Chapter 5 — Acids Bases And Salts

What are Acids and Bases?

Acids and bases are two important classes of substances we meet in daily life. An acid is a substance that tastes sour, can turn blue litmus red, and in water produces hydrogen ions (H+). A base (also called alkali when soluble in water) tastes bitter, feels soapy/slippery, can turn red litmus blue, and in water produces hydroxide ions (OH-).

Key properties

• Acids: sour taste, react with metals to give hydrogen gas, react with carbonates to produce carbon dioxide, conduct electricity in solution (electrolytes).
• Bases: bitter taste, slippery feel, react with acids to form salts and water (neutralisation), conduct electricity in solution (if ionic).

Neutralisation and salts

When an acid reacts with a base, they neutralise each other to form a salt and water. This reaction is very common and is used to treat acidity in the stomach (antacids neutralise excess acid) and in preparing many salts used at home and industry.

Indicators and pH

Indicators such as litmus paper and turmeric show whether a substance is acidic or basic by changing colour. pH is a scale used to express how acidic or basic a solution is: values less than 7 are acidic, 7 is neutral (pure water), and greater than 7 are basic.

Safety note

Many acids and bases are corrosive and can cause burns (e.g., concentrated sulfuric acid, sodium hydroxide). Handle them with care, use protective equipment, and never taste chemicals to identify them.

What are acids?
Acids are substances that taste sour, turn blue litmus paper red, and produce hydrogen ions (H⁺) when dissolved in water. They can be strong (completely ionize in water) or weak (partially ionize).

Common properties of acids

• Turn blue litmus red.
• Taste sour (e.g., lemon, vinegar).
• Conduct electricity in aqueous solution (because of ions).
• React with certain metals to produce hydrogen gas.
• React with carbonates to produce carbon dioxide, and with bases to form salt and water (neutralization).

How acids behave in water
In water an acid donates H⁺ ions. For example, hydrochloric acid (HCl) ionizes completely: HCl → H⁺ + Cl⁻. A weak acid such as acetic acid (CH₃COOH) ionizes partially: CH₃COOH ⇌ CH₃COO⁻ + H⁺.

Reactions to remember

• With metals (above hydrogen in reactivity series): Acid + Metal → Salt + Hydrogen gas (e.g., 2HCl + Zn → ZnCl₂ + H₂).
• With carbonates: Acid + Carbonate → Salt + Water + CO₂ (e.g., 2HCl + CaCO₃ → CaCl₂ + H₂O + CO₂).
• Neutralization with bases: Acid + Base → Salt + Water (e.g., HCl + NaOH → NaCl + H₂O). This reaction is usually exothermic.

pH and strength
The pH scale (0–14) measures how acidic or basic a solution is: values below 7 are acidic, 7 is neutral, and above 7 are basic. Strong acids have low pH (close to 0–2) because they release many H⁺ ions; weak acids have higher pH (closer to 3–6).

Everyday examples and uses
Many everyday substances are acids: lemon juice (citric acid), vinegar (acetic acid), stomach acid (HCl) helps digestion, and sulfuric acid in car batteries. Acids are used in medicine, food preservation, cleaning agents, and industry.

Safety note: Concentrated acids are corrosive and must be handled with care (use gloves, goggles and proper ventilation).

What is a base? A base is a substance that can accept hydrogen ions (H+) or produce hydroxide ions (OH-) in water. Bases that dissolve in water and give OH- ions are called alkalis. Examples: sodium hydroxide (NaOH), potassium hydroxide (KOH), calcium hydroxide (Ca(OH)2), and ammonia solution (NH3 in water).

Important properties of bases/alkalis

• They produce OH- ions in aqueous solution: base + H2O → OH- (or dissociation for ionic bases).
• Turn red litmus paper blue. Phenolphthalein turns pink in basic solutions.
• Taste bitter and feel slippery (but do NOT taste or touch chemicals in practice — many are corrosive).
• Conduct electricity when dissolved in water (electrolytes).
• React with acids to form a salt and water (neutralization).

Strong vs weak bases

• Strong bases (e.g., NaOH, KOH) dissociate completely in water to give many OH- ions.
• Weak bases (e.g., NH3 in water) partly ionize and give fewer OH- ions.

pH of bases

Bases have pH > 7 (up to 14). The higher the pH, the stronger the base. pH is related to hydroxide ion concentration by pOH = -log[OH-] and pH = 14 - pOH.

Neutralization (acid + base)

When an acid reacts with a base, they neutralize each other to form a salt and water. Example: HCl + NaOH → NaCl + H2O

Uses and safety

• Uses: manufacture of soaps and detergents (NaOH), drain cleaners (NaOH), fertilizers, glass and paper making (NaOH, KOH), antacids (Mg(OH)2), treatment of water (Ca(OH)2).
• Safety: strong alkalis are corrosive to skin and eyes. Use gloves and goggles and follow safety rules.

