CHEMICAL EQUATIONS AND REACTIONS

Chemical equations represent the reactants and products involved in a chemical reaction, following the law of conservation of mass, which states that matter cannot be created or destroyed .balancing of chemical equations is essential to ensure that the number of atoms for each element is equal on both sides of the equation. This process involves adjusting the coefficients of the reactants and products to achieve this balance, allowing for accurate stoichiometric calculations in chemical calculations.

How do we know that a chemical reaction has occured ?

• Change in colour

• Change in state

• Evolution of gas

• Change in temperature

Chemical Formula Construction and Valency Review

To write accurate chemical equations, understanding valency and ion charges is essential.

Positive Ions (Cations):

• Valency +1: Sodium ($Na^{+1}$), Potassium ($K^{+1}$), Silver ($Ag^{+1}$), Copper (I) or Cuprous ($Cu^{+1}$).

• Valency +2: Magnesium ($Mg^{2+}$), Calcium ($Ca^{2+}$), Zinc ($Zn^{2+}$), Iron (II) or Ferrous ($Fe^{2+}$), Copper (II) or Cupric ($Cu^{2+}$), Lead (II) ($Pb^{2+}$), Barium ($Ba^{2+}$).

• Valency +3: Aluminium ($Al^{3+}$), Iron (III) or Ferric ($Fe^{3+}$).

Negative Ions (Anions):

• Valency -1: Chloride ($Cl^{-1}$), Bromide ($Br^{-1}$), Nitrate ($NO_3^{-1}$), Hydroxide ($OH^{-1}$).

• Valency -2: Oxide ($O^{2-}$), Carbonate ($CO_3^{2-}$), Sulfate ($SO_4^{2-}$).

Formula Writing Examples:

• Sodium Chloride: $Na^{+1}$ and $Cl^{-1}$ cancel out to form $NaCl$.

• Aluminium Chloride: $Al^{3+}$ and $Cl^{-1}$ cross-multiply to form $AlCl_3$.

• Ferrous Sulfate: $Fe^{2+}$ and $SO_4^{2-}$ cancel out to form $FeSO_4$.

• Lead Nitrate: $Pb^{2+}$ and $NO_3^{-1}$ cross-multiply to form $Pb(NO_3)_2$.

Balancing Chemical Equations

• Definition: Ensuring the number of atoms of each element is equal on both the reactant and product sides of the equation.

• The Law of Conservation of Mass: Mass can neither be created nor destroyed in a chemical reaction. Therefore, the total mass of the reactants must equal the total mass of the products. This is the primary reason balancing is required.

• The Hierarchical Balancing Method:

  1. Metals: Balance metals first (e.g., $Na, K, Ag, Cu, Zn, Ca, Fe, Al, Pb, Mn, Ba$).

  2. Non-Metals: Balance non-metals next (e.g., $Cl, Br, S, N, C, P$), excluding oxygen and hydrogen.

  3. Oxygen: Balance oxygen atoms.

  4. Hydrogen: Balance hydrogen atoms last.

  5. Repeat: If any adjustment disturbs a previously balanced element, restart the process from step 1.

Characteristics of Chemical Reactions (FACTS)

A chemical reaction can be identified by certain observable changes, summarized by the mnemonic FACTS:

• Formation of a Precipitate: An insoluble solid forms in the solution.

• Altered Color: A change in the color of the substances.

• Concentration/Evolution of a Gas: Bubbles or fumes indicate gas production.

• Temperature Change: The reaction vessel becomes hotter (exothermic) or colder (endothermic).

• State Change: Substance changes from solid to liquid, liquid to gas, etc.

Combination Reactions

• Definition: A reaction where two or more reactants combine to form a single product ($A + B \rightarrow C$).

• Case Study 1: Burning of Magnesium Ribbon:

  • Equation: $2Mg(s) + O_2(g) \rightarrow 2MgO(s) + Heat$

  • Observations: Burns with a dazzling white flame. A white powder (Magnesium Oxide) is formed.

  • Classification: Both a combination reaction and an exothermic reaction (releases heat).

• Case Study 2: Formation of Slaked Lime:

  • Equation: $CaO(s) + H_2O(l) \rightarrow Ca(OH)_2(aq) + Heat$

  • Reactants: Quick Lime (Calcium Oxide).

  • Products: Slaked Lime (Calcium Hydroxide).

  • Notes: Reaction is vigorous. Water must be added slowly to avoid the beaker bursting due to high heat.

