diff --git "a/lancedb1/cbse_.lance/data/ba080219-f6f4-4264-958a-eebed06d25a6.lance" "b/lancedb1/cbse_.lance/data/ba080219-f6f4-4264-958a-eebed06d25a6.lance" deleted file mode 100644--- "a/lancedb1/cbse_.lance/data/ba080219-f6f4-4264-958a-eebed06d25a6.lance" +++ /dev/null @@ -1,4890 +0,0 @@ -chemistry-notes-CBSE CBSE CBSE CBSE class-10-chapter-1.txt -Chemical Reactions and Equations -Introduction to Chemical Reactions and Equations -Physical and chemical changes -Chemical change - one or more new substances with new physical and chemical properties -are formed. -Example: Fe(s)  +  CuSO 4(aq) →FeSO 4(aq) +Cu(s)  -       (Blue)                      (Green)        -Here, when copper sulphate reacts with iron, two new substances, i.e., ferrous sulphate and -copper are formed. -Physical change - change in colour or state occurs but no new substance is formed. -Example: Water changes to steam on boiling but no new substance is formed(Even though -steam and water look different when they are made to react with a piece of Na, they react -the same way and give the exact same products). This involves only change in state (liquid -to vapour).  -Observations that help determine a chemical reaction -A chemical reaction can be determined with the help of any of the following observations: -a) Evolution of a gas -b) Change in temperature -c) Formation of a precipitate -d) Change in colour -e) Change of state -Chemical reaction -Chemical reactions are chemical changes in which reactants transform into products by -making or breaking of bonds(or both) between different atoms. -Types of chemical reactionsTaking into consideration different factors, chemical reactions are grouped into multiple -categories. -Few examples are: -●Combination -●Decomposition -●Single Displacement -●Double displacement -●Redox -●Endothermic -●Exothermic -●Precipitation -●Neutralisation -Chemical Reactions and Equations I -Word equation -A  word equation is a chemical reaction expressed in words rather than chemical -formulas. It helps identify the reactants and products in a chemical reaction. -For example,  -Sodium + Chlorine → Sodium chloride -The above equation means: "Sodium reacts with chlorine to form sodium chloride."  -Symbols of elements and their valencies -A symbol is the chemical code for an element. Each element has one or two letter atomicchemistry-notes-CBSE CBSE CBSE class-10-chapter-1.txt -formulas. It helps identify the reactants and products in a chemical reaction. -For example,  -Sodium + Chlorine → Sodium chloride -The above equation means: "Sodium reacts with chlorine to form sodium chloride."  -Symbols of elements and their valencies -A symbol is the chemical code for an element. Each element has one or two letter atomic -symbol, which is the abbreviated form of its name. -Valency is the combining capacity of an element. It can be considered as the number of -electrons lost, gain or shared by an atom when it combines with another atom to form a -molecule. -Writing chemical equations -Representation of a chemical reaction in terms of symbols and chemical formulae of the -reactants and products is known as a chemical equation. -Zn(s) +dil.H2SO 4(aq) →ZnSO 4(aq) +H2(↑) - (Reactants)   (Products) -• For solids, the symbol is "(s)". -• For liquids, it is "(l)". -• For gases, it is "(g)".• For aqueous solutions, it is "(aq)". -• For gas produced in the reaction, it is represented by "(↑)". -• For precipitate formed in the reaction, it is represented by "(↓)". -Balancing of a Chemical Reaction -Conservation of mass -According to the law of conservation of mass, no atoms can be created or destroyed in a -chemical reaction, so the number of atoms for each element in the reactants side has to -balance the number of atoms that are present in the products side. -In other words, the total mass of the products formed in a chemical reaction is equal to the -total mass of the reactants participated in a chemical reaction. -Balanced chemical equation -The chemical equation in which the number of atoms of each element in the reactants side -is equal to that of the products side is called a balanced chemical equation. -Steps for balancing chemical equations -Hit and trial method: While balancing the equation, change the coefficients (the numbers in -front of the compound or molecule) so that the number of atoms of each element is same -on each side of the chemical equation.chemistry-notes-CBSE CBSE CBSE class-10-chapter-1.txt -is equal to that of the products side is called a balanced chemical equation. -Steps for balancing chemical equations -Hit and trial method: While balancing the equation, change the coefficients (the numbers in -front of the compound or molecule) so that the number of atoms of each element is same -on each side of the chemical equation.  -Short-cut technique for balancing a chemical equation -Example: -aCaCO 3+bH3PO 4→cCa 3(PO 4)2+dH2CO 3 -Set up a series of simultaneous equations, one for each element. -Ca: a=3c -C:   a=d -O:   3a+4b=8c+3d -H:   3b=2d -P:    b=2c -Let's set c=1 -Then a=3 and -d=a=3 -b=2c=2So a=3; b=2; c=1; d=3 -The balanced equation is -3CaCO 3+ 2H 3PO 4→Ca3(PO 4)2+ 3H 2CO 3 -Chemical Reactions and Equations II -Types of chemical reactions -Taking into consideration different factors, chemical reactions are grouped into multiple -categories. -Few examples are: -●Combination -●Decomposition -●Single Displacement -●Double displacement -●Redox -●Endothermic -●Exothermic -●Precipitation -●Neutralisation -Combination reaction -In a combination reaction, two elements or one element and one compound or two -compounds combine to give one single product. -H2+Cl2→ 2HCl -element + element → compound -2CO +O2→ 2CO 2 -compound + element → compound -NH 3+HCl →NH 4Cl -compound + compound → compound -Decomposition reaction -A single reactant decomposes on the application of heat or light or electricity to give two or -more products. -Types of decomposition reactions: -a. Decomposition reactions which require heat - thermolytic decomposition or thermolysis. -Thermal decomposition of HgO -b. Decomposition reactions which require light - photolytic decomposition or photolysis. -Photolytic decomposition of H2O2 -c. Decomposition reactions which require electricity - electrolytic decomposition or -electrolysis. -Electrolytic decomposition of H 2O -Displacement reaction -More reactive element displaces a less reactive element from its compound or solution.i)Zn(s) +CuSO 4(aq) →ZnSO 4(aq) +Cu(s)chemistry-notes-CBSE CBSE CBSE class-10-chapter-1.txt -Photolytic decomposition of H2O2 -c. Decomposition reactions which require electricity - electrolytic decomposition or -electrolysis. -Electrolytic decomposition of H 2O -Displacement reaction -More reactive element displaces a less reactive element from its compound or solution.i)Zn(s) +CuSO 4(aq) →ZnSO 4(aq) +Cu(s) -ii)Cu(s) + 2AgNO 3(aq) →Cu(NO 3)2(aq) + 2Ag (s) -Double displacement reaction -An exchange of ions between the reactants takes place to give new products. -For example, Al 2(SO4)3(aq) + 3Ca( OH)2(aq) → 2Al (OH)3(aq) + 3CaSO 4(s) -Precipitation reaction -An insoluble compound called precipitate forms when two solutions containing soluble salts -are combined.  -For example, Pb( NO 3)2(aq) + 2KI (aq) → 2KNO 3(aq) +PbI 2(↓)(s)(yellow ) -Redox reaction -Oxidation and reduction take place simultaneously. -Oxidation: Substance loses electrons or gains oxygen or loses hydrogen. -Reduction: Substance gains electrons or loses oxygen or gains hydrogen. -Oxidising agent - a substance that oxidises another substance and self-gets reduced. -Reducing agent - a substance that reduces another substance and self-gets oxidised. -Examples: -1.Fe(s) +CuSO 4(aq) →FeSO 4(aq) +Cu(s)       (Blue)                (Green) -Fe→Fe+2+ 2e −  (oxidation ) ; Fe - reducing agent. -Cu+2+ 2e − →Cu(s) (reduction ) ; Cu - oxidising agent. -2.ZnO +C→Zn+CO -ZnO reduces to Zn → reduction -C oxidises to CO → oxidation -ZnO - Oxidising agent -C - Reducing agent -Endothermic and exothermic reaction -Exothermic reaction - heat is evolved during a reaction. Most of the combination reactions -are exothermic. -Al+Fe2O3→Al2O3+Fe+heat -CH 4+ 2O 2→CO 2+ 2H 2O+heat -Endothermic - Heat is required to carry out the reaction. -6CO 2+ 6H 2O+Sunlight →C6H12O6+ 6O 2 -       Glucose -Most of the decomposition reactions are endothermic. -Corrosion -Gradual deterioration of a material, usually a metal, by the action of moisture, air or -chemicals in the surrounding environment. -Rusting:chemistry-notes-CBSE CBSE CBSE class-10-chapter-1.txt -are exothermic. -Al+Fe2O3→Al2O3+Fe+heat -CH 4+ 2O 2→CO 2+ 2H 2O+heat -Endothermic - Heat is required to carry out the reaction. -6CO 2+ 6H 2O+Sunlight →C6H12O6+ 6O 2 -       Glucose -Most of the decomposition reactions are endothermic. -Corrosion -Gradual deterioration of a material, usually a metal, by the action of moisture, air or -chemicals in the surrounding environment. -Rusting: -4Fe(s) + 3O 2(from  air) +xH 2O(moisture ) → 2Fe 2O3.xH 2O(rust) -Corrosion of copper: -Cu(s) +H2O(moisture ) +CO 2(from  air) →CuCO 3.Cu(OH)2(green ) -Corrosion of silver: -Ag(s) +H2S(from  air) →Ag2S(black) +H2(g) -Rancidity -It refers to oxidation of fats and oils in food that is kept for a long time. It gives foul smell -and bad taste to food. Rancid food causes stomach infection on consumption. -Prevention: -(i) Use of air-tight containers(ii) Packaging with nitrogen -(iii) Refrigeration -(iv) Addition of antioxidants or preservativescbse-CBSE CBSE CBSE class-10-science-notes-chapter-2-acids-bases-and-salts.txt -Acids, Bases and Salts -Introduction to Acids, Bases and Salts -Classification of matter -On the basis of -a) composition -  elements, compounds and mixtures -b) state - solids, liquids and gases -c) solubility - suspensions, colloids and solutions -Types of mixtures - homogeneous and heterogeneous -Types of compounds - covalent and ionic -What Is an Acid and a Base? -Ionisable and non-ionisable compounds -An ionisable compound when dissolved in water or in its molten state, dissociates into ions -almost entirely. Example: NaCl, HCl, KOH, etc. -A non-ionisable compound does not dissociate into ions when dissolved in water or in its -molten state. Example: glucose, acetone, etc. -Arrhenius theory of acids and bases -Arrhenius acid - when dissolved in water, dissociates to give H+(aq) or H 3O+ ion. -Arrhenius base - when dissolved in water, dissociates to give OH− ion. -Examples Acids  -Hydrochloric acid (HCl ) -Sulphuric acid  (H 2SO 4) -Nitric acid (HNO 3) -Bases  -Sodium hydroxide (NaOH) -Potassium hydroxide (KOH) -Calcium