Chem Lec Finals PDF

Redox Reactions (Oxidation-Reduction) Definition Redox reactions involve the transfer of electrons between species, causing changes in oxidation states. ○ Oxidation: Loss of electrons (increase in oxidation state). ○ Reduction: Gain of electrons (decrease in...

Redox Reactions (Oxidation-Reduction) Definition Redox reactions involve the transfer of electrons between species, causing changes in oxidation states. ○ Oxidation: Loss of electrons (increase in oxidation state). ○ Reduction: Gain of electrons (decrease in oxidation state). Oxidation and reduction occur simultaneously, with equal electron transfer: ○ Oxidizing Agent (OA): Accepts electrons (reduced). ○ Reducing Agent (RA): Donates electrons (oxidized). Significance of Redox Reactions 1. Understanding Reaction Mechanisms: ○ Tracks electron transfer, oxidation state changes, and reaction types (e.g., combustion, redox). 2. Predicting Reaction Outcomes: ○ Determines products, reaction spontaneity, and energy changes. 3. Practical Applications: ○ Synthesis, industrial processes, and pollutant removal. 4. Analytical Techniques: ○ Includes titration, electrochemistry, and spectroscopy. 5. Safety: ○ Recognizes hazards and ensures safe handling. Importance of Redox 1. Energy Production: ○ Cellular respiration and photosynthesis. 2. Corrosion: ○ Material degradation (e.g., rusting). 3. Environmental Processes: ○ Water purification, soil formation. 4. Industrial Uses: ○ Electroplating, batteries. Basic Concepts 1. Oxidation: Loss of electrons. 2. Reduction: Gain of electrons. 3. Oxidizing Agent: Gains electrons. 4. Reducing Agent: Donates electrons. Characteristics of Redox Reactions 1. Electron Transfer: Essential for the process. 2. Oxidation State Changes: Tracks electron movement. 3. Chemical Bonding: Formation or breaking of bonds. Identifying Redox Reactions 1. Look for oxidation state changes. 2. Identify electron transfer. 3. Determine OA and RA. Indicators of Redox Reactions 1. Color changes. 2. Gas evolution. 3. Precipitate formation. 4. Heat release. Oxidation Numbers (ON) Represent hypothetical charges for electron tracking. Rules: 1. Neutral atoms: ON = 0 (e.g., O₂, H₂). 2. Monatomic ions: ON = ion charge (e.g., Na⁺ = +1). 3. Oxygen: ON = -2 (except in peroxides, -1). 4. Hydrogen: ON = +1 (except in hydrides, -1). 5. Sum of ON in compounds = compound charge. Common Redox Examples Example 1: Methane Combustion Reaction: CH₄ + 2O₂ → CO₂ + 2H₂O OA: O₂ (reduced to H₂O). RA: CH₄ (oxidized to CO₂). Example 2: Zinc and Copper(II) Sulfate Reaction: Zn + CuSO₄ → ZnSO₄ + Cu OA: Cu²⁺ (reduced to Cu). RA: Zn (oxidized to Zn²⁺). Strong Oxidizing Agents 1. O₂, Cl₂, F₂, KMnO₄, HNO₃, H₂O₂. 2. Used in combustion, organic synthesis, and environmental processes. Strong Reducing Agents 1. Metals: Na, Mg, Zn. 2. Nonmetals: H₂, C, S. 3. Compounds: LiAlH₄, NaBH₄. 4. Used in synthesis, pollutant removal, and biological processes. Summary 1. OIL RIG: Oxidation Is Loss, Reduction Is Gain. 2. Oxidizing agents gain electrons, causing reduction. 3. Reducing agents lose electrons, causing oxidation. 4. Balance equations to identify OA/RA and electron flow. Electrochemistry 1. Electrolysis Definition: A process where electrical energy is used to induce a chemical change, often decomposing substances or plating metals. Components: ○ Electrolytic Cell: Consists of a container, electrolyte, cathode (negative electrode), and anode (positive electrode). ○ Process: Reduction occurs at the cathode: Positive ions gain electrons. Oxidation occurs at the anode: Negative ions lose electrons. ○ Used for electroplating (e.g., silver utensils) and extraction of metals like magnesium and calcium. Electroplating: Example: Coating steel with tin for tin cans. Process: ○ Cathode: Object to be plated. ○ Anode: Metal to be deposited. ○ Electrolyte: Solution containing