Physics I Module of the MCAT Self Prep eCourse: Lesson 12: Redox Reactions (Pro)
What is the oxidation number for any atom by itself?
0
What is the ε° for concentration cells?
0 And therefore ΔG is also 0
What is the cell potential, ε, at equilibrium?
0 Explanation: ΔG = 0 at equilibrium, so ΔG = -nFε 0 = -nFε ε must be 0
True or False: All batteries are influenced by temperature changes.
True
What is the oxidation half-reaction that occurs in a nickel-cadmium battery?
occurs at the anode E°red = -0.86 V
A spontaneous reaction in a voltaic cell will have a ___________ cell potential.
positive
True or False: ΔG° and ε° will have opposite signs.
True
True or False: The salt bridge is composed of inert salt.
True Explanation: inert ions will not react with the electrodes or ions in the solution. They will only precipitate/trickle into the solution to balance out the newly created charges.
True or False: In an electrolytic cell, the half-reactions do not need to be separated into different compartments.
True Explanation: the desired reaction is *non-spontaneous*
True or False: Oxidation *always* occurs at the anode and reduction *always* occurs at the cathode.
True ✴︎doesn't matter if it's a galvanic or electrolytic cell (though the charges on each electrode *do* switch)
What is the equation for the reaction quotient, Q? When is it used?
When conditions deviate from equilibrium Remember, only the species *in solution* are included in the expression. Pure *solids* and *liquids* will *not* be included, while gases and aqueous solutions *will* be.
When will H₂SO₄ lose both its protons?
When it reacts with a base, such as ammonia In water, it will only readily lose the first proton. The Kₐ for the loss of the second proton is only 10⁻², and only 10% of the H₂SO₄ molecules in a 1 M solution lose a second proton.
Does temperature have any effect on electrochemical cells/batteries?
YES
For a Ni-Cd battery, the reaction occurs in a ____________ solution.
basic ✴︎so, when balancing, you pretend like it's an acidic solution and then change to a basic solution: After adding H⁺ to both sides, add OH⁻ to balance them out. Some will form water.
In general, an oxidizing agent can oxidize any reducing agent that lies _____________ it on the standard reduction potential table.
below ✷so Pb²⁺ can oxidize Al(s), but not Cu(s)
For electrolytic cells, the electrode with the more negative reduction potential is the ____________.
cathode *NOTE* - reduction still occurs here, but with the compound that is less naturally likely to be reduced (spontaneously would be oxidized), so it is is nonspontaneous
No matter the type of cell, cations are always attracted to the ____________.
cathode Reduction always occurs at the cathode and in a galvanic cell, the cathode receives electrons from the anode, and thus attracts cations because it looks more negative. In an electrolytic cell, the cathode is - and is attached to the - terminal of the voltage source, where it attracts *cations* from the solution
For galvanic cells, the electrode with the more positive reduction potential is the ____________.
cathode Explanation: it will be spontaneously reduced
Cations ALWAYS migrate towards the ____________.
cathode Mnemonic: *Cation to cathode* Explanation: In a galvanic cell, the cathode is +, so this may be counterintuitive, but read this: "Cations (e.g. K⁺) are flowing *from the salt bridge* toward the cathode to *replace* the positive charge of the Cu²⁺ ions that got consumed. Easier to remember for electrolytic cells because cathode is negative. So please just remember this convention 😊
For all electrochemical cells, the direction of the current runs from ____________ ⟶ _____________.
cathode ⟶ anode
In a galvanic cell, the concentration of copper solution is ____________. (Use a Daniel cell to apply this concept).
decreasing Explanation: Cu is the cathode, and the metal electrode is gaining mass as Cu²⁺ ions in solution are being reduced to Cu(s), which flow out of solution and deposit onto the electrode. *less ions in solution = decreased concentration*
For a polyatomic ion, the sum of the individual oxidation numbers must equal...
its overall charge
In a concentration cell, which side is oxidation occurring?
less concentrated Explanation: we want both to be equal, so to get this side more concentrated, we want ions to flow into solution. This happens if the metal electrode is oxidized to cations.