hydroxide (Ca(OH)2) -Bronsted Lowry theoryA Bronsted acid is a H+(aq) ion donor. -A Bronsted base is a H+(aq) ion acceptor. -Example -In the reaction:  HCl(aq) +NH 3(aq) →NH+ -4(aq) +Cl−(aq) -HCl - Bronsted acid and Cl− - its conjugate acid -NH 3  - Bronsted base and NH+ -4 - its conjugate acid -Physical test -Given are two possible physical tests to identify an acid or a base. -a. Taste -An acid tastes sour whereas a base tastes bitter. -The method of taste is not advised as an acid or a base could be contaminated or corrosive. -b. Effect on indicators by acids and bases -An indicator is a chemical substance which shows a change in its physical properties, -mainly colour or odour when brought in contact with an acid or a base. -Below mentioned are commonly used indicators and the different colours they exhibit:  -a) Litmus -In neutral solution - purple -In acidic solution - red -In basic solution - bluecbse-CBSE CBSE CBSE class-10-science-notes-chapter-2-acids-bases-and-salts.txt -b. Effect on indicators by acids and bases -An indicator is a chemical substance which shows a change in its physical properties, -mainly colour or odour when brought in contact with an acid or a base. -Below mentioned are commonly used indicators and the different colours they exhibit:  -a) Litmus -In neutral solution - purple -In acidic solution - red -In basic solution - blue -Litmus is also available as strips of paper in two variants - red litmus and blue litmus. -An acid turns a moist blue litmus paper to red. -A base turns a moist red litmus paper to blue. -b) Methyl orange -In neutral solution - orange -In acidic solution - red -In basic solution - yellow -c) Phenolphthalein -In neutral solution - colourless -In acidic solution - remains colourless -In basic solution - pink -Acid Base ReactionsReactions of acids and bases -a) Reaction of acids and bases with metals -Acid + active metal →  salt + hydrogen + heat -2HCl   +Mg→MgCl 2+H2(↑) -Base + metal → salt + hydrogen + heat -2NaOH   +Zn→Na 2ZnO 2+H2(↑) -A more reactive metal displaces the less reactive metal from its base. -2Na +Mg(OH)2→ 2NaOH +Mg -b) Reaction of acids with metal carbonates and bicarbonates -Acid + metal carbonate or bicarbonate  →  salt + water + carbon dioxide. -2HCl   +  CaCO 3→CaCl 2  +  H 2O  +  CO 2 -H2SO 4  +  Mg (HCO 3)2→MgSO 4  +  2H 2O  +  2CO 2 -Effervescence indicates liberation of  CO 2 gas. -c) Neutralisation reaction -1. Reaction of metal oxides and hydroxides with acids -Metal oxides or metal hydroxides are basic in nature. -Acid + base → salt + water + heat -H2SO 4  +  MgO →MgSO 4  +  H 2O -2HCl +Mg(OH)2→MgCl 2+ 2H 2O -2. Reaction of non-metal oxides with bases -Non-metal oxides are acidic in nature -Base + Non-metal oxide  →  salt + water + heat -2NaOH +CO 2→Na 2CO 3+H2O -Water -Acids and bases in water -When added to water, acids and bases dissociate into their respective ions and help in -conducting electricity. -Difference between a base and an alkali -Base-Bases undergo neutralisation reaction with acids.cbse-CBSE CBSE CBSE class-10-science-notes-chapter-2-acids-bases-and-salts.txt -2. Reaction of non-metal oxides with bases -Non-metal oxides are acidic in nature -Base + Non-metal oxide  →  salt + water + heat -2NaOH +CO 2→Na 2CO 3+H2O -Water -Acids and bases in water -When added to water, acids and bases dissociate into their respective ions and help in -conducting electricity. -Difference between a base and an alkali -Base-Bases undergo neutralisation reaction with acids. -They are comprised of metal oxides, metal hydroxides, metal carbonates and metal -bicarbonates. -Most of them are insoluble in water.  -Alkali -  -An alkali is an aqueous solution of a base, (mainly metallic hydroxides). -It dissolves in water and dissociates to give   OH− ion. -All alkalis are bases, but not all bases are alkalis. -Hydronium ion -Hydronium ion is formed when a hydrogen ion accepts a lone pair of electrons from the -oxygen atom of a water molecule, forming a coordinate covalent bond. -Formation of a hydronium ion -Dilution -Dilution is the process of reducing the concentration of a solution by adding more solvent -(usually water) to it. -It is a highly exothermic process. -To dilute an acid, the acid must be added to water and not the other way round. -Strength of acids and basesStrong acid or base : When all molecules of given amount of an acid or a base dissociate -completely in water to furnish their respective ions, H+(aq)  for acid and OH−(aq) for base). -Weak acid or base: When only a few of the molecules of given amount of an acid or a base -dissociate in water to furnish their respective ions, H+(aq) for acid and OH−(aq) for base).  -Dilute acid: contains less number of H+(aq) ions per unit volume. -Concentrated acid: contains more number of H+(aq) ions per unit volume.Universal indicator -A universal indicator has pH range from 0 to 14 that indicates the acidity or alkalinity of a -solution. -A neutral solution has pH=7 -pH -       pH= −log 10[H+] -In pure water,  [H+] = [OH−] = 10−7 mol/L. Hence, the pH of pure water is 7. -The pH scale ranges from 0 to 14.cbse-CBSE CBSE CBSE class-10-science-notes-chapter-2-acids-bases-and-salts.txt -Concentrated acid: contains more number of H+(aq) ions per unit volume.Universal indicator -A universal indicator has pH range from 0 to 14 that indicates the acidity or alkalinity of a -solution. -A neutral solution has pH=7 -pH -       pH= −log 10[H+] -In pure water,  [H+] = [OH−] = 10−7 mol/L. Hence, the pH of pure water is 7. -The pH scale ranges from 0 to 14. -If pH < 7 - acidic solution -If pH > 7-  basic solution -pH scale -Importance of pH in everyday life1. pH sensitivity of plants and animals -Plants and animals are sensitive to pH. Crucial life processes such as digestion of food, -functions of enzymes and hormones happen at a certain pH value. -2. pH of a soil -The pH of a soil optimal for the growth of plants or crops is 6.5 to 7.0. -3. pH in the digestive system -The process of digestion happens at a specific pH in our stomach which is 1.5 - 4. -The pH of the interaction of enzymes, while food is being digested, is influenced by HCl in -our stomach.  -4. pH in tooth decay -Tooth decay happens when the teeth are exposed to an acidic environment of pH -5.5 and below.   -5. pH of self-defense by animals and plants -Acidic substances are used by animals and plants as a self-defense mechanism. For example, -bee and plants like nettle secrete a highly acidic substance for self-defense. These secreted -acidic substances have a specific pH. -Manufacture of Acids and Bases -Manufacture of acids and bases -a) Non-metal oxide + water → acid -SO2(g) +H2O(l) →H2SO3(aq) -SO3(g) +H2O(l) →H2SO4(aq) -4NO 2(g) + 2H 2O(l) +O2(g) → 4HNO 3(aq) -Non-metal oxides are thus referred to as acid anhydrides. -b) Hydrogen + halogen → acid -H2(g) +Cl2(g) → 2HCl (g) -HCl(g) +H2O(l) →HCl(aq) -c) Metallic salt + conc. sulphuric acid → salt + more volatile acid -2NaCl (aq) +H2SO4(aq) →Na2SO4(aq) + 2HCl (aq) -2KNO 3(aq) +H2SO4(aq) →K2SO4(aq) + 2HNO 3(aq)d) Metal + oxygen → metallic oxide (base) -4Na( s) +O2(g) → 2Na 2O(s) -2Mg (s) +O2(g) → 2MgO (s) -e) Metal + water → base or alkali + hydrogen -Zn(s) + H2O(steam ) → ZnO(s)+ H 2(g)cbse-CBSE CBSE CBSE class-10-science-notes-chapter-2-acids-bases-and-salts.txt -H2(g) +Cl2(g) → 2HCl (g) -HCl(g) +H2O(l) →HCl(aq) -c) Metallic salt + conc. sulphuric acid → salt + more volatile acid -2NaCl (aq) +H2SO4(aq) →Na2SO4(aq) + 2HCl (aq) -2KNO 3(aq) +H2SO4(aq) →K2SO4(aq) + 2HNO 3(aq)d) Metal + oxygen → metallic oxide (base) -4Na( s) +O2(g) → 2Na 2O(s) -2Mg (s) +O2(g) → 2MgO (s) -e) Metal + water → base or alkali + hydrogen -Zn(s) + H2O(steam ) → ZnO(s)+ H 2(g) -f) Few metallic oxides + water → alkali -Na 2O(s) +H2O(l) → 2NaOH (aq) -g) Ammonia + water → ammonium hydroxide -NH 3(g) +H2O(l) →NH 4OH(aq) -Salts -Salts -A salt is a combination of an anion of an acid and a cation of a base. -Examples - KCl ,NaNO 3,CaSO 4,etc. -Salts are usually prepared by neutralisation reaction of an acid and a base. -Common salt -Sodium Chloride (NaCl) is referred to as common salt because it’s used all over the world for -cooking. -Family of salts -Salts having the same cation or anion belong to the same family. For example, NaCl, KCl, -LiCl. -pH of salts -A salt of a strong acid and a strong base will be neutral in nature. pH = 7 (approx.). -A salt of a weak acid and a strong base will be basic in nature. pH > 7. -A salt of a strong acid and a weak base will be acidic in nature. pH < 7. -The pH of a salt of a weak acid and a weak base is determined by conducting a pH test. -Preparation of Sodium hydroxide  -Chemical formula - NaOH -Also known as - caustic sodaPreparation (Chlor-alkali process): -Electrolysis of brine (solution of common salt, NaCl) is carried out.At anode: Cl 2 is released -At cathode: H2 is released -Sodium hydroxide remains in the solution. -Bleaching powder -Chemical formula - Ca(OCl)Cl or CaOCl 2 -Preparation - Ca( OH)2(aq) +Cl2(g) →CaOCl 2(aq) +H2O(l) -On interaction with water - bleaching powder releases chlorine which is responsible for -bleaching action. -Baking soda -Chemical name - Sodium hydrogen carbonate -Chemical formula - NaHCO 3 -Preparation (Solvay process) -  -a. Limestone is heated:  CaCO 3→CaO +CO 2cbse-CBSE CBSE CBSE class-10-science-notes-chapter-2-acids-bases-and-salts.txt -Bleaching powder -Chemical formula - Ca(OCl)Cl or CaOCl 2 -Preparation - Ca( OH)2(aq) +Cl2(g) →CaOCl 2(aq) +H2O(l) -On interaction with water - bleaching powder releases chlorine which is responsible for -bleaching action. -Baking soda -Chemical name - Sodium hydrogen carbonate -Chemical formula - NaHCO 3 -Preparation (Solvay process) -  -a. Limestone is heated:  CaCO 3→CaO +CO 2 -b. CO_2 is passed through a concentrated solution of sodium chloride and ammonia : -NaCl (aq) +NH 3(g) +CO 2(g) +H2O(l) →NaHCO 3(aq) +NH 4Cl(aq) -Uses: -1. Textile industry -2. Paper industry -3. Disinfectant -Washing soda -Chemical name  - Sodium carbonate decahydrate. -Chemical formuala - \(Na_2CO_3 \) -Preparation: By heating NaHCO 3 -2NaHCO 3(s) →Na 2CO 3(s) +CO 2(g) +H2O(g) -Na 2CO 3(s)  +  10H 2O(l)  →  Na 2CO 3.10H 2O(s)Uses -1. In glass, soap and paper industries -2. Softening of water -3. Domestic cleaner -Crystals of salts -Certain salts form crystals by combining with a definite proportion of water. The water that -combines with the salt is called water of crystallisation. -Plaster of parisGypsum ,  CaSO 4.2H 2O (s) on  heating  at 100°C  (373K ) gives  CaSO 4.H2O and  H2O -CaSO 4.H2O is plaster of paris. -CaSO 4.H2O means two formula units of CaSO 