Comprehensive NCERT Class X Science Notes

Chemical Reactions and Equations

• Chemical Reactions and Changes:

  • Chemical Change: A process where the chemical nature and identity of the initial substance change, forming new substances with different properties.

  • Daily Life Examples of Chemical Changes:

    • Milk left at room temperature during summers turns sour.

    • An iron tawa, pan, or nail is left exposed to humid atmosphere (rusting).

    • Grapes get fermented.

    • Food is cooked.

    • Food gets digested in our body.

    • Respiration process in living organisms.

  • Observations Indicating a Chemical Reaction:

    • Change in state.

    • Change in color.

    • Evolution of a gas.

    • Change in temperature.

• Key Experimental Activities:

  • Magnesium Ribbon Burning (Activity 1.1):

    • Procedure: Clean a magnesium ribbon about $2\,cm$ long by rubbing it with sandpaper (to remove the protective layer of basic magnesium carbonate). Hold it with tongs and burn it using a spirit lamp or burner, collecting the ash in a watch-glass.

    • Observation: Magnesium ribbon burns with a dazzling white flame and changes into a white powder.

    • Chemical Identity: The white powder formed is magnesium oxide ($MgO$), formed by the reaction between magnesium and oxygen present in air:             $2Mg(s) + O_2(g) \rightarrow 2MgO(s)$

    • Caution: Wear protective eye glass as the flame emits intense ultraviolet light.

  • Lead Nitrate and Potassium Iodide (Activity 1.2):

    • Lead nitrate solution ($Pb(NO_3)_2$) reacts with potassium iodide solution ($KI$) to form a yellow precipitate of lead iodide ($PbI_2$) and potassium nitrate ($KNO_3$):             $Pb(NO_3)_2(aq) + 2KI(aq) \rightarrow PbI_2(s) + 2KNO_3(aq)$

  • Zinc Granules with Acid (Activity 1.3):

    • Zinc granules in a conical flask react with dilute hydrochloric acid ($HCl$) or dilute sulphuric acid ($H_2SO_4$).

    • Observation: Evolution of hydrogen gas ($H_2$) and the conical flask becomes hot due to heat generation (exothermic reaction).             $Zn(s) + H_2SO_4(aq) \rightarrow ZnSO_4(aq) + H_2(g)$

• Chemical Equations and Balancing:

  • Skeletal Chemical Equation: An unbalanced chemical equation where the mass is not the same on both sides of the equation (e.g., $Mg + O_2 \rightarrow MgO$).

  • Balanced Chemical Equation: An equation where the total mass of elements present in products equals the total mass of elements present in reactants, complying with the Law of Conservation of Mass (mass can neither be created nor destroyed in a chemical reaction).

  • Step-by-step Balancing Method (Hit and Trial Method):

    • Unbalanced Equation: $Fe + H_2O \rightarrow Fe_3O_4 + H_2$

    • Step 1: Draw boxes around each formula; do not change anything inside boxes.

    • Step 2: Count the number of atoms of each element on LHS and RHS.

    • Step 3: Balance the element with the maximum number of atoms ($O$ in $Fe_3O_4$). Multiply $H_2O$ by $4$ on LHS: $Fe + 4H_2O \rightarrow Fe_3O_4 + H_2$

    • Step 4: Balance $H$ atoms. Multiply $H_2$ by $4$ on RHS: $Fe + 4H_2O \rightarrow Fe_3O_4 + 4H_2$

    • Step 5: Balance $Fe$ atoms. Multiply $Fe$ by $3$ on LHS: $3Fe + 4H_2O \rightarrow Fe_3O_4 + 4H_2$

    • Step 6: Verify the number of atoms on both sides ($3\,Fe$, $8\,H$, $4\,O$ on both sides).

    • Step 7: Write physical states: gaseous $(g)$, liquid $(l)$, solid $(s)$, and aqueous $(aq)$.             $3Fe(s) + 4H_2O(g) \rightarrow Fe_3O_4(s) + 4H_2(g)$             (Note: $H_2O(g)$ denotes steam).

    • Specific reaction conditions (temperature, pressure, catalyst) are written above/below the arrow (e.g., $CO(g) + 2H_2(g) \xrightarrow{340\,atm} CH_3OH(l)$).

• Types of Chemical Reactions:

  1. Combination Reaction: A reaction where two or more substances combine to form a single product.

     • Slaking of Lime: Calcium oxide (quicklime) reacts vigorously with water to produce slaked lime (calcium hydroxide), releasing a large amount of heat:             $CaO(s) + H_2O(l) \rightarrow Ca(OH)_2(aq) + \text{Heat}$

     • Whitewashing: Slaked lime $Ca(OH)_2$ applied to walls reacts slowly with $CO_2$ in air to form a thin layer of calcium carbonate ($CaCO_3$) giving a shiny finish after 2–3 days. (Chemical formula for marble is also $CaCO_3$):             $Ca(OH)_2(aq) + CO_2(g) \rightarrow CaCO_3(s) + H_2O(l)$

     • Other Examples: Burning of coal: $C(s) + O_2(g) \rightarrow CO_2(g)$; Formation of water: $2H_2(g) + O_2(g) \rightarrow 2H_2O(l)$.

  2. Exothermic Reactions: Reactions in which heat is released along with the formation of products.

     • Burning of natural gas: $CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(g)$

     • Respiration: Digestion breaks carbohydrates into glucose ($C_6H_{12}O_6$), which combines with oxygen in cells to provide energy:             $C_6H_{12}O_6(aq) + 6O_2(aq) \rightarrow 6CO_2(aq) + 6H_2O(l) + \text{Energy}$

     • Decomposition of vegetable matter into compost.