• Application: White Washing Walls:

  • Slaked lime is applied to walls. It reacts slowly with atmospheric carbon dioxide.

  • Equation: $Ca(OH)_2(aq) + CO_2(g) \rightarrow CaCO_3(s) + H_2O(l)$.

  • Result: Calcium Carbonate forms over 2-3 days, providing a shiny finish to the wall.

Decomposition Reactions

• Definition: A reaction where a single reactant breaks down into two or more simpler products. It is the reverse of a combination reaction. These reactions are endothermic (require energy to break bonds).

1. Thermal Decomposition (Thermolysis):

• Uses heat energy for decomposition.

• Ferrous Sulfate Crystals:

  • Step 1: $FeSO_4 \cdot 7H_2O(s) \xrightarrow{Heat} FeSO_4(s) + 7H_2O(g)$ (Green crystals turn white as they lose water of crystallization).

  • Step 2: $2FeSO_4(s) \xrightarrow{Heat} Fe_2O_3(s) + SO_2(g) + SO_3(g)$.

  • Observations: White powder turns reddish-brown (Ferric Oxide). A pungent, suffocating smell of burning sulfur (like a matchstick) is observed due to $SO_2$ and $SO_3$.

• Lead Nitrate:

  • Equation: $2Pb(NO_3)_2(s) \xrightarrow{Heat} 2PbO(s) + 4NO_2(g) + O_2(g)$.

  • Observations: Colorless/White Lead Nitrate produces Yellow Lead Oxide, Brown fumes of Nitrogen Dioxide ($NO_2$), and Oxygen gas. A characteristic cracking sound is heard.

2. Electrolytic Decomposition (Electrolysis):

• Uses electrical energy for decomposition.

• Electrolysis of Water:

  • Equation: $2H_2O(l) \xrightarrow{Electric\,Current} 2H_2(g) + O_2(g)$.

  • Setup: 6V Battery, Plastic mug, Graphite electrodes.

  • Cathode (Negative): Hydrogen gas collects here. Hydrogen is combustible and burns with a pop sound.

  • Anode (Positive): Oxygen gas collects here. Oxygen supports combustion (re-kindles a glowing splinter).

  • Mnemonic: PAO ($P$ositive $A$node $O$xygen).

  • Ratio: The volume ratio of $H_2$ to $O_2$ is $2:1$. The mass ratio is $1:8$.

  • Note: A few drops of acid (like $H_2SO_4$) are added to pure water because pure water is a poor conductor of electricity.

3. Photochemical Decomposition (Photolysis):

• Uses light energy for decomposition.

• Silver Chloride: $2AgCl(s) \xrightarrow{Sunlight} 2Ag(s) + Cl_2(g)$.

• Silver Bromide: $2AgBr(s) \xrightarrow{Sunlight} 2Ag(s) + Br_2(g)$.

• Observations: White $AgCl$ turns grey as silver metal is formed.

• Application: Used in Black and White Photography.

• Storage: These chemicals are stored in dark or black bottles to prevent accidental decomposition by light.

Displacement Reactions

• Definition: A reaction where a more reactive element displaces a less reactive element from its compound.

• Reactivity Series (Descending Order):

  • Potassium ($K$), Sodium ($Na$), Calcium ($Ca$), Magnesium ($Mg$), Aluminium ($Al$), Zinc ($Zn$), Iron ($Fe$), Lead ($Pb$), [Hydrogen ($H$)], Copper ($Cu$), Mercury ($Hg$), Silver ($Ag$), Gold ($Au$).

• Iron Nail in Copper Sulfate Experiment:

  • Equation: $Fe(s) + CuSO_4(aq) \rightarrow FeSO_4(aq) + Cu(s)$.

  • Observations: The blue color of Copper Sulfate fades to light green/pale green (Ferrous Sulfate). The iron nail becomes coated with a reddish-brown layer of Copper.

• Zinc and Sulfuric Acid:

  • Equation: $Zn(s) + H_2SO_4(aq) \rightarrow ZnSO_4(aq) + H_2(g)$.

  • Notes: Zinc displaces Hydrogen. The reaction is highly exothermic.

Double Displacement and Precipitation Reactions

• Definition: A reaction where there is an exchange of ions between the reactants to form new compounds ($AB + CD \rightarrow AD + CB$).

• Precipitation Reaction: A specific double displacement reaction where two aqueous solutions react to form an insoluble solid (precipitate).

