diff --git "a/lancedb/cbse.lance/data/3f75d07a-b551-455e-8ab5-a84fa739b0cc.lance" "b/lancedb/cbse.lance/data/3f75d07a-b551-455e-8ab5-a84fa739b0cc.lance" new file mode 100644--- /dev/null +++ "b/lancedb/cbse.lance/data/3f75d07a-b551-455e-8ab5-a84fa739b0cc.lance" @@ -0,0 +1,4930 @@ +chemistry-notes-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 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 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 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 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 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 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 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 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 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 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 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 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 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 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 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 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 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 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 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 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 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