  3. Decomposition Reaction: A reaction in which a single reactant breaks down to yield two or more simpler products.

     • Thermal Decomposition (decomposition carried out by heating):

       • Heating ferrous sulphate crystals: Green $FeSO_4 \cdot 7H_2O$ crystals lose water upon heating, changing color, and then decompose to ferric oxide ($Fe_2O_3$), sulphur dioxide ($SO_2$), and sulphur trioxide ($SO_3$):                 $2FeSO_4(s) \xrightarrow{\text{Heat}} Fe_2O_3(s) + SO_2(g) + SO_3(g)$

       • Decomposition of limestone: $CaCO_3(s) \xrightarrow{\text{Heat}} CaO(s) + CO_2(g)$. Quicklime ($CaO$) is used in cement manufacturing.

       • Decomposition of lead nitrate: Heating $Pb(NO_3)_2$ produces brown fumes of nitrogen dioxide ($NO_2$):                 $2Pb(NO_3)_2(s) \xrightarrow{\text{Heat}} 2PbO(s) + 4NO_2(g) + O_2(g)$

     • Electrolytic Decomposition (Electrolysis):

       • Electrolysis of water: Passage of electric current through water breaks it into hydrogen gas at the cathode and oxygen gas at the anode in a $2:1$ volume ratio:                 $2H_2O(l) \xrightarrow{\text{Electricity}} 2H_2(g) + O_2(g)$

     • Photolytic Decomposition (decomposition by light):

       • White silver chloride ($AgCl$) turns grey in sunlight due to decomposition into silver and chlorine:                 $2AgCl(s) \xrightarrow{\text{Sunlight}} 2Ag(s) + Cl_2(g)$

       • Silver bromide ($AgBr$) decomposes similarly (used in black-and-white photography):                 $2AgBr(s) \xrightarrow{\text{Sunlight}} 2Ag(s) + Br_2(g)$

  4. Endothermic Reactions: Reactions in which energy is absorbed in the form of heat, light, or electricity.

  5. Displacement Reaction: A chemical reaction in which a more reactive element displaces a less reactive element from its compound solution.

     • Iron nail in copper sulphate solution: The blue color of $CuSO_4$ fades to light green, and a brown coating of copper deposits on the iron nail:             $Fe(s) + CuSO_4(aq) \rightarrow FeSO_4(aq) + Cu(s)$

     • Zinc displacing copper: $Zn(s) + CuSO_4(aq) \rightarrow ZnSO_4(aq) + Cu(s)$

     • Lead displacing copper: $Pb(s) + CuCl_2(aq) \rightarrow PbCl_2(aq) + Cu(s)$

  6. Double Displacement Reaction: A reaction where there is an exchange of ions between reactants, often forming an insoluble precipitate.

     • Sodium sulphate and barium chloride: Formation of a white insoluble precipitate of barium sulphate ($BaSO_4$):             $Na_2SO_4(aq) + BaCl_2(aq) \rightarrow BaSO_4(s) + 2NaCl(aq)$

     • Precipitation Reaction: Any reaction that produces an insoluble salt (precipitate).

  7. Oxidation and Reduction (Redox) Reactions:

     • Oxidation: A process involving the gain of oxygen or loss of hydrogen.

     • Reduction: A process involving the loss of oxygen or gain of hydrogen.

     • Example 1: Copper heated in air forms black copper(II) oxide ($2Cu + O_2 \xrightarrow{\text{Heat}} 2CuO$). When $H_2$ gas is passed over heated $CuO$, the black coating turns brown as copper is formed:             $CuO + H_2 \xrightarrow{\text{Heat}} Cu + H_2O$             (Here, $CuO$ is reduced to $Cu$, and $H_2$ is oxidized to $H_2O$).

     • Example 2: $ZnO + C \rightarrow Zn + CO$ ($ZnO$ reduced, $C$ oxidized).

     • Example 3: $MnO_2 + 4HCl \rightarrow MnCl_2 + 2H_2O + Cl_2$ ($HCl$ oxidized to $Cl_2$, $MnO_2$ reduced to $MnCl_2$).

• Corrosion and Rancidity:

  • Corrosion: The process in which metals are slowly eaten away by the action of air, moisture, or chemical substances on their surface.

    • Iron: Forms reddish-brown flaky rust ($Fe_2O_3 \cdot xH_2O$).

    • Silver: Develops a black coating of silver sulphide ($Ag_2S$) due to reaction with $H_2S$ in air.

    • Copper: Develops a green coating of basic copper carbonate ($CuCO_3 \cdot Cu(OH)_2$) due to reaction with moist $CO_2$ in air.

  • Rancidity: The oxidation of fats and oils in food when exposed to air, causing a change in smell and taste.

    • Prevention Methods: Adding antioxidants to foods containing fats/oils, keeping food in airtight containers, or flushing food packaging (like potato chips) with unreactive nitrogen gas ($N_2$) to prevent oxidation.

Acids, Bases and Salts

• Indicators and Basic Properties:

  • Acids: Sour in taste, turn blue litmus paper red, release $H^+(aq)$ ions in water.

  • Bases: Bitter in taste, soapy to touch, turn red litmus paper blue, release $OH^-(aq)$ ions in water.

  • Litmus: A natural indicator; a purple dye extracted from Lichen (a plant belonging to division Thallophyta). Neutral litmus solution is purple.

  • Other Natural Indicators: Turmeric (yellow, turns reddish-brown in basic medium, turns back yellow on washing with water), red cabbage leaves, colored petals of flowers like Hydrangea, Petunia, and Geranium.

  • Synthetic Indicators: Methyl orange, Phenolphthalein (colorless in acidic, pink in basic solution).

  • Olfactory Indicators: Substances whose odor changes in acidic or basic media (e.g., onion, vanilla essence, clove oil). Vanilla, onion, and clove retain odor in acid but lose odor in $NaOH$ solution.

• Chemical Properties of Acids and Bases:

  1. Reaction of Metals with Acids and Bases:

     • $\text{Acid} + \text{Metal} \rightarrow \text{Salt} + \text{Hydrogen gas}$

     • Zinc + Sulphuric Acid: $Zn(s) + H_2SO_4(aq) \rightarrow ZnSO_4(aq) + H_2(g)$

     • Hydrogen Test: Gas bubbles passed through soap solution form soap bubbles. On bringing a burning candle near a bubble filled with hydrogen, it burns with a characteristic pop sound.

