Acids Bases and Salts — Maharashtra (SSC) Class 9 Science Solutions (Free)

Acids are substances that release hydrogen ions (H⁺) when dissolved in water. Examples: hydrochloric acid (HCl), sulphuric acid (H₂SO₄), citric acid (in lemons). Bases are substances that release hydroxide ions (OH⁻) when dissolved in water (an alkali is a base that is soluble in water). Examples: sodium hydroxide (NaOH), ammonia (NH₃). Why the properties occur: - Sour taste (acids) / bitter taste (bases): these tastes are the tongue's direct sensory response to the concentration of H⁺ or OH⁻ ions present — it is literally how our taste receptors detect these ions. - Litmus colour change: litmus is a natural dye whose molecular structure changes shape depending on the surrounding H⁺ concentration. In acid (high H⁺), the molecule takes a form that reflects red light; in base (low H⁺, high OH⁻), it takes a form that reflects blue light — the colour change is a direct visual signal of the ion concentration. - Neutralisation: when an acid and a base are mixed, the H⁺ ions from the acid combine with the OH⁻ ions from the base to form neutral water (H⁺ + OH⁻ → H₂O), which is why acid + base → salt + water is called a neutralisation reaction — the characteristic ions of both are consumed.

pH is a numerical scale (0 to 14) that measures how acidic or basic a solution is, defined as pH = −log₁₀[H⁺], where [H⁺] is the concentration of hydrogen ions in mol/L. - pH < 7 → acidic (more H⁺ ions) - pH = 7 → neutral (pure water) - pH > 7 → basic/alkaline (fewer H⁺ ions, more OH⁻) Given [H⁺] = 10⁻³ mol/L: pH = −log₁₀(10⁻³) = −(−3) = 3 Since pH = 3 is less than 7, the solution is acidic. In fact, since each whole step down the pH scale represents a 10-fold increase in H⁺ concentration, a pH of 3 means the solution has 10,000 times more H⁺ ions than a neutral solution at pH 7.

In any aqueous solution at room temperature, pH and pOH always add up to 14: pH + pOH = 14 Given pH = 4: pOH = 14 − 4 = 10 To find [OH⁻], use pOH = −log₁₀[OH⁻], which rearranges to [OH⁻] = 10⁻ᵖᴼᴴ: [OH⁻] = 10⁻¹⁰ mol/L This confirms the solution is strongly acidic — a very low [OH⁻] of 10⁻¹⁰ mol/L, far below the neutral value of 10⁻⁷ mol/L.

Salts are ionic compounds formed when the H⁺ ion of an acid is replaced by a metal ion (or ammonium ion) — most commonly through the neutralisation reaction: Acid + Base → Salt + Water. Examples: - NaCl (common salt), from HCl + NaOH - CuSO₄ (copper sulphate), from H₂SO₄ + Cu(OH)₂ - KNO₃ (potassium nitrate), from HNO₃ + KOH Are all salts neutral? No — a salt's pH depends on the STRENGTH of the acid and base it came from: - Salt of a strong acid + strong base (e.g. NaCl, from HCl + NaOH) → neutral, pH ≈ 7. - Salt of a strong acid + weak base (e.g. NH₄Cl, from HCl + NH₃) → acidic, pH < 7. - Salt of a weak acid + strong base (e.g. Na₂CO₃, from H₂CO₃ + NaOH) → basic, pH > 7. So "salt" does not automatically mean neutral — it is a common misconception that all salts behave like table salt.

Indicators are substances that change colour depending on whether they are in an acidic or a basic solution, letting us identify which one we're dealing with without needing lab equipment. Three natural indicators: 1. Litmus (extracted from lichen) — turns red in acid, blue in base. The most commonly used indicator in school labs, usually as litmus paper. 2. Turmeric — its yellow colour stays yellow in acidic or neutral solutions but turns reddish-brown in a basic solution; this is why a turmeric stain on cloth often turns reddish when treated with soap (soap is mildly basic). 3. Red cabbage juice — a versatile natural indicator that turns red/pink in strongly acidic solutions, purple in neutral solutions, and green to yellow in basic solutions, making it useful for showing a wider range on the pH scale than litmus alone. (Methyl orange and phenolphthalein are also common indicators, but these are synthetic, not natural.)

Full equation: HCl + NaOH → NaCl + H₂O What happens at the ionic level: Both HCl and NaOH are strong electrolytes, so in water they exist almost entirely as separate ions: HCl → H⁺ + Cl⁻, and NaOH → Na⁺ + OH⁻. The actual reaction that occurs is between the H⁺ and OH⁻ ions only: H⁺ + OH⁻ → H₂O The Na⁺ and Cl⁻ ions do not react with each other — they remain in solution as "spectator ions," and simply stay together as dissolved NaCl once the water evaporates. This reaction is exothermic (releases heat), which is why mixing a strong acid and a strong base in the lab often makes the container noticeably warm. Since the acid and base are both strong, the resulting salt (NaCl) is neutral, with pH ≈ 7.

NaCl (sodium chloride): NEUTRAL. Formed from a strong acid (HCl) and a strong base (NaOH) — neither ion (Na⁺ or Cl⁻) reacts further with water, so the solution stays at pH ≈ 7. NH₄Cl (ammonium chloride): ACIDIC. Formed from a strong acid (HCl) and a weak base (NH₃). The ammonium ion (NH₄⁺) partially reacts with water (hydrolyses), releasing extra H⁺ ions: NH₄⁺ + H₂O ⇌ NH₃ + H₃O⁺ — this extra H⁺ makes the solution acidic, pH < 7. Na₂CO₃ (sodium carbonate, washing soda): BASIC. Formed from a weak acid (H₂CO₃, carbonic acid) and a strong base (NaOH). The carbonate ion (CO₃²⁻) hydrolyses in water, releasing extra OH⁻ ions: CO₃²⁻ + H₂O ⇌ HCO₃⁻ + OH⁻ — this extra OH⁻ makes the solution basic, pH > 7, which is why washing soda solution feels slippery, a common indicator of a basic solution.

Passing CO₂ through lime water (a solution of calcium hydroxide, Ca(OH)₂): CO₂ + Ca(OH)₂ → CaCO₃↓ + H₂O Calcium carbonate (CaCO₃) is insoluble in water, so it forms as a white precipitate, turning the clear lime water milky/turbid. This colour change is the standard laboratory test used to confirm the presence of CO₂ gas. What happens with EXCESS CO₂: If CO₂ continues to be passed well beyond this point, the white precipitate reacts further and dissolves: CaCO₃ + CO₂ + H₂O → Ca(HCO₃)₂ Calcium bicarbonate, Ca(HCO₃)₂, is soluble in water, so the milky solution turns clear again. This is an important two-stage test: cloudiness confirms CO₂ is present, and the cloudiness later clearing on continued bubbling confirms that CO₂ specifically (rather than some other gas) is what caused it, since the reversal only happens with excess CO₂.

At the point of exact neutralisation, the moles of H⁺ from the acid equal the moles of OH⁻ from the base (since HCl and NaOH react in a 1:1 ratio). Step 1: Find moles of HCl Moles = Volume (L) × Concentration (mol/L) Moles of HCl = 0.025 L × 0.1 mol/L = 0.0025 mol Step 2: At neutralisation, moles of NaOH = moles of HCl = 0.0025 mol Step 3: Find the concentration of NaOH Concentration = Moles / Volume (L) = 0.0025 mol / 0.020 L = 0.125 mol/L The concentration of the NaOH solution is 0.125 M.