Ch 01 - Chemical Reactions and Equations10th Science

A Physical Change involves no new chemical substance formation. The change is usually reversible and involves only a state or phase transition (e.g., melting ice, or the sublimation of camphor from solid directly to gas).

A Chemical Change forms entirely new substances with new chemical properties. The bonds between atoms are broken and rearranged (e.g., burning LPG, digestion of food, curdling of milk).

• Cleaning: To remove the protective layer of Magnesium Oxide / Basic Magnesium Carbonate formed due to atmospheric exposure, ensuring a steady reaction.
• Observation: It burns with a dazzling white flame.
• Product: Forms a white powder called Magnesium Oxide (MgO).

A chemical reaction can be identified by any of the following physical observations:

• Change in state
• Change in colour
• Evolution of a gas
• Change in temperature

Variables & Units:

• Mg = Magnesium (Solid)
• O_2 = Oxygen gas from air (Gas)
• MgO = Magnesium Oxide (White powder, Solid)

Conditions/Notes:

• The ribbon burns with a dazzling white flame.
• This is also an example of an oxidation and combination reaction.

• Skeletal Equation: Represents the chemical change using formulas but is unbalanced (total mass is not equal on both sides). Example: $Mg + O_2 \rightarrow MgO$
• Balanced Equation: Has an equal number of atoms of each element on both the reactant and product sides, strictly adhering to the Law of Conservation of Mass.

• The Law: Mass can neither be created nor destroyed in a chemical reaction.
• Implication: The total mass of the elements present in the reactants must strictly equal the total mass of the elements in the products.
• Therefore, the number of atoms of each element remains constant before and after the reaction, demanding that we balance the equation.

• Hit-and-Trial Method: A method of balancing equations by making multiple trials to equalize atoms using the smallest whole number coefficients.
• Strategy: Always start with the compound that contains the maximum number of atoms (whether it is a reactant or product).
• Within that compound, balance the element with the highest number of atoms first.

• (s) = Solid state (also used for precipitates).
• (l) = Liquid state.
• (g) = Gaseous state.
• (aq) = Aqueous solution (substance is dissolved in water).
• Note: Reaction conditions like temperature, pressure, or catalysts are typically written above or below the reaction arrow.

• Definition: A reaction in which two or more reactants combine to form a single product.
• General Form: $A + B \rightarrow AB$
• Key Note: These reactions are frequently highly exothermic as new chemical bonds are formed, releasing energy.

• Definition: Reactions in which heat is released along with the formation of products.
• Examples:
  1. Burning of Natural Gas (Methane): $CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O + \text{Heat}$
  2. Cellular Respiration in our bodies.
  3. Decomposition of vegetable matter into compost.

Variables & Units:

• CaO = Calcium Oxide / Quicklime (Solid)
• H_2O = Water (Liquid)
• Ca(OH)_2 = Calcium Hydroxide / Slaked Lime (Aqueous)

Conditions/Notes:

• This is a highly exothermic combination reaction.
• It produces a hissing sound and a large amount of heat.

• Definition: A reaction in which a single reactant breaks down to yield two or more simpler products.
• General Form: $AB \rightarrow A + B$
• Types of Decomposition:
  1. Thermal: Driven by heat.
  2. Electrolytic: Driven by electricity.
  3. Photolytic: Driven by light.
• Note: These are inherently Endothermic, as breaking bonds requires energy.

Variables & Units:

• Ca(OH)_2 = Slaked Lime (Aqueous)
• CO_2 = Carbon Dioxide from the atmosphere (Gas)
• CaCO_3 = Calcium Carbonate (Solid)

Conditions/Notes:

• The reaction takes 2 to 3 days to form a thin layer of $CaCO_3$, providing a shiny finish to the walls.
• Fact: $CaCO_3$ is also the chemical formula for Marble.

• Definition: Reactions in which energy is absorbed from the surroundings.
• Characteristic: The energy required can be in the form of heat, light, or electricity to break the bonds of the reactants.
• Examples: All decomposition reactions, photosynthesis, and dissolving ammonium chloride in water.

• Definition: A reaction in which a more reactive element displaces a less reactive element from its compound.
• General Form: $A + BC \rightarrow AC + B$
• Principle: It strictly relies on the Reactivity Series of metals. An element can only kick out elements below it in the series.

Variables & Units:

• C_6H_{12}O_6 = Glucose (Aqueous)
• O_2 = Oxygen (Aqueous)
• CO_2 = Carbon Dioxide (Aqueous)
• H_2O = Water (Liquid)

Conditions/Notes:

• Respiration is fundamentally an exothermic process because it breaks down glucose to release the energy our bodies need to function.

