CHAPTER-1 CHEMICAL REACTIONS AND EQUATIONS

Topic-1 Chemical Reaction and Equations

Concepts Covered: Chemical reaction and examples • Skeletal and balanced chemical equation • Steps to balance a chemical equation.

A chemical reaction

• A chemical reaction is a process in which the original substance(s) loses its nature and identity and forms new substance(s) with different properties.
• Breaking of the chemical bonds and formation of new chemical bonds is responsible for the occurrence of a chemical reaction.
• The substances which take part in a chemical reaction are called Reactants.
• The substances which are formed in a chemical reaction are called Products.

Examples of chemical reaction:

1. Digestion of food
2. Respiration
3. Rusting of iron
4. Burning of magnesium ribbon
5. Formation of curd

A chemical reaction can be identified by either of the following observations:

Chemical equations

• A chemical equation is the symbolic representation of a chemical reaction in the form of symbols and formulae.
• It is a way to represent the chemical reaction in a concise and informative way.
• Example (word equation): Magnesium + Oxygen → Magnesium oxide
• Chemical equation: \[ 2\mathrm{Mg + O_2 \rightarrow 2MgO} \]

Writing a chemical equation

1. Symbols and formulae of reactants are written on the left-hand side with a plus (+) sign between them.
2. Symbols and formulae of products are written on the right-hand side with a plus (+) sign between them.
3. An arrow sign (→) is put between the reactants and the products.
4. Physical states of the reactants and products are also mentioned.

Skeletal chemical equation: A chemical equation which simply represents the symbols and formulas of reactants and products. Example: \[ \mathrm{Mg + O_2 \rightarrow MgO} \]

Balanced Equation: The equation in which atoms of various elements on both sides are equal in accordance with the law of conservation of mass.

Examples:

• \[ \mathrm{CO(g) + 2H_2(g) \xrightarrow{340\,atm} CH_3OH(l)} \]
• \[ \mathrm{6CO_2(g) + 6H_2O(l) \xrightarrow{sunlight} C_6H_{12}O_6(aq) + 6O_2(g)} \] (Glucose)

Example – Balancing a chemical equation (Hit and Trial method):

Step 1: Write the chemical equation and draw boxes around each formula.
\[ \mathrm{Fe + H_2O \rightarrow Fe_3O_4 + H_2} \]

Step 2: Count the number of atoms of each element on both sides.

Step 3: Equalize the element with maximum atoms (oxygen): \[ \mathrm{Fe + 4H_2O \rightarrow Fe_3O_4 + H_2} \]

Step 4: Equalize all atoms: \[ \mathrm{3Fe + 4H_2O \rightarrow Fe_3O_4 + 4H_2} \]

Step 5: Write physical states: \[ \mathrm{3Fe(s) + 4H_2O(g) \rightarrow Fe_3O_4(s) + 4H_2(g)} \]

Step 6: Write necessary conditions of temperature, pressure or catalyst if required.

Topic-2 Types of Chemical Reactions

Concepts Covered: Combination, Decomposition, Displacement, Double displacement, Redox, Oxidation & Reduction, Exothermic & Endothermic reactions.

I. Combination Reaction: Two or more reactants combine to form a single product.

• Burning of coal: \[ \mathrm{C(s) + O_2(g) \rightarrow CO_2(g)} \]
• Formation of water: \[ \mathrm{2H_2(g) + O_2(g) \rightarrow 2H_2O(l)} \]
• \[ \mathrm{CaO(s) + H_2O(l) \rightarrow Ca(OH)_2(aq) + Heat} \] (Quick lime → Slaked lime)

Exothermic Reactions: Heat is released along with formation of products.

• Burning of natural gas: \[ \mathrm{CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(g) + Heat} \]
• Respiration: \[ \mathrm{C_6H_{12}O_6(aq) + 6O_2(g) \rightarrow 6CO_2(aq) + 6H_2O(l) + energy} \]

II. Decomposition Reaction: A compound splits into two or more simpler substances. \[ A \rightarrow B + C \]

(a) Thermal decomposition (by heating)

• \[ \mathrm{2FeSO_4(s) \xrightarrow{Heat} Fe_2O_3(s) + SO_2(g) + SO_3(g)} \] (Green → Red-brown)
• \[ \mathrm{CaCO_3(s) \xrightarrow{Heat} CaO(s) + CO_2(g)} \]

(b) Electrolytic Decomposition (by electricity)
\[ \mathrm{2H_2O(l) \xrightarrow{Electric\,current} 2H_2(g) + O_2(g)} \]

(c) Photolytic Decomposition (by sunlight)

• \[ \mathrm{2AgCl(s) \xrightarrow{Sunlight} 2Ag(s) + Cl_2(g)} \]
• \[ \mathrm{2AgBr(s) \xrightarrow{Sunlight} 2Ag(s) + Br_2(g)} \]

Endothermic Reaction: Reactions which require energy in the form of heat, light or electricity to break reactants.

