Current Technology for Electric Vehicles

The only ZEVs at this point are electric vehicles. Several car manufacturers sell electric vehicles to fleet buyers, and kits are available for individuals who want to convert ordinary cars into electric vehicles. For example, the Honda Insight and the Toyota Prius are hybrid gasoline/electric vehicles that do not require external charging. Although they are superior to ordinary gasoline vehicles in terms of both gas mileage and emissions, they are not ZEVs. Battery technology is the key to making electric vehicles. Some commercially available and most home-built electric vehicles use lead-acid batteries, similar to ordinary car batteries. To understand how to improve electric vehicles, we must first understand how batteries work.

Batteries in Electric Vehicles

Batteries are one of the most common applications of electrochemistry. Batteries harness oxidation-reduction, or redox, reactions to produce energy. Redox reactions involve the transfer of electrons from one substance to another. The electrochemical cell is the basis of every battery. In an electrochemical cell, electrons are transferred between two substances, even though the reactants are not in physical contact with each other. This electron transfer produces a current. If the cell is galvanic, the reaction proceeds spontaneously and the resulting current can be used to do work such as driving an electrical motor, or illuminating a flashlight. The tin(II)/ferricyanide cell you constructed in Experiment 5, "Electrochemistry: Cells and Redox Reactions" is a good example of a cell-based electrochemical reaction.

All redox reactions can be separated into two half-reactions. Figure 6 is a schematic of a lead-acid battery (an ordinary car battery). In the lead-acid battery, Pb is oxidized to Pb²⁺ at the anode (negative battery terminal) as shown in Equation 10 and Pb⁴⁺ is reduced to Pb²⁺ at the cathode (positive battery terminal) as shown in Equation 11.

Battery

Figure 6

This schematic of a lead-acid battery shows the oxidation-reduction reaction between lead and lead (IV) oxide, mediated by sulfuric acid. Pb is the anode and PbO₂ is the cathode.

Pb (s) + SO₄^2- (aq) ↔ PbSO₄(s) + 2e-

e^o = 0.356 V

(10)

Reduction half-reaction:

PbO₂ (s) + SO₄^2-(aq) + 4H⁺(aq) + 2e- ↔

PbSO₄ (s) + 2H₂O (l)

e^o= 1.685 V

(11)

During discharge (when the battery is being used to generate electricity), the sum of the two half-reactions is the following cell reaction:

Pb (s) + PbO₂ (s) + 2 H₂SO₄ (aq) ↔

2 PbSO₄ (s) + 2 H₂O (l)

e^o= 2.04 V

(12)

Notice that as the cell reaction proceeds, PbSO₄ precipitate forms on both electrodes.

For the lead-acid battery, e^o= 2.04 volts. (See Box 2, below, for a review of electrochemical potentials.) In the lead-acid battery,

because all the reactants and products other than H₂SO₄ are solids or liquids that do not appear in the equilibrium expression. This means that the actual battery voltage depends on the concentration of sulfuric acid, which is consumed as the battery discharges. (During discharge when e > 0, the cell is galvanic and the reaction occurs spontaneously, since ΔG < 0.) Common car batteries are twelve-volt batteries, because they are actually six lead-acid cells (each of which has e^o= 2.04 volts as seen in Equation 12, above) connected in series and housed in a single container.

Box 2

Electrical Potentials in Lead-Acid Batteries

Every half-reaction has a standard electrical potential, designated as e^o. By convention, we write the half-reactions as reductions; i.e., the addition of electrons to the oxidized form of the substance. The two half reactions that comprise the redox reaction in a lead-acid battery are the following (written as reduction reactions):

PbO₂(s) + SO₄^2- (aq) + 4H⁺ (aq) + 2e- ↔

PbSO₄(s) + 2H₂O (l)

e^o= 1.685 V

(13)

PbSO₄(s) + 2e- ↔ Pb(s) + SO₄^2-(aq)

e^o= -0.356 V

(14)

e^o is a measure of the tendency of the reactant in the half-reaction to accept electrons relative to H⁺ + e- → ¹/₂H₂ (every species in its standard state). We obtain the e^o for the whole cell or reaction by adding together the standard potentials for the half reactions.

e^o_cell= e^o_cathode+ e^o_anode

(15)

Note that the anode half-reaction must be written in the reverse direction from Equation 14 above (as an oxidation reaction), and that the sign of e^ochanges when the reaction is reversed. Rewriting Equations 13 and 14, and adding them together gives:

PbO₂(s) + SO₄^2- (aq) + 4H⁺ (aq) + 2e- ↔

PbSO₄(s) + 2H₂O (l)

e^o= 1.685 V

(13)

Pb (s) + SO₄^2- (aq) ↔ PbSO₄(s) + 2e-

e^o= 0.356 V

(16)

e^o_cell= 1.685 + 0.356 = 2.04 V

(17)

Recall, a potential of 2.04 V will only be measured if all the reactants and products are in their standard state, which means a concentration of 1 M for aqueous solutions. The Nernst Equation (Equation 18, below) allows us to calculate the potential when we have conditions other than the standard state.

(18)

Here, R is the gas constant, 8.314 J/mol-K; T is the Kelvin temperature; n is the number of moles of electrons transferred in the cell reaction; F is the value of the Faraday constant, 9.65 x10⁴C/mol (i.e., the absolute value of the electric charge of one mole of electrons); and Q is the reaction quotient, a quantity of the same form as the equilibrium constant but employing activities actually present in the cell rather than those at equilibrium.

To recharge the lead-acid battery, the redox reaction (Equation 12) must be run in reverse. The electrical potential for this reverse reaction is negative (e^o< 0); therefore, the reaction is nonspontaneous. This means that energy is needed to drive the reaction (i.e., recharge the battery). Notice that during recharging the PbSO₄ precipitate on the electrodes is oxidized and reduced to PbO₂ and Pb, respectively. Once recharging is complete, the forward reaction can run spontaneously again and create electrical work.

The main drawbacks to lead-acid batteries is that they are heavy and must be recharged over a period of hours. Depending on the weight of the vehicle and driving conditions, twenty batteries may be required to power the car for forty miles before recharging. The batteries weigh up to eighty pounds each, and they take up quite a bit of space. Although the weight and volume of the internal combustion engine is eliminated, electric vehicles are still heavier than their gasoline counterparts and often sacrifice trunk space for battery storage. In addition, the discharged batteries must be plugged into a battery charger for six or more hours for recharging.