     • Base + Metal: $2NaOH(aq) + Zn(s) \rightarrow Na_2ZnO_2(aq) + H_2(g)$ (Sodium zincate formed; such reactions do not occur with all metals).

  2. Reaction with Metal Carbonates and Metal Hydrogen Carbonates:

     • $\text{Metal Carbonate / Hydrogen Carbonate} + \text{Acid} \rightarrow \text{Salt} + \text{Water} + \text{Carbon Dioxide}$

     • $Na_2CO_3(s) + 2HCl(aq) \rightarrow 2NaCl(aq) + H_2O(l) + CO_2(g)$

     • $NaHCO_3(s) + HCl(aq) \rightarrow NaCl(aq) + H_2O(l) + CO_2(g)$

     • Carbon Dioxide / Lime Water Test: Passing $CO_2$ gas through lime water ($Ca(OH)_2$) turns it milky due to insoluble calcium carbonate formation:             $Ca(OH)_2(aq) + CO_2(g) \rightarrow CaCO_3(s) + H_2O(l)$

     • Passing excess $CO_2$: Milkiness disappears due to soluble calcium hydrogen carbonate formation:             $CaCO_3(s) + H_2O(l) + CO_2(g) \rightarrow Ca(HCO_3)_2(aq)$

     • Limestone, chalk, and marble are different crystalline forms of calcium carbonate ($CaCO_3$).

  3. Neutralisation Reaction:

     • The reaction between an acid and a base to produce salt and water:             $\text{Base} + \text{Acid} \rightarrow \text{Salt} + \text{Water}$             $NaOH(aq) + HCl(aq) \rightarrow NaCl(aq) + H_2O(l)$

  4. Reaction of Metallic Oxides with Acids:

     • $\text{Metal Oxide} + \text{Acid} \rightarrow \text{Salt} + \text{Water}$

     • $CuO(s) + 2HCl(aq) \rightarrow CuCl_2(aq) + H_2O(l)$ (Solution turns blue-green due to copper(II) chloride formation).

     • Since metallic oxides react with acids to yield salt and water, metallic oxides are basic oxides.

  5. Reaction of Non-Metallic Oxides with Bases:

     • $\text{Non-metallic Oxide} + \text{Base} \rightarrow \text{Salt} + \text{Water}$

     • $Ca(OH)_2(aq) + CO_2(g) \rightarrow CaCO_3(s) + H_2O(l)$

     • Since non-metallic oxides react with bases to yield salt and water, non-metallic oxides are acidic oxides.

• Acids and Bases in Water Solution:

  • Acids produce hydronium ions ($H_3O^+$) or hydrogen ions ($H^+(aq)$) only in the presence of water:         $HCl + H_2O \rightarrow H_3O^+ + Cl^-$         (Free $H^+$ ions cannot exist alone; they exist as $H^+(aq)$ or $H_3O^+$).

  • Dry $HCl$ gas does not turn dry blue litmus paper red because no $H^+$ ions are produced without moisture.

  • Bases generate hydroxide ions ($OH^-$) in water:         $NaOH(s) \xrightarrow{H_2O} Na^+(aq) + OH^-(aq)$         $KOH(s) \xrightarrow{H_2O} K^+(aq) + OH^-(aq)$         $Mg(OH)_2(s) \xrightarrow{H_2O} Mg^{2+}(aq) + 2OH^-(aq)$

  • Alkalis: Bases that are soluble in water (e.g., $NaOH$, $KOH$, $Ca(OH)_2$). Alkalis are soapy to touch, bitter, and corrosive.

  • Dilution Process: Mixing an acid or base with water results in a decrease in the concentration of ions ($H_3O^+$ or $OH^-$) per unit volume.

    • Warning: Mixing concentrated acid (like $H_2SO_4$ or $HNO_3$) with water is highly exothermic. Always add acid slowly to water with constant stirring, never water to acid, as intense local heating may cause the splash of acidic mixture or glass container breakage.

• Strength of Acid and Base Solutions & pH Scale:

  • pH Scale: A scale for measuring hydrogen ion concentration in a solution. The $p$ in pH stands for potenz (German for power).

    • pH ranges from $0$ (very acidic) to $14$ (very alkaline); $7$ represents a neutral solution.

    • $\text{Neutral Solution}: pH = 7$

    • \text{Acidic Solution}: pH < 7 (Lower pH means higher $H^+$ concentration).

    • \text{Basic Solution}: pH > 7 (Higher pH means higher $OH^-$ concentration).

  • Universal Indicator: A mixture of several indicators showing different colors at different pH values on the pH scale.

  • Importance of pH in Everyday Life:

    • Biological sensitivity: Human body functions within a narrow pH range of $7.0$ to $7.8$. When pH of rain water is less than $5.6$, it is called acid rain, which lowers river water pH and threatens aquatic life.

    • Soil pH: Optimal plant growth requires a specific pH range. Acidic soil is treated with quicklime ($CaO$), slaked lime ($Ca(OH)_2$), or chalk ($CaCO_3$).

    • Digestive system: Stomach produces hydrochloric acid ($HCl$) of pH ~$1.2$ to digest food. Excess acid during indigestion causes pain/irritation; treated with antacids (mild bases like sodium hydrogen carbonate $NaHCO_3$ or Milk of Magnesia $Mg(OH)_2$) to neutralize excess acid.

    • Tooth decay: Tooth decay starts when mouth pH falls below $5.5$. Tooth enamel, made of calcium hydroxyapatite (a crystalline form of calcium phosphate, $Ca_5(PO_4)_3OH$), is the hardest substance in the body. It does not dissolve in water but corrodes when pH $< 5.5$. Bacteria produce acids by degrading sugar/food particles. Prevented by brushing with basic toothpaste.

