TLDR;
Alright students, welcome to another chemistry session! Today, we're diving deep into electrochemistry, a chapter many of you have been waiting for. We'll explore its importance, key topics, and real-world applications. Get ready to transform your understanding of how electrochemistry surrounds us!
- Electrochemistry is used in batteries, cars, phones, drones, lights, and corrosion prevention.
- The chapter is divided into electrolytic conductance (numerical) and electrochemical cells (conceptual).
- Galvanic cells have spontaneous reactions (delta G is negative), while electrolytic cells are non-spontaneous (delta G is positive).
Electrochemistry in Daily Life [0:46]
Electrochemistry is all around us, even if we don't realize it. Think about the batteries in your remotes and video game controllers. Cars also use electrochemistry. Even the phones and laptops you're watching this lecture on rely on it. Drones, lights, and even the rusting of old cars involve electrochemical principles. Electroplating, like on "golden spoons," and various types of batteries also use electrochemistry. Once you start looking, you'll see electrochemistry everywhere, changing how you understand the world.
Two Major Parts of Electrochemistry [4:36]
Electrochemistry is mainly divided into two parts. First, electrolytic conductance, which tells you if a compound will conduct electricity and how well. This part is mostly numerical, so remember your formulas. Second, electrochemical cells, which are more conceptual. Numerical questions are asked, but the main focus is on understanding the concepts.
Chemical Energy and Electrochemical Cells [6:06]
Chemical energy isn't just enthalpy (delta H); it's Gibbs free energy (delta G). Electrochemical cells come in two types: galvanic and electrolytic. Galvanic cells have spontaneous reactions, meaning they happen on their own, and electrolytic cells have non-spontaneous reactions, needing an external force. For galvanic cells, delta G is negative, while for electrolytic cells, it's positive.
Galvanic vs Electrolytic Cells: The Smartphone Analogy [9:15]
Think of your smartphone. When you charge it, you're using an external force (electricity), making it a non-spontaneous process. Electrical energy is converted into chemical energy. This is like an electrolytic cell. When your phone is charged and you're using it, it's a spontaneous process. The chemical energy stored in the phone is converted into electrical energy, like a galvanic cell.
Electrolytic and Galvanic Cells Summarized [14:55]
Let's summarize. Electrolytic cells are non-spontaneous, have a positive delta G, and convert electrical energy into chemical energy. Galvanic cells are spontaneous, have a negative delta G, and convert chemical energy into electrical energy. Remember these points!
Electrolytic Cell Apparatus [17:03]
An electrolytic cell needs a battery because it has non-spontaneous processes. The electrode connected to the positive terminal is positive, and the one connected to the negative terminal is negative. In a setup with NaCl, Na+ goes to the negative electrode, and Cl- goes to the positive electrode. Oxidation (loss of electrons) happens at the anode, and reduction (gain of electrons) happens at the cathode.
Understanding Anode and Cathode [21:45]
Anode is where oxidation happens, and cathode is where reduction happens. If Na+ gains an electron at the negative terminal, it's reduction, so that's the cathode. If Cl- loses an electron at the positive terminal, it's oxidation, so that's the anode. In an electrochemical cell, the cathode is negatively charged, and the anode is positively charged.
Key Points About Electrodes [26:17]
Oxidation happens at the anode, and reduction happens at the cathode. Electrodes are metallic plates where electrons are exchanged. These exchanges always happen on the surface of the electrodes, not in the solution. We always assume we're using equal solutions.
Galvanic Cell: Setup and Spontaneous Processes [27:38]
Galvanic cells, also known as voltaic or Daniel cells, have spontaneous processes and don't need a battery. Delta G is negative, and chemical energy is converted into electrical energy. The setup involves two containers with different solutions, like ZnSO4 and CuSO4. Zinc wants to become Zn2+ by losing two electrons, and copper wants to accept two electrons to become copper solid.
Finding Cathode and Anode in Galvanic Cell [31:27]
Zinc losing electrons is oxidation, so the zinc electrode is the anode. Copper gaining electrons is reduction, so the copper electrode is the cathode. Zinc from the electrode goes into the solution, making the zinc electrode thinner. Copper ions from the solution deposit on the copper electrode, making it thicker.
Electron Flow and Salt Bridge [34:36]
Zinc gives electrons to the electrode, making it negatively charged (anode). Copper takes electrons from the electrode, making it positively charged (cathode). Electrons flow from anode to cathode, while current flows from cathode to anode. A salt bridge, made of KCl and agar-agar, neutralizes the solutions by bringing electro-neutrality.
Summarizing Galvanic and Electrolytic Cells Signs [41:00]
In a galvanic cell, the anode is negatively charged, and the cathode is positively charged. Electrolytic cells have the opposite signs. If you remember the sign of one electrode in one cell, you can figure out the signs of the other three.
Practice Questions: Spontaneity and Correct Statements [42:27]
A cell reaction is spontaneous if delta G is negative. In galvanic cells, the anode is negative. Molten sodium chloride conducts electricity because of free ions. Galvanic cells convert chemical energy to electrical energy. The function of a salt bridge is to maintain electrical neutrality.
Cell Representation: Oxidation and Reduction [46:23]
Cells are represented with a salt bridge (two vertical lines). The cathode reaction is on the left, and the anode reaction is on the right. Ions are written inside, and solids are at the ends. For example, Cu2+ (aq) | Cu (s) || Zn2+ (aq) | Zn (s). Remember, ions are "scared" and can't be at the end.
