Chemistry HYE PDF

1. Arrhenius Theory Definition: ○ Acid: Produces H+ ions in water. ○ Base: Produces OH− ions in water. Example: ○ Acid: HCl→H+Cl− ○ Base: NaOH→Na+ OH− Advantages: ○ Simple and easy to understand. ○ Focuses on aqueous solutions. Limitations: ○...

1. Arrhenius Theory Definition: ○ Acid: Produces H+ ions in water. ○ Base: Produces OH− ions in water. Example: ○ Acid: HCl→H+Cl− ○ Base: NaOH→Na+ OH− Advantages: ○ Simple and easy to understand. ○ Focuses on aqueous solutions. Limitations: ○ This only applies to water-based (aqueous) solutions. ○ It cannot explain substances like NH3 (a base thdoesn'tn’t release OH−). 2. Brønsted-Lowry Theory Definition: ○ Acid: Proton (H+) donor. ○ Base: Proton (H+) acceptor. Example: ○ Acid: HCl+H2O→H3O+ Cl− ○ Base: NH3+H2O→NH4+ OH− Advantages: ○ Works in both aqueous and non-aqueous systems. ○ Explain the behaviour of bases like NH3. Limitations: ○ Focuses only on proton transfer. ○ Does not explain reactions involving non-protonic acids and bases (e.g., BF3). 3. Lewis Theory Definition: ○ Acid: Electron pair acceptor. ○ Base: Electron pair donor. Example: ○ NH3+BF3→H3NBF3 ◆ NH3: Donates an electron pair (base). ◆ BF3BF_3BF3: Accepts an electron pair (acid). Advantages: ○ The broadest definition includes all acid-base reactions. ○ Explains reactions where no protons are involved. Limitations:  ○ More abstract and less intuitive. ○ Does not prioritise H+ ions. Key Comparison Theory Focus Advantage Limitation Arrhenius H+ and OH− Simple, water- Limited to ions based reactions aqueous solutions Brønsted- Proton (H+) Includes non- Excludes non- Lowry transfer aqueous protonic reactions reactions Lewis Electron pair Broad and Abstract, harder transfer versatile to visualize Properties of Acids and Bases 1. Definitions: Strong and Weak Acids/Bases Strong Acid: Completely ionises (dissociates) in water. Example: HCl→H+ Cl− Weak Acid: Partially ionises in water. Example: CH3COOH⇋H+ CH3COO− Strong Base: Completely dissociates in water to release OH− ions. Example: NaOH→Na+ OH− Weak Base: Partially ionises in water. Example: NH3+H2O⇋NH4+ OH- 2. Key Properties of Strong vs. Weak Acids/Bases Property Strong Acid/Base Weak Acid/Base Degree of Ionization 100% ionized in water. Partial ionization (low percentage). Electrical High (more ions in Low (fewer ions in Conductivity solution). solution). pH Strong acids: Very low Higher pH for acids pH (e.g., 1–2). (e.g., 4–6). Strong bases: Very Lower pH for bases high pH (e.g., 12–14). (e.g., 8–10). Reactivity Very reactive (reacts Slower reactions with quickly). other substances. 3. Examples of Strong and Weak Acids/Bases Strong Acids: ○ Hydrochloric acid (HCl) HClHCl ○ Sulfuric acid (H2SO4) ○ Nitric acid (HNO3)  Weak Acids: ○ Acetic acid (CH3COOH) ○ Carbonic acid (H2CO3) ○ Hydrofluoric acid (HF) Strong Bases: ○ Sodium hydroxide (NaOH) ○ Potassium hydroxide (KOH) Weak Bases: ○ Ammonia (NH3) ○ Aluminum hydroxide (Al(OH)3) 4. How to Distinguish Between Strong and Weak Acids/Bases 1. pH Measurement: ○ Strong acids have a lower pH; strong bases have a higher pH. 2. Electrical Conductivity: ○ Strong acids/bases conduct electricity better due to more ions. 3. Reaction Rate: ○ Strong acids/bases react faster with metals, carbonates, or other substances. 