    • Self-defense by animals/plants: Bee sting injects an acid causing pain/swelling (relieved by applying mild base like baking soda). Ant sting / Nettle leaf sting injects methanoic acid causing burning pain; traditional remedy is rubbing leaves of the dock plant (Rumex), which often grows near nettle plants.

• Naturally Occurring Acids:

  • Vinegar: Acetic acid

  • Orange / Lemon: Citric acid

  • Tamarind: Tartaric acid

  • Tomato: Oxalic acid

  • Sour milk (Curd): Lactic acid

  • Ant sting / Nettle sting: Methanoic acid

• Salts and Chemical Compounds from Common Salt ($NaCl$):

  • pH of Salts:

    • Strong Acid + Strong Base $\rightarrow$ Neutral Salt ($pH = 7$, e.g., $NaCl$, $Na_2SO_4$).

    • Strong Acid + Weak Base $\rightarrow$ Acidic Salt ($pH < 7$, e.g., $NH_4Cl$).

    • Weak Acid + Strong Base $\rightarrow$ Basic Salt ($pH > 7$, e.g., $Na_2CO_3$, $CH_3COONa$).

  • Common Salt ($NaCl$): Obtained from sea water or mined as rock salt (brown colored due to impurities).

  • Sodium Hydroxide ($NaOH$) / Chlor-Alkali Process:

    • Electrolysis of an aqueous solution of sodium chloride (called brine):             $2NaCl(aq) + 2H_2O(l) \xrightarrow{\text{Electricity}} 2NaOH(aq) + Cl_2(g) + H_2(g)$

    • Chlorine gas ($Cl_2$) liberated at the anode (used for water treatment, PVC, CFCs, disinfectants, pesticides).

    • Hydrogen gas ($H_2$) liberated at the cathode (used for fuels, margarine, ammonia for fertilisers).

    • Sodium hydroxide ($NaOH$) formed near the cathode (used for de-greasing metals, soaps, detergents, paper making, artificial fibres).

  • Bleaching Powder ($CaOCl_2$):

    • Produced by the action of chlorine on dry slaked lime:             $Ca(OH)_2 + Cl_2 \rightarrow CaOCl_2 + H_2O$

    • Uses: Bleaching cotton/linen in textile industry, wood pulp in paper factories, disinfecting drinking water, oxidizing agent in chemical industries.

  • Baking Soda ($NaHCO_3$ - Sodium Hydrogen Carbonate):

    • Produced using sodium chloride:             $NaCl + H_2O + CO_2 + NH_3 \rightarrow NH_4Cl + NaHCO_3$

    • Mild non-corrosive basic salt. On heating during cooking:             $2NaHCO_3 \xrightarrow{\text{Heat}} Na_2CO_3 + H_2O + CO_2(g)$

    • Uses:

      1. Making baking powder (mixture of baking soda + mild edible acid like tartaric acid). When heated or mixed in water, $CO_2$ gas released makes cake/bread rise, turning them soft and spongy:                 $NaHCO_3 + H^+ \rightarrow CO_2 + H_2O + \text{Sodium salt of acid}$

      2. Ingredient in antacids.

      3. Used in soda-acid fire extinguishers.

  • Washing Soda ($Na_2CO_3 \cdot 10H_2O$ - Sodium Carbonate Decahydrate):

    • Recrystallisation of sodium carbonate yields washing soda:             $Na_2CO_3 + 10H_2O \rightarrow Na_2CO_3 \cdot 10H_2O$

    • Basic salt.

    • Uses: Glass, soap, and paper industries; manufacturing sodium compounds like borax ($Na_2B_4O_7 \cdot 10H_2O$); cleaning agent for domestic purposes; removing permanent hardness of water.

  • Water of Crystallisation:

    • The fixed number of water molecules present in one formula unit of a salt in crystalline state.

    • Copper sulphate: $CuSO_4 \cdot 5H_2O$ (Blue). Heating drives off water, turning it white. Re-moistening restores blue color.

    • Gypsum: $CaSO_4 \cdot 2H_2O$ (contains 2 molecules of water of crystallisation).

  • Plaster of Paris ($CaSO_4 \cdot \frac{1}{2}H_2O$ - Calcium Sulphate Hemihydrate):

    • Obtained by heating gypsum at $373\,K$ ($100^\circ C$):             $CaSO_4 \cdot 2H_2O \xrightarrow{373\,K} CaSO_4 \cdot \frac{1}{2}H_2O + 1\frac{1}{2}H_2O$

    • On mixing with water, it sets into a hard solid mass of gypsum:             $CaSO_4 \cdot \frac{1}{2}H_2O + 1\frac{1}{2}H_2O \rightarrow CaSO_4 \cdot 2H_2O$

    • Note: Two formula units of $CaSO_4$ share one molecule of water ($2CaSO_4 \cdot H_2O$).

    • Uses: Plastering fractured bones, making toys, decorative materials, casts, smooth surfaces.

Metals and Non-Metals

• Physical Properties of Metals and Non-Metals:

  • Metals:

    • Metallic Lustre: Clean shiny surface in pure state.

    • Hardness: Generally hard (varies from metal to metal).

    • Malleability: Can be beaten into thin sheets (Gold and silver are the most malleable metals).

    • Ductility: Can be drawn into thin wires (Gold is the most ductile metal: $1\,g$ of gold can be drawn into a wire of about $2\,km$ length).

    • Thermal Conductivity: Good conductors of heat with high melting points. Silver ($Ag$) and Copper ($Cu$) are the best conductors. Lead ($Pb$) and Mercury ($Hg$) are comparatively poor conductors of heat.

    • Electrical Conductivity: Good conductors of electricity (electric wires coated with PVC or rubber-like insulating material).

    • Sonorousness: Produce a sound on striking a hard surface.

  • Non-Metals:

    • Examples: Carbon, sulphur, iodine, oxygen, hydrogen, etc.

    • Exist as solids or gases, except Bromine which is a liquid non-metal.

  • Exceptions in Physical Properties:

    • All metals except Mercury are solids at room temperature.