• Definition: A reaction in which there is an exchange of ions between the two reactant compounds.
• General Form: $AB + CD \rightarrow AD + CB$
• Mechanism: The cations and anions of two different compounds switch places, forming two entirely new compounds. One of these is frequently an insoluble solid (precipitate).

• Oxidation: The gain of Oxygen OR the loss of Hydrogen during a reaction.
• Reduction: The loss of Oxygen OR the gain of Hydrogen during a reaction.
• Redox Reaction: When oxidation and reduction occur simultaneously in a reaction (one reactant gets oxidised while the other gets reduced).

Variables & Units:

• FeSO_4 = Ferrous Sulphate (Solid)
• Fe_2O_3 = Ferric Oxide (Solid)
• SO_2 = Sulphur Dioxide (Gas)
• SO_3 = Sulphur Trioxide (Gas)

Conditions/Notes:

• Green crystals of $FeSO_4 \cdot 7H_2O$ lose water when heated, turning white, then decompose to reddish-brown iron (III) oxide.
• Accompanied by the characteristic smell of burning sulphur.

Variables & Units:

• CaCO_3 = Calcium Carbonate / Limestone (Solid)
• CaO = Calcium Oxide / Quicklime (Solid)
• CO_2 = Carbon Dioxide (Gas)

Conditions/Notes:

• CaO (Quicklime) is heavily used in the manufacture of cement.

• Cathode (Negative Electrode): Hydrogen gas ($H_2$) is collected.
• Anode (Positive Electrode): Oxygen gas ($O_2$) is collected.
• Volume Ratio: The volume of Hydrogen collected is double the volume of Oxygen (2:1 ratio).
• Reason: Water ($H_2O$) is composed of two parts hydrogen and one part oxygen by volume.

Variables & Units:

• Pb(NO_3)_2 = Lead Nitrate (Solid)
• PbO = Lead Oxide (Solid)
• NO_2 = Nitrogen Dioxide (Gas)
• O_2 = Oxygen (Gas)

Conditions/Notes:

• Crucial Observation: The emission of thick brown fumes, which are Nitrogen Dioxide ($NO_2$) gas.

Variables & Units:

• AgBr = Silver Bromide (Solid)
• Ag = Silver (Solid)
• Br_2 = Bromine (Gas)

Conditions/Notes:

• The pale yellow $AgBr$ turns grey due to the formation of silver metal.
• Application: This specific reaction (along with $AgCl$) is strictly used in black and white photography.

Variables & Units:

• Fe = Iron (Solid nail)
• CuSO_4 = Copper Sulphate (Aqueous, Blue)
• FeSO_4 = Iron(II) Sulphate (Aqueous, Light Green)
• Cu = Copper (Solid, Brown)

Conditions/Notes:

• Observation 1: The blue colour of $CuSO_4$ fades to light green ($FeSO_4$).
• Observation 2: A brownish coating of copper metal deposits on the iron nail.
• Works strictly because Iron (Fe) is more reactive than Copper (Cu).

Variables & Units:

• Na_2SO_4 = Sodium Sulphate (Aqueous)
• BaCl_2 = Barium Chloride (Aqueous)
• BaSO_4 = Barium Sulphate (Solid precipitate)
• NaCl = Sodium Chloride (Aqueous)

Conditions/Notes:

• Observation: A white precipitate of $BaSO_4$ is immediately formed.
• Any reaction that produces an insoluble solid (precipitate) is also called a Precipitation Reaction.

Variables & Units:

• CuO = Copper(II) Oxide (Solid, Black)
• H_2 = Hydrogen (Gas)
• Cu = Copper (Solid, Brown)
• H_2O = Water (Liquid)

Conditions/Notes:

• Reduced Substance: $CuO$ (loses oxygen to become $Cu$).
• Oxidised Substance: $H_2$ (gains oxygen to become $H_2O$).
• Observation: The black coating of Copper(II) Oxide turns brown as pure Copper is regenerated.

• Corrosion: Gradual degradation of metals by the attack of moisture, acids, or atmospheric gases (e.g., Rusting of iron, black coating on silver, green coating on copper).
• Rancidity: The oxidation of fats and oils in food, resulting in a foul smell and bad taste.
• Prevention of Rancidity:
  1. Adding Antioxidants.
  2. Storing food in airtight containers.
  3. Flushing bags with Nitrogen gas (e.g., in potato chip bags).