In a galvanic cell, the zinc electrode is ____________ mass. (Use a Daniel cell to apply this concept).
losing Explanation: Zn metal (anode) is oxidized to Zn²⁺ ions which go into solution
Ni-Cd batteries have ____________ internal resistance.
low
How do lead-storage batteries rate on the *energy density* scale?
low ✴︎require a heavier amount of battery material compared to other batteries
In a concentration cell, which side is reduction occurring?
more concentrated Explanation: we want both to be equal, so to get this side less concentrated, we want ions in solution to get out and deposit onto the electrode. This happens if the ions are reduced to elemental solid form.
A spontaneous reaction in a voltaic cell will have a ___________ change in free energy (ΔG°).
negative
In a electrolytic cell, the cathode is ____________.
negative
In a galvanic cell, the anode is ____________.
negative
For electrolytic cells, the difference of the reduction potentials of the 2 half-reactions is ___________.
negative Explanation: anode has the more +ε° this time, so according to the equation: Ecell = Ered,cat - Ered,an subtracting a larger # from a smaller # gives you a - value ✴︎also makes intuitive sense: galvanic cells are spontaneous (ΔG > 0 and ε° < 0)
The natural logarithm of any number between 0 and 1 is _____________.
negative ✴︎so, conceptually, if Keq is < 1, the E°cell will be negative (nonspontaneous).
Most AA and AAA batteries are ____________.
nickel-cadmium batteries
What is the reduction half-reaction that occurs in a nickel-cadmium battery?
occurs at the cathode E°red = 0.49 V
As standard reduction potential becomes more positive, the species is a better ____________ agent.
oxidizing *it is better at being reduced, so it oxidizes something else*
surge current
periods of above-average current transiently released early in the discharge cycle it wanes rapidly until a stable current is achieved
In a galvanic cell, the cathode is ____________.
positive
In an electrolytic cell, the anode is ____________.
positive
The change in free energy (ΔG) for an electrolytic cell is ____________.
positive
For galvanic cells, the difference of the reduction potentials of the 2 half-reactions is ___________.
positive Explanation: cathode has the more + ε°, so according to the equation: Ecell = Ered,cat - Ered,an subtracting a smaller # from a larger # gives you a + value ✴︎also makes intuitive sense: galvanic cells are spontaneous (ΔG < 0 and ε° > 0)
The more ____________ the value of the standard cell reduction potential ε°, the more likely the substance is to get reduced.
positive ✴︎So for a redox reaction involving copper and zinc, copper will be reduced.
The natural logarithm of any number greater than 1 is _____________.
positive ✴︎so, conceptually, if Keq is > 1, the E°cell will be positive (spontaneous)
What type of device cn be used to measure the emf of a cell?
potentiometer Explanation: It is a special type of voltmeter that draws no current and gives a more accurate reading of the difference in potential between the two electrodes.
If Q < K, the reaction will favor ___________.
products (shift right; forward)
If Q > K, the reaction will favor ___________.
reactants (shift left; backward)
Lead storage batteries are ____________ batteries that function as both galvanic and electrolytic cells.
rechargeable
By convention, the standard reduction potential are written as if the compound is being ____________.
reduced, *not* oxidized
As standard reduction potential becomes more negative, the species is a better ____________ agent.
reducing *it is better at being oxidized, so it reduces something else*
electrode
strips of metal or other conductive material placed in an electrolyte solution *where oxidation and reduction take place*
standard hydrogen electrode (SHE)
the half-cell consisting of an inert platinum electrode immersed in 1 M HCl with hydrogen gas at 1 atm bubbling through the solution 2 H+ and 2 e⁻ *used as the standard of a cell potential of zero*
Flip to see the equation for calculating EMF that I prefer
MUST FLIP THE OXIDATION PORTION IF BOTH ARE GIVEN AS REDUCTION POTENTIALS
What is the equation for *moles of electrons transferred during reduction*?
Mⁿ⁺ + n e⁻ ⟶ M (s) ^ 1 mole of metal (M) solid will be produced if n moles of electrons are supplied to one mole of Mⁿ⁺, where ⁿ is the charge on the metal ion.
Which batteries have higher energy density, lead-storage or nickel-cadmium?