III. Displacement Reaction: More reactive element displaces less reactive element from its salt solution.

• \[ \mathrm{Fe(s) + CuSO_4(aq) \rightarrow FeSO_4(aq) + Cu(s)} \] (Iron nail becomes brownish; blue colour changes to dirty green)
• \[ \mathrm{Zn(s) + CuSO_4(aq) \rightarrow ZnSO_4(aq) + Cu(s)} \]

IV. Double Displacement Reaction: Mutual exchange of ions between two compounds.
\[ \mathrm{Na_2SO_4(aq) + BaCl_2(aq) \rightarrow BaSO_4(s) + 2NaCl(aq)} \] (Also called precipitation reaction)

V. Oxidation and Reduction

• Oxidation: Loss of electrons / gain of oxygen
• Reduction: Gain of electrons / loss of oxygen

\[ \mathrm{2Cu + O_2 \xrightarrow{Heat} 2CuO} \] (Oxidation)
\[ \mathrm{CuO + H_2 \rightarrow Cu + H_2O} \] (CuO reduced, H₂ oxidised → Redox reaction)

Important Redox equation:
\[ \mathrm{MnO_2(s) + 4HCl(conc.) \rightarrow MnCl_2(aq) + Cl_2(g) + 2H_2O} \] (MnO₂ is reduced)

Some usually asked equations for balancing:

• \[ \mathrm{2CO(g) + O_2(g) \rightarrow 2CO_2(g)} \]
• \[ \mathrm{ZnCO_3 \rightarrow ZnO + CO_2} \]
• \[ \mathrm{2FeSO_4(s) \rightarrow Fe_2O_3(s) + SO_2(g) + SO_3(g)} \]
• \[ \mathrm{Pb(NO_3)_2 + 2KI \rightarrow 2KNO_3 + PbI_2} \]
• \[ \mathrm{CaO(s) + H_2O \rightarrow Ca(OH)_2 + Heat} \]
• \[ \mathrm{NaCl + AgNO_3 \rightarrow AgCl + NaNO_3} \]
• \[ \mathrm{Ca + 2HNO_3 \rightarrow Ca(NO_3)_2 + H_2\uparrow} \]
• \[ \mathrm{Mg + 2HNO_3 \rightarrow Mg(NO_3)_2 + H_2\uparrow} \]
• \[ \mathrm{2Al + 3H_2SO_4 \rightarrow Al_2(SO_4)_3 + 3H_2\uparrow} \]
• \[ \mathrm{Na_2CO_3 + 2HCl \rightarrow 2NaCl + H_2O + CO_2} \]
• \[ \mathrm{Ca(OH)_2 + CO_2 \rightarrow CaCO_3 + H_2O} \]
• \[ \mathrm{Zn + H_2SO_4 \rightarrow ZnSO_4 + H_2\uparrow} \]
• \[ \mathrm{Zn + 2HCl \rightarrow ZnCl_2 + H_2\uparrow} \]
• \[ \mathrm{4Zn + 10HNO_3 \rightarrow 4Zn(NO_3)_2 + 5H_2O + N_2O} \]
• \[ \mathrm{Zn + 2NaOH \rightarrow Na_2ZnO_2 + H_2\uparrow} \]

Effects of oxidation reactions in everyday life

1. Corrosion: Deterioration of metals by action of air, moisture, chemicals etc. It is a redox reaction.

• Rusting of iron: Iron objects left in moist open air get coated with reddish brown powder (rust = hydrated iron(III) oxide \[ \mathrm{Fe_2O_3\cdot xH_2O} \]).
• Green coating on copper and black coating on silver are other examples.

Effects: Damage to ships, car bodies, bridges, railings; wastage of metals and money.

Prevention of Rusting: Painting, oiling, greasing, galvanising, chromium coating.

2. Rancidity: Slow oxidation of oil and fat in food materials resulting in foul odour and taste. Prevention: Store in refrigerator, air-tight containers, add anti-oxidants, flush with nitrogen.

CHAPTER-2 ACIDS, BASES AND SALTS

Topic-1 Acids and Bases

Concepts Covered: Definition of acids and bases • Properties • Indicators • pH scale and importance of pH.

Acids

• Substances that furnish H⁺ ions in aqueous solution. Sour in taste. Turn blue litmus red.
• Examples: H₂SO₄, CH₃COOH, HNO₃
• Dilute solution – low concentration; Concentrated solution – high concentration.
• Strong acids dissociate completely (H₂SO₄, HCl); Weak acids dissociate partially (citric acid, acetic acid).

Naturally occurring acids:

Bases

• Bitter in taste, soapy to touch, turn red litmus blue, give OH⁻ ions in aqueous solution.
• Examples: NaOH, KOH
• Strong bases ionise completely; Weak bases ionise partially (Mg(OH)₂, Cu(OH)₂).

Types of Indicators:

Chemical Properties of Acids and Bases

1. Reaction with Metals

Test for H₂: Burning candle near the gas bursts with a pop sound.

2. Reaction with Metal Carbonates / Hydrogen Carbonates

• With acids: Acid + Metal carbonate/hydrogen carbonate → Salt + CO₂ + H₂O
  \[ 2\mathrm{HCl + Na_2CO_3 \rightarrow 2NaCl + CO_2 + H_2O} \]
  \[ \mathrm{HCl + NaHCO_3 \rightarrow NaCl + CO_2 + H_2O} \]
• With bases: No reaction.

