To what extent does the concentration of sulfuric acid in a lead-acid cell affect the voltage at the end of a discharge cycle and the accumulation of lead sulfate on a lead electrode at standard thermodynamic conditions?

This research examines how sulfuric acid content in a lead-acid cell impacts lead electrode sulfation and cell voltage at the end of a discharge cycle (otherwise stated as the discharge time). This yields an optimum voltage and sulfation concentration.

Lead-acid batteries, invented by Gaston Planté in 1860, are the earliest rechargeable batteries still in use (Yamaguchi). Lead-acid batteries power car motors and emergency electrical storage. Lead-acid batteries and contemporary batteries must have a long ageing life and slow discharge time. Battery life and parameter optimization research continues. To meet battery demands, it's necessary to study electrochemistry parameters.

 

Electrochemistry combines my interests in electrical and chemical engineering. Since I use remotes, phones, and other devices at home, I saw how important batteries are. I investigated the problem of mobile phones discharging and dying quickly for a battery of my choice. Realizing I could modify lead-acid cell parameters piqued my curiosity in slowing discharge time to improve battery life.

Secondary (rechargeable) lead-acid batteries power electrical systems. Lead-acid cells are voltaic and electrolytic, unlike primary batteries. Industrial lead-acid batteries have cells with lead electrodes and diluted sulfuric acid as electrolytes (Battery Basics) (H₂SO₄). The lead electrodes are lead and lead dioxide. The lead dioxide electrode reduces, whereas the lead electrode oxidizes. Each cell stores 2.1 volts.

 

Since lead-acid batteries are rechargeable, discharging (chemical energy to electrical energy), reactions differ from charging replies (converting electrical energy into chemical energy).

 

Sulfate and hydrogen ions form from dilute sulfuric acid. Due to electrical resistance, lead in the electrode forms lead (II) ions and two electrons. Lead sulfate is produced on the anode when acid sulfate ions react with lead (II) ions. The information (II) ions release electrons that pass through the wire to the cathode (positive electrode). Anode reaction: (Lugo).

 

Oxidation Half-Equation: Pb (s) + SO₄²- (aq) → PbSO₄ (s) + 2e^-

 

Electrons draw electrolyte hydrogen ions to the cathode. Hydrogen ions react with electrode lead dioxide oxygen to generate water and lead (II) ions. Lead ions react with electrolyte sulfate ions to cause cathode-surface lead sulfate. Lugo's half-equation describes this cathode reaction:

 

Reduction Half-Equation: PbO₂ (s) + 4H+ (aq) + SO4₂^- (aq) + 2e^- → PbSO₄ (s) + 2H₂O (l)

 

Combining the oxidation and reduction half-equations creates a discharge redox reaction. Figure 1 (Singh) displays the redox reaction figure.

 

Pb (s) + PbO₂ (s) + 2H₂SO₄ (aq) → 2PbSO₄ (s) + 2H₂O (l)

A lead-acid battery's function is to store charge to be utilized as a power source. Charging a lead-acid battery in an electrolytic reaction reverses the reaction during discharging. An external DC power source is necessary for setting.

 

The electrons are compelled to move to the cathode (negative electrode) when a power supply is linked between the two electrodes, as shown in Figure 2 (Singh). See how the polarity of the cathode and anode has changed. Lead (II) sulfate transforms into lead and sulfate ions when the electrons are forced to the negative electrode. A reduction half-equation is used to depict this process (Lugo):

 

Reduction Half-Equation: PbSO₄ (s) + 2e^- → Pb (s) + SO₄^2- (aq)

 

As hydroxide ions are preferred to be discharged, the anode's positive charge allows water to break down into oxygen (O₂) and hydrogen ions while the reaction occurs at the cathode.

