You’ve gotta love chemistry– it’s everywhere! It’s what makes us go, and what makes things around us go. Every chemical reaction either requires energy or produces energy– our entire civilization depends on these concepts. For those of us in the battery industry, metrics such as voltage, energy, and power are common indicators of performance. This 2 part blog series will examine electrochemical reactions in terms of battery chemistries.
Let’s start at the beginning: in 1800, Alessando Volta invented the first battery, also called a voltaic pile: alternating sheets of Zn and Cu separated by cardboard and stacked in H2SO4 or brine produced a steady electric current with 0.76 V. This voltage originates from Zn oxidation, but the anodic reaction is the formation of H2, which can have obvious consequences of fire or explosion. If the pile is contained in a closed vessel, H2 bubbles build up on the plates and block contact with the electrolyte, and the battery output declines rapidly. This was proof of concept but not practical: let’s move on.
Oxidation/reduction, ionic diffusion, crystal structures, conductivity– all topics in chemistry texts that impact battery operation. We’ll start with the most fundamental property of battery electrochemistry, oxidation/reduction reactions, the driving force behind energy storage. This is the origin of cell voltage, and remember, voltage x capacity (mAh/g) = energy (mWh/g). Whether primary (non-rechargeable) or secondary (rechargeable), redox reactions are common to all cell types and chemistries.
In a previous blog, The Role of Battery Materials and Electrode Fabrication in Cell Performance, we presented various anodic and cathodic reactions in common batteries.
Primary Battery Chemistry
Let’s dig a little deeper into the chemistry, starting with the ubiquitous alkaline battery, AKA manganese-zinc. This technology features a compressed MnO2 cathode, using graphite as a binder, and a powdered Zn/9M KOH slurry anode. The half reactions and potentials are:
2MnIVO2 + H2O + 2e– ⇒ MnIII2O3 + 2OH– +1.28V
Zn + 2OH– ⇒ ZnO + H2O + 2e– +0.15V
Simple algebraic addition yields the overall cell reaction and voltage:
Zn + 2MnO2 ⇒ ZnO + Mn2O3 +1.43 V
Although the graphite and KOH solution offer excellent conductivity, the disordered MnO2 is a mix of two crystal phases, and contains circa 2% interstitial water, which slows the reduction reaction above. This is why alkaline cells are good at providing steady energy (for example, 0.5 hours or more at constant rate) but less satisfactory for high current pulse operation, since the Zn+2 ion diffuses relatively slowly into the cathode. Nominal cell voltage is 1.5V, working voltage under load is 1.1-1.3V, and end-of-life is about 0.8V; fast discharge increases impedance and pushes the voltage curve down.
The precursor to alkaline chemistry was the Leclanché cell, which employed a saturated NH4Cl solution electrolyte rather than KOH and produced nearly 1.4V. The redox reaction is:
Zn + 2MnO2 + 2NH4Cl ⇒ ZnCl2 + Mn2O3 + H2O + 2NH3
This technology was a favorite in the telegraph and telephone industries of the late 19th century, before the electric grid was widely available. Unfortunately, the depletion of NH4Cl rapidly amplified internal resistance, which lowered voltage and limited the cell’s working life.
For consumers seeking twice the energy and willing to spend more money, there are lithium primary batteries. These cells are available in several flavors, not all suitable for the general public. The driving force in such cells is the 3.0 V half-reaction of Li ionization, but the downside is the presence of hazardous Li metal. Most common are Li-MnO2 3.3 V batteries, chemically related to Mn-Zn but with a Li metal anode and electrolyte of LiClO4 in an organic solvent mix. The full reaction yields LiMnO2, not Mn2O3, which reduces the anticipated (ca 4.3 V) electrochemical output. Li+ diffusion is very rapid, explaining why this system is excellent for pulse operations (such as flash cameras, for those of you old enough to remember them). Despite high self-discharge above room temperature, Li-MnO2 technology has more than 75% of the Li primary battery market.
Other Li primaries include Li-CFx, Li-FeS2, and Li-SOCl2, and a multitude of less-common or specialty Li chemistries. The redox reactions are:
Li + CFx ⇒ LiF + C 3.1 V (A)
Li + FeS2 ⇒ LiFeS2 1.8 V (B)
4Li + 2SOCl2 ⇒ 4LiCl + S + SO2 3.65 V (C)
Let’s examine each in turn and explore the pluses and minuses.
Cell A features an inert cathode prepared by subjecting graphite to fluorine gas under high pressure. While thermodynamically the reaction with Li is very energetic, Li diffusion into the cathode is slow (that is, kinetically controlled), and these high-reliability cells are best for low-current backup in small devices for military and aerospace operations. Unlike Li-MnO2, Li-CFx cells are stable to 800C and have very long shelf lives, with excellent energy density – 1000 Wh/L.
Cell B features an older chemistry, but still found in some AA and AAA batteries. Cathodes are formed much the same as in Mn-Zn cells: a compressed mix of FeS2 and graphite. Exhibiting a nominal 1.5V, such cells are drop-in replacements for alkaline batteries, and exhibit better high rate (power) capability and very low self-discharge. Although FeS2 is inexpensive, Li is not, which is why this niche technology will not replace alkaline cells.
Finally, cell C contains a high voltage, high energy (3.5 V, 1200 Wh/L, twice that of alkaline cells) system that is not for consumer use, due to the toxicity of the liquid cathode, SOCl2, and the likelihood of explosion if overcharged. On the plus side, the SOCl2/LiAlCl4 electrolyte remains liquid to nearly -600C, enabling cell operation in extreme cold. Very expensive and with a reasonably long shelf life, these batteries find use in portable devices and memory backup for the military. Totaling all the primary battery chemistries, they make up