Chemistry Cheatsheet

Cheatsheet Content

1. Materials for Energy Devices Semiconductors: Bridging Conductors and Insulators What they are: Materials that conduct electricity better than insulators but not as well as conductors. Their conductivity can be controlled. How they work: Their atoms have electron "holes" or extra electrons that can move when energy is applied. Types: Organic Semiconductors: Made from carbon-based molecules or polymers. Electrons move through overlapping atomic orbitals called $\pi$-orbitals. Examples: Used in flexible screens (OLEDs) and solar cells (photovoltaics). Inorganic Semiconductors: Not carbon-based (e.g., Silicon, Germanium, Gallium Arsenide). They have a fixed, crystalline structure. Examples: Found in computer chips (transistors, diodes). p-type Semiconductors (Hole Conductors): These materials conduct electricity mainly through "holes" (missing electrons that act like positive charges). Pentacene: A common organic p-type semiconductor. Properties: Excellent at carrying positive charge (holes). Absorbs visible light strongly. Sensitive to air (oxygen) and moisture, which can degrade it. Uses: Organic transistors (OFETs), flexible electronics. Structure: A flat, ladder-like molecule made of 5 benzene rings. Pentacene Structure (Conceptual) n-type Semiconductors (Electron Conductors): These materials conduct electricity mainly through extra electrons. Perfluoropentacene (PFP): A modified version of pentacene where hydrogen atoms are replaced by fluorine. Properties: Attracts electrons strongly (high electron affinity). Good electron mobility (electrons move easily). Uses: N-type components in transistors, memory devices. Energy Storage Devices: Batteries What they are: Devices that store chemical energy and convert it into electrical energy. They are made of one or more "cells." Key Parts of a Battery Cell: Anode (Negative Electrode): Where oxidation happens (loses electrons). Cathode (Positive Electrode): Where reduction happens (gains electrons). Electrolyte: A chemical medium (liquid or gel) that allows ions (charged atoms/molecules) to move between the anode and cathode. Separator: A porous barrier that keeps the anode and cathode from touching (preventing short circuits) but lets ions pass through. Battery Classification: Primary (Non-rechargeable) Batteries: Used once and discarded. The chemical reactions are irreversible. Examples: Standard AA/AAA dry cells, lithium-manganese dioxide batteries. Secondary (Rechargeable) Batteries: Can be charged and discharged many times. The chemical reactions are reversible. Examples: Car batteries (lead-acid), laptop/phone batteries (Li-ion). Reserve Batteries: Stored with key components (like the electrolyte) kept separate until needed, giving them a very long shelf life. Important Characteristics: Capacity (measured in Amp-hours, $Ah$): How much current the battery can deliver over a certain time. A higher capacity means it lasts longer. Power Density (measured in Watts per kilogram, $W/kg$): How much power the battery can deliver relative to its weight or volume. High power density means it can deliver a lot of power quickly. Cycle Life: How many times a rechargeable battery can be charged and discharged before its performance significantly degrades. Shelf Life: How long a battery can be stored without losing significant charge or performance. Lithium-ion Battery (LIB): The Rechargeable Workhorse How it works: Lithium ions ($Li^+$) move back and forth between the anode and cathode during charging and discharging. This is called "intercalation" – the ions insert themselves into the electrode materials without changing their basic structure. Key Components: Anode: Usually graphite (a form of carbon). Cathode: Often lithium cobalt oxide ($LiCoO_2$) or other lithium metal oxides. Electrolyte: A lithium salt (like $LiPF_6$) dissolved in an organic solvent. Charging Process (Putting energy in): At the cathode: Lithium ions leave the cathode material. $LiCoO_2 \rightarrow Li_{(1-x)}CoO_2 + xLi^+ + xe^-$ At the anode: Lithium ions and electrons move into the graphite. $xC_6 + xLi^+ + xe^- \rightarrow xLiC_6$ Overall: $LiCoO_2 + xC_6 \rightarrow Li_{(1-x)}CoO_2 + xLiC_6$ (Lithium moves from cathode to anode) Discharging Process (Getting energy out): At the anode: Lithium ions leave the graphite. $xLiC_6 \rightarrow xC_6 + xLi^+ + xe^-$ At the cathode: Lithium ions and electrons move back into the cathode material. $Li_{(1-x)}CoO_2 + xLi^+ + xe^- \rightarrow LiCoO_2$ Overall: $Li_{(1-x)}CoO_2 + xC_6 \rightarrow LiCoO_2 + xC_6$ (Lithium moves from anode to cathode, generating current) Uses: Smartphones, laptops, electric vehicles (EVs), grid-scale energy storage. Supercapacitors: Fast Energy Storage What they are: Devices that store electrical energy by physically separating ions at the surface of electrode materials, rather than through chemical reactions (like batteries). They charge and discharge much faster than batteries. How they work: When voltage is applied, ions from the electrolyte move to the surfaces of the highly porous electrodes, forming an "electrical double-layer." This physical separation of charges stores energy. Types: Electrical Double-Layer Capacitors (EDLCs): Store energy purely by physical charge separation. Pseudocapacitors: Store energy through fast, reversible surface redox reactions in addition to charge separation. Hybrid Capacitors: Combine features of both EDLCs and pseudocapacitors, or combine with battery-like electrodes. Ultra-small Asymmetric Supercapacitor: A common type that uses two different electrode materials. Key Parts: Negative Electrode: Often an EDLC material like activated carbon (stores charge physically). Positive Electrode: Often a pseudocapacitive material like manganese dioxide ($MnO_2$) (stores charge physically and with quick surface reactions). Electrolyte: An ion-conducting solution (e.g., salt water or organic solutions). Charging Process (Example with $MnO_2$ and Carbon in $Na^+$ electrolyte): At the positive electrode ($MnO_2$): $MnO_2 + Na^+ + e^- \rightarrow MnO_2Na$ (Sodium ions interact with the $MnO_2$ surface). At the negative electrode (Carbon): $C + Na^+ + e^- \rightarrow C||Na^+$ (Sodium ions accumulate at the carbon surface). Discharging Process (Getting energy out): The reverse of charging. Ions move away from the electrode surfaces, releasing energy. Energy Conversion Devices Photovoltaic (PV) Cells (Solar Cells): Turning Light into Electricity How they work (Photovoltaic Effect): When light hits certain materials (semiconductors), it creates electron-hole pairs. These are then separated by an internal electric field, causing electrons to flow and generate electricity. Basic Structure: A sandwich of n-type and p-type semiconductor layers (a "p-n junction"), with metal contacts to collect the current. An anti-reflective coating helps absorb more light. Advantages: Clean, renewable, silent, minimal maintenance once installed. Disadvantages: High initial cost, only work when there's light, need battery storage for nighttime/cloudy days, require space. MEMS-Based Energy Harvesters: Capturing Ambient Energy What they are: Tiny (Micro-Electro-Mechanical Systems) devices that convert small amounts of ambient energy (like vibrations, heat, or radio waves) from the environment into usable electrical power. 1. Piezoelectric Harvesters: Principle: Certain materials (piezoelectrics) generate an electric charge when mechanically stressed or deformed. How it works: Vibrations or pressure cause the material to deform, shifting its internal charges and creating a voltage. 2. Electrostatic Harvesters: Principle: Use changes in electrical capacitance to generate power. How it works: Mechanical motion (e.g., vibration) moves two conductive plates closer or farther apart, changing their capacitance and generating voltage. ($C = \frac{\epsilon A}{d}$, where $C$ is capacitance, $A$ is area, $d$ is distance). 3. Electromagnetic Harvesters: Principle: Based on Faraday's Law of Induction – a changing magnetic field induces an electric current. How it works: Movement of a coil through a magnetic field (or vice versa) creates a changing magnetic flux, which generates electricity. Uses: Powering small, low-power devices like wireless sensors (IoT), wearable electronics, medical implants. 