 

2H₂O (l) → O₂ (g) + 4H⁺ (aq) + 4e^-

 

On the anode, the oxygen from the equation above interacts with the lead sulfate:

 

PbSO₄ (s) + O₂ (g) + 2e^- → PbO₂ (s) + SO₄^2- (aq)

 

An oxidation half-equation is created at the anode by combining the two equations above:

 

Oxidation Half-Equation: PbSO₄ (s) + 2H₂O (l)→ PbO₂ (s) + SO₄^2- (aq) + 4H⁺ (aq) + 2e^-

 

The following is the redox reaction that occurs during charging when the reduction and oxidation half-equations are combined:

 

PbSO₄ + 2H₂O → Pb + PbO₂ + 2H₂SO₄

For batteries to be practical, they must have good battery life. The number of cycles a battery experiences before giving out is called aging. One complete discharge and charge time is referred to as a cycle. Sulfation, or the buildup of lead sulfate on the electrodes, is one of the significant causes of battery failure.

 

Lead sulfate dissolves during the charging phase and recombines with lead and sulfate ions. Yet because not all of the chemical energy generated is used to dissolve the lead sulfate, not all of it is converted back to its constituent parts during the electrolysis reaction (Gandhi). Some energy is put into it by creating oxygen gas, and some are lost as heat. Another topic is whether lead sulfate is generated chemically or electrochemically. This indicates that there are two types of sulfation, one of which is "reversible" and the other of which is "irreversible" (Yamaguchi). Reversible sulfation refers to the ability to dissolve and reincorporate lead sulfate crystals back into the electrolyte and electrode after they have formed on the surface of an electrode.

 

When sulfate ions cannot be broken down into constituent parts, irreversible sulfation occurs. Chemical reactions, not electrical ones, result in irreversible sulfation (Yamaguchi). As shown in Figure 3, when a cell is depleted and left uncharged for a while, chemical reactions take place. The lead sulfate ions on the surface of the electrode combine with sulfuric acid to produce this chemical reaction.

The lead sulfate on the battery could theoretically all be broken down into constituent parts. Practically speaking. However, this is implausible because there is typically downtime between discharging and charging.

A typical lead-acid battery has 4.5 to 6.0 mol dm^-3 of material (Berera). The electrolyte content impacts the cell potential and, thus, the discharge duration because a higher likelihood results in a longer discharge time.

 

A lead-acid battery's voltage will rise with a higher electrolyte concentration, resulting in a more considerable electrode potential difference. This impact is partially the result of equilibrium reactions (Gaidis). According to Section 3.1, sulfuric acid is split into sulfate and hydrogen ions. However, because of the diprotic nature of sulfuric acid, the removal of the second hydrogen atom will trigger an equilibrium reaction:

 

H₂SO₄ + H₂O → HSO₄^- + H₃O⁺

 

HSO₄^- + H₂O ⇌ SO₄^2- + H₃O⁺

 

A higher H₂SO₄ concentration will accelerate the forward reaction since it will tilt the equilibrium in favor of the product side. A cell with a greater electrode voltage exhibits a more forward response. The further a reaction is from the electrode in a lead-acid battery, the longer the discharge time, and the more balance.

 

The Nernst Equation also demonstrates the effect of concentration on voltage. On electrochemical cells, the electrode potential of one half-cell can be calculated using the Nernst equation. The electrode potential of the lead or lead dioxide electrodes can also be found in the case of lead-acid batteries. In the following calculations, it will be assumed that the lead-acid cell is functioning as a voltaic cell because this study focuses on the discharge rate. As stated below (Berera), the Nernst Equation:

 

E = \(-\frac{RT}{nF}ln(Q)\)

 

E = Electrode Potential of one half-cell

 

R = Universal Gas Constant

 

F = Faraday’s Constant

 

T = Absolute Temperature

 

n = amount of electrons transferred

 

Q = Reaction quotient

 

Since the spontaneity of an electrochemical cell is influenced by electrode potential, this equation relates to the Gibbs Free Energy equations. ΔH stands for enthalpy change, and S for entropy change.