Test for CO₂: Lime water turns milky. \[ \mathrm{Ca(OH)_2 + CO_2 \rightarrow CaCO_3 + H_2O} \]. Excess CO₂ dissolves the precipitate: \[ \mathrm{CaCO_3 + CO_2 + H_2O \rightarrow Ca(HCO_3)_2} \]

3. Neutralisation Reaction
\[ \mathrm{Acid + Base \rightarrow Salt + Water} \]
\[ \mathrm{H^+(aq) + OH^-(aq) \rightarrow H_2O(l)} \]
Example: \[ \mathrm{NaOH(aq) + HCl(aq) \rightarrow NaCl(aq) + H_2O(l)} \]

Dilution: Always add acid to water (never water to acid). It is an exothermic process.

Common property of all acids: They contain hydrogen and produce H⁺ (or H₃O⁺) ions in water.

\[ \mathrm{HCl(aq) \rightarrow H^+(aq) + Cl^-(aq)} \] \[ \mathrm{HCl + H_2O \rightarrow H_3O^+ + Cl^-} \]

Reaction of acids with metal oxides: Metal oxide + Acid → Salt + Water
\[ \mathrm{CuO + 2HCl \rightarrow CuCl_2 + H_2O} \] (black CuO → blue-green solution)

Strength of acids and bases

• Strong acids: Complete dissociation (HCl, H₂SO₄, HNO₃)
• Weak acids: Partial dissociation (CH₃COOH)
• Strong bases: Complete ionisation (NaOH, KOH)
• Weak bases: Partial dissociation (NH₄OH, Ca(OH)₂)

pH Scale
\[ \mathrm{pH = -\log[H^+]} \] or \[ \mathrm{pH = -\log[H_3O^+]} \]
pH = 7 → Neutral pH < 7 → Acidic pH > 7 → Basic
On diluting an acid, pH increases; on diluting a base, pH decreases.

Topic-2 Salts, their Properties and Uses

Salts: Ionic compounds formed by neutralisation of an acid and a base. Electrically neutral.

Types of Salts:

• Neutral salts (strong acid + strong base) → pH ≈ 7 (e.g. NaCl)
• Acidic salts (strong acid + weak base) → pH < 7 (e.g. NH₄Cl)
• Basic salts (strong base + weak acid) → pH > 7 (e.g. Na₂CO₃)

Common Salt (NaCl)
Preparation: \[ \mathrm{NaOH + HCl \rightarrow NaCl + H_2O} \]
Electrolysis (Chlor-alkali process): \[ \mathrm{2NaCl(aq) + 2H_2O(l) \rightarrow 2NaOH(aq) + Cl_2(g) + H_2(g)} \]
Uses: Food, preservative, manufacture of Na and Cl₂.

Sodium hydroxide (NaOH): Used in paper, soap, detergents, degreasing metals. Cl₂ used in water treatment, PVC, pesticides. H₂ used as rocket fuel and for hydrogenation.

Bleaching Powder (CaOCl₂)
\[ \mathrm{Cl_2 + Ca(OH)_2 \rightarrow CaOCl_2 + H_2O} \]
Uses: Bleaching cotton/linen, wood pulp, oxidising agent, disinfecting drinking water.

Baking Soda (NaHCO₃)
\[ \mathrm{NaCl + H_2O + CO_2 + NH_3 \rightarrow NH_4Cl + NaHCO_3} \]
On heating: \[ \mathrm{2NaHCO_3 \xrightarrow{\Delta} Na_2CO_3 + H_2O + CO_2} \]
Uses: Baking powder, antacid, soda-acid fire extinguishers.

Washing Soda (Na₂CO₃·10H₂O)
Recrystallisation of sodium carbonate.
Uses: Glass, soap, paper industry, manufacture of borax, cleaning agent, removing permanent hardness of water.

Plaster of Paris (CaSO₄·½H₂O)
Heating gypsum at 373 K: \[ \mathrm{CaSO_4·2H_2O \rightarrow CaSO_4·½H_2O + 1½H_2O} \]
On mixing with water it sets: \[ \mathrm{CaSO_4·½H_2O + 1½H_2O \rightarrow CaSO_4·2H_2O} \]
Uses: Supporting fractured bones, making toys, decorative materials, statues.

Water of crystallisation: Fixed number of water molecules present in one formula unit of a salt (hydrated salts). Examples: CuSO₄·5H₂O (blue vitriol), Na₂CO₃·10H₂O, CaSO₄·2H₂O (gypsum).

CHAPTER-3 METALS AND NON-METALS

Topic-1 Properties of Metals and Non-Metals

Physical Properties:

Chemical Properties of Metals

(A) Reaction with Air / Oxygen: Metal + O₂ → Metal oxide

• \[ \mathrm{2Cu + O_2 \rightarrow 2CuO} \] (black)
• \[ \mathrm{4Al + 3O_2 \rightarrow 2Al_2O_3} \]
• \[ \mathrm{2Mg + O_2 \rightarrow 2MgO} \]

Na and K react so vigorously that they catch fire if kept in open → kept immersed in kerosene. Surfaces of Mg, Al, Zn, Pb are covered with a thin oxide layer. Au and Ag do not react with oxygen.

Amphoteric Oxides: React with both acids and bases.
\[ \mathrm{Al_2O_3 + 6HCl \rightarrow 2AlCl_3 + 3H_2O} \]
\[ \mathrm{Al_2O_3 + 2NaOH \rightarrow 2NaAlO_2 + H_2O} \]

(B) Reaction with Water: Metal + Water → Metal oxide / hydroxide + H₂

• Na and K react vigorously with cold water.
• Mg reacts with hot water: \[ \mathrm{Mg + 2H_2O \rightarrow Mg(OH)_2 + H_2} \]
• Al and Zn react with steam.
• Ag, Au, Cu, Pb do not react with water.