metal ions. Table: Common Electroplating Processes Metal Anode Electrolyte Applications Copper Cu CuSO₄ solution Electrotyping Silver Ag AgCN, KCN, K₃CO₃ solutions Tableware, jewelry Chromium Pb Cr₂O₃, H₂SO₄ solutions Automotive parts Tin Sn SnSO₄, sulfuric acid Tin cans 2. Electrochemical Cells Definition: Produce electrical energy from spontaneous redox reactions. Components: ○ Two half-cells with electrodes and electrolytes connected by a salt bridge. ○ Example: Zn (oxidation at anode) and Cu (reduction at cathode) in a galvanic cell. Standard Reduction Potentials (SRP): ○ Metals with high SRP are strong oxidizing agents (e.g., Cu). ○ Metals with low SRP are strong reducing agents (e.g., Zn). 3. Practical Cells (Batteries) Primary Cells:  Non-rechargeable; e.g., dry cells. ○ Anode: Zinc. ○ Cathode: Carbon. ○ Electrolyte: NH₄Cl, ZnCl₂. Secondary Cells: Rechargeable; e.g., lead storage batteries and nickel-cadmium batteries. ○ Lead Battery: Used in vehicles. Anode: Pb. Cathode: PbO₂. Electrolyte: H₂SO₄. ○ Nickel-Cadmium Battery: Anode: Cd. Cathode: NiO₂. Electrolyte: KOH. Common in power tools and mobile devices. Fuel Cells: Generate electricity continuously with reactant supply. ○ Example: H₂-O₂ fuel cells in spacecraft. ○ Reaction: Anode: H₂ → 2H⁺ + 2e⁻. Cathode: O₂ + 4H⁺ + 4e⁻ → 2H₂O. ○ Efficient but costly. 4. Corrosion Definition: Deterioration of metals through electrochemical processes. ○ Example: Rusting of iron. Anode: Fe → Fe²⁺ + 2e⁻. Cathode: O₂ + 4H⁺ + 4e⁻ → 2H₂O. Overall: 4Fe + 3O₂ + 6H₂O → 4Fe(OH)₃. Accelerated by: ○ Electrolytes (e.g., seawater). ○ Acidic environments and pollutants. Protection: ○ Cathodic Protection: Use of sacrificial anodes like Zn or Mg. ○ Formation of passive layers (e.g., ZnCO₃ on zinc, CuO on copper). Chemistry of Engineering Materials Maria Rhodalita M. Chiong, RCHE, MSES What is Chemistry? Chemistry is the science that deals with matter: its composition, structure, properties, changes, and the laws governing these changes. What are Engineering Materials? Engineering materials are materials used in the construction of manmade structures and components. Examples of Engineering Materials 1. Ferrous Metals: ○ Stainless Steel: Pipes, bars, cookware, screws, wires, utensils ○ Ductile Cast Iron: Pipes, ranger, castings 2. Non-Ferrous Metals: ○ Aluminum: Foils, cans, aerospace parts ○ Zinc-Aluminum Alloy: Coated sheets, incense burners ○ Nickel, Copper, Tin, Lead, Titanium: Various applications Primary Function of Engineering Materials To withstand applied loading without breaking or exhibiting excessive deflection. Importance of Material Science and Engineering Enables engineers to develop and select materials for specialized applications: 1. Mechanical Engineers: High-temperature and wear-resistant materials. 2. Electrical Engineers: Materials for faster electrical devices with minimal power loss. 3. Aerospace/Automobile Engineers: Materials with high strength-to-weight ratios. 4. Civil Engineers: Strong materials like steel and reinforced concrete. 5. Chemical Engineers: Corrosion-resistant materials. Best Material for Engineering Purposes The best material serves the desired objective at minimal cost. Factors to Consider: 1. Availability  2. Suitability for working conditions 3. Cost Classification of Engineering Materials 1. General Properties 2. Nature of Materials 3. Application Types of Engineering Materials 1. Metals ○ Ferrous Metals: Stainless steel, cast iron, wrought iron ○ Non-Ferrous Metals: Aluminum, copper, zinc, tin, lead, titanium 2. Non-Metallic Materials ○ Synthetic Materials ○ Natural Materials Non-Spontaneous Redox Reactions Non-spontaneous redox reactions require external energy to proceed and are characterized by unfavorable thermodynamic or kinetic conditions. Thermodynamic Conditions ΔG > 0: Positive Gibbs Free Energy indicates the reaction is non-spontaneous and requires energy input. E°cell < 0: Negative standard cell potential shows that reduction at the cathode is unfavorable. Kinetic Barriers 1. High Activation Energy: Significant energy is required to initiate the reaction. 