    • Gallium ($Ga$) and Caesium ($Cs$) have very low melting points; they melt on the palm of the hand.

    • Iodine is a non-metal but it is lustrous.

    • Carbon exists in different forms called allotropes:

      • Diamond: Allotrope of carbon, hardest natural substance known, extremely high melting and boiling point.

      • Graphite: Allotrope of carbon, good conductor of electricity.

    • Alkali metals (Lithium, Sodium, Potassium): Very soft, can be cut easily with a knife, low densities and low melting points.

• Chemical Properties of Metals:

  1. Reaction of Metals with Oxygen:

     • $\text{Metal} + \text{Oxygen} \rightarrow \text{Metal oxide}$

     • $2Cu + O_2 \xrightarrow{\text{Heat}} 2CuO$ (Copper(II) oxide, black oxide).

     • $4Al + 3O_2 \rightarrow 2Al_2O_3$ (Aluminium oxide).

     • Amphoteric Oxides: Metal oxides that react with both acids as well as bases to produce salt and water (e.g., $Al_2O_3$, $ZnO$):             $Al_2O_3 + 6HCl \rightarrow 2AlCl_3 + 3H_2O$             $Al_2O_3 + 2NaOH \rightarrow 2NaAlO_2 + 3H_2O$ (Sodium aluminate)

     • Most metal oxides are insoluble in water, but some dissolve to form alkalis:             $Na_2O(s) + H_2O(l) \rightarrow 2NaOH(aq)$             $K_2O(s) + H_2O(l) \rightarrow 2KOH(aq)$

     • Reactivity with $O_2$:

       • Sodium ($Na$) and Potassium ($K$) react vigorously, catching fire if kept in open air. Stored immersed in kerosene oil.

       • Magnesium, Aluminium, Zinc, Lead ($Mg, Al, Zn, Pb$) form a thin protective oxide layer at ordinary temperatures to prevent further oxidation.

       • Iron ($Fe$) does not burn on heating, but iron filings burn vigorously when sprinkled in the flame.

       • Copper ($Cu$) does not burn, but hot metal is coated with black $CuO$.

       • Silver ($Ag$) and Gold ($Au$) do not react with oxygen even at high temperatures.

     • Anodising: An electrochemical process of forming a thick protective aluminium oxide layer on aluminium articles to prevent corrosion.

  2. Reaction of Metals with Water:

     • $\text{Metal} + \text{Water} \rightarrow \text{Metal oxide} + \text{Hydrogen}$

     • $\text{Metal oxide} + \text{Water} \rightarrow \text{Metal hydroxide}$

     • Sodium and Potassium: React violently with cold water; evolved $H_2$ catches fire immediately (exothermic):             $2K(s) + 2H_2O(l) \rightarrow 2KOH(aq) + H_2(g) + \text{Heat energy}$             $2Na(s) + 2H_2O(l) \rightarrow 2NaOH(aq) + H_2(g) + \text{Heat energy}$

     • Calcium: Reacts less violently with cold water; heat released is insufficient for $H_2$ to catch fire. Calcium floats because bubbles of $H_2$ gas stick to its surface:             $Ca(s) + 2H_2O(l) \rightarrow Ca(OH)_2(aq) + H_2(g)$

     • Magnesium: Does not react with cold water; reacts with hot water to form $Mg(OH)_2$ and $H_2$, and floats due to hydrogen gas bubbles.

     • Aluminium, Iron, Zinc ($Al, Fe, Zn$): Do not react with cold or hot water; react with steam:             $2Al(s) + 3H_2O(g) \rightarrow Al_2O_3(s) + 3H_2(g)$             $3Fe(s) + 4H_2O(g) \rightarrow Fe_3O_4(s) + 4H_2(g)$

     • Lead, Copper, Silver, Gold ($Pb, Cu, Ag, Au$): Do not react with water at all.

  3. Reaction of Metals with Acids:

     • $\text{Metal} + \text{Dilute acid} \rightarrow \text{Salt} + \text{Hydrogen}$

     • Reactivity order with $HCl$: Mg > Al > Zn > Fe. Copper does not react with dilute $HCl$ (no hydrogen gas bubbles, temperature remains unchanged).

     • Nitric Acid ($HNO_3$): Hydrogen gas is not evolved when a metal reacts with $HNO_3$ because $HNO_3$ is a strong oxidizing agent. It oxidizes $H_2$ produced to water ($H_2O$) and itself gets reduced to nitrogen oxides ($NO_2, NO, N_2O$).

     • Exception: Magnesium ($Mg$) and Manganese ($Mn$) react with very dilute $HNO_3$ to evolve $H_2$ gas:             $Mg + 2HNO_3 \rightarrow Mg(NO_3)_2 + H_2$

     • Aqua Regia (Royal water): A freshly prepared mixture of concentrated hydrochloric acid ($HCl$) and concentrated nitric acid ($HNO_3$) in the ratio $3:1$. It can dissolve gold and platinum, though neither acid can do so alone. Highly corrosive, fuming liquid.

  4. Reaction of Metals with Solutions of Other Metal Salts (Displacement):

     • $\text{Metal A} + \text{Salt solution of B} \rightarrow \text{Salt solution of A} + \text{Metal B}$

     • A more reactive metal displaces a less reactive metal from its compound in solution.

• The Reactivity (Activity) Series:

  • Arrangement of metals in order of their decreasing activities:

    • $\text{Potassium } (K)$ - Most reactive

    • $\text{Sodium } (Na)$

    • $\text{Calcium } (Ca)$

    • $\text{Magnesium } (Mg)$

    • $\text{Aluminium } (Al)$

    • $\text{Zinc } (Zn)$

    • $\text{Iron } (Fe)$

    • $\text{Lead } (Pb)$

    • $\text{[Hydrogen] } (H)$

    • $\text{Copper } (Cu)$

    • $\text{Mercury } (Hg)$

    • $\text{Silver } (Ag)$

    • $\text{Gold } (Au)$ - Least reactive

• How Metals and Non-Metals React (Ionic Bonding):

  • Reactivity of elements is explained as a tendency to attain a completely filled valence shell (noble gas electronic configuration).

