Introduction to Electrochemistry

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Introduction to Electrochemistry

• Electrochemistry studies the interconversion of chemical and electrical energy.
• It includes converting chemical energy into electrical energy and vice versa.
• Devices like batteries and electrochemical cells utilize these energy conversions.
• Chemical energy is used in batteries, then batteries can be recharged by converting electrical energy back to chemical energy.
• Electrolytic cells use electrical energy to drive non-spontaneous chemical reactions.

Electrochemical Cells

• Electrochemical cells convert chemical energy into electrical energy.
• These cells use spontaneous chemical reactions to generate electricity.
• The Daniel cell is a specific type of electrochemical cell.
• The Daniel cell demonstrates the basic principles of how electrochemical cells work.

The Daniel Cell: Setup

• A Daniel cell consists of two containers.
• One container holds a zinc sulfate (ZnSO₄) solution with a zinc (Zn) electrode.
• The other container holds a copper sulfate (CuSO₄) solution with a copper (Cu) electrode.
• The two solutions are connected by a salt bridge.

Components and Function

• A zinc rod is placed in the zinc sulfate solution.
• A copper rod is placed in the copper sulfate solution.
• A Galvanometer can be connected to the circuit to meausre the flow of current.
• The two electrodes are connected via a galvanometer to allow electron flow.
• A switch is included to control activation of the cell.

Oxidation and Reduction

• Zinc is more reactive than copper, so it undergoes oxidation.
• Oxidation occurs at the anode, where zinc atoms lose electrons to form Zn²⁺ ions.
• Anode is where oxidation occurs.
• Electrons released by zinc flow through the external circuit towards the copper electrode.
• At the cathode, copper ions (Cu²⁺) gain electrons and are reduced to copper metal.
• Reduction occurs at the cathode.

Half-Reactions

• Oxidation half-reaction: Zn → Zn²⁺ + 2e⁻
• Reduction half-reaction: Cu²⁺ + 2e⁻ → Cu
• Zinc metal is oxidized into Zn²⁺ ions, causing the zinc electrode to corrode.
• Copper ions (Cu²⁺) are reduced into Copper metal, leading to a deposit of cooper on the electrode.

Salt Bridge

• The salt bridge connects the two half-cells.
• The two main roles of the salt bridge are to:
  • Complete the circuit by allowing ion flow between the half-cells.
  • Maintain electrical neutrality in both half-cells by preventing charge build-up.
• The salt bridge typically contains an inert electrolyte like KCl or NaNO₃.

Overall Reaction

• The overall cell reaction: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
• This is a redox reaction where zinc displaces copper from its salt solution.
• The reaction is carried out in two separate containers unlike usual displacement reactions.
• Enables the flow of electrons from one container to another.
• This electron flow produces an electric current.

Cell Representation

• The standard notation for the Daniel cell is: Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)
• The anode (oxidation half-cell) is written on the left.
• The cathode (reduction half-cell) is written on the right.
• A double vertical line (||) represents the salt bridge.

Electron Flow and Polarity

• Electrons flow from the zinc electrode (anode) to the copper electrode (cathode).
• In an electrochemical cell, the anode is considered negative, and the cathode is positive.
• Anode terminals are negative.
• Cathode terminals are positive.

Electrode Potential

• Anode electrode charge is negative, releasing Zn²⁺ ions into the zinc sulfate solution, creating an electrical imbalance.
• Cathode electrode charge is positive, and S0₄- molecules increase the cathode solution.
• This difference in charge leads to a potential difference called the electrode potential.
• The electrode's ability to push or pull the electron is the electrode potential.
• Electrical potential is established between the electrode and the solution.
• Depends on concentration, temperature, and ion activity.
• Standard conditions, typically indicated with the superscript °, maintain constant temperature, pressure, and concentration.

Standard Electrode Potential

• When conditions are standard, the electrode potential becomes the standard electrode potential.
• Standard conditions include 1M concentration, 298K temperature, and 1 atm pressure.

Measuring Single Electrode Potentials

• Measuring the potential of a single electrode is difficult, as it requires two half-cells to function.
• Use a reference electrode to measure individual electrode potentials.
• Need to measure the ability to push or pull the electrons.

Standard Hydrogen Electrode (SHE)

• The SHE is a reference electrode used to measure the standard electrode potentials of other half-cells.
• The SHE consists of a platinum electrode immersed in 1M HCl solution.
• Hydrogen gas at 1 atm pressure is bubbled through the solution.
• The SHE is assigned a standard electrode potential of zero volts.

Standard Hydrogen Electrode (SHE) - Reaction

• The half-reaction in SHE is: 2H⁺(aq) + 2e⁻ ⇌ H₂(g)
• Hydrogen is the standard.
• It is assigned a value of zero, acting as the baseline to compare any other half-cell connected to it.

SHE Example

• If a SHE is connected to another half-cell, a voltmeter will show a reading.
• The reading corresponds to the other half-cell due to standardizing and referencing against the Hydrogen Electrode, which is set to zero.
• If a zinc electrode is connected to a SHE displaying 0.74V, we attribute total voltage to Zinc, because Hydrogen's contribution is zero.

Oxidation vs. Reduction Potential

• Electrod potential can be written as oxidation or reduction potential.
• Magnitude wll remain, but sign will change.
• Oxidation potential is the tendency of a species to lose electrons.
• Reduction potential is the tendency of a species to gain electrons.
• The oxidation potential is the reverse of the reduction potential.
• Oxidation and reduction potential are basically the same, but opposite signs.

Electrochemical Series

• It's a list that organizes electrochemical potential.
• Electrochemical series arranges standard reduction potentials in decreasing order.
• It allows predicting the spontaneity of redox reactions.

