Chemical Reactions & Equations Full Chapter in Animation | Class 10 Science Chapter 1 | CBSE 2026

[00:00:04][music] >> Chemical reactions and equations.

[00:00:10]In our day-to-day life, we see different changes around us.

[00:00:14]Like, if milk is left at room temperature during summer, it becomes sour and curdling takes place.

[00:00:21]If an iron tawa is exposed to moisture, it rusts.

[00:00:26]If grapes are fermented, they smell like alcohol.

[00:00:30]If food is cooked, its texture and taste changes.

[00:00:34]When food is digested, it is converted into useful nutrients.

[00:00:39]When we respire, food gets oxidized and releases heat and energy.

[00:00:44]In all these cases, the chemical nature of the initial substances is changed.

[00:00:50]These kind of changes are called chemical changes.

[00:00:54]A chemical change takes place due to a chemical reaction.

[00:00:59]Let us see some examples of chemical reactions in lab setup.

[00:01:05]When a magnesium ribbon is burnt in presence of oxygen, it burns with a dazzling white flame and white ash is obtained.

[00:01:13]This ash is magnesium oxide.

[00:01:17]When some potassium iodide solution is added to lead nitrate, an yellow precipitate is formed.

[00:01:23]This yellow precipitate is nothing but lead iodide.

[00:01:27]In the same way, when a few zinc granules are added to dilute HCl or H2SO4, hydrogen gas is evolved.

[00:01:36]Heat is also released in this reaction.

[00:01:40]In the above reactions, we observed change in state, change of color, evolution of gas, change in temperature.

[00:01:49]Along with change in chemical nature of the initial substance, we also observe any of these changes in a chemical reaction.

[00:01:57]So, these observations help us to identify a chemical reaction.

[00:02:02]Chemical equation.

[00:02:04]Magnesium, when burnt in presence of oxygen, it produces magnesium oxide.

[00:02:10]This is the verbal description of a chemical reaction.

[00:02:14]This can be presented in the short form, that is in the form of a word equation.

[00:02:19]Magnesium plus oxygen gives rise to magnesium oxide.

[00:02:26]In this equation, the starting materials magnesium and oxygen are called reactants, and the final substance formed magnesium oxide is the product.

[00:02:37]The reactants and products are separated by an arrow.

[00:02:41]This arrow head points towards the products and shows the direction of the reaction.

[00:02:47]Reactants are always written on left-hand side of the arrow, and products are written on the right-hand side.

[00:02:55]There is a plus sign between the two reactants and between the products.

[00:03:01]Writing a chemical equation using chemical formulas.

[00:03:06]Magnesium plus oxygen gives rise to magnesium oxide.

[00:03:11]Let us write this equation using chemical formulas.

[00:03:16]Mg for magnesium and O2 for oxygen. Mg plus O2 gives rise to MgO.

[00:03:23]Here the number of atoms on the left-hand side and right-hand side are not equal.

[00:03:28]That means this is an unbalanced equation.

[00:03:32]Such equations are called skeletal equations.

[00:03:36]Then, how to balance a skeletal equation?

[00:03:39]Let's see.

[00:03:40]Balancing a chemical equation.

[00:03:43]The law of conservation of mass states that in a closed system, the mass of the reactants must equal with the mass of the products for any given chemical reaction.

[00:03:54]It implies that atoms are neither created nor destroyed in a chemical reaction, but they are rearranged it to form different substances.

[00:04:03]That means the number of atoms of each element remains the same before and after a chemical reaction.

[00:04:10]So, we need to balance a skeletal equation to present it properly.

[00:04:16]Fe + H2O gives rise to Fe3O4 + H2.

[00:04:22]Now, let us balance this chemical equation.

[00:04:25]Step one.

[00:04:26]When balancing a chemical equation, begin with enclosing each chemical formula within boxes. Do not change anything inside these boxes while working to balance the equation.

[00:04:38]Now, tabulate the number of atoms of different elements present in this unbalanced equation.

[00:04:44]Iron, number of atoms in the reactants is one and in the products it is three.