What is an indicator?

An indicator is a substance that changes colour depending on whether it is in an acidic or basic solution. Indicators help us identify whether a solution is acidic, basic (alkaline) or neutral. They can be natural (plant extracts) or synthetic (prepared chemicals).

How indicators work (simple idea)

Most indicators are weak acids or weak bases that exist in two forms: the protonated form (HIn) and the deprotonated form (In–). Each form has a different colour. In acidic solutions the HIn form dominates; in basic solutions the In– form dominates. The visible colour depends on the ratio of these two forms.

Common indicators and their colour changes

• Litmus: turns red in acids and blue in bases. (Neutral → purple)
• Phenolphthalein: colourless in acidic/neutral solutions, pink in basic solutions (≈ pH 8.2–10).
• Methyl orange: red in strong acids and yellow in neutral/weakly basic solutions (≈ pH 3.1–4.4).
• Universal indicator: a mixture giving a continuous range of colours across pH 1–14 (red → orange → yellow → green → blue → violet).
• Natural indicator example: Red cabbage juice changes through a wide range of colours with pH (useful as a classroom indicator).

Uses

• School and laboratory tests to detect acids and bases quickly.
• Monitoring pools, aquariums and drinking water (pH control).
• Soil testing for agriculture (acidic or alkaline soils).
• Testing food and household substances (vinegar, soap, antacids).

Limitations

• Some indicators work only over a limited pH range; choosing the right indicator is important.
• Indicators give qualitative or semi-quantitative information (colour change), not precise pH values unless used with a calibrated universal indicator or pH meter.

What is pH? pH is a number that tells how acidic or basic (alkaline) a solution is. The pH scale commonly runs from 0 to 14. A pH less than 7 means the solution is acidic, pH equal to 7 is neutral, and pH greater than 7 is basic.

Why pH works: pH measures the concentration of hydrogen ions (H+) in a solution. Because H+ concentrations vary over many orders of magnitude, pH uses a logarithmic scale: pH = -log₁₀[H⁺]. That means each whole pH value differs by a factor of 10 in H+ concentration (for example, pH 3 has ten times more H+ than pH 4).

Neutrality and temperature note: At 25 °C (standard lab temperature), pure water has [H⁺] = 1×10⁻⁷ M so pH = 7. Water also has hydroxide ions [OH⁻], and at 25 °C the product [H⁺][OH⁻] = 1×10⁻¹⁴ (called Kw). Using this we get the useful relation pH + pOH = 14 at 25 °C. (Kw and the pH of pure water change slightly with temperature.)

Strong vs weak acids/bases: Strong acids/bases almost fully dissociate in water and give pH values near the ends of the scale; weak acids/bases partially dissociate and give pH values closer to neutral. pH is important in chemistry, biology (e.g., blood pH), environment (acid rain), agriculture (soil pH), and daily life (food, cleaning products).

Safety: Many acids and bases are corrosive—do not taste or touch unknown chemicals. Use indicators (litmus, universal indicator paper) or pH meters to measure pH safely.

Definition: Neutralisation is a chemical reaction in which an acid and a base react to form a salt and water. In this process the acidic property (H+) and basic property (OH−) are neutralised.

Chemical idea: At the ionic level, hydrogen ions (H+) from the acid combine with hydroxide ions (OH−) from the base to form water:
H+ + OH− → H2O

General equation:
Acid + Base → Salt + Water

Examples of balanced reactions: HCl + NaOH → NaCl + H2O; Ca(OH)2 + 2HCl → CaCl2 + 2H2O.

Properties: Neutralisation is usually an exothermic reaction (releases heat). The pH of the mixture moves toward 7 (neutral) as neutralisation proceeds. An indicator (like phenolphthalein, litmus, methyl orange) is used to detect the end point: colour changes show when neutralisation is achieved.

Importance and applications: Neutralisation has many everyday uses — treating acidity in the stomach (antacids), removing acidic pollutants, adjusting soil pH in agriculture (liming), cleaning spills, and laboratory titrations to find concentrations of acids or bases.

Safety note: Always handle strong acids and bases with care, using gloves and eye protection.

What is a salt? A salt is an ionic compound formed when the hydrogen ion (H+) of an acid is replaced by a metal or ammonium (NH4+) ion. In simple terms, salts are products of reactions between acids and bases (or metals/metallic compounds) and consist of positive ions (cations) and negative ions (anions).

