Electrochemistry Notes PDF

A. Introduction to Electrochemistry Electrochemistry deals with the relationship between electrical energy and chemical reactions. It is fundamental in understanding how chemical reactions can produce electricity and how electrical energy can drive chemical changes. Key Concepts: 1. Redox Reactio...

A. Introduction to Electrochemistry Electrochemistry deals with the relationship between electrical energy and chemical reactions. It is fundamental in understanding how chemical reactions can produce electricity and how electrical energy can drive chemical changes. Key Concepts: 1. Redox Reactions: The basis of electrochemical processes, involving oxidation (loss of electrons) and reduction (gain of electrons). 2. Electrochemical Cells: Devices that convert chemical energy into electrical energy (galvanic cells or voltaic cells) or use electrical energy to drive a chemical reaction (electrolytic cells). Figure 1: Electrochemical Cell An electrochemical cell can either generate electricity from a spontaneous redox reaction like the galvanic cell or consume electricity to drive a nonspontaneous reaction thus requiring an external electron source like a DC battery or an AC power source like an electrolytic cell. Both galvanic and electrolytic cells consist of two electrodes (an anode and a cathode), which can be made of the same or different metals, and an electrolyte in which the two electrodes are immersed. GALVANIC CELLS Galvanic cells traditionally are used as sources of DC electrical power. A simple galvanic cell may contain only one electrolyte separated by a semi-porous membrane, while a more complex version involves two separate half-cells connected by a salt bridge. The salt bridge contains an inert electrolyte like potassium sulfate whose ions will diffuse into the separate half-cells to balance the building charges at the electrodes. According to the mnemonic “Red Cat An Ox”, oxidation occurs at the anode and reduction occurs at the cathode. Since the reaction at the anode is the source of electrons for the current, the anode is the negative terminal for the galvanic cell. Voltaic cells can be represented as follows for the zinc-copper cell In this representation, a single line ( | ) represents an interface at which a potential develops, that is, an electrode. It is conventional to write the anode half-cell on the left in this notation. ELECTROLYTIC CELLS An electrolytic cell is an electrochemical cell in which the energy from an external power source is used to drive a normally non-spontaneous reaction, i.e. apply a reverse voltage to a voltaic cell. We encounter electrolytic cells during the charging phase of any type of rechargeable battery from the lead-acid battery in automobiles to the lithium-ion battery in smartphones. In comparison to the galvanic cell, the electrodes of an electrolytic cell can be placed in a single compartment containing the molten or aqueous electrolyte. In addition, since the external battery source is what drives the electrons through the circuit, the electrodes will match the positive and negative terminal of the battery. While the anode remains the site of oxidation, it becomes the positive terminal, and the cathode becomes the negative terminal. 3. Electrode Potentials: The potential difference between two electrodes determines the cell’s voltage. 4. Applications: Electrochemistry is used in batteries, corrosion prevention, electroplating, and analytical techniques. Practical Applications: ​ Understanding how batteries work. ​ Measuring ion concentrations in solutions through electrochemical methods. ​ Using electrochemical sensors in various industries. B. Basic Principles of Potentiometry and Voltammetry Potentiometry and voltammetry are important electroanalytical techniques used to measure the potential or current in electrochemical systems. Potentiometry: 1.​ Definition: Measures the potential difference between a reference electrode and an indicator electrode without drawing current. 2.​ Applications: Commonly used in pH measurement, determination of ion concentrations (e.g., fluoride, chloride), and redox potential studies. 3.​ Key Components: o​ Reference Electrode: Provides a stable and known electrode potential regardless of the composition of the solution it is placed in. It serves as a constant reference point against which the potential of another electrode (usually the working electrode) is measured. Common Reference Electrodes: 1.​ Standard Hydrogen Electrode (SHE): o​ Potential: 0.00 V by definition. o​ Rarely used in practice due to its complexity and maintenance. 2.​ Silver/Silver Chloride Electrode (Ag/AgCl): o​ Potential depends on the chloride ion concentration o​ Common in aqueous systems. 3.​ Calomel Electrode (Hg/Hg₂Cl₂): o​ Potential depends on the chloride ion concentration. o​ Used less frequently due to mercury toxicity. o​ Indicator Electrode: Responds to changes in the analyte concentration. Usually, this is an inert noble metal, such as Pt (they do not give any side reaction and they are not toxic in nature), and the potential of the cell is measured vs a reference electrode. o​ Solution: Contains the analyte of interest 4. Key Principles: 1. Measurement of Potential Difference without Current Flow o​ The principle of potentiometry involves measuring the equilibrium potential between two electrodes. o​ Example: In pH measurement, a glass electrode (working electrode) and a reference electrode (e.g., Ag/AgCl) are used to measure the hydrogen ion concentration.  