2. Corrosion Science and E-waste Management Corrosion Chemistry: The Degradation of Metals What it is: The natural process where refined metals convert into a more chemically stable form such as oxides, hydroxides, or sulfides. It's essentially the destruction of a metal due to a reaction with its environment. Electrochemical Theory of Corrosion (Focusing on Iron Rusting): Corrosion often happens like a tiny battery (galvanic cell) forming on the metal surface. Key Steps: Anode (Oxidation): The metal loses electrons and turns into ions. This is where the metal actually corrodes. For iron: $Fe \rightarrow Fe^{2+} + 2e^-$ (Iron loses electrons and becomes an iron ion). Cathode (Reduction): Electrons travel through the metal to another spot where they react with something in the environment (like oxygen or water). This spot does not corrode. If oxygen is present (most common for iron): Water and oxygen gain electrons to form hydroxide ions. $H_2O + \frac{1}{2}O_2 + 2e^- \rightarrow 2OH^-$ If oxygen is absent (e.g., in acidic solutions): Hydrogen ions gain electrons to form hydrogen gas. $2H^+ + 2e^- \rightarrow H_2 \uparrow$ Rust Formation (for Iron): The $Fe^{2+}$ ions from the anode then react with the $OH^-$ ions from the cathode to form iron hydroxide, which further reacts with oxygen to become rust ($Fe_2O_3 \cdot nH_2O$). Common Types of Corrosion: 1. Galvanic Corrosion (Two Different Metals): Happens when two different metals are connected and exposed to an electrolyte (like saltwater). The more "active" metal (the one that wants to lose electrons more easily) corrodes preferentially. Example: Steel corroding faster when in contact with copper. 2. Differential Aeration Corrosion (Different Oxygen Levels): Occurs when different parts of the *same* metal are exposed to different amounts of oxygen. The area with *less* oxygen becomes the anode and corrodes. Waterline Corrosion: A metal partially submerged in water corrodes most intensely just below the waterline, where oxygen is scarce. Pitting Corrosion: Small, deep holes form on the metal surface, often under a dirt particle. The area under the particle has less oxygen and becomes an anode. Corrosion Control: Protecting Metals I) Metallic Coatings (Barrier Protection): Covering a base metal with a layer of another metal. Galvanization: Coating iron or steel with a thin layer of zinc. Zinc is more active than iron, so it corrodes first (sacrificially) to protect the iron. Anodizing of Aluminum: An electrochemical process that thickens the natural oxide layer ($Al_2O_3$) on aluminum. This hard, protective oxide layer prevents further corrosion. III) Cathodic Protection (Making it a Cathode): A method to prevent corrosion by making the entire metal structure act as the cathode in an electrochemical cell. A cathode doesn't corrode. Sacrificial Anode Method: Connecting the metal to be protected (e.g., an underground pipeline) to a more active metal (like magnesium or zinc). The active metal acts as the anode and corrodes away, "sacrificing" itself to protect the main structure. Impressed Current Method: An external power source is used to supply electrons to the protected metal, making it a cathode. An inert anode (like graphite) is also buried nearby. This is used for larger structures. Corrosion Penetration Rate (CPR): How Fast Metal Degrades What it is: A measure of how quickly a metal is being eaten away by corrosion. It can be expressed as loss of thickness per year (e.g., millimeters per year, mmpy) or weight loss. Factors Affecting CPR: The type of metal, the environment it's in, and the properties of any corrosion products formed. Formula: $CPR = \frac{K \cdot W}{D \cdot A \cdot T}$ $W$: Weight loss of the metal. $D$: Density of the metal. $A$: Surface area exposed to corrosion. $T$: Time of exposure. $K$: A constant to get the units right (e.g., 87.6 for mmpy). How it's tested (Weight Loss Method): A pre-weighed metal sample is exposed to a corrosive environment for a set time, then cleaned and re-weighed. The weight difference is used to calculate the CPR. 3. Electrochemistry Electrochemical Cells: Chemical Energy $\leftrightarrow$ Electrical Energy What they are: Devices that convert chemical energy into electrical energy (or vice versa) through redox reactions. Redox Reactions (Reduction-Oxidation): The core of electrochemistry. Oxidation: A chemical species loses electrons and its oxidation number increases. Happens at the anode . Reduction: A chemical species gains electrons and its oxidation number decreases. Happens at the cathode . Types of Electrochemical Cells: Galvanic (Voltaic) Cell: Generates electricity from spontaneous chemical reactions (e.g., a battery