Acids Bases and Salts — Karnataka (SSLC) Class 10 Science Solutions (Free)

Acids: substances that release hydrogen ions (H⁺) in water. Characterised by a sour taste, turning blue litmus red, and a pH below 7. Example: hydrochloric acid (HCl). Bases: substances that release hydroxide ions (OH⁻) in water. Characterised by a bitter taste, a slippery feel, turning red litmus blue, and a pH above 7. Example: sodium hydroxide (NaOH). Salts: ionic compounds formed when an acid neutralises a base — the general reaction is Acid + Base → Salt + Water. Most common salts are close to neutral (pH ≈ 7), though this depends on the strength of the acid and base that formed them, as covered later in this chapter. Example: sodium chloride (NaCl, common table salt).

pH is a numerical scale, running from 0 to 14, that measures how acidic or basic a solution is, based on its concentration of hydrogen ions (H⁺). It is mathematically defined as pH = −log₁₀[H⁺]. - Acidic: pH below 7. Common examples — vinegar (pH ≈ 3), lemon juice (pH ≈ 2). - Neutral: pH exactly 7. Pure water and a dissolved solution of sodium chloride (NaCl) are both neutral. - Basic (alkaline): pH above 7. Common examples — baking soda solution (pH ≈ 8.3), bleach (pH ≈ 13). Each whole step on the pH scale represents a TEN-FOLD change in H⁺ concentration, not a simple linear difference — which is why lemon juice (pH 2) is not just "slightly more acidic" than vinegar (pH 3), but roughly ten times higher in H⁺ concentration.

HCl + NaOH → NaCl + H₂O The reaction produces a salt (sodium chloride, NaCl) and water. What actually happens: the H⁺ ions released by the acid combine directly with the OH⁻ ions released by the base to form neutral water molecules (H⁺ + OH⁻ → H₂O). The Na⁺ and Cl⁻ ions do not react with each other in this process — they remain dissolved in solution and simply form NaCl once the water evaporates. Because both HCl and NaOH are strong (they fully ionise in water), the resulting salt, NaCl, is neutral and fully soluble. This reaction is described as complete neutralisation when the acid and base are combined in exactly the stoichiometric (chemically balanced) proportions needed to fully consume both.

Indicators are substances that change colour depending on whether the solution they're added to is acidic or basic, allowing quick identification without lab equipment. Three natural/common indicators: 1. Litmus: turns RED in acidic solutions, BLUE in basic solutions. 2. Turmeric: stays YELLOW in acidic and neutral solutions, but turns REDDISH-BROWN in basic (alkaline) solutions — this is exactly why a turmeric stain on cloth often turns a reddish tinge when treated with soap, since ordinary soap is mildly basic. 3. Phenolphthalein: remains COLOURLESS in acidic solutions, but turns PINK in basic solutions — this sharp, easy-to-see colour change makes it a common choice in laboratory titrations. Together, these indicators let you not only detect whether a substance is acidic or basic, but often also give a rough sense of HOW strongly, since the intensity of the colour change can vary with pH.

Strong acids ionise (break apart into ions) COMPLETELY when dissolved in water — essentially every acid molecule releases its H⁺ ion. Examples: hydrochloric acid (HCl), sulphuric acid (H₂SO₄), nitric acid (HNO₃). Weak acids ionise only PARTIALLY — at any given moment, only a fraction of the acid molecules have released their H⁺ ion, while the rest remain un-ionised, existing in a dynamic equilibrium. Examples: acetic acid (in vinegar), citric acid (in lemons and other citrus fruits). Acetic acid's partial ionisation is represented as a reversible reaction: CH₃COOH ⇌ CH₃COO⁻ + H⁺ Practical consequence: for solutions of the SAME concentration, a strong acid produces a much higher concentration of free H⁺ ions than a weak acid, giving the strong acid both a lower pH and significantly better electrical conductivity (since conductivity depends on the concentration of free ions available to carry current).

Preparation of slaked lime: water is added to quicklime (calcium oxide, CaO): CaO + H₂O → Ca(OH)₂ This reaction is strongly exothermic — it releases a substantial amount of heat, enough that the mixture can visibly steam or even boil if the reaction proceeds quickly. Reaction with carbon dioxide (the limewater test): Ca(OH)₂ + CO₂ → CaCO₃ + H₂O Calcium carbonate (CaCO₃) is insoluble in water, so it forms as a white precipitate, turning clear limewater milky/turbid — this colour change is the standard laboratory test used to confirm the presence of CO₂ gas. (If excess CO₂ continues to be passed, the precipitate can further react to form soluble calcium bicarbonate, Ca(HCO₃)₂, and the solution clears again — a useful additional check that the gas really is CO₂.) The calcium carbonate produced from this basic reaction is also a key raw material used in manufacturing cement.

A salt is an ionic compound formed by the neutralisation of an acid and a base, made up of a cation (typically the metal or ammonium ion from the base) and an anion (typically the acid's remaining part after losing its H⁺). Salts are classified into three types, based on the relative strength of the acid and base that produced them: 1. Normal (neutral) salt — formed from a STRONG acid and a STRONG base. Example: NaCl, from HCl (strong acid) + NaOH (strong base). Its solution is neutral, pH ≈ 7. 2. Acidic salt — formed from a STRONG acid and a WEAK base. Example: NH₄Cl (ammonium chloride), from HCl (strong acid) + NH₃ (weak base). Its solution is mildly acidic, pH < 7, because the ammonium ion partially reacts with water to release extra H⁺. 3. Basic salt — formed from a WEAK acid and a STRONG base. Example: NaHCO₃ (sodium bicarbonate, baking soda), from H₂CO₃ (weak acid, carbonic acid) + NaOH (strong base). Its solution is mildly basic, pH > 7, because the bicarbonate ion partially reacts with water to release extra OH⁻. All salts, regardless of this classification, conduct electricity when dissolved in water or melted, since in both states their ions become free to move and carry electric current.