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Chapter 19: Problem 16

What reactions occur when a lead storage battery is recharged?

Short Answer

Expert verified
The lead sulfate (PbSO4) on both plates is converted back into lead (Pb) and lead dioxide (PbO2) with concurrent regeneration of sulfuric acid.

Step by step solution

01

Understanding Lead-Acid Battery Discharge

A lead storage battery consists of lead dioxide (PbO2) as the positive plate and sponge lead (Pb) as the negative plate, with a sulfuric acid (H2SO4) electrolyte. During discharge, the overall chemical reaction is: \[ \text{Pb} + \text{PbO}_2 + 4\text{H}^+ + 2\text{SO}_4^{2-} \rightarrow 2\text{PbSO}_4 + 2\text{H}_2\text{O} \] This converts the lead and lead dioxide into lead sulfate (PbSO4) and water, releasing electrical energy.
02

Reversing the Reaction During Recharge

When the lead storage battery is recharged, the discharge reaction above is reversed. Electrical energy is supplied to convert lead sulfate back into lead at the negative electrode and lead dioxide at the positive electrode. The reverse reaction is: \[ 2\text{PbSO}_4 + 2\text{H}_2\text{O} \rightarrow \text{Pb} + \text{PbO}_2 + 4\text{H}^+ + 2\text{SO}_4^{2-} \] This regenerates the original reactants of the battery, restoring its potential to store energy.
03

Importance of Sulfuric Acid Concentration

During both discharge and recharge, the concentration of sulfuric acid changes. In recharge, as lead sulfate is reconverted, sulfuric acid is regenerated in the electrolyte solution. Maintaining the right concentration of \( \text{H}_2\text{SO}_4 \) is crucial as it affects the battery's ability to function efficiently.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Chemical Reactions
A lead-acid battery operates on the basis of key chemical reactions that transform chemical energy into electrical energy. During the discharge process, the battery generates electrical energy by initiating a specific reaction. This involves lead and lead dioxide plates transitioning into lead sulfate and water.
The discharge reaction can be described by the equation:
  • Lead
  • Lead dioxide
  • Sulfuric acid
These react to produce lead sulfate and water, effectively releasing electrical energy in the process. During recharge, electrical energy supplied to the battery causes this reaction to reverse, converting lead sulfate and water back into lead, lead dioxide, and sulfuric acid. This allows the battery to store energy once again.
Electrochemistry
Electrochemistry plays a pivotal role in the functioning of a lead-acid battery. This branch of science studies the relationship between chemical reactions and electrical energy.
In a lead-acid battery, electrochemical reactions at the electrodes convert chemical energy to electrical energy, through the movement of electrons and ions in the electrolyte.
  • At the positive electrode, electrons are released when lead dioxide reacts with hydrogen ions and sulfate ions.
  • At the negative electrode, lead reacts with the sulfate ions, allowing for electron flow through the external circuit, generating electricity.
During recharge, these processes are reversed, with the battery needing external electrical energy to drive the reaction back to the original state, thereby storing energy for future use.
Sulfuric Acid Concentration
The concentration of sulfuric acid within the electrolyte is vital for the efficient operation of a lead-acid battery. Sulfuric acid acts as the medium that enables the flow of ions necessary for the electrochemical reactions to occur.
During the discharge process, some of the sulfuric acid is consumed, forming lead sulfate and water. This change in concentration can impact the overall efficiency and power output of the battery.
  • As the battery discharges, sulfuric acid concentration decreases.
  • During recharge, the acid concentration increases as lead sulfate is converted back.
Maintaining an optimal concentration of sulfuric acid ensures that the battery not only recharges effectively but also sustains its energy storage capacity and longevity.
Battery Recharge Process
Recharging a lead-acid battery is a process of reversing the chemical reactions that occurred during discharge, which requires external electrical energy. This energy flows into the battery and initiates the conversion of lead sulfate and water back into lead, lead dioxide, and sulfuric acid.
The recharge process is crucial for restoring the battery's energy storage capability. Adequate and timely recharges are necessary to prevent irreversible lead sulfate buildup, which can diminish the battery鈥檚 performance over time.
  • Ensure a compatible charger is used to suit the battery's voltage and capacity.
  • Monitor recharge times, avoiding overcharging, which can lead to overheating and damage.
By understanding and managing the recharge process appropriately, the battery's lifespan can be maximized and its efficiency sustained.