Using the Electrochemical Series

• The electrochemical series helps determine which substances are easily oxidized or reduced.
• The table should not be used on rote, but conceptually.
• Substances at the top are strong reducing agents (easily oxidized).
• Substances at the bottom are strong oxidizing agents (easily reduced).

Application of Electrochemical Series

• To Predict the Spontaneity
  • It facilitates quick assessment of probable redox reactions.
• To Identify Oxidizing and Reducing Agents
  • Elements with high oxidation potentials tend to act as reducing agents, whereas those with high reduction potentials are oxidizing agents.

Calculating Cell Potential

• A cell is made up of 2 half-cells.
• For example, if Nickel is coupled with Oxygen, which becomes the Anode and which becomes the Cathode?
• Need to figure out the electric potential in terms of a whole cell (not just individual half cells)

Standard Cell Potential

• Standard cell potential (E°cell) is the potential difference between two half-cells under standard conditions.
• E°cell = E°cathode - E°anode
• Always make sure that when calculating you are always using Reduction Potential, not Oxidation Pontential.
• The cathode has a higher reduction potential.
• Anode potential is lower than cathode.
• E°cell is always positive.

Example: Magnesium and Tin Cell (Mg/Sn)

• The cell consists of a magnesium electrode and a tin electrode.
• E°(Mg²⁺/Mg) = -2.37 V
• E°(Sn²⁺/Sn) = -0.14 V
• Identify which reactions are Oxidation where Magnesium is the reducing agent and Tin is for Oxidation.
• E°cell equals E° cathode - E° anode.
• Equals -0.14 - (-2.37). The difference in potential is 2.23V.

Example: Chlorine and Copper Cell (Cl/Cu)

• The cell consists of a Chlorine electrode and a Copper electrode.
• E°(Cl/Cl⁻) = 1.36 V
• E°(Cu²⁺/Cu) = 0.34 V
• Identify which reactions are Oxidation where Chlorine is the reducing agent and Copper is for Oxidation.
• E°cell equals E° cathode - E° anode.
• Equals 1.36 - (0.34). The difference in potential is 1.02V.

Nernst Equation Introduction

• The Nernst equation relates cell potential to standard cell potential and reaction quotient.
• Standard conditions are rare, so Nernst Equation can be used in non-standard conditions to determine cell potential.
• Non-standard conditions involve any variations in system temperature and concentrations from standard metrics

Generalized Reaction

• aA + bB ⇌ cC + dD

Deriving the Nernst Equation

• Ecell = E°cell - (2.303RT/nF)log(Q)
  • E°cell represents the cell when in standard conditions
  • R is the gas constant,
  • T is temperature in Kelvin,
  • n is the number of moles of electrons transferred,
  • F is Faraday's constant, and
  • Q is the reaction quotient.
• Reaction quotient is Q = [C]^c[D]^d / [A]^a[B]^b

Simplified Nernst Equation (at 298K)

• At 25 degrees Celsius or 298 K, the Nernst equation simplifies to:
• Ecell = E°cell - (0.0591/n)log(Q)
  • E°cell represents the cell when in standard conditions
  • n is the number of moles of electrons transferred
  • Q is the reaction quotient.
• This simplified form is useful and removes the necessity to calculate constants.
• Remember this constant happens only at 298K or 25C!

Nernst Equation Example: Zn/Cu Cell

• Zn(s) + Cu²⁺(aq) ⇌ Zn²⁺(aq) + Cu(s)
• Ecell = E°cell - (0.0591/n)log(Q) becomes
• Ecell = E°cell - (0.0591/2)log([Zn²⁺] / [Cu²⁺])
  • n equals 2, because there are 2 electrons, and zinc is 2+
  • E°cell represents the cell when in standard conditions
  • [Zn²⁺] is going to above, becaues it is a product
  • [Cu²⁺] is going to bellows, becaues it is a reagent

Nernst Equation Example: Chlorine (Cl₂)

• Cl₂(g) + 2e⁻ -> 2Cl⁻(aq)
• Ecell = E°cell - (0.0591/2)log([Cl⁻] / [Cl₂])
  • [Cl₂] There is nothing, so we raise to the power of 0
  • [Cl⁻] there are products, so products go on top

Gibbs Free Energy Introduction

• Relates cell potential to spontaneity and equilibrium.
• Gibbs free energy (ΔG) indicates the spontaneity of a reaction.
• ΔG = -nFEcell, where n is the number of moles of electrons transferred, F is Faraday's constant, and Ecell is the cell potential.
  • Gibbs Free Energy measures useable energy.

Standard Gibbs Free Energy

• ΔG° = -nFE°cell, where ΔG° is the standard Gibbs free energy and E°cell is the standard cell potential.
• Need to use constants for gibbs free energy!
• Charge is always negative when calculating Free Energy.

Equilibrium Constant

• At equilibrium, Ecell = 0 and ΔG = 0.
• ΔG° = -RT lnK, where K is the equilibrium constant.

Nernst Equation and Equilibrium

• E°cell = (0.0591/n)logK
• Reorganize Nernst equation to calculate K
• E°cell = (RT/nF)lnK

Overview of Conductivity and Molar Conductivity

• R=ρl/A, where R is resistance, ρ (rho) is resistivity, l is length, and A is the area of the conductor.
• Resistance is in Ohams.

Electrochemistry involves converting chemical and electrical energy, seen in batteries and electrochemical cells. Batteries use chemical energy which can be restored by converting electrical energy back. Electrolytic cells use electrical energy to facilitate non-spontaneous reactions, such as those in a Daniel cell.