Ni-Cd, which also provide a higher surge current
Show the reactions that occur when a fully-charged Ni-Cd battery is discharging.
Oxidation: anode Cd (s) + 2 OH⁻ (aq) ⟶ Cd(OH)₂ + 2 e⁻ Reduction: cathode 2NiO(OH) + 2H₂O + 2 e⁻ ⟶ 2Ni(OH)₂ + 2OH⁻ Net reaction: 2NiO(OH) + Cd + 2H₂O ⟶ 2Ni(OH)₂ + Cd(OH)₂ *Result* - both electrodes get plated with their respected solid products, which like the lead-storage, are not soluble in water and thus are available for the reverse reaction.
Write the cell diagram for the discharging state of a lead-acid battery.
Pb (s) | H₂SO₄ (4M) || H₂SO₄ (4M) | PbO₂ (s) ✴︎galvanic: electrodes plate with lead sulfate (PbSO₄) and *dilute* the acid electrolyte: ￫ Water is produced when H⁺ from the acid react with oxide ions produced at the cathode to *form* water. End up with: PbSO₄ | H₂SO₄ (dilute) || H₂SO₄ (dilute) | PbSO₄
Write the cell diagram for the charging state of a lead-acid battery.
PbSO₄ | H₂SO₄ (dilute) || H₂SO₄ (dilute) | PbSO₄ ✴︎electrolytic circuit: an external source reverses the electroplating process and concentrates the acid solution: ￫ electrons are driven onto the Pb plate and pulled from the PbO₂ plate. ￫This process breaks the chemical bond between the lead and the sulfate ions, releasing that sulfate (SO₄²⁻) from the electrodes back into the solution, resulting in a *higher concentration* of sulfuric acid to water. ￫ During the charging process, some electrolysis takes place, which splits water into hydrogen and oxygen gas (amount of water decreases) Back to original: Pb (s) | H₂SO₄ (4M) || H₂SO₄ (4M) | PbO₂ (s)
What is the oxidation number for a monoatomic ion?
its charge
When Q > 1, how does ε compare to ε°?
ε < ε°
For a cell with the following half-reactions: Anode: SO₂ + 2H₂O ⟶ SO₄²⁻ + 4H⁺ + 2 e⁻ Cathode: Pd²⁺ + 2 e⁻ ⟶ Pd How would decreasing the pH of the solution inside the cell affect the electromotive force (emf)? ⁺⁺ A. The emf would decrease. B. The emf would remain the same. C. The emf would increase. D. The emf would become zero.
(A) Explanation: A change in pH directly correlates to the [H⁺]. Decreasing the pH increases the [H⁺], which means the concentration of products has increased. Going from S⁺⁴O₂⁻² ⟶ S⁺⁸O₄⁻² is oxidation and the reverse is reduction. So if we have increased products, the reaction will shift to the left, thus increasing the reduction reaction (more SO₂ reactant). According to the equation, E°cell = E°red, cathode - E°red, anode An increase in the "E°red, anode" term would decrease the emf (subtracting a larger number). Plus, more H⁺ means more interaction b/w protons and electrons, so it's harder to liberate electrons (cell emf ↓ then)
Which of the following compounds is LEAST likely to be found in the salt bridge of a galvanic cell? A. NaCl B. SO₃ C. MgSO₃ D. NH₄NO₃
(B) Explanation: salt bridges contain inert electrolytes, and ionic compounds (A, C, D) are strong electrolytes as they dissociate completely in solution. Option B is a covalent compound and is not an electrolyte, and it would not dissolve in solution
If the surface area of electrode material in an electrochemical cell is tripled, what else is necessarily tripled? I. E°cell II. Current III. Keq A. I only B. II only C. I and II only D. II and III only
(B) II only Explanation: If electrode material increased, the surface area participating in the redox reactions is increased, and more electrons are released, thus current would increase. Potential (E°cell) depends only on the identity of the *electrodes*, not the amount, making statement I false Keq similarly depends only on the identity of the *electrolyte solutions* and *temperature*, making statement III incorrect
Which of the following best describes why overcharging a Ni-Cd battery is not detrimental? A. The energy density of a Ni-Cd battery is high, so it can store more charge than other batteries per its mass. B. The electrodes of a Ni-Cd battery can discharge though the circuit when they are fully charged. C. The Ni-Cd battery will stop accepting electrons from an outside source when its electrodes are recharged. D. Ni-Cd batteries have a high surge current and can dissipate the overcharge before damage can occur to electrodes.