4. Titration Curve: ○ Strong acids/bases produce steeper titration curves compared to weak ones. 1. Concentrated and Dilute Acids and Bases Concentrated Acid/Base: Contains many acid or base molecules per solution unit (less water).Example: Concentrated HCl (~12 M). HClHCl Dilute Acid/Base: A low number of acid or base molecules per solution unit (more water).Example: Dilute HCl (~0.1 M). HClHCl Strength vs. Concentration: Strength depends on ionisation; concentration depends on the amount of solute in the solution. ○ Example: Dilute NaOH is still a strong base because it completely ionises. NaOHNaOH 2. Acids and Bases in the Periodic Table - Oxides Types of Oxides:. Acidic Oxides: ○ Non-metal oxides (e.g., CO2, SO2, P2O5). CO2,SO2,P2O5CO_2, SO_2, P_2O_5 ○ React with water to form acids. ○ Example: SO2+H2O→H2SO3.. Basic Oxides: ○ Metal oxides (e.g., Na2 O, MgO). Na2O,MgONa_2O, MgO ○ React with water to form bases. Example: Na2O+H2O→2NaOH.. Amphoteric Oxides: ○ React with both acids and bases. ○ Example: Al 2O3, ZnO.Al2O3+6HCl→2AlCl3+3H2O. Al2O3,ZnOAl_2O_3, ZnO Al2O3+6HCl→2AlCl3+3H2O Al2O3+2NaOH→2NaAlO2+H2O. Neutral Oxides: ○ Do not react with acids or bases. ○ Example: CO,N2O. Trends Across the Periodic Table: Oxides become more acidic across a period (left to right). Example: Na2O (basic) → SiO2 (amphoteric) → SO3 (acidic). Oxides become more basic down an oof for metals. 3. pH Scale and Indicators pH Scale: Measures the acidity or alkalinity of a solution on a scale of 0–14. ○ 7pH7: Basic Indicators: Substances that change colour depending on pH.Examples: ○ Litmus: Red in acid, blue in base. ○ Phenolphthalein: Colorless in acid, pink in base. ○ Methyl orange: Red in acid, yellow in base. 4. Neutralisation Reactions and Titrations Neutralisation Reaction: The acid reacts with a base to produce salt and water. Example: HCl+NaOH→NaCl+H2O. Titrations: A technique to determine unknown concentrations by reacting acids with bases. Steps for Titration:. Add a known volume of acid/base to a conical flask.. Add an indicator (e.g., phenolphthalein).. Slowly add the titrant (acid/base) from a burette while swirling until the endpoint (colour change).. Record the volume of titrant used. Key Formula: C1V1=C2V2 Lab Example:. Reaction: NaOH+HCl→NaCl+H2O.. Calculate the concentration of NaOH using titration data: 5. Summary Table of Acids/Bases in Reactions Reaction Example Result Acid + Metal 2HCl+Mg→MgCl2+H2 Salt + Hydrogen gas Acid + Metal Oxide H2SO4+CuO→CuSO4 Salt + Water +H2O Acid + Metal HCl+Na2CO3→2NaCl+ Salt + Water + Carbon Carbonate H2O+CO2 dioxide Base + Non-metal 2NaOH+CO2→Na2CO Salt + Water Oxide 3+H2O Impact of Acid Deposition on the Environment 1. What is Acid Deposition? Acid deposition is the process by which acidic particles, gases, and precipitation (acid rain, snow, or fog) fall to the ground. Acid rain has an explicit pH lower than 5.6. 2. Causes of Acid Deposition Acid rain is caused by human activities that release sulfur dioxide (SO₂) and nitrogen oxides (NOₓ) into the atmosphere. These gases react with water, oxygen, and other chemicals to form sulfuric acid (H₂SO₄) and nitric acid (HNO₃). Key Sources of SO₂ and NOₓ: Burning of fossil fuels: Coal-fired power plants, industrial factories, and vehicles release SO₂ and NOₓ. Volcanic eruptions: Emit SO₂ naturally. Agricultural activities: Fertiliser use and livestock farming produce ammonia, which can combine with acidic compounds. Reactions Involved:. SO₂+O₂→SO3 SO3+H2O→H2SO4. NOx+H2O→HNO3 3. Effects of Acid Deposition on the Environment. Soil Acidification ○ Acid rain washes away essential nutrients like calcium, magnesium, and potassium from the soil. ○ Releases harmful ions like aluminium, which can damage plant roots.. Water Bodies (Lakes, Rivers, and Streams) ○ Lowered pH levels lead to acidic waters, killing aquatic life like fish, algae, and invertebrates. ○ Reduced biodiversity as sensitive species fail to survive.. Forests ○ Weakens trees by leaching nutrients from the soil and washing away protective coatings on leaves. ○ Higher vulnerability to diseases, pests, and harsh weather.. Buildings and Infrastructure ○ Acid rain corrodes materials like limestone, marble, and metal, accelerating the decay of historical monuments and buildings.. Human Health ○ Increases delicate particulate matter in the air, causing respiratory and cardiovascular issues. 