[00:04:49]Hydrogen, number of atoms in the reactants two, in the products two.

[00:04:54]Oxygen, in the reactants one and in the products four.

[00:04:58]Step three. Sometimes it is easier to balance a chemical equation by focusing on the substance with maximum number of atoms.

[00:05:07]This substance could be on either side of the equation.

[00:05:10]For example, if we have Fe3O4, which has a lot of oxygen atoms.

[00:05:15]If we look at the oxygen atoms on both sides, there are four oxygen atoms on one side, that is on right hand side, and only one oxygen atom on the left hand side.

[00:05:25]So, we'll start by adjusting the number of oxygen atoms to make them equal on both these sides.

[00:05:31]To balance the oxygen atoms, the number of oxygen atoms initial in reactants it is one, that is in H2O, and in products four, that is in Fe3O4.

[00:05:43]So, to balance it, we'll multiply the one with four, 1 into four in the reactants.

[00:05:49]That means we are going to add four to the coefficient of H2O.

[00:05:54]So, to equalize the number of atoms, we can put coefficient four as 4H2O.

[00:06:01]Now, the number of oxygen atoms on both left-hand side and right-hand side are balanced.

[00:06:06]Step four.

[00:06:08]Now, balance the hydrogen atoms.

[00:06:10]To equalize the number of hydrogen atoms, make the number of molecules of hydrogen as four on the right-hand side.

[00:06:17]So, atoms of hydrogen in reactants, it is eight in 4H2O.

[00:06:22]Whereas in products, it is only two. So, to balance it in the products, we are multiplying two with four.

[00:06:29]That means we are adding four as a coefficient for hydrogen atoms.

[00:06:35]So, now the equation becomes like this.

[00:06:37]Fe plus 4H2O gives rise to Fe3O4 plus 4H2. But, still this is an unbalanced equation.

[00:06:45]Here, iron is left unbalanced.

[00:06:48]To equalize iron, we take three atoms of Fe on the left-hand side.

[00:06:53]Let's see.

[00:06:54]Atoms of iron initial in the reactants, only one.

[00:06:59]Whereas in the products, three. Fe3O4.

[00:07:04]So, now in the reactants, we are multiplying one with three. That means we are adding three as a coefficient for Fe.

[00:07:13]Now, it becomes 3Fe plus 4H2O gives rise to Fe3O4 plus 4H2.

[00:07:20]So, now this is a balanced equation.

[00:07:23]So, step six.

[00:07:25]Now, count the number of atoms on both sides to check whether the number is balanced on both sides or not. This method of balancing is called hit and trial method.

[00:07:35]Step seven.

[00:07:36]Now, writing the symbols of physical states.

[00:07:39]To make a chemical equation more informative, the physical states of the reactants and products are mentioned along with the chemical formulae.

[00:07:47]The gaseous, liquid, aqueous, and solid states of reactants and products are represented by the notations G for gases, L for liquid, AQ for aqueous, and S for solids respectively.

[00:08:01]Aqueous is a solution in which water is the solvent.

[00:08:05]Sometimes the reaction conditions such as temperature, pressure, catalyst, etc. for the reaction are indicated above and or below the arrow in the equation.

[00:08:16]For example, CO G means gas plus 2H2 G, that is gas, gives rise to CH3OH L, that means it's liquid. And here on the arrow we can see 340 atm. It indicates the atmospheric pressure in this reaction.

[00:08:35]In the same way in the next reaction, 6CO2 AQ, that is aqueous, plus 12H2O L, that is liquid, gives rise to C6H12O6 aqueous, that is glucose, plus 6O2, that is aqueous, plus 6H2O L, liquid.

[00:08:54]So this reaction takes place in presence of sunlight and chlorophyll. These things are mentioned on and below the arrow mark.

[00:09:03]It is to be noted that usually physical states are not included in a chemical equation unless it is necessary to specify them.

[00:09:11]Types of chemical reactions.

[00:09:13]One, combination reaction.

[00:09:15]A combination reaction is a type of chemical reaction where two or more reactants combine to form a single product. Take some water in a beaker and add calcium oxide to it.