  • Metals: Lose valence electrons to form positive ions (cations), e.g., $Na \rightarrow Na^+ + e^-$.

  • Non-metals: Gain electrons into valence shell to form negative ions (anions), e.g., $Cl + e^- \rightarrow Cl^-$.

  • Ionic (Electrovalent) Compounds: Compounds formed by the transfer of electrons from a metal to a non-metal (held by strong electrostatic forces of attraction).

    • Formation of $NaCl$: $Na (2,8,1) + Cl (2,8,7) \rightarrow [Na]^+\,[:\ddot{\text{Cl}}:]^-$

    • Formation of $MgCl_2$: $Mg (2,8,2) + 2Cl (2,8,7) \rightarrow [Mg]^{2+}\,2[:\ddot{\text{Cl}}:]^-$

  • Properties of Ionic Compounds:

    1. Physical Nature: Solid, hard, and generally brittle; break into pieces when pressure is applied.

    2. Melting and Boiling Points: High melting and boiling points due to strong inter-ionic electrostatic forces requiring considerable energy to break (e.g., $NaCl$ mp $1074\,K$, bp $1686\,K$).

    3. Solubility: Soluble in water (polar solvent), insoluble in organic solvents like kerosene, petrol, etc.

    4. Conduction of Electricity: Conduct electricity in aqueous solution or molten state (ions are free to move); do not conduct electricity in solid state (rigid crystal structure prevents movement of ions).

• Occurrence and Extraction of Metals (Metallurgy):

  • Minerals: Naturally occurring elements/compounds in the earth's crust.

  • Ores: Minerals containing a very high percentage of a particular metal that can be profitably extracted.

  • Gangue: Earthy impurities like soil, sand, etc., present in mined ores.

  • Steps in Metallurgy:

    1. Concentration / Enrichment of Ore: Removal of gangue impurities based on physical/chemical properties.

    2. Extraction of Metals Low in Activity Series (Unreactive metals; oxides reduced by heat alone):

       • Cinnabar ($HgS$, ore of mercury): $2HgS(s) + 3O_2(g) \xrightarrow{\text{Heat}} 2HgO(s) + 2SO_2(g)$; $2HgO(s) \xrightarrow{\text{Heat}} 2Hg(l) + O_2(g)$

       • Copper glance ($Cu_2S$): $2Cu_2S + 3O_2 \xrightarrow{\text{Heat}} 2Cu_2O + 2SO_2$; $2Cu_2O + Cu_2S \xrightarrow{\text{Heat}} 6Cu + SO_2$

    3. Extraction of Metals in Middle of Activity Series ($Zn, Fe, Pb, Cu$; converted to oxides first):

       • Roasting: Heating sulphide ores strongly in the presence of excess air:                 $2ZnS(s) + 3O_2(g) \xrightarrow{\text{Heat}} 2ZnO(s) + 2SO_2(g)$

       • Calcination: Heating carbonate ores strongly in limited air:                 $ZnCO_3(s) \xrightarrow{\text{Heat}} ZnO(s) + CO_2(g)$

       • Reduction: Metal oxides reduced using carbon (coke):                 $ZnO(s) + C(s) \rightarrow Zn(s) + CO(g)$

       • Reduction using Displacement: Highly reactive metals ($Na, Ca, Al$) used as reducing agents:                 $3MnO_2(s) + 4Al(s) \rightarrow 3Mn(l) + 2Al_2O_3(s) + \text{Heat}$

       • Thermite Process: Reduction of iron(III) oxide with aluminium powder. Highly exothermic; iron is produced in molten state, used to join railway tracks or cracked machine parts:                 $Fe_2O_3(s) + 2Al(s) \rightarrow 2Fe(l) + Al_2O_3(s) + \text{Heat}$

    4. Extraction of Metals High in Activity Series ($Na, Mg, Ca, Al$; cannot be reduced by carbon due to high affinity for oxygen):

       • Electrolytic Reduction: Metals extracted by electrolysis of their molten salts.

       • Sodium from molten $NaCl$:                 $\text{At Cathode (Negative electrode)}: Na^+ + e^- \rightarrow Na$                 $\text{At Anode (Positive electrode)}: 2Cl^- \rightarrow Cl_2 + 2e^-$

       • Aluminium is extracted by electrolytic reduction of molten aluminium oxide.

    5. Refining of Metals (Electrolytic Refining):

       • Purification of crude metal (e.g., Copper, Zinc, Tin, Nickel, Silver, Gold).

       • Anode: Block of impure metal.

       • Cathode: Thin strip of pure metal.

       • Electrolyte: Aqueous solution of metal salt (e.g., acidified $CuSO_4$).

       • On passing current, pure metal from anode dissolves into electrolyte and an equivalent amount of pure metal deposits on cathode:                 $\text{Anode}: Cu \rightarrow Cu^{2+} + 2e^-$                 $\text{Cathode}: Cu^{2+} + 2e^- \rightarrow Cu$

       • Anode Mud: Insoluble impurities that settle down at the bottom of the anode.

• Corrosion and Its Prevention:

  • Conditions for Rusting: Both air (oxygen) and water (moisture) are essential for rusting of iron.

  • Methods of Prevention:

    • Painting, oiling, greasing, enameling.

    • Galvanisation: Method of protecting steel/iron from rusting by coating with a thin layer of Zinc. Galvanised articles remain protected even if the zinc coating is scratched.

    • Chrome plating, Anodising.

    • Alloying: Mixing a metal with other metals or non-metals to improve properties.

  • Alloys:

    • Homogeneous mixture of two or more metals, or a metal and a non-metal.

    • Pure Iron: Very soft, stretches easily when hot. Mixed with $0.05\%$ carbon to become hard and strong.

    • Stainless Steel: Iron + Nickel ($Ni$) + Chromium ($Cr$). Hard, does not rust.