2. Sudden Potential Change in potentiometric Titration: o​ During a potentiometric titration, a significant change in the electrode potential occurs near the equivalence point. o​ This sudden potential shift indicates the completion of the reaction, making the point at which the titrant has fully reacted with the analyte. o​ Example: Determining the equivalence point in an acid base titration. 5. Key Equations: Nernst equation relates the electrode potential to the concentration of ions. It is crucial for understanding and predicting the behavior of electrochemical systems, including potentiometry and voltammetry. Voltammetry: 1.​ Definition: o​ an electrochemical technique that measures the current as a function of applied voltage. o​ It provides information about the redox behavior and concentration of analytes in a solution. 2.​ Key Principles: 1. Measurement of the Relationship between Voltage and Current o​ In a voltammetry experiment, the voltage applied to the working electrode is varied, and the resulting current is measured. o​ The current response provides information about the oxidation/reduction reactions of the analyte. o​ Example: Determining the concentration of metal ions through anodic stripping voltammetry. 2. Cyclic Voltammetry and cyclic Voltammogram Interpretation: o​ In cyclic voltammetry (CV), the applied potential is scanned cyclically, and the resulting current is recorded. o​ A cyclic voltammogram is a plot of current vs. potential providing insights into: o​The redox potential of the analyte o​The reversibility of the redox reaction o​The kinetics of electron transfer processes. 3. Key Components: o​ Working Electrode: Where the redox reaction occurs (e.g., glassy carbon, platinum, gold) o​ Reference Electrode: Provides a constant potential for measurement. o​ Counter Electrode: Completes the circuit by allowing current flow. 4. Voltammetric Techniques:  o​ Linear Sweep Voltammetry (LSV): Voltage is increased linearly with time. o​ Cyclic Voltammetry (CV): Voltage is swept forward and backward to study redox processes. o​ Differential Pulse Voltammetry (DPV): Voltage is applied in pulses to improve resolution and sensitivity. o​ Square Wave Voltammetry (SWV): Uses square wave modulation for enhanced sensitivity. o​ Anodic stripping Voltammetry (ASV): Involves pre-concentration of metal ions onto the working electrode by applying a negative potential. : After accumulation, an anodic potential sweep oxidizes the metal ions, generating a current peak proportional to the metal concentration : Highly sensitive for trace metal analysis in environmental and biological samples 3.​ Practical Applications: o​ Environmental analysis (detection of trace metals; monitoring pH levels in environmental samples) o​ Pharmaceutical and biological testing o​ Food quality analysis o​ Electrochemical sensor development C. Basic Knowledge and Principles on Spectrochemical Methods of Analysis Spectrochemical methods involve the interaction of light with matter to determine the concentration of substances in a sample. Types of Spectrochemical Methods: 1.​ UV-Vis Spectroscopy: o​ Definition: Measures a sample's absorption of ultraviolet and visible light. o​ Principle: Based on the Beer-Lambert Law, which relates the absorbance of light to the concentration of the absorbing species in a solution. The equation is A = εbC , where A is absorbance, ε is the molar absorptivity, C is concentration, and b is the path length of light. o​ Applications: Commonly used to determine the concentration of colored compounds and analyze organic molecules, transition metals, and biological samples. o​ Advantages: Non-destructive and suitable for both qualitative and quantitative analysis.  Figure 2: UV-Vis Spectroscopy 2. Atomic Absorption Spectroscopy (AAS): o​ Definition: Measures light absorption by free atoms in the gaseous state. o​ Principle: Involves the use of a light source (usually a hollow cathode lamp) that emits a wavelength specific to the element being analyzed. The sample is atomized in a flame or graphite furnace, and the atoms absorb light at specific wavelengths. o​ Applications: ▪​ Detection of trace metals like lead, cadmium, zinc, and mercury in water, food, and biological samples. ▪​ Environmental monitoring for heavy metal contamination. o​ Advantages: Highly sensitive and selective for metal ions. o​ Limitations: Requires careful calibration and maintenance of the instrument. Figure 3. Atomic Absorption Spectroscopy 3. Fluorescence Spectroscopy: o​ Definition: Measures the emission of light by a substance that has absorbed light or other electromagnetic radiation. o​ Principle: After a molecule absorbs energy, it transitions to an excited state. When the molecule returns to its ground state, it emits light of a  longer wavelength (fluorescence). This emitted light is detected and measured. o​ Applications: ▪​ Used in biological and chemical analysis to detect and quantify proteins, nucleic acids, and other biomolecules. ▪​ Commonly applied in clinical diagnostics, environmental monitoring, and pharmaceutical research. o​ Advantages: ▪​ Highly sensitive, capable of detecting very low concentrations of analytes. ▪​ Non-destructive and can be used for real-time analysis. o​ Limitations: Susceptible to interference from other fluorescent substances in the sample. Key Components of Spectrometers: 1.