(C) First, *overcharging* refers to continuing to try to run current into the battery even when the electrodes are reverted to their original state. Explanation: During the recharge cycle, Ni-Cd cells will accept current from an outside source until the Cd and NiO(OH) electrodes are pure. At this point, the reaction will stop because Cd(OH)₂ runs out and no more electrons can be accepted. 2NiO(OH) + Cd + 2H₂O ⟵ 2Ni(OH)₂ + Cd(OH)₂ ^ note the direction of the arrow *⟵* Choices (A) and (B) are both true statements, but they fail to answer the question. Surge current refers to the initial burst of current seen in some batteries. Choice (D) is incorrect because once charged, the surge current will not increase even if the power source continues to be run because no additional charge will be stored on the electrodes.
An electrolytic cell is filled with water. Which of the following will move toward the cathode of such a cell? I. H⁺ ions II. O²⁻ ions III. Electrons A. I only B. II only C. I and III only D. II and III only
(C) I and III only Explanation: In an electrolytic cell, ionic compounds are broken up into their constituents, with the cations migrating towards the cathode & anions migrating towards the anode. This makes option I. correct. In *all* types of cells, electrons flow from anode to cathode, making option III correct Option II is incorrect because these anions would flow to the anode, and furthermore, O₂⁻ would not likely be the anions in any cell (more likely OH⁻).
Which of the following can alter the emf of an electrochemical cell? A. The mass of the electrodes B. The length of the wire connecting the half-cells C. The overall size of the battery D. The temperature of the solutions in the half-cells
(D) Explanation: E°cell is dependent on the change in free energy of the system through the equation RTlnKeq = nFE°cell. The temperature (T) appears in this equation, so there it is the correct answer.
What is equation that relates standard cell potential (E°cell) to the equilibrium constant, K? (This is the equation at *standard conditions/equilibrium*)
*bottom equation in pic* Explanation: We first set the following equations equal to each other: ΔG = -nFε° ΔG = -RTlnK° -nFε° = -RTlnK° Then solve for ε° ε° = RTlnK°/nF Then, since R and F are both constants, we can simplify (assume temp of 25°C/298K), so all 3 of these variables get combined into the one numerical value. Then, convert to log form.
What is the equation that relates the free energy change of a voltaic cell to the standard cell potential?
*n* is # of mol of electrons *F* is Faraday's constant (96,500 C/mol) *ε°* has units of J/C (V) *ΔG°* will have units of J
nickel metal hydride battery
A battery that uses the same cathode reaction as the NiCad battery (NiO(OH)) but a different anode reaction, the oxidation of hydrogens in a metal alloy. ❤️ red = discharging: Cathode: NiO(OH)(s) + H₂O + e⁻ ⟶ Ni(OH)₂ + OH⁻ Anode: (occurs in the lattice) MH + OH⁻ ⟶ M + H₂O + e⁻ 💜 purple = charging: Cathode: Ni(OH)₂ + OH⁻ ⟶ NiO(OH)(s) + H₂O + e⁻ Anode: (occurs in the lattice) M + H₂O + e⁻ ⟶ MH + OH⁻
Describe the setup of a fully-charged Ni-Cd battery.
A fully charged NiCd cell contains: ￫ a nickel(III) oxide-hydroxide positive electrode plate (cathode) ￫ a cadmium negative electrode plate (anode) ￫ a separator ￫ an alkaline electrolyte (KOH) (in battery, it's placed with the cathode) Ni-Cd batteries usually have a metal case with a sealing plate equipped with a self-sealing safety valve. The positive and negative electrode plates, isolated from each other by the separator, are rolled in a spiral shape inside the case. This is known as the jelly-roll design and allows a Ni-Cd cell to deliver a much higher maximum current than an equivalent size alkaline cell.
Why is the Nernst Equation useful?