[00:09:26]Calcium oxide reacts with water vigorously and forms a precipitate, that is calcium hydroxide.

[00:09:33]Large amount of heat is also released in this reaction.

[00:09:37]In this reaction, two reactants, that is water and calcium oxide, react and forms a single product, that is calcium hydroxide. Such reactions are called as combination reactions.

[00:09:49]The chemical reaction in which two reactants react to form single product is called a combination reaction.

[00:09:57]Some examples of combination reactions are one, burning of coal.

[00:10:02]Carbon plus oxygen gives rise to carbon dioxide.

[00:10:06]So here, two reactants combining together and forming a product, carbon dioxide.

[00:10:12]And second one, formation of water from H2 and O2.

[00:10:17]H2 plus O2 gives rise to 2H2O. Hydrogen plus oxygen gives rise to water. Here also two reactants, hydrogen and oxygen, are combining to form a product, water.

[00:10:29]So in this particular reaction, heat is also released out. The chemical reactions in which heat is released are called exothermic reactions.

[00:10:39]Some examples of exothermic reactions.

[00:10:42]One, burning of natural gas.

[00:10:45]CH4 plus 2O2 gives rise to CO2 plus 2H2O.

[00:10:50]In this reaction, heat is released.

[00:10:53]And one more example for exothermic reaction, cellular respiration. Here the glucose reacts with oxygen and gives carbon dioxide plus water plus energy.

[00:11:03]In this reaction also, heat is released.

[00:11:07]Decomposition reaction.

[00:11:09]A decomposition reaction is a type of chemical reaction in which a single compound breaks down into two or more simpler substances.

[00:11:17]Let us see some examples of it.

[00:11:19]Ferrous sulfate crystals are light green in color.

[00:11:23]On heating, the water in these crystals evaporates, resulting in the formation of an anhydrous ferrous sulfate.

[00:11:30]The color of the crystals changes from light green to white.

[00:11:35]Now, this compound decomposes to different substances. The ferric oxide, sulfur dioxide, and sulfur trioxide.

[00:11:44]Now, let's see another example of decomposition reaction.

[00:11:47]On heating, calcium carbonate decomposes into calcium oxide and carbon dioxide.

[00:11:52]It is an important decomposition reaction used in various industries.

[00:11:57]Another example of decomposition.

[00:12:00]On heating lead nitrate, it decomposes to lead oxide, nitrogen oxide, and hydrogen. We can observe the evolution of nitrogen oxide as brown fumes.

[00:12:11]The decomposition reactions that takes place by heating are called thermal decomposition reactions.

[00:12:17]All the above discussed decomposition reactions are thermal decomposition reactions.

[00:12:22]Now, let's see electrolytic decomposition.

[00:12:26]If the decomposition reaction takes place by passing electricity, such reaction is called electrolytic decomposition reaction. Example, electrolysis of water.

[00:12:38]The chemical equation for this reaction is 2H2O in presence of electricity, 2H2 + O2.

[00:12:45]Now, let us see photolytic decomposition.

[00:12:48]If the decomposition reaction takes place by the exposure to light is called photolytic decomposition.

[00:12:55]When silver chloride is exposed to sunlight, it decomposes into silver and chlorine gas.

[00:13:01]The chemical equation for this reaction is 2AgCl gives rise to 2Ag + Cl2.

[00:13:08]Silver bromide also decomposes by sunlight.

[00:13:11]2AgBr gives rise to 2Ag + Br2 in presence of sunlight.

[00:13:17]The decomposition reactions require energy either in the form of heat, light, or electricity for breaking down the reactants.

[00:13:26]The reactions in which energy is absorbed are known as endothermic reactions.

[00:13:32]Now, let us learn about displacement reaction.

[00:13:36]Gather three iron nails and clean them by rubbing with sandpaper.

[00:13:41]Obtain two test tubes labeled as A and B.

[00:13:44]Pour approximately 10 ml of copper sulfate solution into each test tube.

[00:13:49]Securely tie two iron nails together with a thread and carefully immerse them into the copper sulfate solution in test tube B for 20 minutes.

