7.1 Redox Reactions
7.10 Battery Facts
A battery is the ultimate portable voltaic cell. It is a self-contained little power factory that doesn't spill. Its electric power production is due very active oxidizing and reducing agents. Batteries come in all kinds of shapes, sizes, and potentials (volts). There are many batteries that we buy directly and put in our devices. There are specialized batteries that work "behind the scenes" and you never see them - but you know they are there because there is electric power present.
One of the biggest breakthroughs on batteries was the invention of the dry cell battery. It's an old idea that was patented in 1892. This was the first portable battery that didn't "spill" if you knocked it over or turned it upside down.
How is there no liquid or solution in our modern day batteries? You really don't have to have a true aqueous liquid state in the way we think about aqueous solutions which you can pour and measure out in a graduated cylinder. The electrolyte in the dry cell was really contained in a paste. As you know, pastes don't spill. Pastes don't flow in the way water flows or spills. The electrolyte is dissolved in water, but very little water with lots of electrolyte salt are mixed together to make a paste. After the rest of the cell is constructed, it is sealed and ready for use in all orientations. There are enough free aqueous ions in the paste to easily carry the necessary electrolytic current in order for the battery to operate.
Batteries are generally broken into three types. Drum roll....
The primary cell is a voltaic cell (battery) that is made to be a one-time use battery. This means a primary cell is a non-chargeable cell. You do not recharge primary cells. The chemical reactions (and side reactions) in primary cells really only go in one direction and that is from reactants to products. The reverse reaction does not occur efficiently enough to allow one to regenerate the reactants from the products. Primary cells are meant to be used once and then discarded. Let's list some popular primary cells.
Alkaline Cells: Probably the most familiar of batteries is the alkaline cell. These are typically cells that have a voltage of 1.5 V. These come in various shapes and sizes. Certainly the most popular shape is the familiar cylindrical shape of the "letter" cells. "Letter" meaning the D-cell, C-cell, AA-cell, and AAA-cell. All of these cells share the same exact chemistry. The cathode active material (what gets reduced) is MnO2(s). The anode is zinc metal. MnO2 is a non-conductive oxide so in order to make electrical contact, graphite powder is mixed with the MnO2 to make a paste. The graphite then acts as an extended electrode surface in contact with all the active MnO2. The two half reactions are shown below.
|cathode:||2 MnO2(s) + H2O + 2 e- → Mn2O3(s) + 2 OH-||E°red = +0.15 V|
|anode:||Zn (s) + 2 OH- → ZnO(s) + H2O + 2 e-||E°ox = +1.28 V|
|overall:||2 MnO2(s) + Zn (s) → Mn2O3(s) + ZnO(s)||E°cell = +1.43 V|
Shorthand cell notation for this is:
Notice that the overall standard cell potential for this cell will be +1.43 V. Adjusting for the actual concentrations used in the manufacture of the batteries (using the Nernst equation), you get a battery with somewhere between 1.50 and 1.65 V. The diagram shown here is a cut-away diagram of a typical alkaline cell. Note that the positive terminal in the diagram is just the contact point of a stainless stell can that is in electrical contact with the active cathode material of MnO2. The MnO2 is a powder and is mixed with graphite and aqueous KOH to make a paste. THe graphite helps ensure better electrical contact throughout the compressed powder/paste and is not shown in the shorthand notation. The zinc anode in modern batteries is a compressed Zn powder mixed with aqueous KOH (another paste). In powder form, the zinc has more surface area for electron transfer which allows for higher electric currents. The "current collector" in this case acts just like an inert electrode and simply allows for electrical contact with the zinc anode paste.
The "porous separator" is simply a paper-like material that provides the functionality of a salt bridge. It allows electrolytic current to flow across it, but keeps the reactive components (Zn and MnO2) separated so they do not react directly. Note that the entire cell is in a strong solution (although with so much solid in it that it is a paste) of KOH which is what makes it "alkaline".
The main feature of a secondary cell is that it is rechargeable. This means that the redox reaction is very reversible. Products will readily convert back to reactants with great efficiency once one drives the electrolytic reaction back using the appropriate charging system (potential and current). Being rechargeable means that a battery can be used over and over during its lifetime. A good secondary battery will accept hundreds of recharges. Although the initial cost is more than a primary battery, a secondary battery will usually be more cost effective over the long run than a primary battery. Here is a list of many popular secondary batteries.
