Acids, Bases and Salts

Dry $HCl$ gas does not contain water, so it does not ionise to form hydrogen ions ($H^+$) or hydronium ions ($H_3O^+$). Acidic properties and colour change in litmus paper occur only in the presence of water.

• pH = 2: Strongly acidic ($pH < 7$).
• pH = 10: Basic ($pH > 7$).

Hydrogen gas ($H_2$) is evolved.
Test: When a burning matchstick or splinter is brought near the gas, it burns with a characteristic 'pop' sound.

Acids dissociate in water to produce free mobile hydronium ions ($H_3O^+$) and anions that carry electric current. Glucose ($C_6H_{12}O_6$) and alcohol ($C_2H_5OH$), although containing hydrogen, do not ionise in water to produce free ions and therefore cannot conduct electricity.

The dilution of concentrated acid is a highly exothermic process. If water is added to concentrated acid, the sudden heat generated may cause the acid mixture to splash out and cause chemical burns, or the glass container may break due to excessive local heating. Hence, acid must always be added slowly to water with constant stirring.

• Strong Acids: Acids that undergo complete dissociation in aqueous solution to give a high concentration of $H^+$ ions. Examples: Hydrochloric acid ($HCl$), Sulphuric acid ($H_2SO_4$).
• Weak Acids: Acids that undergo partial dissociation in water to give a lower concentration of $H^+$ ions. Examples: Acetic acid ($CH_3COOH$), Carbonic acid ($H_2CO_3$).

(a) Tooth Decay: Tooth enamel (calcium hydroxyapatite) begins to corrode when mouth pH falls below 5.5 due to acid produced by bacteria fermenting sugars. Using basic toothpaste neutralises excess acid.

(b) Digestion: The stomach produces $HCl$ (pH $\approx 1.2–2.0$) which activates pepsin. Excess acid causes acidity, treated using antacids like $Mg(OH)_2$ (Milk of Magnesia) or $NaHCO_3$.

(c) Self-Defence: Honeybee stings and nettle leaves inject methanoic (formic) acid ($HCOOH$), causing burning pain. Applying mild bases like baking soda ($NaHCO_3$) or dock leaf juice neutralises the sting.

When electricity is passed through an aqueous solution of sodium chloride (brine), it decomposes into sodium hydroxide, chlorine gas, and hydrogen gas.

• $NaOH$ (at cathode): De-greasing metals, soap and detergent manufacturing, paper making.
• $Cl_2$ (at anode): Water treatment (disinfectant), PVC, CFCs, bleaching powder.
• $H_2$ (at cathode): Fuels, margarine, manufacturing ammonia for fertilisers.

Water of crystallisation is the fixed number of water molecules chemically bonded in one formula unit of a salt in crystalline form.

• Hydrated Copper Sulphate (Blue Vitriol): $CuSO_4 \cdot 5H_2O$ (contains 5 water molecules per formula unit; turns white anhydrous $CuSO_4$ upon heating).
• Washing Soda: $Na_2CO_3 \cdot 10H_2O$ (contains 10 water molecules).
• Gypsum: $CaSO_4 \cdot 2H_2O$ (contains 2 water molecules).

(a) Bleaching Powder ($CaOCl_2$):
• Preparation: Action of chlorine gas on dry slaked lime: $Ca(OH)_2 + Cl_2 o CaOCl_2 + H_2O$
• Uses: (1) Bleaching cotton/linen in textile industry, (2) Disinfecting drinking water.

(b) Baking Soda ($NaHCO_3$):
• Preparation: Solvay process: $NaCl + H_2O + CO_2 + NH_3 o NH_4Cl + NaHCO_3$
• Uses: (1) Making baking powder (with mild edible tartaric acid), (2) Antacid ingredient, (3) Soda-acid fire extinguishers.

(c) Plaster of Paris ($CaSO_4 \cdot \frac{1}{2}H_2O$):
• Preparation: Heating gypsum at $373K$ ($100^\circC$): $CaSO_4 \cdot 2H_2O \xr→{373K} CaSO_4 \cdot \frac{1}{2}H_2O + 1\frac{1}{2}H_2O$
• Uses: (1) Setting fractured bones in position, (2) Making statues, toys, and decorative false ceilings.

