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