 

ΔG^Θ = -nFE^Θ

 

ΔG^Θ = ΔH - TΔS

 

Gibbs free energy increases (gets closer to zero) as entropy decreases while the temperature remains constant, increasing the electrode potential. Consequently, it can be said that electrolytes at higher concentrations have less entropy than those at lower concentrations. The value of the reaction quotient is assumed to be at t = 0, as the Nernst Equation determines the highest electrode potential a half-cell has. As higher concentration systems are further from equilibrium, the reaction quotient at the start of the Reaction will be less in value. The electrode potential of a half-cell is raised by a lower reaction quotient value, which also boosts the electrode potential difference. The energy cost of the concentration will be far greater than the increase in electrode potential. Enthalpy is a Gibbs Free Energy formula component, which explains this. Gibbs's free energy is harmful since the lead-acid cell is assumed to be voltaic. The entropy is, therefore, positive, while the enthalpy is negative. As a result, the half-internal cell's energy grows in absolute terms as increasing concentration causes entropy to decrease. As a result, higher concentrations demand more fuel to maintain the electrochemical Reaction. As a result, the maximum engagement that can be used in a lead-acid cell is practically limited.

A single lead-acid cell with a maximum potential of roughly 2 volts will be constructed.

 

The construction of the cell is shown in Figure 4. A DC supply will be linked to the lead electrodes during charging by cables with crocodile clips on them. A smaller propeller will be employed during discharge.

The amount of lead sulfate that forms on the anode's surface will only be measured on the anode (the lead electrode) in this experiment. Since the cathode (the lead dioxide electrode) is prone to corrosion, which is brought on by oxidation ("Delivering Batteries..."), the mass of lead sulfate on the cathode will not be detected. This implies that the lead dioxide electrode's overall mass will decrease over the experiment.

 

In contrast to the norm of 4.5 to 6.0 moldm^-3, a lower concentration will be employed in the approach.

 

This is because, as was explained in Section 3.3, systems with higher concentrations require more energy. Lead-acid cells used in an industry often have electrolyte additives, metal alloys in the electrodes, a lead dioxide electrode that wasn't electrochemically created, and other features that make the lead-acid cells built for this study different from those used in business. This could prevent the lead dioxide electrode from effectively delivering the voltage necessary for the elevated concentration. The concentrations employed in this procedure will thus not be greater than 6.00 mol dm^-3.

 

Lastly, just the discharge cycle of the lead-acid cell's life will be measured for voltage; the charging cycle will not be noted.

Instead of the six lead-acid cells that are typically utilized, the approach just uses one. This was used to make it simpler to control the sulfuric acid concentration and to conserve resources and space.

 

The timing of discharge was calculated using two different methods. The cell was initially discharged until an end voltage of 0.18 V, at which point the time was recorded. The voltmeter's reading being between values and having trouble stopping the timer in time were just two of the sources of inaccuracy. The optimum approach was to discharge the cell for 5 minutes before recording the final voltage.

Because it is corrosive, sulfuric acid should never be exposed to the skin or eyes. When handling sulfuric acid, wear gloves and eye protection. Always place the device beneath a fume hood when discharging and charging the cell. Keep a safe distance away from the cell when it's in use. Lead is hazardous and harmful to the environment. After using the lead electrodes, dispose of them in a lab-available waste container. Never flush lead-containing materials down the toilet or put them in the usual trash.

• The electrodes' surface can be extensively cleaned using sandpaper and ethanol. Record the mass of both electrodes after weighing them.
• With a volumetric cylinder, transfer 200 ml of a 3.00 mol dm^-3 solution to a 300 ml beaker.
• Build the cell by positioning the two lead electrodes according to Figure 4 in Section 3.5. Two wires should be connected to the DC power supply using crocodile clips. Ensure the electrodes are not in contact with one another and the power source is turned off.
• Connect the voltmeter in parallel to the cell using two wires.
• Start the cell's 10-minute charge.
• Connect the crocodile clips to the propeller and unplug the wires from the DC source once 10 minutes have passed. Release for five minutes.
• After five minutes, note the voltage measurement on the voltmeter. Set the clock for three minutes.
• Take out the anode (the electrode that seems smoother) and weigh it on a scale. To protect the equipment, use paper. Reconnect the electrodes using crocodile clips after placing them back inside the beaker. Everything is finished in 3 minutes.
• Start charging for 10 minutes after 3 minutes. Steps 6 through 8 should be repeated two more times.
• Repeat procedures 1 through 9 with the next concentration to be evaluated.