Allows us to calculate a cell potential at *nonstandard* concentrations (away from equilibrium, but temp is still 25°C). ✷like the instantaneous cell potential related to the progress of the reaction: as concentration of reactants and products changes, reaction quotient (Q) changes, thereby changing the cell potential, ε
Why is sulfuric acid (H₂SO₄) considered a strong acid?
Because the Kₐ for the loss of the first proton is much larger than 1 It is *not* because it loses both of its protons when it reacts with water. Its complete dissociation is a two-stage process, with stage 2 (the dissociation of the weak acid HSO₄⁻) having a much lower Kₐ value. http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch17/diprotic.php
When finding the standard reduction potential of a cell, how is the convention different from a situation where you are given simply reduction potentials?
Before, we were given reduction potentials & had to figure out which species would get reduced and which would get oxidized. For the one that would get oxidized, we would flip the sign and then add it to the unchanged reduction potential. ✷ If you follow the new equation, you don't have to flip any signs, but you *must* make sure that both are given as Ered, *not* Eox We are given both Ered, so by subtracting, we are accounting for the fact that the species that got oxidized (anode) should have the opposite sign than we were given.
What is the overall equation for a Ni-Cd battery?
E°cell = 0.49 V - (-0.86 V) = 1.35 V
What is the equation for the *standard electromotive force of a cell (E°cell)*?
E°cell = E°red, cathode - E°red, anode Only valid when *both* half reactions are written as *reductions*
What is the purpose of a salt bridge?
If only a wire were provided for electron flow, the reaction would soon stop because an excess + charge would build up on the anode, and an excess - charge would build up on the cathode. Eventually, these charge accumulations would provide a countervoltage large enough to prevent the redox from occurring any further, so the current would cease. *the salt bridge dissipates this charge gradient and permits the exchange of cations and anions*
What are the biological applications of a concentration cell?
It's like the cell membrane of a neuron. Explanation: Because of the potential difference from the concentration gradient, electrons are driven in the direction that results in the equilibration of the ion gradient. The current will stop when the concentrations of ionic species of each half cell are equal. ^ at this point, the voltage will be zero ✴︎ Use the Nernst equation to calculate the voltage at any point. The actual value depends on both the concentrations and charges of the ions.
What happens to the cell potential, ε, as the reaction progresses?
Since Q *increases*, (more products vs reactants), cell potential *decreases*
What happens during recharging of a Ni-Cd battery?
The reactions proceed from right to left (opposite, nonspontaneous direction).
Why are some Ni-Cd batteries vented?
This contains any generation of oxygen and hydrogen gases until they can recombine back to water. Such generation typically occurs during rapid charge and discharge, and exceedingly at overcharge condition
In general, an oxidizing agent *cannot* oxidize any reducing agent that lies _____________ it on the standard reduction potential table.
above ✷so Pb²⁺ cannot oxidize Cu(s)
energy density
amount of energy a cell can produce relative to the mass of battery material ￫ more mass needed to produce same amount of energy = low energy density ￫ less mass needed to produce same amount of energy = high energy density
For electrolytic cells, the electrode with the more positive reduction potential is the ____________.
anode *NOTE* - oxidation still occurs here, but with the compound that is more naturally reduced, so it is nonspontaneous
For galvanic cells, the electrode with the less positive reduction potential is the ___________.
anode Explanation: it will be spontaneously oxidized
Anions ALWAYS migrate towards the __________
anode Mnemonic: *Anion to anode* Explanation: In a galvanic cell, it's negative, so this may be counterintuitive, but read this: "Anions, negative ions (e.g. Cl⁻), are flowing *from the salt bridge* toward the anode to balance the positive charge of the Zn²⁺ ions produced. Easier to remember for electrolytic cells because anode is positive. So please just remember this convention 😊 ❗️but ELECTRONS *always* flow from anode to cathode!