4 share one molecule of water. -Uses - cast for healing fractures.1 -23 -2 -1 -21 -2cbse-CBSE CBSE CBSE class-10-science-notes-chapter-3-metals-and-non-metals.txt -Corrosion -Alloys -Alloys are homogeneous mixtures of metal with other metals or nonmetals. Alloy formation -enhances the desirable properties of the material, such as hardness, tensile strength and -resistance to corrosion. -Examples of few alloys - -Brass: copper and zinc -Bronze: copper and tin -Solder: lead and tin -Amalgam: mercury and other metal -Corrosion -Gradual deterioration of a material usually a metal by the action of moisture, air or -chemicals in the surrounding environment. -Rusting: -4Fe(s) + 3O 2(from  air) + xH 2O(moisture ) → 2Fe 2O3.xH 2O(rust) -Corrosion of copper: -Cu(s) +H2O(moisture ) +CO 2(from  air) → CuCO 3.Cu(OH)2(green ) -Corrosion of silver: -Ag(s) +H2S(from  air) → Ag2S(black) + H2(g) -Prevention of CorrosionPrevention :  -1. Coating with paints or oil or grease: Application of paint or oil or grease on metal surfaces -keep out air and moisture. -2. Alloying: Alloyed metal is more resistant to corrosion. Example: stainless steel. -3. Galvanization: This is a process of coating molten zinc on iron articles. Zinc forms a -protective layer and prevents corrosion. -4. Electroplating: It is a method of coating one metal with another by use of electric current. -This method not only lends protection but also enhances the metallic appearance. -Example: silver plating, nickel plating. -5. Sacrificial protection: Magnesium is more reactive than iron. When it is coated on the -articles made of iron or steel, it acts as the cathode, undergoes reaction (sacrifice) instead -of iron and protects the articles.Metals and Non-metalsPhysical Properties -Physical Properties of Metals -●Hard and have a high tensile strength -●Solids at room temperature -●Sonorous -●Good conductors of heat and electricity -●Malleable, i.e., can be beaten into thin sheets -●Ductile, i.e., can be drawn into thin wires -●High melting and boiling points (except Caesium (Cs) and Gallium (Ga)) -●Dense, (except alkali metals). Osmium - highest density and lithium - least density -●Lustrouscbse-CBSE CBSE CBSE class-10-science-notes-chapter-3-metals-and-non-metals.txt -●Hard and have a high tensile strength -●Solids at room temperature -●Sonorous -●Good conductors of heat and electricity -●Malleable, i.e., can be beaten into thin sheets -●Ductile, i.e., can be drawn into thin wires -●High melting and boiling points (except Caesium (Cs) and Gallium (Ga)) -●Dense, (except alkali metals). Osmium - highest density and lithium - least density -●Lustrous -●Silver-grey in colour, (except gold and copper) -Non-Metals -Nonmetals are those elements which do not exhibit the properties of metals. -Physical Properties of Nonmetals -Occur as solids, liquids and gases at room temperature -Brittle -Non-malleable -Non-ductile -Non-sonorous -Bad conductors of heat and electricity -Exceptions in Physical Properties -Alkali metals (Na, K, Li) can be cut using a knife. -Mercury is a liquid metal. -Lead and mercury are poor conductors of heat. -Mercury expands significantly for the slightest change in temperature. -Gallium and caesium have a very low melting point -Iodine is non-metal but it has lustre. -Graphite conducts electricity. -Diamond conducts heat and has a very high melting point. -Chemical Properties -Chemical Properties of Metals●Alkali metals (Li, Na, K, etc) react vigorously with water and oxygen or air. -●Mg reacts with hot water. -●Al, Fe and Zn react with steam. -●Cu, Ag, Pt, Au do not react with water or dilute acids. -Reaction of Metals with Oxygen (Burnt in Air) -Metal + Oxygen  →  Metal oxide (basic) -●Na and K are kept immersed in kerosene oil as they react vigorously with air and catch -fire. - 4K(s) +O2(g) → 2K 2O(s) (vigorous reaction) -●Mg, Al, Zn, Pb react slowly with air and form a protective layer that prevents corrosion. -2Mg (s) +O2(g) → 2MgO (s) (Mg burns with a white dazzling light) -4Al(s) + 3O 2(g) → 2Al 2O3(s) -●Silver, platinum and gold don't burn or react with air. -Basic Oxides of Metals -Some metallic oxides get dissolved in water and form alkalis. Their aqueous solution turns -red litmus blue. -Na 2O(s) +H2O(l) → 2NaOH (aq) -K2O(s) +H2O(l) → 2KOH (aq)cbse-CBSE CBSE CBSE class-10-science-notes-chapter-3-metals-and-non-metals.txt -2Mg (s) +O2(g) → 2MgO (s) (Mg burns with a white dazzling light) -4Al(s) + 3O 2(g) → 2Al 2O3(s) -●Silver, platinum and gold don't burn or react with air. -Basic Oxides of Metals -Some metallic oxides get dissolved in water and form alkalis. Their aqueous solution turns -red litmus blue. -Na 2O(s) +H2O(l) → 2NaOH (aq) -K2O(s) +H2O(l) → 2KOH (aq) -Amphoteric Oxides of Metals -Amphoteric oxides are metal oxides which react with both acids as well as bases to form -salt and water. -For example - Al 2O3,ZnO ,PbO ,SnO -Al2O3(s) + 6HCl (aq) → 2AlCl 3(aq) + 3H 2O(l) -Al2O3(s) + 2NaOH (aq) → 2NaAlO 2(aq) +H2O(l) -ZnO (s) + 2HCl (aq) →ZnCl 2(aq) +H2O(l) -ZnO (s) + 2NaOH (aq) →Na 2ZnO 2(aq) +H2O(l) -Reactivity Series -The below table illustrates the reactivity of metals from high order to low order. -Symbol               ElementK Potassium ( Highly Active Metal) -Ba Barium -Ca Calcium -Na Sodium -Mg Magnesium -Al Aluminium -Zn Zinc -Fe Iron -Ni Nickel -Sn Tin -Pb Lead -H Hydrogen -Cu Copper -Hg Mercury -Ag Silver -Au Gold -Pt Platinum -Reaction of Metals with Water or Steam -Metal +Water →Metal  hydroxide  or Metal  oxide +Hydrogen -2Na + 2H 2O(cold) → 2NaOH +H2+heat -Ca+ 2H 2O(cold) →Ca(OH)2+H2 -Mg+ 2H 2O(hot) →Mg(OH)2+H2 -2Al+ 3H 2O(steam ) →Al2O3+ 3H 2 -Zn+H2O(steam ) →ZnO +H2 -3Fe + 4H 2O(steam ) →Fe3O4+ 4H 2 -Reaction of Metals with Acid -Metal +dilute  acid →Salt +Hydrogen  gas -2Na(s) + 2HCl (dilute ) → 2NaCl (aq) +H2(g) -2K(s) +H2SO 4(dilute ) →K2SO 4(aq) +H2(g) -Only Mg and Mn, react with very dilute nitric acid to liberate hydrogen gas.  -Mg(s) + 2HNO 3(dilute ) →Mg(NO 3)2(aq) +H2(g) -Mn(s) + 2HNO 3(dilute ) →Mn(NO 3)2(aq) +H2(g) -Displacement ReactionA more reactive element displaces a less reactive element from its compound or solution. -How Do Metal React with Solution of Other Metal Salts -Metal  A+Salt of  metal  B→Salt of  metal  A+Metal  B -Fe(s) +CuSO 4(aq) →FeSO 4(aq) +Cu(s) -Cu(s) + 2AgNO 3(aq) →Cu(NO 3)(aq) + 2Ag (s) -Reaction of Metals with Bases -Base +metal →salt+hydrogencbse-CBSE CBSE CBSE class-10-science-notes-chapter-3-metals-and-non-metals.txt -Displacement ReactionA more reactive element displaces a less reactive element from its compound or solution. -How Do Metal React with Solution of Other Metal Salts -Metal  A+Salt of  metal  B→Salt of  metal  A+Metal  B -Fe(s) +CuSO 4(aq) →FeSO 4(aq) +Cu(s) -Cu(s) + 2AgNO 3(aq) →Cu(NO 3)(aq) + 2Ag (s) -Reaction of Metals with Bases -Base +metal →salt+hydrogen -2NaOH (aq) +Zn(s) →Na 2ZnO 2(aq) +H2(g) -2NaOH (aq) + 2Al (s) + 2H 2O(l) → 2NaAlO 2(aq) + 2H 2(g) -Extraction of Metals and Non-Metals -Applications of Displacement Reaction -Uses of displacement reaction -1. Extraction of metals -2. Manufacturing of steel -3. Thermite reaction: Al( s) +Fe2O3(s) →Al2O3+Fe(molten ) -The thermite reaction is used in welding of railway tracks, cracked machine parts, etc. -Occurrence of Metals -Most of the elements especially metals occur in nature in the combined state with other -elements. All these compounds of metals are known as  minerals. But out of them, only a few -are viable sources of that metal. Such sources are called ores. -Au, Pt - exist in the native or free state. -Extraction of MetalsMetals of high reactivity - Na, K, Mg, Al. -Metals of medium reactivity - Fe, Zn, Pb, Sn. -Metals of low reactivity - Cu, Ag, Hg -Roasting -Converts sulphide ores into oxides on heating strongly in the presence of excess air. -It also removes volatile impurities. -2ZnS (s) + 3O 2(g) +Heat → 2ZnO (s) + 2SO 2(g) -Calcination -Converts carbonate and hydrated ores into oxides on heating strongly in the presence of -limited air. It also removes volatile impurities. -ZnCO 3(s) +heat →ZnO (s) +CO 2(g) -CaCO 3(s) +heat →CaO (s) +CO 2(g) -Al2O3.2H 2O(s) +heat → 2Al 2O3(s) + 2H 2O(l) -2Fe 2O3.3H 2O(s) +heat → 2Fe 2O3(s) + 3H 2O(l) -Extracting Metals Low in Reactivity SeriesBy self-reduction- when the sulphide ores of less electropositive metals like Hg, Pb, Cu etc., -are heated in air, a part of the ore gets converted to oxide which then reacts with thecbse-CBSE CBSE CBSE class-10-science-notes-chapter-3-metals-and-non-metals.txt -ZnCO 3(s) +heat →ZnO (s) +CO 2(g) -CaCO 3(s) +heat →CaO (s) +CO 2(g) -Al2O3.2H 2O(s) +heat → 2Al 2O3(s) + 2H 2O(l) -2Fe 2O3.3H 2O(s) +heat → 2Fe 2O3(s) + 3H 2O(l) -Extracting Metals Low in Reactivity SeriesBy self-reduction- when the sulphide ores of less electropositive metals like Hg, Pb, Cu etc., -are heated in air, a part of the ore gets converted to oxide which then reacts with the -remaining sulphide ore to give the crude metal and sulphur dioxide. In this process, no -external reducing agent is used. -1.2HgS (Cinnabar) + 3O 2(g) +heat → 2HgO (crude  metal ) + 2SO 2(g) -2HgO (s) +heat → 2Hg (l) +O2(g) -2.Cu2S(Copper pyrite ) + 3O 2(g) +heat → 2Cu 2O(s) + 2SO 2(g) -2Cu 2O(s) +Cu2S(s) +heat → 6Cu(crude  metal ) +SO 2(g) -3.2PbS (Galena) + 3O 2(g) +heat → 2PbO (s) + 2SO 2(g) -PbS(s) + 2PbO (s) → 2Pb (crude  metal ) +SO 2(g) -Extracting Metals in the Middle of Reactivity Series -Smelting - it involves heating the roasted or calcined ore(metal oxide) to a high temperature -with a suitable reducing agent. The crude metal is obtained in its molten state. -Fe2O3+ 3C (coke) → 2Fe + 3CO 2 -Aluminothermic reaction - also known as the Goldschmidt reaction is a highly exothermic -reaction in which metal oxides usually of Fe and Cr are heated to a high temperature with -aluminium. -Fe2O3+ 2Al →Al2O3+ 2Fe +heat -Cr2O3+ 2Al →Al2O3+ 2Cr +heat -Extraction of Metals Towards the Top of the Reactivity Series -Electrolytic reduction: -1. Down’s process:  Molten NaCl is electrolysed in a special apparatus. -At the cathode (reduction) -  -Na+(molten ) +e−→Na(s) -Metal is deposited. -At the anode (oxidation) - -2Cl−(molten ) →Cl2(g) + 2e– -Chlorine gas is liberated. -2. Hall’s process: Mixture of molten alumina and a fluoride solvent usually cryolite, -(Na 3AlF 6) is electrolysed. -At the cathode (reduction) - -2Al3++ 6e–→ 2Al (s)Metal is deposited. -At the anode (oxidation) - -6O2–→ 3O 2(g) + 12e–  -Oxygen gas is liberated. -Enrichment of Orescbse-CBSE CBSE CBSE class-10-science-notes-chapter-3-metals-and-non-metals.txt -Na+(molten ) +e−→Na(s) -Metal is deposited. -At the anode (oxidation) - -2Cl−(molten ) →Cl2(g) + 2e– -Chlorine gas is liberated. -2. Hall’s process: Mixture of molten alumina and a fluoride solvent usually cryolite, -(Na 3AlF 6) is electrolysed. -At the cathode (reduction) - -2Al3++ 6e–→ 2Al (s)Metal is deposited. -At the anode (oxidation) - -6O2–→ 3O 2(g) + 12e–  -Oxygen gas is liberated. -Enrichment of Ores -It means removal of impurities or gangue from ore, through various physical and chemical -processes. The technique used for a particular ore depends on the difference in the -properties of the ore and the gangue. -Refining of Metals -Refining of metals - removing impurities or gangue from crude metal. It is the last step in -metallurgy and is based on the difference between the properties of metal and the gangue. -Electrolytic Refining -Metals like copper, zinc, nickel, silver, tin, gold etc., are refined electrolytically. -Anode – impure or crude metal -Cathode – thin strip of pure metal -Electrolyte – aqueous solution of metal salt -From anode  (oxidation) - metal ions are released into the solution -At cathode (reduction) - equivalent amount of metal from solution is deposited -Impurities deposit at the bottom of the anode. -The Why Questions -Electronic configuration - Group 1 elements - Alkali metals -Element Electronic  configuration -Lithium (Li) 2, 1 -Sodium (Na) 2, 8, 1 -Potassium (K) 2, 8, 8, 1 -Rubidium (Rb) 2, 8, 18, 8, 1 Group 2 elements - Alkaline earth metals -Element Electronic  configuration -Beryllium (Be) 2, 2 -Magnesium (Mg ) 2, 8, 2 -Calcium (Ca) 2, 8, 8, 2 -Stronium (Sr) 2, 8, 18, 8, 2 -How Do Metals and Nonmetals React -Metals lose valence electron(s) and form cations. -Non-metals gain those electrons in their valence shell and form anions. -The cation and the anion are attracted to each other by strong electrostatic force, thus -forming an ionic bond.  -For example: In Calcium chloride, the ionic bond is formed by oppositely charged -calcium and chloride ions.cbse-CBSE CBSE CBSE class-10-science-notes-chapter-3-metals-and-non-metals.txt -How Do Metals and Nonmetals React -Metals lose valence electron(s) and form cations. -Non-metals gain those electrons in their valence shell and form anions. -The cation and the anion are attracted to each other by strong electrostatic force, thus -forming an ionic bond.  -For example: In Calcium chloride, the ionic bond is formed by oppositely charged -calcium and chloride ions. -Calcium atom loses 2 electrons and attains the electronic configuration of the nearest noble -gas (Ar). By doing so, it gains a net charge of +2. -The two Chlorine atoms take one electron each, thus gaining a charge of -1 (each) and attain -the electronic configuration of the nearest noble gas (Ar). -Ionic CompoundsThe electrostatic attractions between the oppositely charged ions hold the compound -together. -Example: MgCl 2,CaO ,MgO ,NaCl ,etc . -Properties of Ionic Compound -Ionic compounds -1. Are usually crystalline solids (made of ions). -2. Have high melting and boiling points. -3. Conduct electricity when in aqueous solution and when melted. -4. Are mostly soluble in water and polar solvents. -Physical Nature -Ionic solids usually exist in a regular, well-defined crystal structures. -Electric Conduction of Ionic Compounds -Ionic compounds conduct electricity in the molten or aqueous state when ions become free -and act as charge carriers. -In solid form, ions are strongly held by electrostatic forces of attractions and not free to -move; hence do not conduct electricity. -For example, ionic compounds such as NaCl does not conduct electricity when solidconduct electricity but when dissolved in water or in molten state, it will -conduct  electricity. -Salt solution conducts electricity -Melting and Boiling Points of Ionic Compounds -In ionic compounds, the strong electrostatic forces between ions require a high amount of -energy to break. Thus, the melting point and boiling point of an ionic compound are usually -very high. -Solubility of Ionic Compoundscbse-CBSE CBSE CBSE class-10-science-notes-chapter-3-metals-and-non-metals.txt -conduct  electricity. -Salt solution conducts electricity -Melting and Boiling Points of Ionic Compounds -In ionic compounds, the strong electrostatic forces between ions require a high amount of -energy to break. Thus, the melting point and boiling point of an ionic compound are usually -very high. -Solubility of Ionic Compounds -Most ionic compounds are soluble in water due to the separation of ions by water. This -occurs due to the polar nature of water. For example, NaCl is a 3-D salt crystal composed of Na+ and Cl− ions bound together -through electrostatic forces of attractions. When a crystal of NaCl comes into contact with -water, the partial positively charged ends of water molecules interact with the Cl− ions, -while the negatively charged end of the water molecules interacts with the Na+ ions. -This ion-dipole interaction between ions and water molecules assist in the breaking of the -strong electrostatic forces of attractions within the crystal and ultimately in the solubility of -the crystal.cbse-CBSE CBSE CBSE class-10-science-notes-chapter-4-carbon-and-its-compounds.txt -Carbon and Its Compounds -Soaps and Detergents:- -Cleansing Action of Soap -When soap is added to water, the soap molecules uniquely orient themselves to