4. Long-Term Effects Global Ecosystems: Reduced biodiversity due to species loss in acidic environments. Economic Impact: High costs of restoring soil and water health, repairing damaged structures, and treating health problems. 5. Mitigation Strategies. Reduce Emissions ○ Use cleaner energy sources like renewable energy (solar, wind, and hydropower). ○ Install scrubbers in factories to capture SO₂ and NOₓ emissions.. International Agreements ○ Protocols like the Clean Air Act and the Gothenburg Protocol aim to reduce transboundary pollution.. Liming of Lakes and Soil ○ Adding limestone neutralises acid in lakes and soil.. Public Awareness ○ Promoting energy efficiency and educating people on reducing fossil fuel consumption. Formation of Soluble and Insoluble Salts 1. What Are Salts? Definition: Salts are ionic compounds formed by the neutralisation reaction between an acid and a base. Example: HCl+NaOH→NaCl+H2O\text{HCl} + \text{NaOH} → \text{NaCl} + \text{H}_2\text{O}HCl+NaOH→NaCl+H2O. Salts can be soluble (dissolve in water) or insoluble (do not dissolve in water). 2. Soluble and Insoluble Salts General Solubility Rules:. Soluble Salts: ○ Salts containing alkali metals (Group 1) or ammonium ions (NH4+\text{NH}_4^+NH4+) are soluble. ○ Salts of nitrates (NO^-_3), acetates (CH3COO−), and chlorates (ClO3−) are soluble. ○ Most chlorides, bromides, and iodides are soluble, except for Ag⁺, Pb²⁺, and Hg₂²⁺. ○ Most sulfates (SO^2−_4) are soluble, except for those of Ba²⁺, Sr²⁺, Pb²⁺, and Ca²⁺ (partially soluble).. Insoluble Salts: ○ Carbonates (CO^2−_3), phosphates (PO^3−_4), and hydroxides (OH−) are generally insoluble, except for alkali metals and NH4+. ○ Sulfides (S^2−) are generally insoluble except for alkali metals, NH4+, and alkaline earth metals. 3. Net Ionic Equations for Precipitation Reactions Definition: A precipitation reaction occurs when two aqueous solutions react to form an insoluble solid (precipitate). Steps to Write Net Ionic Equations:. Write the balanced molecular equation. Example: BaCl2(aq)+Na2SO4(aq)→BaSO4(s)+2NaCl(aq)\. Separate the soluble ionic compounds into their ions (complete ionic equation). Ba2+(aq)+2Cl−(aq)+2Na+(aq)+SO42−(aq)→BaSO4(s)+2Na+(aq) +2Cl−(aq). Remove the spectator ions (ions that remain unchanged on both sides). Net ionic equation: Ba2+(aq)+SO42−(aq)→BaSO4(s) 4. Practical Applications of Solubility Principles. Water Treatment ○ Removal of Hardness: Adding lime or soda ash precipitates insoluble salts like calcium carbonate (CaCO3) and magnesium hydroxide (Mg(OH)2). ○ Flocculation: Precipitating impurities like iron or aluminium salts from water.. Medicine ○ Insoluble salts like barium sulfate (BaSO4) are used for X-ray imaging as they are non-toxic and opaque to X-rays.. Agriculture  ○ Soluble salts (e.g., ammonium nitrate) are used in fertilisers to supply plant nutrients.. Chemical Analysis ○ Precipitation reactions help identify ions in a solution. For example, adding silver nitrate (AgNO3) to a solution can detect halides like Cl− by forming a white silver chloride (AgCl) precipitate.. Industrial Uses ○ Solubility principles extract metals, produce pigments, and remove industrial waste products. 