Banana

Apples (malic acid), oranges (citric & ascorbic acid), and strawberries (citric & malic acid) are distinctly acidic. Ripe bananas have a mild alkaline-forming nature (higher pH relative to citrus fruits) and are known to soothe gastric acidity.

Acid: Lemon juice, Curd; Base: Baking soda, Broccoli

Lemon juice contains citric acid and curd contains lactic acid (sour taste, acidic). Baking soda is sodium hydrogen carbonate ($\text{NaHCO}_3$, basic salt) and broccoli contains basic/alkaline minerals.

Blue litmus paper: no colour change; Red litmus paper: changes colour

Basic solutions turn moist red litmus paper blue, while blue litmus paper remains unaffected (no colour change).

Methyl orange: red; Phenolphthalein: colourless

A solution with $\text{pH} = 2$ is strongly acidic. Methyl orange turns red in acidic medium ($\text{pH} < 3.1$), and phenolphthalein remains completely colourless in acidic medium (turning pink only in alkaline medium $\text{pH} > 8.3$).

X: zinc; Y: hydrogen

Zinc metal reacts with dilute sulphuric acid to produce zinc sulphate salt and liberate hydrogen gas: $$\text{Zn}(s) + \text{H}_2\text{SO}_4(aq) \longrightarrow \text{ZnSO}_4(aq) + \text{H}_2(g) \uparrow$$ Copper does not displace hydrogen from non-oxidising dilute acids.

By a burning candle

Hydrogen gas is highly combustible. When a lighted splint or burning candle is brought near the mouth of the test tube, hydrogen gas burns rapidly with a characteristic "pop" sound.

White precipitate is formed

Carbon dioxide gas reacts with lime water (calcium hydroxide solution, $\text{Ca(OH)}_2$) to precipitate insoluble white calcium carbonate, turning the lime water milky: $$\text{Ca(OH)}_2(aq) + \text{CO}_2(g) \longrightarrow \text{CaCO}_3(s) \downarrow (\text{white ppt}) + \text{H}_2\text{O}(l)$$

X: hydrochloric acid; Y: carbon dioxide

Metal carbonates react with acids (such as $\text{HCl}$) to yield a metallic salt, water, and carbon dioxide gas ($\text{CO}_2$): $$\text{Metal Carbonate} + \text{Acid} \longrightarrow \text{Salt} + \text{H}_2\text{O} + \text{CO}_2 \uparrow$$

$$\text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O}$$

One mole of strong monoprotic acid ($\text{HCl}$) neutralises one mole of strong monobasic alkali ($\text{NaOH}$) to form neutral table salt ($\text{NaCl}$) and water in an equimolar $1:1$ stoichiometric ratio.

Beaker 1: increase; Beaker 2: reduce

In Beaker 1 (acidic $\text{HCl}$), adding alkali $\text{NaOH}$ consumes $\text{H}^+$ ions, increasing the pH towards 7. In Beaker 2 (basic $\text{NaOH}$), adding acid $\text{HCl}$ consumes $\text{OH}^-$ ions, decreasing the pH towards 7.

Salt

Metallic oxides (such as $\text{CuO}$, $\text{MgO}$) are basic in nature. When they react with acids, they undergo neutralisation to produce salt and water: $$\text{CuO}(s) + 2\text{HCl}(aq) \longrightarrow \text{CuCl}_2(aq) + \text{H}_2\text{O}(l)$$

Metal (Basic oxide)

Substances that react with acids to produce salt and water are bases. Metallic oxides are basic oxides, formed by metals (element P = metal). (Note: CBSE resource manual key records option a).

Acid: $\text{H}^+$; Base: $\text{OH}^-$

According to Arrhenius theory, acids ionise in water to liberate hydrogen / hydronium ions ($\text{H}^+ / \text{H}_3\text{O}^+$), whereas bases dissociate to produce hydroxide ions ($\text{OH}^-$). These mobile charge carriers facilitate electrical conductivity.

The distilled water does not have ions present in it

Pure distilled water consists essentially of neutral covalent water molecules with negligible dissociation ($[\text{H}^+] = [\text{OH}^-] = 10^{-7}\text{ M}$). Lacking free mobile ions to carry electric current, it acts as an electrical insulator.