2. Catalyst Inhibition: Lack of a catalyst or its inhibition slows down the reaction. Concentration and Pressure Effects 1. Low Reactant Concentrations: Insufficient reactants prevent the reaction from proceeding. 2. High Product Concentrations: Equilibrium shifts toward reactants. 3. Non-Optimal Pressure: Deviations from standard pressure reduce reaction feasibility. Electrochemical Considerations 1. Non-Standard Electrode Potentials: Factors like concentration, temperature, or pH alter E° values. 2. Electrode Polarization: Resistance at electrode surfaces hinders electron transfer. Examples of Non-Spontaneous Reactions 1. Zn + Cu²⁺ → Zn²⁺ + Cu (E°cell = -1.10 V) 2. Fe + 2H⁺ → Fe²⁺ + H₂ (E°cell = -0.44 V) 3. Pb + Hg²⁺ → Pb²⁺ + Hg (E°cell = -0.91 V) Driving Non-Spontaneous Reactions 1. External Voltage: Electrolysis supplies energy to drive reactions. 2. Catalysts: Lower activation energy to enhance reaction rates. 3. Optimized Conditions: Control temperature, pressure, and reactant concentrations. 4. Reaction Coupling: Combine with spontaneous reactions to enable the process. Reversibility of Non-Spontaneous Reactions Under specific conditions, non-spontaneous reactions can become reversible. Conditions for Reversibility 1. External Energy Input: Applying voltage can drive the reaction backward. 2. Concentration Adjustments: Altering concentrations shifts the equilibrium. 3. Temperature and Pressure Changes: Modify these to favor reversibility. 4. Catalysts: Facilitate reverse reactions. Examples of Reversible Reactions 1. Lead-Acid Battery: Pb + H₂SO₄ ⇌ PbSO₄ + H₂ 2. Lithium-Ion Battery: Li⁺ + FePO₄ ⇌ LiFePO₄ 3. Water Electrolysis: 2H₂O ⇌ 2H₂ + O₂ Spontaneous Redox Reactions Spontaneous redox reactions release energy and occur naturally under favorable thermodynamic and kinetic conditions. Thermodynamic Conditions 1. ΔG < 0: Negative Gibbs Free Energy indicates spontaneity. 2. E°cell > 0: Positive cell potential reflects favorable oxidation and reduction. Kinetic Factors 1. Low Activation Energy: Reaction proceeds easily. 2. Catalyst Presence: Speeds up the reaction. 3. Optimal Conditions: Favorable temperature and pressure. Concentration and Equilibrium Effects 1. High Reactant Concentrations: Drives the reaction forward. 2. Low Product Concentrations: Shifts equilibrium to form products. 3. K > 1: Indicates a spontaneous reaction. Examples of Spontaneous Reactions  1. Zn + Cu²⁺ → Zn²⁺ + Cu (E°cell = +1.10 V) 2. Fe + O₂ → Fe₂O₃ (E°cell = +1.65 V) 3. Al + O₂ → Al₂O₃ (E°cell = +1.68 V) 4. Mg + H₂O → Mg(OH)₂ + H₂ (E°cell = +2.37 V) Applications of Spontaneous Reactions 1. Energy Generation: ○ Combustion reactions power engines. ○ Electrochemical cells drive batteries. 2. Environmental Processes: ○ Corrosion of metals (e.g., rusting of iron). ○ Atmospheric redox reactions (e.g., ozone formation). 3. Biological Systems: ○ Cellular respiration (mitochondrial energy generation). ○ Photosynthesis (chloroplast-based redox reactions). 4. Geological and Aquatic Systems: ○ Oxidation of metals in ores. ○ Sediment redox gradients in water. Comparison of Spontaneous and Non-Spontaneous Reactions Feature Spontaneous Reactions Non-Spontaneous Reactions ΔG Negative Positive E°cell Positive Negative Energy Requirement No external energy required Requires external energy input Examples Combustion, corrosion, batteries Electrolysis, forced redox reactions Nuclear Chemistry Definition: Nuclear chemistry is the branch of chemistry focused on the study of nuclear reactions, radioactivity, and the properties and behavior of atomic nuclei. Scope of Nuclear Chemistry Nuclear chemistry encompasses: 1. Radioactive Decay and Stability: Understanding how and why nuclei decay and determining their stability. 