5. Laboratory Titrations for Soluble Salts Titrations help determine the concentration of acids, bases, or salts in a solution: Example: HCl+NaOH→NaCl+H2O Steps:. Add a standard solution (known concentration) of NaOH to a burette.. Add a known volume of HCl to a conical flask with a few drops of indicator (e.g., phenolphthalein).. Slowly add NaOH to HCl until the indicator changes colour.. Record the volume of NaOH used and calculate the concentration of HCl. Formula: Concentration of acid or base=Moles/Volume (dm3)\ Preparation and Testing of Various Gases Gas Preparation Reaction Test Observatio Method n Hydrogen Reacting Zn (s) Bring a lit Produces a (H2) dilute acid +2HCl (aq)→ splint near “pop” with a metal. ZnCl2(aq) the test sound, Example: +H2(g) tube. confirming zinc and hydrogen. hydrochloric acid. Oxygen Decompositi 2H2O2(aq) Insert a The splint (O2) on of →2H2O (l) glowing relights, hydrogen +O2(g) splint into confirming peroxide the test oxygen. using tube. manganese dioxide as a catalyst. Carbon Reacting CaCO3(s) Bubble the Limewater Dioxide calcium +2HCl (aq)→ gas through turns milky, (CO2 ) carbonate CaCl2(aq) limewater confirming  manganese dioxide as a catalyst. Carbon Reacting CaCO3(s) Bubble the Limewater Dioxide calcium +2HCl (aq)→ gas through turns milky, (CO2 ) carbonate CaCl2(aq) limewater confirming with dilute +H2O (l) (Ca(OH)2 ). carbon hydrochloric +CO2(g) dioxide. acid. Chlorine Reacting MnO2(s) Hold damp Litmus (Cl2) concentrate +4HCl (aq)→ blue litmus paper turns d MnCl2(aq) paper near red, then hydrochloric +2H2O (l) the gas. white acid with +Cl2(g) (bleached), manganese confirming dioxide. chlorine. Ammonia Heating an NH4Cl (s) Hold damp Litmus (NH3 ) ammonium +NaOH (aq) red litmus paper turns salt with an →NaCl (aq) paper near blue, alkali. +H2O (l) the gas. confirming +NH3(g) ammonia. Safety Precautions Gas Precautions Hydrogen (H2 ) Avoid open flames as hydrogen is highly flammable and explosive. Oxygen (O2 ) Ensure proper ventilation as oxygen supports combustion and can increase fire risks. Carbon Dioxide (CO2 ) Avoid inhaling concentrated CO2 as it can cause suffocation. Chlorine (Cl2\2 ) Conduct experiments in a fume hood as chlorine is toxic and corrosive. Ammonia (NH3) Work in a well-ventilated area as ammonia has a pungent smell and can irritate the respiratory system. Reversible Reactions and Equilibrium Reversible Reactions: These chemical reactions proceed in both the forward and reverse directions. For example, N2(g) +3H2(g)⇌2NH3(g)N_2 (g) + 3H_2 (g) \rightleftharpoons 2NH_3 (g)N2(g)+3H2(g)⇌2NH3(g). At equiEquilibriume, the forward reaction rate equals the reverse reaction rate, and the concentrations of reactants and products remain constant. Dynamic Equilibrium Occurs in a closed system where the reaction continues at a molecular level, but macroscopic properties like  concentration and pressure stay constant. For example, the dissolution of sugar in water reaches equilibrium when no more sugar dissolves. Factors Affecting the Position of Equilibrium (Chatelier'sr's Principle) Factor Change Made Effect on Reason Equilibrium Concentration Increase in Shifts to the More reactants reactants right (forward favor the reaction) formation of products to reduce the disturbance. Increase in Shifts to the left More products products (reverse favor the reaction) formation of reactants to restore balance. Temperature Increase in Shifts to the