C and D

On the pH scale at $25^\circ\text{C}$, $\text{pH} = 7$ represents neutral, $\text{pH} < 7$ represents acidic, and $\text{pH} > 7$ represents basic/alkaline. Solutions C ($\text{pH} \approx 9$) and D ($\text{pH} \approx 12$) both have $\text{pH} > 7$ and are alkaline.

Strongest Acid: A; Strongest Base: E

The lowest pH value corresponds to the highest concentration of $\text{H}^+$ ions (Solution A at $\text{pH} \approx 1$ is the strongest acid). The highest pH value corresponds to the greatest concentration of $\text{OH}^-$ ions (Solution E at $\text{pH} \approx 14$ is the strongest base).

5.5–7 (Near neutral / slightly acidic to neutral)

Most common agricultural crops thrive when soil pH is between 5.5 and 7.0, where essential mineral nutrients like nitrogen, phosphorus, and potassium are most bioavailable for root absorption.

Sodium hydrogen carbonate (Baking soda, mild base)

Baking soda solution ($\text{NaHCO}_3$, weak base $\text{pH} \approx 8.4$) safely neutralises acidic bee or insect venom without harming biological tissue. (Note: Although CBSE question text references option b in its raw key, concentrated $\text{NaOH}$ is highly caustic; baking soda is the correct textbook first-aid antacid treatment).

$\text{Na}^+$: positive radical (cation); $\text{Cl}^-$: negative radical (anion)

In sodium chloride ($\text{NaCl}$), the basic radical derived from the base ($\text{NaOH}$) is the positively charged sodium cation $\text{Na}^+$, and the acid radical derived from $\text{HCl}$ is the negatively charged chloride anion $\text{Cl}^-$.

$\text{HCl}$ (Strong Acid) and $\text{NH}_4\text{OH}$ (Weak Base)

A salt formed from a strong acid and a weak base (such as ammonium chloride, $\text{NH}_4\text{Cl}$) undergoes cationic hydrolysis in aqueous solution, yielding excess $\text{H}_3\text{O}^+$ ions and producing an acidic solution ($\text{pH} < 7$, typically around 5.0–5.5).

Decomposition reaction (Electrolytic decomposition)

The chlor-alkali process is an electrolytic decomposition reaction where brine ($2\text{NaCl} + 2\text{H}_2\text{O}$) is split by electric current into simpler substances: sodium hydroxide ($2\text{NaOH}$), chlorine gas ($\text{Cl}_2$), and hydrogen gas ($\text{H}_2$).

$$2\text{NaCl}(aq) + 2\text{H}_2\text{O}(l) \longrightarrow 2\text{NaOH}(aq) + \text{Cl}_2(g) + \text{H}_2(g)$$

Chlorine gas is liberated at the positive anode, hydrogen gas is liberated at the negative cathode, and sodium hydroxide solution forms near the cathode.

$$\text{X: NH}_4\text{Cl} \text{ (Ammonium chloride)}; \text{ Y: NaHCO}_3 \text{ (Sodium hydrogen carbonate)}$$

In the Solvay process, cold saturated brine reacts with carbon dioxide and ammonia to yield ammonium chloride and sparingly soluble sodium hydrogen carbonate (baking soda).

From blue to white

Hydrated copper sulphate ($\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$) is deep blue due to coordinate bonds with 5 water molecules. Heating drives off the water of crystallisation, forming anhydrous copper sulphate ($\text{CuSO}_4$) which is a white powder.

Liquid 2 (pH 12)

Hydrogen ion concentration $[\text{H}^+]$ is inversely proportional to the pH value ($\text{pH} = -\log_{10}[\text{H}^+]$). As pH increases, $[\text{H}^+]$ decreases. Liquid 2 has the highest pH value (12), which corresponds to the lowest concentration of hydrogen ions ($[\text{H}^+] = 10^{-12}\text{ M}$).

Evaluations:

1. Gloves keep the hands warm: No — Safety PPE in a chemistry lab is not for thermal comfort.
2. Forceps provide better grip than bare hands: No — Forceps prevent chemical contamination and skin contact, not mere grip.
3. Gloves protect hands from corrosive liquids: Yes — Strong acids and alkalis are highly caustic and corrosive, capable of causing severe chemical burns on contact with bare skin.