(C) Reaction with Acids: Metal + Dilute acid → Salt + H₂
\[ \mathrm{Mg + H_2SO_4 \rightarrow MgSO_4 + H_2} \] (Cu, Hg, Ag do not react with dilute acids)

(D) Reaction with solutions of other metal salts: More reactive metal displaces less reactive metal.
\[ \mathrm{Fe + CuSO_4 \rightarrow FeSO_4 + Cu} \]

Reactivity / Activity Series (decreasing order): K > Na > Ca > Mg > Al > Zn > Fe > Sn > Pb > H > Cu > Hg > Ag > Au > Pt

Reaction of Non-Metals:

• With oxygen → acidic oxides (e.g. \[ \mathrm{C + O_2 \rightarrow CO_2} \])
• With water → generally no reaction
• With dilute acids → no reaction
• With chlorine / hydrogen → form chlorides / hydrides

Aqua Regia: Mixture of conc. HCl and conc. HNO₃ in the ratio 3:1. It can dissolve gold and platinum.

Reaction between metal and non-metal (Ionic bond formation):
Metals lose electrons to form cations; non-metals gain electrons to form anions; opposite charges attract.

Example – Formation of NaCl:
\[ \mathrm{Na \rightarrow Na^+ + e^-} \] (2,8,1 → 2,8)
\[ \mathrm{Cl + e^- \rightarrow Cl^-} \] (2,8,7 → 2,8,8)
\[ \mathrm{Na^+ + Cl^- \rightarrow NaCl} \]

Topic-2 Ionic compounds, Metallurgy and Corrosion

Ionic Compounds: Formed by transfer of electrons from metal to non-metal.

Properties:

1. Solid and hard, generally brittle.
2. High melting and boiling points.
3. Generally soluble in water, insoluble in kerosene, petrol etc.
4. Conduct electricity in molten and aqueous state (ions free to move) but not in solid state.

Occurrence of Metals

• Minerals: Elements or compounds occurring naturally in the earth’s crust.
• Ores: Minerals containing a high percentage of a particular metal that can be profitably extracted.

Metals of low reactivity (Au, Ag, Pt, Cu) are often found free. Highly reactive metals (K, Na, Ca, Mg, Al) are never found free. Moderately reactive metals (Zn, Fe, Pb) occur as oxides, sulphides or carbonates.

Thermite reaction: \[ \mathrm{Fe_2O_3(s) + 2Al(s) \rightarrow 2Fe(l) + Al_2O_3(s) + Heat} \] (used to join railway tracks).

Extraction of metals (Metallurgy): Enrichment / concentration of ore → Extraction of metal from concentrated ore → Refining of metal.

Corrosion: Deterioration of a metal by chemical reaction with its environment. Examples: black coating on silver (Ag₂S), green coat on copper, rust on iron (\[ \mathrm{Fe_2O_3·xH_2O} \]).

Prevention: Painting, oiling, greasing, galvanising, making alloys.

Alloys: Homogeneous mixtures of metals (or metal + non-metal). Examples: Stainless steel (Fe + Ni + Cr), Brass (Cu + Zn), Bronze (Cu + Sn), Solder (Pb + Sn). Amalgam – alloy containing mercury.

CHAPTER-4 CARBON COMPOUNDS

Topic-1 Carbon and its Properties, Homologous Series and IUPAC Names

Carbon is a non-metal (symbol C). Atomic number 6, electronic configuration 2, 4. It is tetravalent and forms covalent bonds by sharing electrons.

Properties of covalent compounds: Generally liquids or gases, insoluble in water, low melting and boiling points, do not conduct electricity.

Versatile nature of carbon is due to:

• Catenation: Ability to form long chains, branched chains and rings.
• Tetravalency: Four valence electrons available for bonding.

Hydrocarbons: Compounds containing only carbon and hydrogen.

• Saturated (Alkanes) – single bonds only, general formula CₙH₂ₙ₊₂
• Unsaturated – Alkenes (CₙH₂ₙ, double bond), Alkynes (CₙH₂ₙ₋₂, triple bond)
• Aromatic – contain benzene ring

IUPAC naming: Word root (meth, eth, prop, but, pent…) + suffix (-ane, -ene, -yne).

Important functional groups:

Isomerism: Compounds having the same molecular formula but different structural formulae (e.g. butane C₄H₁₀ has two isomers).

Homologous Series: A series of compounds in which successive members differ by –CH₂ or 14 a.m.u. Physical properties vary gradually; chemical properties remain similar.

Chemical Properties of carbon compounds

1. Combustion: \[ \mathrm{CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O + Heat + Light} \] (Saturated – blue non-sooty flame; Unsaturated – yellow sooty flame)
2. Oxidation: Alcohols → carboxylic acids with alkaline KMnO₄ or acidic K₂Cr₂O₇
   \[ \mathrm{CH_3CH_2OH \xrightarrow{alk.\ KMnO_4\ or\ acid\ K_2Cr_2O_7} CH_3COOH} \]
3. Addition Reaction: Unsaturated hydrocarbons add H₂ in presence of Ni or Pd (hydrogenation of oils).
4. Substitution Reaction: Saturated hydrocarbons in presence of sunlight: \[ \mathrm{CH_4 + Cl_2 \xrightarrow{sunlight} CH_3Cl + HCl} \]

Topic-2 Ethanol, Ethanoic acid, Soaps and Detergents

Ethanol (C₂H₅OH)

• Reaction with sodium: \[ \mathrm{2CH_3CH_2OH + 2Na \rightarrow 2CH_3CH_2ONa + H_2} \]
• Esterification: \[ \mathrm{CH_3COOH + C_2H_5OH \xrightarrow{conc.\ H_2SO_4} CH_3COOC_2H_5 + H_2O} \]
• Uses: Soap, cosmetics, beverages, medicines, laboratory reagent.