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Most popular questions from this chapter

Balance the following redox equations. All occur in acid solution. (a) \(\operatorname{sn}(s)+H^{+}(a q) \rightarrow S n^{2+}(a q)+H_{2}(g)\) (b) \(\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(\mathrm{aq})+\mathrm{Fe}^{2+}(\mathrm{aq}) \rightarrow\) \(\mathrm{Cr}^{3+}(\mathrm{aq})+\mathrm{Fe}^{3+}(\mathrm{aq})\) (c) \(\mathrm{MnO}_{2}(\mathrm{s})+\mathrm{Cl}^{-}(\mathrm{aq}) \rightarrow \mathrm{Mn}^{2+}(\mathrm{aq})+\mathrm{Cl}_{2}(\mathrm{g})\) (d) \(\mathrm{CH}_{2} \mathrm{O}(\mathrm{aq})+\mathrm{Ag}^{+}(\mathrm{aq}) \rightarrow \mathrm{HCO}_{2} \mathrm{H}(\mathrm{aq})+\mathrm{Ag}(\mathrm{s})\)

Use cell notation to depict an electrochemical cell based upon the following reaction that is productfavored at equilibrium. \(\mathrm{Fe}^{3+}(\mathrm{aq})+\mathrm{Ag}(\mathrm{s})+\mathrm{Cl}^{-}(\mathrm{aq}) \rightarrow \mathrm{Fe}^{2+}(\mathrm{aq})+\mathrm{AgCl}(\mathrm{s})\)

A Living organisms derive energy from the oxidation of food, typified by glucose. $$ \mathrm{C}_{0} \mathrm{H}_{12} \mathrm{O}_{6}(\mathrm{aq})+6 \mathrm{O}_{2}(\mathrm{g}) \rightarrow 6 \mathrm{CO}_{2}(\mathrm{g})+6 \mathrm{H}_{2} \mathrm{O}(\ell) $$ Electrons in this redox process are transferred from glucose to oxygen in a series of at least 25 steps. It is instructive to calculate the total daily current flow in a typical organism and the rate of energy expenditure (power). (See T. P. Chirpich: Journal of Chemical Education, Vol. \(52,\) p. 99 1975.) (a) The molar enthalpy of combustion of glucose is \(-2800 \mathrm{kJ} / \mathrm{mol}\) -nan. If you are on a typical daily diet of 2400 Cal (kilocalories), what amount of glucose (in moles) must be consumed in a day if glucose is the only source of energy? What amount of \(\mathrm{O}_{2}\) must be consumed in the oxidation process? (b) How many moles of electrons must be supplied to reduce the amount of \(\mathrm{O}_{2}\) calculated in part (a)? (c) Based on the answer in part (b), calculate the current flowing, per second, in your body from the combustion of glucose. (d) If the average standard potential in the electron transport chain is \(1.0 \mathrm{V},\) what is the rate of energy expenditure in watts?

In the presence of oxgyen and acid, two half. reactions responsible for the corrosion of iron are $$ \begin{array}{c} \mathrm{Fe}(\mathrm{s}) \rightarrow \mathrm{Fe}^{2+}(\mathrm{aq})+2 e^{-} \\ \mathrm{O}_{2}(\mathrm{g})+4 \mathrm{H}^{+}(\mathrm{aq})+4 e^{-} \rightarrow 2 \mathrm{H}_{2} \mathrm{O}(\ell) \end{array} $$ Calculate the the standard potential, \(E^{\circ},\) and decide whether the reaction is product-favored at equilibrium. Will decreasing the pH make the reaction less thermodynamically product-favored at equilibrium?

The products formed in the electrolysis of aqueous \(\mathrm{CuSO}_{4}\) are \(\mathrm{Cu}(\mathrm{s})\) and \(\mathrm{O}_{2}(\mathrm{g}) .\) Write equations for the anode and cathode reactions.
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