Colour: Green

Explanation: Pure distilled water has a neutral $\text{pH} = 7$ at $25^\circ\text{C}$ with equal concentrations of hydrogen and hydroxide ions ($[\text{H}^+] = [\text{OH}^-] = 10^{-7}\text{ M}$). On universal indicator pH paper, neutral pH 7 produces a distinct green colour.

Chemical Nature: Salt

Explanation: $\text{HX}$ represents an acid (composed of hydrogen and non-metallic anion $\text{X}^-$) and $\text{YOH}$ represents a base/hydroxide (composed of metal cation $\text{Y}^+$ and hydroxide $\text{OH}^-$). The reaction between an acid and a base produces water ($\text{HOH}$) and an ionic compound formed by the metal cation and non-metal anion ($\text{YX}$), which is a salt.

It gets hard when mixed with water.

Plaster of Paris (calcium sulphate hemihydrate, $\text{CaSO}_4 \cdot \frac{1}{2}\text{H}_2\text{O}$) on mixing with water hydrates and re-crystallises into gypsum ($\text{CaSO}_4 \cdot 2\text{H}_2\text{O}$), setting into a hard solid mass within 10 to 15 minutes.

Evaluations:

1. Contains moisture / water of crystallisation: Yes — Evaporated water condensed into visible liquid droplets on the lid.
2. Changes colour on heating: Yes — The salt changed from hydrated pale green ($\text{FeSO}_4 \cdot 7\text{H}_2\text{O}$) to anhydrous white ($\text{FeSO}_4$).
3. Changes to a completely new chemical compound: No — Because adding water instantly restored the green hydrated salt, the reversible dehydration did not decompose the chemical into iron oxide.

Explanation:

• Powder colour change: The green crystal powder will still lose its water of crystallisation on heating and turn into a whitish anhydrous powder.
• Water droplets: In an uncovered beaker, the released water vapour will escape freely into the surrounding atmosphere instead of condensing into visible droplets on a cover.

Electrical conductivity of aqueous solutions.

Mike was testing whether aqueous acids and bases dissociate into free mobile ions that carry electric current and complete the circuit, causing the lamp to glow.

No, the lamp will not glow.

Reason: Distilled water is pure and consists of un-ionised neutral covalent water molecules. Because it does not contain free mobile ions (electrolytes) to conduct electric charge across the electrodes, no current flows to light the bulb.

$$\text{X} = \text{Hydrogen gas } (\text{H}_2)$$

Reactive metals displace hydrogen from dilute mineral acids, producing the corresponding metal salt and liberating diatomic hydrogen gas: $$\text{Zn}(s) + 2\text{HCl}(aq) \longrightarrow \text{ZnCl}_2(aq) + \text{H}_2(g) \uparrow$$

Acids turn blue litmus to red and bases turn red litmus to blue.

Natural litmus dye extracted from lichens is purple in neutral solution, turning red in acidic conditions ($\text{pH} < 7$) and blue in basic/alkaline conditions ($\text{pH} > 7$).

hydrogen ions ($\text{H}^+ / \text{H}_3\text{O}^+$)

The characteristic acidic properties of any acid in water (sour taste, turning blue litmus red, reactivity with active metals and carbonates) arise specifically from solvated hydronium / hydrogen ions released upon ionisation.

Observation & Explanation:

• Observation: The wet red litmus paper turns blue.
• Explanation: Magnesium oxide is a metallic oxide. It reacts with water on the moist litmus paper to produce magnesium hydroxide, a base: $$\text{MgO}(s) + \text{H}_2\text{O}(l) \longrightarrow \text{Mg(OH)}_2(aq)$$ The basic hydroxide ions ($\text{OH}^-$) turn red litmus blue.

Word Equation:

$$\text{Ammonia} + \text{Nitric acid} \longrightarrow \text{Ammonium nitrate}$$

Chemical Equation: $\text{NH}_3(aq) + \text{HNO}_3(aq) \longrightarrow \text{NH}_4\text{NO}_3(aq)$

Acidic salt.

Ammonium nitrate ($\text{NH}_4\text{NO}_3$) is formed by the neutralisation of a strong acid (nitric acid, $\text{HNO}_3$) with a weak base (ammonium hydroxide, $\text{NH}_4\text{OH}$). Its aqueous solution undergoes cationic hydrolysis, resulting in $\text{pH} < 7$.