Endothermic temperature right (forward reactions (endothermic) reaction) absorb heat, so adding heat favors the forward reaction. Increase in Shifts to the left Exothermic temperature (reverse reactions (exothermic) reaction) release heat, so adding heat favors the reverse reaction. Pressure Increase in Shifts toward Fewer gas (gases) pressure the side with molecules fewer gas reduce molecules pressure, counteracting the change. Decrease in Shifts toward More gas pressure the side with molecules more gas increase molecules pressure to counteract the change. Catalyst Catalyst added No change in A catalyst  molecules pressure, counteracting the change. Decrease in Shifts toward More gas pressure the side with molecules more gas increase molecules pressure to counteract the change. Catalyst Catalyst added No change in A catalyst equilibrium speeds up both position forward and reverse reactions equally, reaching equilibrium faster. Example Applications:. Haber Process: N2+3H2⇌2NH3 ○ High-pressure ffavoursNH3 production due to fewer gas molecules on the product side. ○ Moderate temperature ensures the balance between reaction rate and equilibrium yield.. Contact Process: 2SO2+O2⇌2SO3 ○ High pressure and low temperature maximise SO3 production, but temperature must be optimised to maintain reaction speed. 3. Applying Chatelier'sr's Principle Predicting the Direction of Equilibrium Shift. Increase in Reactants: The equilibrium shifts toward the products to counteract the change.. Increase in Products: The equilibrium shifts toward the reactants to reduce the concentration of products.. Decrease in Reactants: The equilibrium shifts toward the reactants to replenish the lost reactants.. Decrease in Products: The equilibrium shifts toward the products to restore balance.. Temperature Changes: ○ For endothermic reactions (+ΔH): Adding heat shifts equiEquilibriumard products. ○ For exothermic reactions (−ΔH): Adding heat shifts equiEquilibriumard reactants. Reason for Predicted Shifts According to Chatelier'sr's Principle, the system tries to counteract any imposed change to restore equiEquilibriumr example: For example:  ○ If the pressure increases, the system reduces the pressure by favouring the side with fewer gas molecules. ○ If heat is added, the system adjusts by favouring the reaction that absorbs heat. 4. Applying Chatelier'sr's Principle to Real-World Processes. Haber Process (N2+3H2⇌2NH3): ○ Pressure: High pressure favours the production of ammonia (NH3) fewer gas molecules areas on the product side. ○ Temperature: Moderate temperatures balance yield (favours lower temperature) and reaction rate (favours higher temperature). ○ Catalyst: Irospeedsed up the reaction without altering the equilibrium position.. Contact Process (2SO2+O2⇌2SO3): ○ Pressure: High pressure favors the production of sulfur trioxide (SO3) by reducing the number of gas molecules. ○ Temperature: Low temperatures favour the forward exothermic reaction but must be balanced with rate considerations. ○ Catalyst: Vanadium pentoxide (V2O5) is used to accelerate the reaction.. Industrial Carbon Dioxide Removal (CO2+H2O⇌H2CO3): ○ Used to remove excess CO2 in closed environments like submarines. ○ Adding a base shifts the equilibrium carbonate formation to reduce CO2.. Carbonated Beverages (CO2+H2O⇌H2CO3): ○ Bottling under high pressure shifts the equilibrium to dissolving CO2. ○ When the bottle is opened, pressure decreases, and CO2 escapes, shifting the equiEquilibrium Haber Process. 