Half-Cell Representation and Gas Ion Electrodes [51:22]
For a half-cell, only the cathode or anode part is written. For a gas ion electrode, like H2 becoming H+, the representation is: H+ (aq) | H2 (g) | Pt (s). If there's no solid, use an inert electrode like platinum.
More Practice: Representing and Interpreting Cells [54:59]
To represent a cell, find the oxidation and reduction reactions, write the cathode and anode reactions, and then write the cell representation. For example, if H2 becomes 2H+ and Cl2 becomes 2Cl-, the cell representation is Pt (s) | H2 (g) | H+ (aq) || Cl- (aq) | Cl2 (g) | Pt (s).
Identifying Oxidation and Reduction in Cell Reactions [1:00:44]
To find the cell reaction from a cell representation, write the cathode and anode reactions. For example, if the cell is Fe | Fe2+ || Fe3+ | Fe2+, the cathode reaction is Fe3+ + e- -> Fe2+, and the anode reaction is Fe -> Fe2+ + 2e-. Balance the electrons and add the reactions to get the net cell reaction.
Oxidizing and Reducing Agents: Oxidation Potential [1:04:19]
Oxidation potential is the tendency to lose electrons. Reduction potential is the tendency to gain electrons. E(reduction potential) = -E(oxidation potential). Oxidizing agents oxidize others and get reduced themselves, so they have high reduction potential. Reducing agents reduce others and get oxidized themselves, so they have high oxidation potential.
Electricity in Daily Life and EMF of Cells [1:10:54]
Electricity is used everywhere. The EMF of a cell depends on the nature of the chemical reaction, concentration of reactants and products, pressure of gases, and temperature. Standard conditions are 1 molar concentration, 298 Kelvin temperature, and 1 atmosphere pressure.
Calculating EMF of Cells: Standard and Non-Standard Conditions [1:14:12]
To calculate the EMF of a cell under standard conditions, use the formula: E°cell = E°cathode - E°anode. Use reduction potentials for both. A positive E°cell means the reaction is spontaneous. Standard Hydrogen Electrode (SHE) is used as a reference, with its reduction potential set to zero.
Electrochemical Series: Reduction Potential and Reactivity [1:22:04]
The electrochemical series lists elements by their reduction potentials. As you move down the series, reduction potential increases, and oxidation potential decreases. Better oxidizing agents are at the bottom. Elements above can displace elements below them from their aqueous solutions because they have higher oxidation potential.
Applying Electrochemical Series: Reducing Power [1:29:17]
To arrange metals by increasing reducing power, look at their oxidation potentials. The higher the oxidation potential, the better the reducing power. Remember, reducing power is the ability to reduce someone else.
Nernst Equation: EMF at Non-Standard Conditions [1:35:33]
The Nernst equation is used to calculate cell EMF at non-standard conditions: Ecell = E°cell - (0.059/n) log (products/reactants). At equilibrium, Ecell = 0.
Gibbs Free Energy and EMF Relationship [1:47:58]
The relationship between standard Gibbs free energy change and EMF of the cell is: ΔG° = -nFE°cell, where n is the number of electrons exchanged, and F is Faraday's constant (96500 coulombs).
Faraday's Law of Electrolysis: Charge and Moles [1:49:18]
Faraday's law tells us that the charge of one mole of electrons is one Faraday (96500 coulombs). The number of Faradays is equal to the moles of electrons. To solve electrolysis problems, calculate the total charge, find the number of Faradays, and then use mole concepts.
Electrolysis: Preferential Discharge of Ions [1:58:22]
In electrolysis, positive ions go to the cathode and get reduced, while negative ions go to the anode and get oxidized. If there are multiple positive ions, the one with the higher reduction potential gets reduced first. For anions, Cl- is preferred over OH-, and OH- is preferred over SO42-.
Electrolytic Conductance: Resistance and Conductivity [2:05:06]
Resistance is directly proportional to length and inversely proportional to area: R = ρ(L/A). Conductance (G) is the reciprocal of resistance, and conductivity (κ) is the reciprocal of resistivity. Cell constant (G*) is L/A. The relationship is κ = G * G*.
Units of Conductance and Conductivity [2:08:25]
The unit of resistance is ohm (Ω). The unit of conductance is ohm inverse (Ω-1) or Siemens (S). The unit of resistivity is ohm-centimeter (Ω cm). The unit of conductivity is Siemens per centimeter (S cm-1). The unit of cell constant is centimeter inverse (cm-1).
Molar and Equivalent Conductivity [2:14:32]
Molar conductivity (Λm) is κ * 1000 / C, where C is the molarity. Equivalent conductivity (Λeq) is κ * 1000 / N, where N is the normality. The relationship between them is Λm = Λeq * n-factor, where n-factor is the total positive or negative charge.
Strong vs Weak Electrolytes: Kohlrausch's Law [2:18:40]
Strong electrolytes (strong acids, strong bases, and salts) have high conductance. Their molar conductivity can be found by extrapolating the graph to zero concentration. Weak electrolytes have lower conductance, and their graph doesn't extrapolate easily. Kohlrausch's law states that at infinite dilution, each ion contributes independently to the total conductivity.
Applying Kohlrausch's Law: Independent Migration of Ions [2:22:33]
To find the infinite dilution conductance of a weak electrolyte, use strong electrolytes. For example, to find the conductance of H2O, use NaOH, NaCl, and HCl. Add the conductances of NaOH and HCl, and subtract the conductance of NaCl. This gives you the conductance of H2O.