    • Amalgam: An alloy containing Mercury ($Hg$) as one of the constituents.

    • Brass: Alloy of Copper ($Cu$) and Zinc ($Zn$). Non-sonorous, poor electrical conductor.

    • Bronze: Alloy of Copper ($Cu$) and Tin ($Sn$). Poor electrical conductor.

    • Solder: Alloy of Lead ($Pb$) and Tin ($Sn$). Has low melting point; used for welding electrical wires.

    • Purity of Gold: Pure gold is 24 carat (very soft, unsuitable for jewelry). Alloyed with Copper or Silver to make 22 carat gold (22 parts pure gold + 2 parts Cu/Ag).

  • Iron Pillar at Delhi: Near Qutub Minar, over $1600$ years old, $8\,m$ high, weighs $6\,tonnes$. Developed rust-resistance mechanism (formation of thin magnetic oxide film $Fe_3O_4$ on surface).

Carbon and Its Compounds

• Covalent Bonding in Carbon:

  • Carbon atomic number = $6$; electronic configuration = $(2,4)$; valence electrons = $4$.

  • Carbon cannot form ionic bonds:

    • Gaining 4 electrons ($C^{4-}$ anion): Difficult for nucleus with 6 protons to hold 10 electrons.

    • Losing 4 electrons ($C^{4+}$ cation): Requires a huge amount of energy to remove 4 electrons, leaving a cation with 6 protons holding 2 electrons.

  • Carbon overcomes this by sharing its valence electrons with other carbon atoms or atoms of other elements, forming covalent bonds.

  • Simple Covalent Molecules:

    • Hydrogen ($H_2$): Two H atoms share 1 pair of electrons $\rightarrow$ Single bond ($H-H$).

    • Chlorine ($Cl_2$): Single bond ($Cl-Cl$).

    • Oxygen ($O_2$): Shares 2 pairs of electrons $\rightarrow$ Double bond ($O=O$).

    • Nitrogen ($N_2$): Shares 3 pairs of electrons $\rightarrow$ Triple bond ($N \equiv N$).

    • Water ($H_2O$): Single bonds between O and two H atoms ($H-O-H$).

    • Ammonia ($NH_3$): Nitrogen bonded to three H atoms with one lone pair.

    • Methane ($CH_4$): Carbon bonded to four H atoms.

  • Properties of Covalent Compounds:

    • Strong covalent bonds within the molecule, but weak intermolecular forces.

    • Low melting and boiling points (e.g., $CH_4$ mp $90\,K$, bp $111\,K$; $CH_3COOH$ mp $290\,K$, bp $391\,K$).

    • Poor conductors of electricity because electrons are shared and no free charged particles/ions are formed.

• Allotropes of Carbon:

  1. Diamond: Each carbon atom is bonded to four other carbon atoms in a rigid three-dimensional structure. Extremely hard, insulator.

  2. Graphite: Each carbon atom is bonded to three other carbon atoms in the same plane giving a hexagonal array (layers stacked above each other). One bond is a double bond. Smooth and slippery, good conductor of electricity due to free valence electrons.

  3. Buckminsterfullerene ($C_{60}$): Carbon atoms arranged in the shape of a football (geodesic dome designed by Buckminster Fuller).

• Versatile Nature of Carbon:

  • Carbon forms millions of compounds due to two major factors:

    1. Catenation: The unique ability of carbon to form bonds with other carbon atoms, giving rise to large molecules (long straight chains, branched chains, or closed rings). C-C single bond energy is extremely strong and stable.

    2. Tetravalency: Having a valency of 4, carbon can bond with 4 other monovalent atoms or heteroatoms ($O, N, S, Cl$).

• Saturated and Unsaturated Carbon Compounds:

  • Saturated Compounds (Alkanes): Compounds of carbon and hydrogen where carbon atoms are linked only by single bonds ($C-C$). General formula: $C_n H_{2n+2}$. Unreactive/stable.

  • Unsaturated Compounds: Carbon compounds containing one or more double bonds (Alkenes: $C_n H_{2n}$) or triple bonds (Alkynes: $C_n H_{2n-2}$). More reactive than saturated hydrocarbons.

  • Saturated/Unsaturated Hydrocarbons Table:

    • Methane ($CH_4$)

    • Ethane ($C_2H_6$)

    • Propane ($C_3H_8$)

    • Butane ($C_4H_{10}$)

    • Pentane ($C_5H_{12}$)

    • Hexane ($C_6H_{14}$)

  • Structural Isomers: Compounds with identical molecular formula but different structural formulas (e.g., $n$-butane and isobutane for $C_4H_{10}$).

  • Cyclic Hydrocarbons: Saturated (e.g., Cyclohexane $C_6H_{12}$) or Unsaturated (e.g., Benzene $C_6H_6$ containing alternate double bonds).

• Functional Groups and Homologous Series:

  • Heteroatoms: Elements like $Cl, Br, O, N, S$ that substitute hydrogen in hydrocarbon chains.

  • Functional Group: An atom or group of atoms that confers specific chemical properties to the carbon compound regardless of length/nature of carbon chain.

    • Halo (Chloro/Bromo): $-Cl, -Br$

    • Alcohol: $-OH$

    • Aldehyde: $-CHO$

    • Ketone: >C=O

    • Carboxylic Acid: $-COOH$

  • Homologous Series: A series of compounds in which the same functional group substitutes for hydrogen in a carbon chain.

    • Successive members differ by a $-CH_2-$ unit and a molecular mass of $14\,u$

    • Examples: Alkanes ($CH_4, C_2H_6, C_3H_8$), Alcohols ($CH_3OH, C_2H_5OH, C_3H_7OH$).

    • Show similar chemical properties, but gradation in physical properties (melting point, boiling point, density increase with molecular mass).

• Nomenclature of Carbon Compounds (IUPAC):

  • Identify number of carbon atoms (1 = Meth-, 2 = Eth-, 3 = Prop-, 4 = But-, 5 = Pent-, 6 = Hex-).