Name: Hydrated Copper(II) Sulphate (Blue Vitriol)
Chemical Formula: $$\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$$

Explanation: On heating, hydrated copper sulphate loses its 5 chemically bonded molecules of water of crystallisation, converting from the blue pentahydrate crystalline complex into anhydrous copper sulphate ($\text{CuSO}_4$), which is a colourless/white amorphous powder.

pH: Approximately 13 to 14

Sodium hydroxide is a strong base that dissociates completely in water to produce a high concentration of hydroxide ions ($\text{OH}^-$), placing it at the top extreme of the pH scale.

The pH decreases towards 7.

Dilution reduces the concentration of hydroxide ions ($\text{OH}^-$) per unit volume, decreasing alkalinity and therefore lowering the pH value.

Decreases.

Adding more water increases the total volume of the solution while keeping the number of $\text{OH}^-$ ions constant, which directly decreases the concentration of $\text{OH}^-$ ions per unit volume.

Thermal decomposition reaction.

A single chemical reactant breaks down into multiple simpler product substances when thermal energy (heat) is supplied.

Balanced Symbol Equation:

$$2\text{NaHCO}_3(s) \xrightarrow{\Delta} \text{Na}_2\text{CO}_3(s) + \text{H}_2\text{O}(g) + \text{CO}_2(g) \uparrow$$

Balanced Chemical Equation:

$$\text{Na}_2\text{CO}_3(s) + 10\text{H}_2\text{O}(l) \longrightarrow \text{Na}_2\text{CO}_3 \cdot 10\text{H}_2\text{O}(s)$$

Typical pH values:

• A. Pure water: $\text{pH} = 7$ (neutral)
• B. Strong acid: $\text{pH} = 0 \text{ to } 2$ (e.g., concentrated $\text{HCl}$, $\text{H}_2\text{SO}_4$)
• C. Weak base: $\text{pH} = 8 \text{ to } 10$ (e.g., $\text{NH}_4\text{OH}$, $\text{NaHCO}_3$)

Definition:

pH is a measure of the acidic or basic nature of an aqueous solution, defined mathematically as the negative logarithm (to base 10) of the hydrogen ion concentration in moles per litre: $$\text{pH} = -\log_{10}[\text{H}^+]$$ In German, 'p' stands for potenz, meaning power of hydrogen ions.

pH meter (or digital pH probe / sensor)

A digital pH meter measures electrical potential difference across a glass electrode to provide an accurate numerical pH reading to two decimal places.

Formation & Environmental Impacts of Acid Rain:

• Formation: Combustion of fossil fuels in power stations and vehicle engines releases sulfur dioxide ($\text{SO}_2$) and nitrogen oxides ($\text{NO}_x$). These acidic gases react with atmospheric water vapour, oxygen, and sunlight to form sulfuric acid ($\text{H}_2\text{SO}_4$) and nitric acid ($\text{HNO}_3$), lowering rain pH below 5.6.
• Aquatic ecosystems: Acidifies rivers and lakes, leaching toxic aluminium ions from soil into water bodies and killing fish and aquatic life.
• Soil fertility: Strips soils of essential plant nutrients (calcium, magnesium), stunting forest and crop growth.
• Historical monuments: Corrodes carbonate stone structures like the Taj Mahal (marble cancer): $$\text{CaCO}_3(s) + \text{H}_2\text{SO}_4(aq) \longrightarrow \text{CaSO}_4(s) + \text{H}_2\text{O}(l) + \text{CO}_2(g)$$

Red litmus solution

Sodium hydroxide is a strong alkali and turns red litmus solution blue. Water is neutral ($\text{pH} = 7$) and causes no colour change in red litmus.

Explanation:

In the electrochemical reactivity series, copper is more reactive than silver. If silver nitrate solution ($\text{AgNO}_3$) is stored in a copper container, copper displaces silver, corroding holes in the vessel and contaminating the solution: $$\text{Cu}(s) + 2\text{AgNO}_3(aq) \longrightarrow \text{Cu(NO}_3)_2(aq) + 2\text{Ag}(s)$$