[00:13:58]Set aside one nail for comparison later.

[00:14:01]After 20 minutes, remove the iron nails from the copper sulfate solution.

[00:14:06]Compare the intensity of the blue color of copper sulfate solution in test tubes A and B. Also, compare the color of the iron nails dipped in the copper sulfate solution with the one kept aside.

[00:14:18]The following chemical reaction takes place in this activity.

[00:14:22]Iron plus copper sulfate gives rise to iron sulfate plus copper.

[00:14:27]In this reaction, iron has displaced copper from the copper sulfate solution.

[00:14:32]Such a reaction is termed as a displacement reaction.

[00:14:36]Some other examples of displacement reactions.

[00:14:39]Zinc plus copper sulfate gives rise to zinc sulfate plus copper.

[00:14:43]Lead plus copper chloride gives rise to lead chloride plus copper.

[00:14:48]Since zinc and lead are more reactive than copper, they displace copper from their compounds.

[00:14:53]Now, let us see double displacement reaction.

[00:14:57]A double displacement reaction is a type of chemical reaction where two compounds react by exchanging ions to form two new compounds.

[00:15:06]If we mix sodium sulfate and barium chloride solution, we get a white precipitate. In any chemical reaction, if a precipitate is formed, it's called a precipitation reaction.

[00:15:17]In this reaction, sodium chloride reacts with barium chloride and forms barium sulfate and sodium chloride.

[00:15:24]Here the barium sulfate forms the precipitate and the sodium chloride remains in the solution.

[00:15:30]Oxidation and reduction reactions.

[00:15:34]If a substance gains oxygen during reaction, then the substance is said to be oxidized and such reaction is called an oxidation reaction.

[00:15:42]For example, when copper powder is heated in presence of oxygen, the brown colored surface of the copper powder turns black due to the formation of copper two oxide.

[00:15:53]So when copper reacts with oxygen, it forms copper oxide.

[00:15:57]Here, can we get back the copper from copper oxide? Yes, by a reduction reaction.

[00:16:04]If a substance loses oxygen during a reaction, it is said to be reduced and such reaction is called a reduction reaction.

[00:16:12]When hydrogen gas is passed over heated copper oxide, the black surface coating undergoes a transformation turning brown as the reverse reaction occurs and finally copper is obtained.

[00:16:24]This reaction has both oxidation and reduction in it.

[00:16:29]Let's see, copper oxide turning to copper is a reduction reaction, whereas hydrogen turns to water, it is an oxidation reaction.

[00:16:39]During this reaction, copper oxide gets reduced by losing oxygen. The hydrogen is gaining oxygen and is being oxidized.

[00:16:46]So in other words, one reactant gets oxidized while the other gets reduced during a reaction.

[00:16:52]Such reactions are called oxidation reduction reactions or simply redox reactions.

[00:16:58]Some other examples of redox reactions.

[00:17:01]Zinc oxide plus carbon gives rise to zinc plus carbon monoxide. In the same way, manganese oxide plus 4 HCl gives rise to manganese chloride plus 2 H2O plus Cl2.

[00:17:14]Now let's see the effects of oxidation reactions in our day-to-day life.

[00:17:18]Corrosion.

[00:17:19]When a metal is affected by substances around it, such as moisture, acids, etc., it is said to be corroded, and this process is called corrosion.

[00:17:30]When iron is exposed to moisture, it gets coated with a reddish-brown powder.

[00:17:34]This process is called as rusting of iron.

[00:17:37]The black coating on the silver articles and the green coating on the brass items are the examples of corrosion.

[00:17:44]Rancidity. When fats and oils are exposed to air, they gets oxidized. This leads to change in their smell and taste.

[00:17:53]Then those fats are said to be rancid fats, and this process is called rancidity.

[00:17:59]To prevent the oxidation of fatty foods, the following preventive steps are followed. Adding antioxidant chemicals to the foods.

[00:18:07]Two, keeping food in airtight container.

[00:18:10]Three, filling the chips packets with nitrogen gas.

[00:18:14]This is all about the chemical reactions and equations.

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