10³ in hydrogen-ion concentration. For basicity, the relevant hydroxide-ion concentration changes oppositely: pH 14 corresponds to pOH 0, while pH 11 corresponds to pOH 3. Thus [OH⁻] falls from about 1 M to 10⁻³ M, a thousandfold decrease. The solution remains basic, but much less strongly basic; A is correct.

Equal volume does not mean equal chemical amount. The base may have a higher concentration, may be stronger, or may provide more moles of effective OH⁻ than the acid provides moles of H⁺. After neutralisation, excess OH⁻ remains, making the mixture basic. Neutralisation does occur, but it may be incomplete with respect to the base. Thus A is correct; the other statements are chemically false.

The governing relation is pH = −log₁₀[H⁺]. Therefore, an increase of one pH unit means a tenfold decrease in hydrogen-ion concentration. The change from pH 2 to pH 5 is 5 − 2 = 3 units, so [H⁺] decreases by 10³ = 1000 times. Equivalently, the concentration changes from 10⁻² to 10⁻⁵ mol L⁻¹. Hence option C is correct; A and B count too few units, while D reverses the direction.

A pH of 10 indicates that the salt solution is basic. In the usual school-level classification, a salt formed from a weak acid and a strong base undergoes hydrolysis of its anion, producing OH⁻ ions and hence a basic solution; sodium acetate is a familiar example. A strong acid–weak base salt is generally acidic, a strong acid–strong base salt is approximately neutral, and option D does not describe salt formation. Therefore B is correct.

For basic solutions, $\text{pOH}=14-\text{pH}$ and $[OH^-] \propto 10^{-\text{pOH}}$.
Initial: $\text{pOH}=14-13=1$.
Final: $\text{pOH}=14-10=4$.
So
\[
\frac{[OH^-]_{\text{initial}}}{[OH^-]_{\text{final}}}=10^{(4-1)}=10^3=1000.
\]
Thus the OH$^-$ concentration (basicity) decreased by a factor of $1000$.