discharging). Chemical energy $\rightarrow$ Electrical energy. Electrolytic Cell: Uses electrical energy to drive non-spontaneous chemical reactions (e.g., recharging a battery, electroplating). Electrical energy $\rightarrow$ Chemical energy. Electrode Potential: The electrical potential difference that develops between a metal electrode and its surrounding solution of ions. Electromotive Force (EMF) / Cell Potential ($E_{cell}$): The total voltage produced by an electrochemical cell. It's the difference between the cathode's potential and the anode's potential. $E_{cell} = E_{cathode} - E_{anode}$. Standard Electrode Potential ($E^0$): The electrode potential measured under standard conditions (25°C, 1 M concentration for solutions, 1 atm pressure for gases). Nernst Equation: Electrode Potential under Non-Standard Conditions What it does: It tells us how the electrode potential changes when concentrations or pressures are not at standard conditions. Equation: $E = E^0 - \frac{2.303RT}{nF} \log[M^{n+}]$ $E$: Non-standard electrode potential. $E^0$: Standard electrode potential. $R$: Gas constant. $T$: Temperature in Kelvin. $n$: Number of electrons transferred in the reaction. $F$: Faraday constant. $[M^{n+}]$: Concentration of the metal ions. Simplified at 25°C (298 K): $E = E^0 - \frac{0.0591}{n} \log[M^{n+}]$ Electrolyte Concentration Cell: Generating Electricity from Concentration Differences What it is: A special type of galvanic cell where both electrodes are identical, but they are immersed in solutions of the *same* electrolyte at *different* concentrations. How it works: The cell generates voltage as it tries to equalize the concentrations. The more dilute solution will act as the anode (oxidation), and the more concentrated solution as the cathode (reduction). Cell Potential: $E_{cell} = \frac{0.0591}{n} \log \frac{C_2}{C_1}$ (at 298K, where $C_2$ is the higher concentration and $C_1$ is the lower concentration). Types of Electrodes: The Interface of Reaction What they are: Conductors that make electrical contact with a non-metallic part of a circuit (like an electrolyte), where electron transfer (redox reactions) occurs. Common Types: Metal-Metal Ion Electrode: A metal rod dipped into a solution containing its own ions (e.g., Zinc metal in $ZnSO_4$ solution). Metal-Insoluble Salt Electrode: A metal coated with one of its sparingly soluble salts, immersed in a solution containing the anion of that salt (e.g., Calomel electrode: Mercury metal with mercurous chloride, $Hg_2Cl_2$, in KCl solution). Gas Electrode: An inert metal (like platinum) through which a gas is bubbled, while it's immersed in a solution containing ions related to the gas (e.g., Standard Hydrogen Electrode, SHE). Redox Electrode: An inert metal (Pt or Au) dipped into a solution containing both oxidized and reduced forms of a species (e.g., $Pt$ in a solution of $Fe^{2+}$ and $Fe^{3+}$ ions). Ion-Selective Electrode: A special electrode whose potential depends specifically on the concentration of a particular ion in the solution (e.g., Glass electrode for pH measurement). Reference Electrodes: Stable Benchmarks What they are: Electrodes with a very stable and known potential, used to measure the potential of other "indicator" electrodes. Primary Reference Electrode: Standard Hydrogen Electrode (SHE) Definition: A platinum electrode bathed in hydrogen gas at 1 atm, immersed in a 1 M $H^+$ solution. Its potential is *defined* as 0 Volts. Limitations: Difficult to set up and maintain precisely in a lab; sensitive to impurities. Secondary Reference Electrode: Calomel Electrode Definition: A more practical, commonly used reference electrode. It consists of mercury ($Hg$) in contact with mercurous chloride ($Hg_2Cl_2$, called calomel) and a potassium chloride ($KCl$) solution. Potential: Its potential depends on the concentration of the $KCl$ solution (e.g., saturated $KCl$ gives a stable potential). Uses: Widely used in labs for measuring single electrode potentials and in titrations. Glass Electrode: Measuring pH What it is: A type of ion-selective electrode specifically designed to measure the concentration of hydrogen ions ($H^+$), which determines pH. Construction: Consists of a thin, pH-sensitive glass membrane at the tip, enclosing a solution of known $HCl$ concentration and an internal reference