• Case 1: Sodium Sulfate and Barium Chloride:

  • Equation: $Na_2SO_4(aq) + BaCl_2(aq) \rightarrow BaSO_4(s) + 2NaCl(aq)$.

  • Observation: A white precipitate of Barium Sulfate forms.

  • Ions in solution: After the precipitate settles, the remaining ions in the aqueous solution are $Na^{+}$ and $Cl^{-}$.

• Case 2: Lead Nitrate and Potassium Iodide:

  • Equation: $Pb(NO_3)_2(aq) + 2KI(aq) \rightarrow PbI_2(s) + 2KNO_3(aq)$.

  • Observation: A yellow precipitate of Lead Iodide is formed.

Oxidation, Reduction, and Redox Reactions

• Oxidation: Gain of oxygen or loss of hydrogen.

• Reduction: Loss of oxygen or gain of hydrogen.

• Redox Reaction: A reaction where oxidation and reduction occur simultaneously.

• Oxidizing Agent: The substance that undergoes reduction (provides oxygen or accepts hydrogen).

• Reducing Agent: The substance that undergoes oxidation (accepts oxygen or provides hydrogen).

• Copper Oxidation:

  • Reaction: $2Cu + O_2 \xrightarrow{Heat} 2CuO$. Reddish-brown copper turns black Copper Oxide.

  • Reversal: Passing Hydrogen gas over heated $CuO$ restores the color: $CuO + H_2 \xrightarrow{Heat} Cu + H_2O$.

Corrosion and Rancidity

• Corrosion: The slow eating away of metal surfaces by the action of air, moisture, or chemicals. It is an oxidation process.

  • Rusting of Iron: Formation of reddish-brown hydrated iron oxide.

  • Tarnishing of Copper: Formation of a green coating (Basic Copper Carbonate).

  • Tarnishing of Silver: Silver reacts with $H_2S$ in the air to form black Silver Sulfide ($Ag_2S$).

• Rancidity: The oxidation of fats and oils in food, resulting in an unpleasant smell and taste.

  • Prevention:

    1. Adding Antioxidants.

    2. Storage in air-tight containers.

    3. Flushing food packaging with Nitrogen gas (an inert gas that prevents oxidation).

Physical vs. Chemical Changes

• Chemical Change: Results in the formation of new substances. Reversibility is difficult. (Examples: Burning of paper, rusting, digestion, respiration, fermentation of grapes, souring of milk).

• Physical Change: No new chemical substances are formed. Involves changes in state, shape, or size. (Examples: Boiling of water, melting of ice, melting of wax, crushing paper).

Questions & Discussion

• Question: Why should a magnesium ribbon be cleaned with sandpaper before burning in air?

• Answer: To remove the protective layer of Magnesium Oxide ($MgO$) that may have formed on its surface due to reaction with moist air, ensuring it burns effectively.

• Question: Why is it advised to keep the magnesium ribbon away from eyes while burning?

• Answer: It emits a dazzling white light (UV radiation) so bright that it can damage the eyes or cause temporary blindness.

• Question: Why is the amount of gas collected in one test tube during the electrolysis of water double the amount collected in the other?

• Answer: According to the chemical formula of water ($H_2O$), there are two parts of hydrogen for every one part of oxygen. Thus, the volume of hydrogen collected at the cathode is double the volume of oxygen collected at the anode.

• Question: Why are some decomposition reactions called endothermic?

• Answer: Because they require the absorption of energy in the form of heat, light, or electricity to break the chemical bonds of the reactants.

• Question: Identify the substances oxidized and reduced in the reaction: $MnO_2 + 4HCl \rightarrow MnCl_2 + 2H_2O + Cl_2$.

• Answer: $HCl$ is oxidized to $Cl_2$ (loss of hydrogen). $MnO_2$ is reduced to $MnCl_2$ (loss of oxygen). $MnO_2$ is the oxidizing agent and $HCl$ is the reducing agent.

• Question: What is the law that suggests the balancing of a chemical equation?

• Answer: The Law of Conservation of Mass, which states that mass can neither be created nor destroyed in a chemical reaction.

• Question: Why does the color of copper sulfate solution change when an iron nail is dipped in it?

• Answer: Because Iron, being more reactive than Copper, displaces Copper from the solution to form Ferrous Sulfate ($FeSO_4$), which is pale green, while the blue color of $CuSO_4$ disappears.