Ethanoic acid (CH₃COOH)

• Vinegar = 5–8 % solution; Glacial acetic acid = pure.
• Reacts with Na₂CO₃ / NaHCO₃ giving CO₂ (brisk effervescence).
• Neutralisation: \[ \mathrm{CH_3COOH + NaOH \rightarrow CH_3COONa + H_2O} \]

Soap and Detergents

• Soap is sodium or potassium salt of long-chain carboxylic acid (e.g. C₁₇H₃₅COONa).
• Saponification = alkaline hydrolysis of ester.
• Soaps work only in soft water; detergents work in both hard and soft water.
• Soap molecule has hydrophilic (ionic) head and hydrophobic (hydrocarbon) tail → forms micelles that clean dirt/oil.
• Hard water forms scum with soap (Ca/Mg salts); detergents do not form scum.

CHAPTER-5 LIFE PROCESSES

Topic-1 Nutrition

All living organisms perform life processes (growth, excretion, respiration, circulation, reproduction) for survival. Energy is obtained from carbon-based food sources through nutrition.

Modes of nutrition:

• Autotrophic: Organism prepares its own food by photosynthesis (green plants, some bacteria).
• Heterotrophic: Takes food from another organism (animals, fungi). Types: Holozoic, Saprophytic, Parasitic.

Main events of Photosynthesis:

1. Absorption of light energy by chlorophyll.
2. Conversion of light energy to chemical energy and splitting of water into hydrogen and oxygen.
3. Reduction of CO₂ to form carbohydrates.

Site of photosynthesis is the chloroplast. Opening and closing of stomata is controlled by the turgidity of guard cells.

Nutrition in human beings: Alimentary canal + associated glands (salivary glands, gastric glands, liver, pancreas).

Small intestine is the site of complete digestion of carbohydrates, proteins and fats. Bile (secreted by liver, stored in gall bladder) emulsifies fats. Villi help in absorption of digested food into blood.

Topic-2 Respiration

Respiration involves (i) Breathing (gaseous exchange) and (ii) Breakdown of food to release energy.

Breakdown of glucose:

• Glycolysis (cytoplasm) → Pyruvic acid + Energy
• In presence of oxygen (mitochondria) → CO₂ + H₂O + large amount of energy (Aerobic)
• In absence of oxygen → Lactic acid (in muscles) or Ethanol + CO₂ (in yeast) + small amount of energy (Anaerobic)

Human Respiratory system: Nostrils → Nasal cavity → Pharynx → Larynx → Trachea → Bronchi → Bronchioles → Alveoli. Exchange of gases occurs at the alveoli. Haemoglobin carries oxygen from lungs to tissues.

Breathing rate in aquatic organisms is much faster than in terrestrial organisms because the amount of dissolved oxygen in water is low.

Topic-3 Circulation and Transportation

Human Circulatory System: Consists of blood, lymph, blood vessels (arteries, veins, capillaries) and heart.

• Human heart is four-chambered (right atrium, right ventricle, left atrium, left ventricle).
• Double circulation: Pulmonary circulation (heart ↔ lungs) and Systemic circulation (heart ↔ rest of body).
• Blood pressure is the force that blood exerts on the walls of blood vessels (normal < 120/80 mm Hg).

Transportation in plants:

Transpiration is the loss of water as vapour from aerial parts. Translocation is the transport of food.

Topic-4 Excretion

Excretion is the removal of harmful metabolic wastes (urea, uric acid) from the body.

Excretory system of humans: Pair of kidneys, pair of ureters, urinary bladder and urethra. Nephron is the basic filtration unit of the kidney.

Urine formation involves three steps:

1. Glomerular filtration
2. Tubular reabsorption
3. Secretion

About 180 litres of filtrate is formed daily but only about 2 litres is excreted as urine.

Haemodialysis: Process of purifying blood by an artificial kidney in case of kidney failure.

In plants, excretion of O₂, CO₂ and water takes place through stomata by transpiration.

CHAPTER-6 CONTROL AND CO-ORDINATION

Topic-1 Control and Co-ordination in Plants

Plants respond to stimuli by movements. Two types:

(A) Growth-dependent (Tropic movements)

• Phototropism – towards light (shoots positive, roots negative)
• Geotropism – towards gravity (roots positive, shoots negative)
• Chemotropism – towards chemicals (pollen tube towards ovule)
• Hydrotropism – towards water

(B) Growth-independent movements: Immediate response (e.g. drooping of leaves of “Touch-me-not” plant – thigmotropism).

Plant Hormones (Phytohormones):

Topic-2 Control and Co-ordination in Animals

Control and co-ordination is brought about by the Nervous system and the Endocrine system.