No matter the type of cell, anions are always attracted to the ____________.
anode Oxidation always occurs at the anode and in a galvanic cell, electrons are sent off to the cathode. Thus, having lost electrons, the anode will attract anions from the salt bridge In an electrolytic cell, the anode is + and is attached to the + terminal of the voltage source, where it attracts anions from the solution
Cell diagrams are written from ____________ to _______________.
anode ⟶ cathode
For *all* electrochemical cells, the movement of electrons is from ____________ ⟶ _____________.
anode ⟶ cathode
What equation can be used to determine amount of gas liberated during electrolysis?
electrodeposition equation
￫ When recharging, a lead-storage battery is part of a ____________ circuit. ￫ What would happen to the net equation of the discharging lead-storage battery and its E°cell?
electrolytic Equations and electrode charge designations are opposite because an external source reverses the electroplating process and *concentrates* the acid solution So, at the end of this electrolytic process, we end up with the fully-charged battery again, ready to act as a voltaic cell and discharge.
Charging is a ____________ state.
electrolytic Explanation: requires energy (nonspontaneous)
What type of reaction is the electrolysis of molten NaCl?
electrolytic nonspontaneous when *discharging*: overall, molten NaCl is decomposed into Cl₂ (g) and Na (l). - Na⁺ ions migrate towards the cathode (-), where they are reduced to Na (l). At this temperature, sodium is a liquid and is less dense than molten NaCl so it floats to the top for easy removal - Cl⁻ ions migrate towards the anode (+), where they are oxidized to Cl₂ (g)
Free energy is an _____________ property.
extensive ✷it depends on the amount you're dealing with 📸 pink
One _____________ is equivalent to the amount of charge contained in one mole of electrons.
faraday (F): 96,485 C ⟶ rounded to 1 × 10⁵ C
In a galvanic cell, the copper electrode is ____________ mass. (Use a Daniel cell to apply this concept).
gaining Explanation: Cu²⁺ ions (solution) are reduced and form more Cu metal (cathode).
Discharging is a ________________ state.
galvanic Explanation: releases energy (spontaneous)
What is the electrodeposition equation and what does it help elucidate?
helps determine the # of moles of element being deposited on a plate *M* is the amount of metal ion deposited at a specific electrode *I* is current *t* is time *n* is # of electron equivalents for a specific metal ion *F* is Faraday's constant ✴︎✴︎equation can also be used to determine amount of gas liberated during electrolysis mnemonic: calculating *moles of M*etal, *It* is *N*ot *F*un.
A reduction potential tell us what?
how likely a compound is to be reduced ✷ The more positive the value, the more it "wants" to be reduced.
In a galvanic cell, the concentration of zinc solution is ____________. (Use a Daniel cell to apply this concept).
increasing Explanation: Zn is the anode, and the metal electrode is losing mass as it's being oxidized to Zn²⁺ ions, which flow into the solution *more ions in solution = greater concentration*
Voltage is an ____________ property.
intensive ✷it doesn't matter how much you're dealing with, the value is always the same if you have to multiply a half-reaction by 2, *don't* multiply the voltage by 2 📸 green
Why are lead-storage batteries rechargeable?
the products don't leave the electrodes (PbSO₄ is insoluble and thus is readily available for the recharging process).
electromotive force (emf)
the voltage or electrical potential difference of the cell describes the current either GENERATED during a reaction (spontaneous) or NEEDED (nonspontaneous) to drive a reaction.
How do you find the voltage of a concentration cell?
use the Nernst equation When finding Q, take the concentration of the metal on the *less* concentrated side over the concentration of the metal on the *more* concentrated side. Standard cell potential will be 0
What is the definition of standard cell potential?
voltage measured when the cell is operating under standard conditions *standard conditions*: all solids are in pure form, solutions are all at 1M concentrations, and temperature is 25°C.
When Q = 1, how does ε compare to ε°?
ε = ε°
When Q < 1, how does ε compare to ε°?
ε > ε°
What is the relationship between Keq and ε°cell?
• When Keq > 1 ⟶ ε°cell > 1 (spontaneous) • When Keq < 1 ⟶ ε°cell < 1 (nonspontaneous) • When Keq = 0 ⟶ ε°cell = 0 (concentration cell)
What is the oxidation number for hydrogen?
⁺¹
What is the oxidation number for oxygen? Are there any exceptions to this?
⁻² yes, for diatomic oxygen 0 and hydrogen peroxide: ⁻¹
What is the oxidation number for fluorine?