form -spherical shape micelles. -The non-polar hydrophobic part or tail of the soap molecules attracts the dirt or oil part of -the fabric, while the polar hydrophilic part or head, (−COO−Na+, remains attracted to water -molecules. -The agitation or scrubbing of the fabric helps the micelles to carry the oil or dirt particles -and detach them from the fibres of the fabric.Hard Water -Hard water contains salts of calcium and magnesium, principally as bicarbonates, chlorides, -and sulphates. When soap is added to hard water, calcium and magnesium ions of hard -water react with soap forming insoluble curdy white precipitates of calcium and magnesium -salts of fatty acids. -2C17H35COONa +MgCl 2→ (C17H35COO )2Mg+ 2NaCl -2C17H35COONa +CaCl 2→ (C17H35COO )2Ca+ 2NaCl -These precipitates stick to the fabric being washed and hence, interfere with the cleaning -ability of the soap. Therefore, a lot of soap is wasted if water is hard. -Covalent Bonds -Difficulty of Carbon to Form a Stable Ion -To achieve the electronic configuration of nearest noble gas, He, if the carbon atom loses -four of its valence electrons, a huge amount of energy is involved. C4+ ion hence formed will -be highly unstable due to the presence of six protons and two electrons. -If the carbon atom gains four electrons to achieve the nearest electronic configuration of -the noble gas, Ne, C4− ion will be formed. But again, a huge amount of energy is required. -Moreover, in C4+ ion it is difficult for 6 protons to hold 10 electrons. Hence, to satisfy its -tetravalency, carbon shares all four of its valence electrons and forms covalent bonds. -Ionic BondIonic bonding involves the transfer of valence electron/s, primarily between a metal and a -nonmetal. The electrostatic attractions between the oppositely charged ions hold the -compound together. -Ionic compounds:cbse-CBSE CBSE CBSE class-10-science-notes-chapter-4-carbon-and-its-compounds.txt -tetravalency, carbon shares all four of its valence electrons and forms covalent bonds. -Ionic BondIonic bonding involves the transfer of valence electron/s, primarily between a metal and a -nonmetal. The electrostatic attractions between the oppositely charged ions hold the -compound together. -Ionic compounds: -1. Are usually crystalline solids (made of ions) -2. Have high melting and boiling points -3. Conduct electricity when melted -4. Are mostly soluble in water and polar solvents -Covalent Bond -A covalent bond is formed when pairs of electrons are shared between two atoms. It is -primarily formed between two same nonmetallic atoms or between nonmetallic atoms with -similar electronegativity. -Lewis Dot Structure -Lewis structures are also known as Lewis dot structures or electron dot structures. -These are basically diagrams with the element's symbol in the centre. The dots around it -represent the valence electrons of the element.  -Lewis structures of elements with atomic number 5-8 -Covalent Bonding in H2, N2 and O2 -Formation of a single bond in a hydrogen molecule: -Each hydrogen atom has a single electron in the valence shell. It requires one more to -acquire nearest noble gas configuration (He). -Therefore, both the atoms share one electron each and form a single bond.Formation of a double bond in an oxygen molecule: -Each oxygen atom has six electrons in the valence shell (2, 6). It requires two electrons to -acquire nearest noble gas configuration (Ne). -Therefore, both the atoms share two electrons each and form a double bond. -Formation of a triple bond in a nitrogen molecule: -Each nitrogen atom has five electrons in the valence shell (2, 5). It requires three electrons -to acquire nearest noble gas configuration (Ne). -Therefore, both atoms share three electrons each and form a triple bond. -Single, Double and Triple Bonds and Their Strengths -A single bond is formed between two atoms when two electrons are shared between them, -i.e., one electron from each participating atom.cbse-CBSE CBSE CBSE class-10-science-notes-chapter-4-carbon-and-its-compounds.txt -to acquire nearest noble gas configuration (Ne). -Therefore, both atoms share three electrons each and form a triple bond. -Single, Double and Triple Bonds and Their Strengths -A single bond is formed between two atoms when two electrons are shared between them, -i.e., one electron from each participating atom. -It is depicted by a single line between the two atoms. -A double bond is formed between two atoms when four electrons are shared between them, -i.e., one pair of electrons from each participating atom. It is depicted by double lines -between the two atoms. -A triple bond is formed between two atoms when six electrons are shared between them, -i.e., two pairs of electrons from each participating atom. It is depicted by triple lines -between the two atoms.Bond strength: -- The bond strength of a bond is determined by the amount of energy required to break a -bond. -- The order of bond strengths when it comes to multiple bonds is: Triple bond>double -bond>single bond -- This is to signify that the energy required to break three bonds is higher than that for two -bonds or a single bond.Bond length: -- Bond length is determined by the distance between nuclei of the two atoms in a bond. -- The order of bond length for multiple bonds is: Triple bond