Nervous system: System of conducting tissues that receives stimulus and transmits it to other parts of the body. Structural and functional unit is the neuron (cyton/cell body, dendrites, axon, myelin sheath, synapse).

Functioning of neuron: Dendrites → Cell body → Axon → Nerve endings → Synapse → Dendrite of next neuron.

• Sensory neurons transmit impulses towards the central nervous system.
• Motor neurons transmit information from the brain to effector organs (muscles and glands).

Reflex action: Quick, sudden and immediate response of the body to a stimulus (e.g. knee jerk, withdrawal of hand on touching a hot object). The pathway is called the reflex arc.

Human brain: Main coordinating centre. Protected by the cranium and three membranes (meninges). CSF acts as a cushion.

• Forebrain (cerebrum): Main thinking part, controls voluntary actions, stores memory, receives sensory impulses, centre of hunger.
• Midbrain: Controls involuntary actions.
• Hindbrain:
  • Cerebellum – posture, balance, precision of voluntary actions
  • Medulla – involuntary actions (blood pressure, salivation, vomiting)
  • Pons – regulation of respiration

There are 12 pairs of cranial nerves and 31 pairs of spinal nerves.

Endocrine system: Ductless glands secrete hormones (chemical messengers) that act on target organs.

Feedback mechanisms ensure that hormones are secreted in precise quantities at the right time.

CHAPTER-7 REPRODUCTION

Topic-1 Asexual Reproduction and Vegetative Propagation

Reproduction is the process by which living organisms produce new individuals similar to themselves. It ensures continuity of life. DNA is the hereditary material; its replication causes variation.

Asexual Reproduction: Single parent, no gametes formed, offspring identical to parent. Modes:

• Fission: Binary (Amoeba) and Multiple (Plasmodium)
• Fragmentation: Body breaks into fragments that grow into new individuals (Spirogyra)
• Regeneration: Ability to give rise to new individuals from body parts (Planaria, Hydra)
• Budding: Small outgrowth (bud) detaches and becomes a new organism (Hydra)
• Vegetative Propagation: New plants from vegetative parts (roots, stems, leaves). Natural methods (potato, ginger, Bryophyllum) and artificial methods (grafting, cutting, layering, tissue culture)
• Spore Formation: Spores covered by thick walls germinate under favourable conditions (Rhizopus)

Topic-2 Sexual Reproduction in Plants

Flower is the reproductive part. Stamens (male) produce pollen; Carpels (female) contain ovules. Flowers may be unisexual or bisexual.

Pollination: Transfer of pollen from anther to stigma (self or cross). Agents: wind, water, animals. Cross-pollination introduces variations.

Fertilisation: Fusion of male and female gametes to form zygote. Flowering plants show double fertilisation (syngamy + triple fusion).

Post-fertilisation changes: Ovule → seed; Ovary → fruit; other floral parts may fall off.

Topic-3 Reproduction in Human Beings

Puberty: Period of sexual maturation when production of gametes begins. Common changes + sex-specific changes (breast development & menstruation in girls; facial hair & voice change in boys).

Male Reproductive System: Testes (in scrotum – lower temperature needed for sperm production) produce sperms and testosterone. Sperms travel through epididymis → vas deferens → urethra. Accessory glands (seminal vesicles, prostate) add fluid to form semen.

Female Reproductive System: Ovaries produce eggs and hormones. Oviduct (fallopian tube) is the site of fertilisation. Uterus is the site of implantation and development of the foetus. Placenta provides nutrition and exchange of materials. Menstruation is the breakdown and removal of the uterine lining if fertilisation does not occur.

Reproductive Health & Contraception

• STDs: Gonorrhoea, syphilis, genital warts, HIV-AIDS
• Methods of contraception: Barrier (condoms, diaphragm, cervical cap), Chemical (oral pills), IUCD (Copper-T), Surgical (vasectomy in males, tubectomy in females)

CHAPTER-8 HEREDITY AND EVOLUTION

Variations arise mainly during sexual reproduction (slight inaccuracies in DNA copying + crossing over). Beneficial variations help the species survive and are selected by nature, leading to evolution.

Mendel’s work on Inheritance: Gregor Johann Mendel (Father of Genetics) performed experiments on garden pea (Pisum sativum) and studied seven pairs of contrasting characters.

Monohybrid cross: Phenotypic ratio in F₂ = 3:1; Genotypic ratio = 1:2:1

Dihybrid cross: Phenotypic ratio in F₂ = 9:3:3:1

Mendel’s Laws:

• Law of Dominance
• Law of Segregation
• Law of Independent Assortment

Alleles are alternate forms of a gene. A dominant allele expresses itself in the presence or absence of the recessive allele; a recessive allele expresses only in the homozygous condition.

Sex determination in humans: 22 pairs of autosomes + 1 pair of sex chromosomes. Female = XX, Male = XY. An egg fertilised by an X-carrying sperm results in a girl; fertilisation by a Y-carrying sperm results in a boy. Thus the sex of the child is determined by the type of sperm.

CHAPTER-9 LIGHT – REFLECTION AND REFRACTION

Topic-1 Reflection of Light, Images Formed by Spherical Mirrors

Laws of Reflection:

1. The incident ray, the reflected ray and the normal all lie in the same plane at the point of incidence.
2. Angle of incidence = Angle of reflection.

Real image – rays actually meet (can be obtained on screen). Virtual image – rays appear to meet (cannot be obtained on screen).