⁻¹
What is the oxidation number for the halogens *most of the time*? When would they be different?
⁻¹ Exceptions: if they are bonded to F or O
How is a concentration cell distinct from a galvanic cell? Why is the fact useful?
⇨ The electrodes are chemically identical = same reduction potential. ⇨ Therefore, current is generated as a function of *concentration gradient* not reduction potential. There *must* be different concentrations of ions in the two compartments for there to be a measurable voltage and current. ✴︎all the rest is just like a galvanic cell
Show the configuration of a lead storage battery when it is functioning as an electrolytic cell (after it has been fully discharged).
⇾ *2 PbSO₄ electroplated lead electrodes* ⇾ Connected by *dilute [H₂SO₄]* ➷Current is used to drive a nonspontaneous reaction (recharge the battery) with the solid PbSO₄ ✷ The products of the reactions at the anode and cathode are insoluble (PbSO₄ in each case). This means that these substances are readily available to participate in the reverse reactions that recharge the cell!
Show the configuration of a lead storage battery when it is functioning as a voltaic cell (fully charged). What are the oxidation states and half reactions that occur? What will happen to the anode and cathode?
⇾ *Pb anode* ⇾ *PbO₂ cathode* (porous) ⇾ Connected by *conductive [4M H₂SO₄]* *Original oxidation states*: Pb⁰ (s) Pb⁺⁴O₂⁻² Oxidation half-reaction: Pb (s) + HSO₄⁻ (aq) ⟶ PbSO₄ (s) + H⁺ + 2 e⁻ E°red = -0.356 V *New oxidation state*: Pb⁺²(SO₄)⁻² Reduction half-reaction: PbO₂ (s) + SO₄²⁻ (aq) + 4H⁺ + 2 e⁻ ⟶ PbSO₄ (s) + 2H₂O E°red = 1.685 V *New oxidation state*: Pb⁺²(SO₄)⁻² *Overall* - Discharging caused electrodes to plate with PbSO₄, and it also diluted the acid electrolyte (*see water that was produced ↑*) *Anode* - Pb is oxidized to Pb²⁺, which then attracts SO₄²⁻ to form PbSO₄. Electrons and protons released. *Cathode* - electrons and protons are consumed here. The porous nature allowed the electrolyte (H₂SO₄) to solvate the cathode into lead and oxide ions. Then, the H⁺ in solution reacted with the O₂²- to produce water. The remaining SO₄²⁻ react with Pb to produce electroplated PbSO₄. ➷Drives a spontaneous redox reaction that produces a current
What is the standard reduction potential of a lead-storage battery?
≅ 2 V Solution: Half Reaction @ anode (oxidation): Pb (s) + HSO₄⁻ (aq) ⟶ PbSO₄ (s) + H⁺ + 2 e⁻ E°red = -0.356 V + Half Reaction @ cathode (reduction): PbO₂ (s) + SO₄²⁻ (aq) + 4H⁺ + 2 e⁻ ⟶ PbSO₄ (s) + 2H₂O E°red = 1.685 V ________________________________________________________________________ Pb (s) + PbO₂ (s) + 2H₂SO₄ (aq) ⟶ 2PbSO₄ (s) + 2H₂O E°cell = E°red, cathode - E°red, anode E°cell = 1.685 V - (-0.356 V) = 2.041 V
What is the shorthand notation for a galvanic cell?
✷ | single lines represent phase boundaries ✷ || double line represents salt bridge ✷ anode always on the left ✷ cathode always on the right
Why have nickel-metal hydride (NiMH) batteries replaced Ni-Cd batteries?
￫ even more energy density ￫ more cost effective ￫ way less toxic Essentially, a metal hydride is used instead of a pure metal anode.
What are the advantages of Ni-Cd batteries?
￫ higher energy density ￫ large surge current (large current early in discharge cycle, good for devices that require rapid response, i.e. remote) ￫ low internal resistance
What are the downsides of Ni-Cd batteries?
￫ limited in application ￫ expensive ￫ toxic ￫ even though they are high capacity, they lose capacity very quickly *thus, they have been replaced with nickel-metal hydride (NiMH) batteries.