​ Light Source: Provides the necessary electromagnetic radiation, such as a deuterium lamp for UV light or a tungsten lamp for visible light. 2.​ Monochromator: Selects a specific wavelength of light to be directed at the sample. 3.​ Sample Holder: Holds the sample in a cuvette, ensuring consistent path length. 4.​ Detector: Measures the intensity of transmitted or emitted light and converts it into an electrical signal for analysis. Applications of Spectrochemical Methods: ​ Water Quality Monitoring: Detecting pollutants and heavy metals in water samples. ​ Food Safety Analysis: Identifying contaminants or verifying the authenticity of food products. ​ Pharmaceutical Quality Control: Ensuring the correct concentration of active ingredients in drugs. ​ Biological Research: Analyzing DNA, RNA, and proteins in molecular biology. Practical Tips for Using Spectrochemical Methods: A. Always calibrate the spectrometer before use. B. Use clean and clear cuvettes to avoid contamination and interference. ​ Dirty cuvettes (e.g., fingerprints) can skew absorbance readings. C. Select the appropriate wavelength for the analyte to maximize accuracy. When using a spectrophotometer to analyze a solution, selecting the appropriate wavelength is crucial to ensuring accurate and reliable results. Here's why and how this is done: 1.​ Maximum Sensitivity: At certain wavelengths, the analyte absorbs light most strongly. Selecting this wavelength ensures the spectrophotometer can detect even small changes in analyte concentration. 2.​ Accuracy of Measurements: At the maximum absorbance (called the λmax), the Beer-Lambert Law is most effective because the relationship between absorbance and concentration is linear. 3.​ Minimizing Interference: By choosing the λmax, you reduce the impact of interfering substances that might absorb at other wavelengths. Step-by-Step Process for Wavelength Selection: 1.​ Understand the Absorbance Spectrum: o​ The absorbance spectrum is a plot of absorbance (y-axis) versus wavelength (x-axis). o​ Different molecules have unique absorbance spectra based on their electronic structure. Peaks in the spectrum indicate wavelengths where the molecule absorbs the most light. o​ Example: A solution of chlorophyll shows peaks in the blue (~430 nm) and red (~660 nm) regions, corresponding to its primary absorption of light for photosynthesis. 2. Choose the λmax: o​ The λmax is the wavelength corresponding to the highest peak in the absorbance spectrum. o​ Advantages of using λmax: 1.​ Greater Sensitivity: Measurements at λmax ensure even small changes in concentration produce measurable changes in absorbance.  2.​ Improved Accuracy: The linearity of the Beer-Lambert Law is more reliable near the λmax, reducing errors in concentration determination. o​ Practical Example: 1.​ For a blue solution, such as Cu²⁺ in ammonia, λmax is often in the orange range (~600 nm), which is complementary to its observed color. 2.​ For yellow solutions, the λmax might be in the violet range. 3. Set the Instrument: Once the λmax is identified, set the spectrophotometer to that wavelength for subsequent calibration and sample analysis. 4. Avoid Other Wavelengths: o​ At wavelengths away from the λmax: 1.​ Absorbance values are lower, leading to poorer sensitivity. 2.​ The Beer-Lambert Law may deviate from linearity. 3.​ Noise or interference from other species can dominate, affecting accuracy. Example ​ Suppose you're analyzing a sample of a colored solution, such as a solution containing Fe³⁺ ions with thiocyanate. The Fe(SCN)²⁺ complex has a λmax at 480 nm. By choosing this wavelength: o​ The spectrophotometer is most sensitive to the concentration of Fe(SCN)²⁺. o​ Interference from other species absorbing outside this range is minimized. D. Handle the instrument carefully and follow standard operating procedures. Why is Wavelength Selection Part of Calibration? During calibration, a series of standard solutions with known concentrations of the analyte are measured to establish a calibration curve (absorbance vs. concentration). Accurate calibration requires: ​ Selecting the best wavelength (λmax) to ensure precise and sensitive measurements. ​ Avoiding wavelengths where the absorbance signal is weak or where interfering species may absorb. Factors to Consider When Selecting λmax:  1.​ Absorbance of the Analyte: o​ Ensure the analyte absorbs strongly at the selected wavelength. 2.​ Interference from Other Species: o​ Minimize overlap with other components in the solution. 3.​ Instrument Limitations: o​ Some spectrophotometers have constraints on the wavelength range (e.g., UV-Vis instruments may operate only from 200–800 nm). Practical Example with Beer-Lambert Law ​ Scenario: You’re analyzing a Fe³⁺-thiocyanate complex. ​ Procedure: 1.​ Prepare the sample and scan the absorbance spectrum from 350–700 nm. 2.​ Identify the λmax (e.g., 480 nm for Fe(SCN)²⁺ complex). 3.​ Set the spectrophotometer to 480 nm and measure the absorbance of standards and unknown samples. ​ Result: o​ A calibration curve is constructed, and unknown concentrations are determined using the curve. Advanced Considerations 1.​ Shoulders in the Spectrum: o​ Some spectra have broad or multiple peaks (e.g., proteins). Choose the most prominent or specific peak to the analyte. 2.​ Spectral Bandwidth: o​ Ensure the instrument's spectral resolution matches the narrowness of the λmax. 3.​ Stray Light: o​ Avoid regions with low absorbance to minimize the impact of stray light, which can cause deviations.