4. Haber Process Balanced Equation for Ammonia Production N2(g)+3H2(g)⇌2NH3(g) ΔH=−92 kJ/mol Optimal Conditions for the Haber Process Condition Explanation Temperature Moderate temperature ( 400−500°C): Higher temperatures increase the rate but reduce yield due to the exothermic nature of the reaction. Pressure High pressure ( 200atm): Favours the formation of ammonia since there are fewer gas molecules on the product side.  temperatures increase the rate but reduce yield due to the exothermic nature of the reaction. Pressure High pressure ( 200atm): Favours the formation of ammonia since there are fewer gas molecules on the product side. Catalyst Iron catalyst with potassium/ aluminum oxide promoters: Speeds up the reaction without affecting equilibrium. Importance of Ammonia in Agriculture and Industry. Agriculture: ○ Ammonia is key in producing fertilisers like ammonium nitrate and urea, boosting crop yields. ○ Essential for feeding the growing global population.. Industry: ○ Used in the production of explosives, plastics, and dyes. ○ Essential in refrigeration systems and water treatment. 5. Environmental and Economic Impact of the Haber Process Environmental Concerns. Greenhouse Gas Emissions: ○ Hydrogen is typically derived from natural gas, releasing large amounts of CO2.. Eutrophication: ○ Excess fertilisers from ammonia runoff into water bodies cause algal blooms, depleting oxygen and harming aquatic life.. Energy Consumption: ○ The process is highly energy-intensive, contributing to fossil fuel depletion. Economic Significance. Global Food Security: ○ Ensures large-scale food production, supporting global agriculture.. Industrial Growth: ○ Ammonia serves as a base for various industries, driving economic development. Trade-offs Benefits: Boosts agricultural productivity and industrial output. Drawbacks: Environmental harm due to emissions, energy use, and ecological imbalance. Solution: To minimise impacts, focus on sustainable methods like green hydrogen and precision farming. 6. Contact Process Balanced Equations for Sulfuric Acid Production. Combustion of Sulfur: S(s)+O2(g)→SO2(g). Oxidation of Sulfur Dioxide: 2SO2(g)+O2(g)⇌2SO3(g) ΔH=−196kJ/mol.. Formation of Sulfuric Acid: SO3(g)+H2O(l)→H2SO4(aq) Optimal Conditions for the Contact Process Condition Explanation Temperature Moderate temperature ( 450°C): Balances reaction rate and yield for the exothermic reaction. Pressure Low to moderate pressure (1−2atm): Sufficient for yield without high costs, as the reaction already favors products. Catalyst Vanadium(V) oxide (V2O5 ): Lowers activation energy, increasing the rate of SO2 oxidation. Importance of Sulfuric Acid. Industries: ○ Manufacturing fertilisers (e.g., superphosphate and ammonium sulfate). ○ Production of detergents, paints, and plastics.. Other Uses: ○ Oil refining, mineral processing, and wastewater treatment. ○ Essential in batteries (lead-acid batteries in vehicles). 7. Environmental and Economic Impact of the Contact Process Potential Environmental Hazards of Sulfuric Acid Production. Air Pollution: ○ If not controlled, sulfur dioxide (SO2) emissions contribute to acid rain, harming ecosystems and infrastructure.. Acid Rain: ○ SO2 reacts with water vapour to form sulfuric acid, which causes water bodies and soil acidification and impacts plant and animal life.. Waste Disposal: ○ Sulfuric acid by-products must be appropriately handled to avoid contamination of the environment. Measures to Mitigate the Environmental Impact. Desulfurisation of Gases: ○ Scrubbers are used to remove sulfur dioxide from exhaust gases before release.. Catalyst Recovery: ○ Recycle