Lead-Acid Cells The lead-acid (Pb-acid) battery is a workhorse of a battery. It is also known as a car battery. It is robust, rechargeable, and of course - heavy (thank you lead). Lead has three different oxidation states in this battery. The two reactants are the extremes for lead which are +4 for the lead in PbO2 (the cathode), and zero (0) for metallic lead (Pb, the anode). BOTH the reduction and oxidation make the same product, PbSO4, which has lead in the +2 oxidation state. The half reactions and standard potentials are given below.
|cathode:||PbO2(s) + HSO4- + 3 H+ + 2 e- → PbSO4(s) + 2 H2O||E° = +1.685 V|
|anode:||Pb (s) + HSO4- → PbSO4(s) + H+ + 2 e-||E° = +0.356 V|
|overall:||Pb (s) + PbO2(s) + 2 HSO4- + 2 H+ → 2 PbSO4(s) + 2 H2O||E° = +2.041 V|
Shorthand cell notation for this is:
One thing unique about this voltaic cell is the fact that there is no salt bridge. This is because all the active components (those that change oxidations states) are solids and they do not migrate through the solution. The electrolyte is sulfuric acid where you have H+ cations and HSO4– anions. Each of the solids is coated onto the Pb(s) electrode. The Pb(s) on the right (the cathode) is really never depleted - only the PbO2 is depleted and replaced with a coating of PbSO4(s).
A fuel cell tends to utilize a fuel much like a combustion engine - natural gas, gasoline, kerosene, or better than all that, hydrogen. All of those fuels react readily with oxygen in combustion reactions. All combustion reactions are very spontaneous as well. The goal though in a fuel cell is to force the electrons to be transferred through the wiring and NOT have a runaway fire or explosion like in a combustion engine. This can be tricky but it is possible and has been around for decades. It is only recently though that the cost of these fuel cells is finally coming down to a reasonable cost such that a consumer (as opposed to an entire government) could purchase one - albeit still on the high end.
We will just focus on the H2/O2 fuel cell. The net overall reaction should be quite familiar:
2H2 + O2 → 2H2O
Pretty staight forward except that the H2 and O2 must be in two separate half-cells. Plus, the fact that they are both gases makes things a bit more complicated than it is with good ol' solid metals and gel-like electrolytes. Also, that net reaction is a bit deceiving in that there ARE other necessary reactants that just happen to cancel out in the net ionic equation. There are two ways to go with the H2/O2 fuel cell: acidic or basic. Let's look at each of them.
Hailed as the environmentally cleanest fuel cell, the H2/O2 fuel cell produces water as its product. NASA has used such cells since the early days of the space program and could even use the water that is produced to provide the needed water for the astronauts on the spaceships. Oxygen gas is a good oxidizing agent and is therefore used at the cathode - itself being reduced to hydroxide ion. Hydrogen acts as the reducing agent and is the primary reactant at the anode with itself being oxidized to water. Below is a table showing the two half reactions (in alkaline solution - generally a concentrated solution of KOH) and the overall net reaction and their potentials.
|cathode:||O2 + 2H2O + 4e- → 4OH–||E° = +0.40 V|
|anode:||2H2 + 4OH- → 4H2O + 4e-||E° = +0.83 V|
|overall (net):||2H2 + O2 → 2H2O||E° = +1.23 V|
modified from Wikipedia
The same overall reaction can be carried out in acidic conditions. The half reactions are slightly different, but the overall net reaction and the potential is the same.
|cathode:||O2 + 4H+ + 4e- → 2H2O||E° = +1.23 V|
|anode:||2H2 → 4H+ + 4e-||E° = 0.00 V|
|overall (net):||2H2 + O2 → 2H2O||E° = +1.23 V|
Do you want more info on fuel cells? Search the internet. There are lots of sources with good information - the amount of information can be overwhelming so expect that. Much of the information is a combination of marketing and the science, but you can see what is available and what is being used in the industry. Here are two links to a sites with a modern take on the H2/O2(air) fuel cell.
The Fuel Cell Explained from Pragma Industries.
Are cars being made with fuel cells? Yes. Check out Toyota's site for the Mirai: Toyota Mirai.