Plane mirror: Virtual, erect, same size as object, laterally inverted, as far behind the mirror as the object is in front.

Spherical mirrors:

• Concave mirror – reflecting surface recessed inward; can form real and virtual images.
• Convex mirror – reflecting surface bulges outward; always forms virtual, erect and diminished images.

Important terms: Pole (P), Centre of curvature (C), Principal axis, Focus (F), Radius of curvature R = 2f, Aperture.

Rules for ray diagrams (concave mirror):

1. A ray parallel to the principal axis passes through the principal focus after reflection.
2. A ray passing through the principal focus becomes parallel to the principal axis after reflection.
3. A ray passing through the centre of curvature is reflected back along the same path.
4. A ray incident obliquely to the principal axis is reflected obliquely making equal angle.

Image formation by a concave mirror:

Image formation by a convex mirror: Always virtual, erect and diminished (between P and F or at F for object at infinity).

Mirror Formula: \[ \dfrac{1}{v} + \dfrac{1}{u} = \dfrac{1}{f} \]

Magnification: \[ m = \dfrac{h_i}{h_o} = -\dfrac{v}{u} \] (negative → real image; positive → virtual image)

Topic-2 Refraction, Lenses and Power of Lens

Refraction of light: Change in the path of light when it travels from one medium to another (caused by change in speed of light).

• Rarer → denser medium: ray bends towards the normal (i > r)
• Denser → rarer medium: ray bends away from the normal (i < r)

Laws of Refraction (Snell’s law):

1. Incident ray, refracted ray and normal lie in the same plane.
2. \[ \dfrac{\sin i}{\sin r} = n \] (constant for a given pair of media and colour of light)

Refractive index: \[ n = \dfrac{c}{v} \] (absolute refractive index). Refractive index of diamond is 2.42 (highest).

Lens Formula: \[ \dfrac{1}{v} - \dfrac{1}{u} = \dfrac{1}{f} \]

Magnification: \[ m = \dfrac{h_i}{h_o} = \dfrac{v}{u} \]

Power of a lens: \[ P = \dfrac{1}{f} \] (f in metres). SI unit = dioptre (D). 1 D = power of a lens of focal length 1 m.

Image formation by a convex lens:

A concave lens always forms a virtual, erect and diminished image.

CHAPTER-10 HUMAN EYE AND COLOURFUL WORLD

Topic-1 Human Eye, Defects of Vision and Corrections

The human eye is a natural optical device that forms an inverted real image on the retina.

Parts of the eye:

• Cornea: Outermost transparent part; provides most of the refraction.
• Iris: Dark muscular diaphragm that controls the size of the pupil.
• Pupil: Central aperture that regulates the amount of light entering the eye.
• Lens: Convex lens that focuses the image on the retina.
• Ciliary muscles: Change the shape and focal length of the lens (accommodation).
• Retina: Light-sensitive membrane containing rods and cones.
• Optic nerve: Carries electrical signals from the retina to the brain.

Far point of a normal eye = infinity; Near point (least distance of distinct vision) = 25 cm. Range of vision = 25 cm to infinity. Accommodation is the ability of the eye lens to adjust its focal length.

Defects of vision and corrections:

• Myopia (Nearsightedness): Distant objects not seen clearly; image formed in front of the retina. Correction – concave lens.
• Hypermetropia (Farsightedness): Nearby objects not seen clearly; image formed behind the retina. Correction – convex lens.
• Presbyopia: Age-related loss of accommodation; near point recedes. Correction – bifocal lenses or convex lens.

Topic-2 Dispersion of Light and Scattering of Light

Dispersion: Splitting of white light into its constituent colours (VIBGYOR) on passing through a glass prism. Different colours undergo different deviations.

If a second identical prism is placed in an inverted position, the colours recombine to form white light.

Atmospheric refraction: Bending of light on passing through the Earth’s atmosphere. Causes twinkling of stars and makes the Sun appear to rise about two minutes earlier and set about two minutes later.

Scattering of light: Very small particles scatter light of shorter wavelengths better (reason for blue colour of the sky). Larger particles scatter light of all wavelengths equally (white appearance of clouds).

CHAPTER-11 ELECTRICITY

Topic-1 Electric Current, Ohm’s Law

Electric current: Rate of flow of electric charge through any cross-section of a conductor. \[ I = \dfrac{Q}{t} \]. SI unit = Ampere (A). It is a scalar quantity. Conventional direction of current is opposite to the direction of flow of electrons.

Potential difference / Voltage: \[ V = \dfrac{W}{Q} \]. SI unit = Volt (V). 1 V = 1 J C⁻¹.

Ohm’s Law: The current through a conductor between two points is directly proportional to the voltage across the two points provided external conditions remain constant. \[ V = IR \] or \[ I \propto V \]. The V–I graph is a straight line through the origin; its slope equals resistance R.

Resistance (R): Property of a conductor to resist the flow of charges. SI unit = Ohm (Ω). 1 Ω = 1 V / 1 A.

Resistance of a uniform metallic conductor is:

• Directly proportional to its length
• Inversely proportional to its area of cross-section
• Depends on the nature of the material and temperature

Resistivity (ρ): Resistance of a wire of unit length and unit cross-sectional area. SI unit = Ω m. Metals and alloys have low resistivity (10⁻⁸ to 10⁻⁶ Ω m); insulators have high resistivity (10¹² to 10¹⁷ Ω m).