and reuse catalysts like vanadium oxide to reduce waste.. Water Treatment: ○ Treating sulfuric acid waste to neutralise its acidic properties before disposal.. Emission Controls: ○ Install filtration and treatment systems to capture and neutralise toxic emissions. Economic Importance of Sulfuric Acid. Critical Industrial Chemical: ○ Sulfuric acid is essential in producing fertilisers, detergents, and chemicals.. Global Manufacturing: ○ It supports industries like petrochemicals, paper, textiles, and pharmaceuticals.. Agricultural Sector: ○ Used in producing phosphate fertilisers and in soil treatment to improve crop yields.. Economic Growth: ○ Sulfuric acid drives economic activities by providing a backbone for many manufacturing processes. 8. Comparing Industrial Processes: Haber and Contact Process Aspect Haber Process Contact Process Primary Product Ammonia (NH3 ) Sulfuric Acid (H2SO4) Balanced Chemical N2+3H2⇌2NH3 S+O2⇌SO2 Equation Key Raw Materials Nitrogen (N2 ) from Sulfur (S) and oxygen the air, Hydrogen (H2) (O2) from natural gas or electrolysis Optimal Temperature 400-500°C 450°C Optimal Pressure 200 atm 1-2 atm Catalyst Iron with potassium/ Vanadium(V) oxide aluminum oxide (V2O) promoters Equilibrium Favors the production The reaction favors Consideration of ammonia at lower products, but temperatures but moderate pressure is requires high pressure used to keep cost to increase yield. down and still achieve good yields. Main Application Fertilizers (ammonium Fertilizers nitrate, urea), (superphosphate), explosives, and detergents, paints,  requires high pressure used to keep cost to increase yield. down and still achieve good yields. Main Application Fertilizers (ammonium Fertilizers nitrate, urea), (superphosphate), explosives, and detergents, paints, industrial chemicals. and petroleum refining. Environmental CO2 emissions from Sulfur dioxide Concerns hydrogen production emissions contributing and energy use, to acid rain. greenhouse gases. Economic Major role in global Sulfuric acid is a Significance agriculture (fertilizer critical industrial production) and chemical used in many industrial growth. industries, contributing to economic activities. 9. Practical Skills and Analysis Conduct Simple Experiments to Demonstrate Reversible Reactions Experiment Example 1: Copper Sulfate and Water ○ Observation: Dissolve copper(II) sulfate (CuSO4) in water. The solution turns blue. Evaporate the water, and the blue solid crystallises. On heating, the blue solid turns white (anhydrous copper sulfate), demonstrating the reversibility of the reaction. ○ Equation: CuSO4(aq)⇌CuSO4(s)+H2O(l) Experiment Example 2: Ammonium Chloride ○ Observation: Ammonium chloride (NH4Cl) is heated in a test tube, and the vapour condenses on the more remarkable part of the test tube. The solid and vapour phases of ammonium chloride show a reversible reaction. ○ Equation: NH4Cl(s)⇌NH3(g)+HCl(g) Analyse Data from Equilibrium Experiments to Determine Equilibrium Constants. Set-Up: ○ Set up a reversible reaction in a closed container. Record the concentration of reactants and products over time as the reaction progresses. ○ Measure the concentrations at equiEquilibriumuilibrium Expression:. Equilibrium Expression: ○ Use the equilibrium concentrations of reactants and products to calculate the equilibrium constant (Kc) using the formula: Kc=[Reactants][Products]Wheree square brackets represent the molar concentrations of each substance at equiEquilibriumterpretation:. Interpretation: ○ If Kc>1, products are favoured at equiEquilibrium Kc