2. Nuclear Reactions: Includes fission, fusion, and neutron-induced reactions. 3. Nuclear Synthesis and Transmutation: Creating new elements or isotopes through nuclear processes. 4. Radiation Detection and Protection: Measuring radiation and ensuring safety from harmful exposure. 5. Nuclear Energy Applications: Power generation through nuclear fission and potentially nuclear fusion. 6. Nuclear Medicine and Radiopharmaceuticals: Using radioactive isotopes for diagnostics and treatment. 7. Environmental and Analytical Aspects: Examining the impact of nuclear processes and using isotopes for analytical techniques. Types of Nuclear Reactions 1. Nuclear Fission A heavy nucleus splits into two or more lighter nuclei. Releases a significant amount of energy and additional neutrons. Applications: ○ Nuclear power plants (controlled fission for electricity). ○ Atomic bombs (uncontrolled chain reactions). 2. Nuclear Fusion Two light nuclei combine to form a heavier nucleus. Releases tremendous energy, as seen in stars. Applications: ○ The focus of ongoing research for clean, abundant energy. ○ Powers thermonuclear weapons (fusion bombs). 3. Radioactive Decay Unstable nuclei spontaneously emit particles or radiation to become stable. Types of Decay: ○ Alpha decay: Emission of helium nuclei. ○ Beta decay: Conversion of a neutron into a proton (or vice versa) with electron or positron emission. ○ Gamma decay: Emission of high-energy photons. Applications: ○ Medical imaging (e.g., PET scans). ○ Carbon dating for determining the age of artifacts. Key Principles of Nuclear Chemistry 1. Energy Release Nuclear reactions release vastly more energy than chemical reactions due to the strong nuclear force. 2. Mass-Energy Equivalence According to Einstein's equation (E=mc²), a small amount of mass is converted into a large amount of energy in nuclear reactions. 3. Nuclear Waste Byproducts of nuclear reactions are often radioactive and require careful management and disposal to minimize environmental and health risks. Applications of Nuclear Chemistry 1. Nuclear Power Plants: ○ Controlled nuclear fission generates electricity. ○ Provides a low-carbon energy source but produces long-lived radioactive waste. 2. Nuclear Medicine: ○ Radioisotopes are used for diagnosis (e.g., technetium-99m) and treatment (e.g., iodine-131 for thyroid disorders). 3. Scientific Research: ○ Studying fundamental nuclear properties and behaviors. 4. Carbon Dating: ○ Determining the age of fossils and artifacts using radioactive isotopes like carbon-14. 5. Industrial Applications: ○ Radiation is used in materials testing, sterilization, and power generation in remote areas (e.g., RTGs in space exploration). Additional Considerations Nuclear Weapons Utilize uncontrolled nuclear chain reactions (fission or fusion). Have immense destructive power but carry significant ethical and security concerns. Nuclear Fusion Power Potential Benefits: ○ Clean, nearly limitless energy source. ○ Minimal radioactive waste compared to fission. Challenges: ○ Achieving and maintaining the extreme temperatures and pressures required for fusion.  ○ Research is ongoing (e.g., ITER project). Environmental Impact Radioactive materials can contaminate the environment if not properly managed. Safe disposal methods and advances in nuclear waste recycling are critical. Summary Nuclear chemistry involves understanding the immense power of atomic nuclei, which can be harnessed for beneficial purposes like energy generation and medical advancements or pose significant risks, such as nuclear weapons and radioactive waste. Mastery of its principles enables responsible and sustainable use, helping address global energy challenges while mitigating potential hazards.