How salts are formed

• Neutralization: Acid + Base → Salt + Water. Example: HCl + NaOH → NaCl + H2O.
• Acid + Metal: Metal + Acid → Salt + Hydrogen gas. Example: Zn + H2SO4 → ZnSO4 + H2↑.
• Acid + Metal Carbonate/Bicarbonate: Produces salt, carbon dioxide and water. Example: CaCO3 + 2HCl → CaCl2 + CO2↑ + H2O.
• Precipitation (Double displacement): Two soluble salts react to form an insoluble salt (precipitate). Example: AgNO3 + NaCl → AgCl(s) + NaNO3.

Types of salts

• Neutral salts: Formed from strong acid + strong base (e.g., NaCl). Solution pH ≈ 7.
• Acidic salts: Formed when a polyprotic acid is partially neutralized (e.g., NaHCO3). Solution pH < 7.
• Basic salts: Formed when a weak acid is neutralized by a strong base leaving OH− (e.g., CH3COONa can give slightly basic solution). Solution pH > 7.

Properties of salts

• Usually crystalline solids with definite melting and boiling points.
• Conduct electricity when molten or dissolved in water (because ions are free to move).
• May be soluble or insoluble in water depending on the salt (e.g., NaCl soluble; AgCl insoluble).
• Sensory and chemical properties differ: many are tasteable (table salt is salty), some are poisonous, some are used as nutrients.

Uses — Table salt (NaCl) for food, fertilizers (KNO3, NH4NO3), medicines, preservatives, industrial raw materials (Na2SO4, CaCO3), water softening, and laboratory reagents.

Indicators and pH — Litmus, phenolphthalein and methyl orange help identify acids, bases and neutralization points. pH of a salt solution depends on the strengths of the parent acid and base.

Acids and bases react with many substances to give characteristic products. These reactions are important in daily life, industry and the laboratory. The main types of reactions covered in Class 7 are:

• Acid + Metal: An acid reacts with certain metals to form a salt and hydrogen gas. This happens with metals above hydrogen in the reactivity series (e.g., Zn, Mg, Fe).
• Acid + Metal Carbonate / Hydrogen Carbonate: Acids react with carbonates and hydrogen carbonates to produce a salt, carbon dioxide (CO2) and water.
• Acid + Metal Oxide: An acid reacts with a metal oxide to form a salt and water.
• Acid + Base (Alkali): An acid reacts with a base or alkali to give a salt and water—this is called neutralization.
• Reactions with Indicators: Acids and bases produce colour changes in indicators (litmus, phenolphthalein, methyl orange), which helps identify them.

Key points:

• Products are usually salt + gas (for metals and carbonates) or salt + water (for oxides and bases).
• Neutralization (acid + base) changes pH toward neutral (pH 7) and is used in antacids, treating acidic soil, and industrial waste neutralization.
• Some reactions (acid + carbonate) are used to produce CO2 in schools (vinegar + baking soda).

Overview
Laboratory experiments in the chapter 'Acids, Bases and Salts' help students observe properties of acids and bases, how they interact with indicators, metals and carbonates, and how neutralisation produces salts and water. Familiar experiments include the litmus test, using natural indicators, testing reactions with metals and carbonates, and simple neutralisation.

Common Experiments (purpose, procedure, observations)

• Litmus Test (identify acid or base)
  1. Place a few drops of the test solution on a strip of blue and red litmus paper.
  2. Observation: An acid turns blue litmus red; a base turns red litmus blue; neutral solutions do not change litmus.
• Natural Indicator (red cabbage)
  1. Boil chopped red cabbage in water and filter to get purple indicator solution.
  2. Add test solutions: it turns red/pink with acids, green/yellow with bases.
• Reaction of Acids with Metals (e.g., Zn or Mg)
  1. Add small pieces of metal to dilute hydrochloric acid in a test tube.
  2. Observation: Effervescence (bubbles) due to hydrogen gas; metal dissolves over time.
  3. Conclusion: Acid reacts with active metals to produce hydrogen and a salt.
• Reaction with Carbonates (e.g., CaCO3 / marble chips)
  1. Add marble chips to dilute HCl; collect evolved gas and test with lime water.
  2. Observation: Bubbles of CO2 are produced; lime water turns milky (formation of CaCO3).
• Neutralisation (antacid test)
  1. Take a measured volume of dilute HCl; add crushed antacid (or NaHCO3 solution) gradually.
  2. Use indicator to detect end point. Observation: Acid becomes neutral, indicator shows neutral colour; effervescence if carbonate antacid used.

What to record in the laboratory
State the aim, materials, procedure, observations (colour changes, gas evolution, temperature change if any), balanced equations, and conclusion. Note control variables (concentration, volume, temperature) and safety steps.

Safety precautions

• Wear goggles, gloves, and apron.
• Use dilute acids/bases; never smell directly; waft if needed.
• Handle metals and glassware carefully; neutralise and dispose solutions as instructed.

Common errors to avoid: Using concentrated acids, adding acid to base instead of base to acid in some experiments (when preparing solutions), not rinsing test tubes between tests, and using too much indicator.