electrode (often Ag/AgCl). How it works: When the glass membrane is immersed in a solution, $H^+$ ions exchange across the membrane, creating a potential difference that is directly proportional to the pH of the external solution. Equation: $E_G = E_G^0 - 0.0591 pH$ (The voltage $E_G$ changes by 0.0591 V for every pH unit change at 25°C). Conductometry: Measuring Solution Conductivity Principle: Measures the ability of an ionic solution to conduct electricity. More ions and more mobile ions mean higher conductivity. Conductance ($C$): The inverse of resistance ($R$). $C = \frac{1}{R}$. Conductivity ($\kappa$): A measure of how well a specific volume of solution conducts electricity. It depends on the concentration and type of ions. Conductometric Titration: A titration method where the equivalence point (the point where the reaction is complete) is determined by monitoring the change in solution conductivity as a titrant is added. How it works: As one ion is replaced by another during the reaction, the total conductivity of the solution changes. A plot of conductivity vs. titrant volume shows a distinct change in slope at the equivalence point. Example (Titrating a Weak Acid with a Strong Base): Initially, the weak acid has low conductivity. As base is added, the weak acid reacts to form a salt, increasing conductivity. After the equivalence point, excess strong base is added, which has very mobile ions, causing a sharp increase in conductivity. Potentiometric Estimations: Titration by Voltage Principle: A titration method where the potential (voltage) of an indicator electrode is measured relative to a reference electrode as the titrant is added. Advantage: More accurate and useful for colored or turbid solutions where visual indicators don't work well. Example (Estimating Ferrous Ammonium Sulphate, FAS, with Potassium Dichromate): This is a redox titration where $Fe^{2+}$ is oxidized to $Fe^{3+}$. Setup: A platinum electrode (indicator) and a calomel electrode (reference). Measurement: The voltage of the platinum electrode changes significantly as the ratio of $Fe^{3+}$ to $Fe^{2+}$ changes during the titration. Endpoint: A sharp jump in the measured voltage indicates the equivalence point, where all the $Fe^{2+}$ has been converted to $Fe^{3+}$. E-waste Management: Dealing with Electronic Waste What it is: Discarded electrical or electronic equipment that is no longer useful. Sources: Everything from old phones and laptops to refrigerators and light bulbs. Why it's a Problem: Environmental Harm: Contains toxic materials (lead, mercury, cadmium) that can leach into soil and water, polluting ecosystems. Health Risks: Improper handling exposes workers and communities to hazardous substances. Resource Depletion: Contains valuable materials (gold, silver, copper) that are lost if not recycled. Why Management is Needed: Protects the environment and human health. Recovers valuable resources, reducing the need for new mining. Saves energy compared to producing new materials from scratch. Extraction of Gold from E-waste (Bioleaching Method): What it is: A method that uses microorganisms (bacteria) to dissolve metals from e-waste, making it easier to recover valuable elements like gold. It's an eco-friendlier alternative to traditional harsh chemical methods. Simplified Steps: Preparation: E-waste is collected, shredded, and non-metal parts are removed. Bacterial Action: Special bacteria (e.g., Acidithiobacillus ferrooxidans ) are introduced in an acidic solution. These bacteria produce powerful oxidizing agents ($Fe^{3+}$ ions) and sulfuric acid. Base Metal Dissolution: The $Fe^{3+}$ ions and acid produced by bacteria dissolve common metals like copper. Example: $Cu + 2Fe^{3+} \rightarrow Cu^{2+} + 2Fe^{2+}$ (Copper dissolves) Gold Dissolution: Another type of bacterium (e.g., Chromobacterium violaceum ) can produce natural cyanide. This cyanide reacts with gold to form a soluble gold-cyanide complex, bringing the gold into solution. Example: $4Au + 8CN^- + O_2 + 2H_2O \rightarrow 4[Au(CN)_2]^- + 4OH^-$ Gold Recovery: Once gold is dissolved, it can be recovered from the solution using methods like: Cementation: Adding zinc dust, which displaces gold from the solution, causing gold to precipitate out. Electrodeposition: Using electricity to plate the dissolved gold onto a cathode.