Variables & Units:

• Fe = Iron (Solid)
• H_2O = Water (Gas/Steam)
• Fe_3O_4 = Iron(II,III) oxide / Magnetic iron oxide (Solid)
• H_2 = Hydrogen (Gas)

Conditions/Notes:

• Iron does not react with cold or hot water; it strictly requires steam.
• This is a classic textbook example used to teach the 'Hit-and-Trial' balancing method.

Variables & Units:

• Pb(NO_3)_2 = Lead Nitrate (Aqueous)
• KI = Potassium Iodide (Aqueous)
• PbI_2 = Lead Iodide (Solid precipitate)
• KNO_3 = Potassium Nitrate (Aqueous)

Conditions/Notes:

• Crucial Observation: A striking yellow precipitate of Lead Iodide ($PbI_2$) is formed immediately upon mixing.
• Frequently tested in practical-based board questions.

Variables & Units:

• Zn = Zinc (Solid granules)
• H_2SO_4 = Sulphuric Acid (Aqueous, dilute)
• ZnSO_4 = Zinc Sulphate (Aqueous)
• H_2 = Hydrogen (Gas)

Conditions/Notes:

• Observation 1: Bubbles of Hydrogen gas are seen around the zinc granules.
• Observation 2: The flask becomes warm, indicating it is an exothermic reaction.
• $H_2$ gas can be tested by bringing a burning candle near it; it burns with a characteristic 'pop' sound.

• Oxidising Agent: The substance that gives oxygen or removes hydrogen. It gets reduced itself during the reaction.
• Reducing Agent: The substance that gives hydrogen or removes oxygen. It gets oxidised itself during the reaction.
• Rule of Thumb: The substance oxidised acts as the reducing agent; the substance reduced acts as the oxidising agent.

Variables & Units:

• CO_2 = Carbon Dioxide (Aqueous/Dissolved)
• H_2O = Water (Liquid)
• C_6H_{12}O_6 = Glucose (Aqueous)
• O_2 = Oxygen (Aqueous)

Conditions/Notes:

• Must occur in the presence of Sunlight and Chlorophyll.
• This is fundamentally an endothermic process, converting solar energy into chemical energy.

Variables & Units:

• MnO_2 = Manganese Dioxide (Solid)
• HCl = Hydrochloric Acid (Aqueous)
• MnCl_2 = Manganese(II) Chloride (Aqueous)
• Cl_2 = Chlorine (Gas)

Conditions/Notes:

• Reduced: $MnO_2$ loses oxygen to become $MnCl_2$.
• Oxidised: $HCl$ loses hydrogen to become $Cl_2$.
• $MnO_2$ acts as the oxidising agent, while $HCl$ acts as the reducing agent.

Variables & Units:

• Ba(OH)_2 = Barium Hydroxide (Solid)
• NH_4Cl = Ammonium Chloride (Solid)
• BaCl_2 = Barium Chloride (Aqueous)
• NH_3 = Ammonia (Gas)

Conditions/Notes:

• Crucial Observation: This is a rare example of a reaction where mixing two solids absorbs so much heat from the surroundings that the bottom of the test tube feels extremely cold to the touch.

Variables & Units:

• CO = Carbon Monoxide (Gas)
• H_2 = Hydrogen (Gas)
• CH_3OH = Methanol (Liquid)

Conditions/Notes:

• This reaction strictly requires a high pressure of 340 atm, illustrating how specific conditions (temperature, pressure, catalysts) are written directly above or below the yield arrow.

Variables & Units:

• ZnO = Zinc Oxide (Solid)
• C = Carbon / Coke (Solid)
• Zn = Zinc (Solid metal)
• CO = Carbon Monoxide (Gas)

Conditions/Notes:

• Reduced: $ZnO$ loses oxygen to form $Zn$.
• Oxidised: $C$ gains oxygen to form $CO$.
• Carbon acts as a strong reducing agent. This principle is widely used in the metallurgical extraction of metals from their ores.

• Spectator Ions: Ions that exist in the same form on both the reactant and product sides of a chemical equation. They do not actively participate in the chemical change.
• Net Ionic Equation: An equation that omits spectator ions, showing only the species that actually react to form a precipitate or a new compound.
• Example: In $Na_2SO_4(aq) + BaCl_2(aq) \rightarrow BaSO_4(s) + 2NaCl(aq)$, $Na^+$ and $Cl^-$ are spectator ions. The net ionic equation is strictly $Ba^{2+}(aq) + SO_4^{2-}(aq) \rightarrow BaSO_4(s)$.