Conclusion
Laboratory activities demonstrate practical differences between acids and bases, show chemical reactions (gas evolution, neutralisation), and teach safe experimental technique and observation recording. They link theoretical concepts (pH, indicators, reaction types) with everyday substances.

Overview: Acids and bases are chemicals that can react strongly with other substances. Knowing how to handle them safely, how they affect the environment, and how they are used in everyday life is important.

Safety rules (simple and important):

• Always wear goggles and gloves when handling concentrated acids or bases.
• Never taste or smell chemicals directly. Work in a well-ventilated area.
• When diluting acids, add acid to water slowly (not water to acid) to avoid splashing and sudden heating.
• Keep chemicals out of reach of children and label containers clearly.
• Neutralise small spills (acid with a weak base like sodium bicarbonate; base with a weak acid like vinegar) before cleanup, then rinse with plenty of water. For large spills, follow school/industrial safety protocols and notify a teacher or supervisor.

Environmental effects:

• Acid rain: gases such as sulfur dioxide (SO2) and nitrogen oxides (NOx) released from burning fossil fuels dissolve in rainwater to form acids. Acid rain can damage plants, soil, buildings and aquatic life.
• Soil and water pH changes: Acidic or highly basic wastes change the pH of soil and water, harming plants, microbes and fish. Most crops and freshwater animals prefer near-neutral pH (around 6.5–8).
• Waste disposal: Industrial acidic or basic wastes must be neutralised before release; untreated discharge pollutes rivers and harms ecosystems.

Practical applications (everyday uses):

• Cleaning: Vinegar (ethanoic acid) dissolves mineral deposits; many soaps and detergents are basic.
• Food: Citric acid (in fruits) adds sourness; baking soda (sodium bicarbonate) is used in cooking and cleaning.
• Medicine: Antacids (e.g., Mg(OH)2, CaCO3) neutralise excess stomach acid.
• Agriculture: Lime (CaO or Ca(OH)2) is added to acidic soils to raise pH and improve crop growth.
• Water treatment: Lime and other bases are used to neutralise acidic water and remove impurities; neutralisation keeps water safe for use.

Key idea — Neutralisation: Acids react with bases to form a salt and water. This is used in labs, industry and at home (antacids, cleaning, treatment of wastes).

Overview: This topic summarises the important concepts from the chapter "Acids, Bases and Salts" for quick revision: definitions, properties, reactions, indicators, pH concept, common salts and their uses, and safety precautions.

• Definitions:
  • Acid: Substance that tastes sour, turns blue litmus to red and produces H+ ions in water (e.g., HCl, CH3COOH).
  • Base (Alkali): Substance that tastes bitter, feels soapy, turns red litmus to blue and produces OH− ions in water (e.g., NaOH, Ca(OH)2). An alkali is a base soluble in water.
  • Salt: Product formed when an acid reacts with a base, metal or carbonate (e.g., NaCl, CaSO4).
• Key properties:
  • Acids: sour, corrosive, conduct electricity (when aqueous), react with metals to give H2, react with carbonates to give CO2.
  • Bases: bitter, slippery, corrosive, conduct electricity (when aqueous), react with acids to form salt + water (neutralisation).
  • Neutralisation: Acid + Base → Salt + Water.
• Common reactions (types):
  • Acid + Metal → Salt + Hydrogen (e.g., 2HCl + Zn → ZnCl2 + H2).
  • Acid + Carbonate → Salt + Carbon dioxide + Water (e.g., 2HCl + CaCO3 → CaCl2 + CO2 + H2O).
  • Acid + Base (Neutralisation) → Salt + Water (e.g., HCl + NaOH → NaCl + H2O).
• Indicators and colour changes:
  • Blue litmus → Red in acid; Red litmus → Blue in base.
  • Methyl orange: Red in acid, yellow in base. Phenolphthalein: Colourless in acid/neutral, pink in basic solutions.
• pH concept (summary):
  • pH is a measure of how acidic or basic a solution is. pH scale ranges from 0 to 14: acids <7, neutral = 7, bases >7.
  • Strong acids have low pH (close to 0); strong bases have high pH (close to 14).
• Important salts and uses:
  • NaCl (table salt) – food seasoning and preservative.
  • NaHCO3 (baking soda) – antacid, baking, cleaning.
  • CaCO3 (chalk, limestone) – building material, antacid (in some forms).
  • CaSO4·2H2O (gypsum) – plaster, cement.
• Safety and precautions:
  • Always handle acids and bases with care: use gloves and goggles in the lab, do not taste, dilute acids by adding acid to water (not water to acid).

Revision tips: Memorise indicator colour changes, common reactions (acid+metal, acid+carbonate, neutralisation), pH ranges and a few typical examples of acids, bases and salts used in daily life.