  • Identify presence of functional group (Suffix: -ane, -ene, -yne, -ol, -al, -one, -oic acid; Prefix: chloro-, bromo-).

  • If functional group suffix begins with a vowel ($a, e, i, o, u$), drop final $-e$ from alkane name (e.g., Propane - $e$ + ol = Propanol).

• Chemical Properties of Carbon Compounds:

  1. Combustion: Carbon and its compounds burn in oxygen to form $CO_2$, water, heat, and light:         $C + O_2 \rightarrow CO_2 + \text{Heat} + \text{Light}$         $CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O + \text{Heat} + \text{Light}$         $CH_3CH_2OH + 3O_2 \rightarrow 2CO_2 + 3H_2O + \text{Heat} + \text{Light}$

     • Saturated hydrocarbons burn with a clean blue flame. Unsaturated hydrocarbons burn with a yellow smoky flame (soot).

     • Incomplete combustion of saturated hydrocarbons (due to limited air supply) also produces soot.

  2. Oxidation: Alcohols are converted to carboxylic acids by strong oxidizing agents like alkaline potassium permanganate ($KMnO_4$) or acidified potassium dichromate ($K_2Cr_2O_7$):         $CH_3CH_2OH \xrightarrow{\text{Alkaline } KMnO_4 + \text{Heat} \text{ or Acidified } K_2Cr_2O_7 + \text{Heat}} CH_3COOH$

  3. Addition Reaction: Unsaturated hydrocarbons add hydrogen in the presence of catalysts such as Nickel ($Ni$) or Palladium ($Pd$) to give saturated hydrocarbons:         $R_2C=CR_2 + H_2 \xrightarrow{Ni} R_2CH-CHR_2$

     • Industrial Application: Hydrogenation of vegetable oils (unsaturated fats, liquid, healthy) to form vegetable ghee (saturated fats, solid, unhealthy).

  4. Substitution Reaction: Saturated hydrocarbons react with chlorine in the presence of sunlight, replacing hydrogen atoms one by one:         $CH_4 + Cl_2 \xrightarrow{\text{Sunlight}} CH_3Cl + HCl$

• Important Carbon Compounds: Ethanol and Ethanoic Acid:

  • Ethanol ($CH_3CH_2OH$):

    • Liquid at room temperature, pleasant odor, burning taste, miscible in water in all proportions.

    • Good solvent; used in medicines like tincture iodine, cough syrups, tonics.

    • Reactions:

      1. Reaction with Sodium: Evolves hydrogen gas and forms sodium ethoxide:                 $2Na + 2CH_3CH_2OH \rightarrow 2CH_3CH_2ONa + H_2(g)$

      2. Dehydration: Heating ethanol at $443\,K$ ($170^\circ C$) with excess concentrated sulphuric acid ($H_2SO_4$) (dehydrating agent) converts it to ethene:                 $CH_3CH_2OH \xrightarrow{\text{Conc. } H_2SO_4, 443\,K} CH_2=CH_2 + H_2O$

    • Harmful Effects: Depresses central nervous system, impairs coordination, causes drowsiness. Pure ethanol (absolute alcohol) is lethal. Consumption of methanol ($CH_3OH$) causes liver damage and optic nerve damage leading to blindness/death.

    • Denatured Alcohol: Ethanol mixed with poisonous copper sulphate/pyridine to prevent industrial misuse.

  • Ethanoic Acid ($CH_3COOH$ / Acetic Acid):

    • $5-8\%$ solution of ethanoic acid in water is called vinegar (used as preservative).

    • Melting point of pure ethanoic acid is $290\,K$ ($17^\circ C$); freezes in winter forming glacial acetic acid.

    • Weak acid compared to mineral acids ($HCl$) because it dissociates partially.

    • Reactions:

      1. Esterification: Reacts with ethanol in presence of an acid catalyst to form sweet-smelling ester (ethyl ethanoate):                 $CH_3COOH + CH_3CH_2OH \xrightarrow{\text{Acid}} CH_3COOCH_2CH_3 + H_2O$                 (Esters are used in perfumes and flavoring agents).

      2. Saponification: Esters react with sodium hydroxide ($NaOH$) to give back alcohol and sodium salt of carboxylic acid (used in making soap):                 $CH_3COOCH_2CH_3 + NaOH \rightarrow CH_3COONa + CH_3CH_2OH$

      3. Reaction with Base: $CH_3COOH + NaOH \rightarrow CH_3COONa + H_2O$

      4. Reaction with Carbonates and Hydrogencarbonates: Forms sodium acetate, water, and brisk effervescence of $CO_2$:                 $2CH_3COOH + Na_2CO_3 \rightarrow 2CH_3COONa + H_2O + CO_2$                 $CH_3COOH + NaHCO_3 \rightarrow CH_3COONa + H_2O + CO_2$

• Soaps and Detergents:

  • Soap: Sodium or potassium salt of long-chain carboxylic acids (e.g., sodium stearate).

  • Structure of Soap Molecule:

    • Ionic Head ($-\text{COO}^- Na^+$): Hydrophilic (water-attracting), soluble in water.

    • Hydrocarbon Tail: Hydrophobic (water-repelling / oil-attracting), soluble in oil/grease/dirt.

  • Micelle Formation: In water, soap molecules align radially with hydrophobic tails directed inside towards dirt/oil droplet and hydrophilic heads facing outward towards water. Clusters of such orientation form micelles.

  • Cleansing Action: Micelles trap oily dirt in the center; stirring washes away the suspended dirt in water as an emulsion.

  • Hard Water & Scum: Hard water contains $Ca^{2+}$ and $Mg^{2+}$ salts. Soaps react with these ions to form an insoluble gummy white precipitate called scum, wasting soap.

  • Synthetic Detergents: Sodium salts of sulphonic acids or ammonium salts with chlorides/bromides. Do not form scum with hard water ($Ca^{2+}$ and $Mg^{2+}$ ions do not form insoluble precipitates with detergents). Used in shampoos and clothes cleaning powders.