Ammeter (low resistance) is connected in series; Voltmeter (high resistance) is connected in parallel.

Topic-2 Resistance in Series & Parallel, Electric Power and Heating Effect

Series combination: \[ R_s = R_1 + R_2 + R_3 + \dots \] (same current flows through each resistor)

Parallel combination: \[ \dfrac{1}{R_p} = \dfrac{1}{R_1} + \dfrac{1}{R_2} + \dfrac{1}{R_3} + \dots \] (same potential difference across each resistor)

Joule’s law of heating: Heat produced in a conductor is directly proportional to the square of current, resistance and time. \[ H = I^2Rt \] (or \[ H = VIt \])

Practical applications: electric iron, room heaters, water heaters, electric bulb (tungsten filament), electric fuse.

Electric Power: Rate at which electric energy is consumed or dissipated. \[ P = VI = I^2R = \dfrac{V^2}{R} \]. SI unit = Watt (W). Commercial unit of electric energy = kilowatt-hour (kWh). 1 kWh = 3.6 × 10⁶ J.

CHAPTER-12 MAGNETIC EFFECTS OF ELECTRIC CURRENT

Topic-1 Magnetic Effects

A magnet has two poles (North and South). Like poles repel; unlike poles attract. A freely suspended bar magnet aligns itself in the north-south direction.

Magnetic field: Region around a magnet in which its magnetic force can be experienced. SI unit = Tesla (T). It is a vector quantity.

Magnetic field lines:

• Emerge from North pole and enter South pole
• Are closed curves
• Never intersect each other
• Closer lines indicate stronger field

Right Hand Thumb Rule: If the thumb of the right hand points in the direction of current, the curled fingers give the direction of the magnetic field.

Magnetic field due to:

• Straight current-carrying conductor → concentric circles
• Circular loop → concentric circles that become larger and larger; almost straight at the centre
• Solenoid → similar to a bar magnet

Strength of magnetic field ∝ current, ∝ number of turns, ∝ 1/distance from the conductor.

Force on a current-carrying conductor in a magnetic field: \[ F = IBl\sin\theta \]

Fleming’s Left Hand Rule: Stretch the thumb, forefinger and middle finger of the left hand mutually perpendicular. Forefinger → magnetic field, Middle finger → current, Thumb → force (motion).

Topic-2 Domestic Electric Circuits

AC (Alternating Current): Reverses direction periodically (50 Hz in India). DC (Direct Current): Flows in one direction only.

Domestic supply in India: 220 V, 50 Hz. Three wires:

• Live wire (red insulation) – positive
• Neutral wire (black insulation) – negative
• Earth wire (green insulation) – safety wire that protects from electric shock

Safety devices:

• Electric fuse: Thin wire of low melting point and high resistance connected in series with the live wire. Protects against short-circuiting and overloading.
• Earth wire: Provides a low-resistance path for leakage current.

CHAPTER-13 OUR ENVIRONMENT

Topic-1 Ecosystem and Food Chain

Environment includes both biotic (living) and abiotic (non-living) components. Their interaction forms an ecosystem.

Types of ecosystem: Natural (forest, pond, lake, ocean) and Artificial (crop field, aquarium, garden).

Components of an ecosystem:

• Producers (Autotrophs): Green plants and algae that prepare food by photosynthesis.
• Consumers (Heterotrophs): Herbivores, carnivores, omnivores and parasites.
• Decomposers: Micro-organisms that break down dead organic matter into simple inorganic substances.

Food chain: Sequence of organisms through which food energy flows. Example: Grass → Grasshopper → Frog → Snake → Eagle.

Trophic levels: Producer (1st) → Herbivore (2nd) → Carnivore (3rd) → Top carnivore (4th).

Food web: Network of interconnected food chains.

Energy flow: Unidirectional. Only about 10 % of the energy is transferred from one trophic level to the next (10 percent law). The remaining 90 % is used in life processes or lost as heat. Therefore food chains are generally short (3–4 trophic levels).

Biological magnification: Concentration of harmful chemicals increases at successive trophic levels. Maximum concentration is found in the top consumer (often humans).

Topic-2 Biodegradable & Non-Biodegradable Substances, Ozone Depletion

Biodegradable wastes: Substances that can be decomposed by micro-organisms (fruit and vegetable peels, paper, cow-dung, cotton, jute etc.).

Non-biodegradable wastes: Substances that cannot be decomposed by micro-organisms (plastic, polythene, metals, synthetic fibres, radioactive wastes, pesticides etc.).

Methods of waste disposal:

• Biogas plant
• Sewage treatment plant
• Landfill
• Composting
• Recycling
• Reuse

Biodegradable and non-biodegradable wastes should be discarded in separate dustbins.

Ozone (O₃): Molecule formed by three atoms of oxygen. It is continuously formed and destroyed in the stratosphere by the action of UV radiation. Ozone layer shields the Earth’s surface from harmful ultraviolet radiation.

Ozone depletion: Reduction in the concentration of the ozone layer, mainly caused by chlorofluorocarbons (CFCs) and nitrogen monoxide (NO). This leads to the formation of the ozone hole (especially over Antarctica). More UV rays reach the Earth, causing harmful effects.

Protection of ozone layer: Stop the release of CFCs, reduce the use of air-conditioners, follow international agreements such as the Montreal Protocol (1987).