Compare the specific characteristics of alkaline manganese-zinc, salt and primary lithium sells.

Lithium sells (IEC code: L) are a type of primary battery dependent upon the reaction between zinc and manganese(IV) oxide (Zn/MnO₂). A rechargeable alkaline battery allows reuse of specially designed cells. The alkaline battery gets its name because it has an alkaline electrolyte of potassium hydroxide, instead of the acidic ammonium chloride or zinc chloride electrolyte of the zinc-carbon batteries. Chemistry. In an alkaline battery, the negative electrode is zinc and the positive electrode manganese (IV) oxide. The alkaline electrolyte of potassium hydroxide is not part of the reaction, only the zinc and manganese (IV) oxide are consumed during discharge. The alkaline electrolyte of potassium hydroxide remains, as there are equal amounts of OH consumed and produced. The half-reactions are: Zn_(s) + 2OH_(aq) ZnO_(s) + H₂O_(l) + 2e [e° = -1.28 V]; 2MnO_2(s) + H₂O_(l) + 2e Mn₂O_3(s) + 2OH_(aq) [e° = +0.15 V]. Overall reaction: Zn_(s) + 2MnO_2(s) ZnO_(s) + Mn₂O_3(s) [e° = +1.43 V]. Capacity. Several sizes of button and coin cells. Some are alkaline and others are silver oxide. Two 9v batteries were added as a size comparison. Enlarge to see the size code markings. Capacity of an alkaline battery is greater than an equal size Leclanché cell or zinc-chloride cell because the manganese (IV) oxide is purer and denser, and space taken up by internal components such as electrodes is less. An alkaline cell can provide between three and five times capacity. The capacity of an alkaline battery is strongly dependent on the load. An AA-sized alkaline battery might have an effective capacity of 3000 mAh at low drain, but at a load of 1 ampere, which is common for digital cameras, the capacity could be as little as 700 mAh. The voltage of the battery declines steadily during use, so the total usable capacity depends on the cut-off voltage of the application. Unlike Leclanche cells, the alkaline cell delivers about as much capacity on intermittent or continuous light loads. On a heavy load, capacity is reduced on continuous discharge compared with intermittent discharge, but the reduction is less than for Leclanche cells. A zinc–carbon battery is a dry cell battery that delivers a potential of 1.5 volts between a zinc metal electrode and a carbon rod from an electrochemical reaction between zinc and manganese dioxide mediated by a suitable electrolyte. It is usually conveniently packaged in a zinc can which also serves as the anode with a negative potential, while the inert carbon rod is the positive cathode. General purpose batteries may use an aqueous paste of ammonium chloride as electrolyte, possibly mixed with some zinc chloride solution. Heavy duty types use a paste primarily composed of zinc chloride. Zinc–carbon batteries were the first commercial dry batteries, developed from the technology of the wet Leclanché cell. They made flashlights and other portable devices possible, because the battery can function in any orientation. They are still useful in low drain or intermittent use devices such as remote controls, flashlights, clocks or transistor radios. Zinc–carbon dry cells are single-use primary cells. Chemical reactions.In a zinc–carbon dry cell, the outer zinc container is the negatively charged terminal. The zinc is oxidised according to the following half reactions: Anode (marked -) Zn(s) Zn²⁺(aq) + 2 e [E° = 0.7626 V]; Cathode (marked +) 2 MnO₂(s) + 2 e + 2 NH₄Cl(aq) Mn₂O₃(s) + 2 NH₃(aq) + H₂O(l) + 2 Cl [E° +0.5 V]. Other side-reactions are possible, but the overall reaction in a zinc–carbon cell can be represented as: Zn(s) + 2 MnO₂(s) + 2 NH₄Cl(aq) Mn₂O₃(s) + Zn(NH₃)₂Cl₂ (aq) + H₂O(l). If zinc chloride is substituted for ammonium chloride as the primary electrolyte, the anode reaction remains he same but the cathode reaction is: MnO₂(s) + H₂O(l) + e MnO(OH)(s) + OH(aq) and the overall reaction: 4 Zn(s) + 8 MnO₂(s) + ZnCl₂(aq) + 9 H₂O(l) 8 MnO(OH)(s) + Zn(OH)Cl(aq) + 5 H₂O +4ZnO.

Zinc and alkaline are two types of batteries, although their names is sort of a misnomer since even alkaline batteries have zinc in their composition. What separated alkaline batteries is the type of elec-trolyte it used. Most Zinc batteries used an acidic electrolyte com-posed of ammonium chloride while alkaline batteries use potassium hydroxide, which is a basic electrolyte.The difference between these two types of batteries that would directly impact the ordinary user is battery capacity. Although zinc chloride batteries have a higher capacity compared to the older zinc carbon battery, the capacity of alkaline batteries is many folds over the capacity of both types of zinc batteries. This instantly translates to longer time periods before needing to replace the batteries on your device.

A lithium-ion battery or Li-ion battery is a type of rechargeable battery in which lithium ions move from the negative electrode to the positive electrode during discharge and back when charging. Li-ion batteries use an intercalated lithium compound as one electrode material, compared to the metallic lithium used in a non-rechargeable lithium battery. The electrolyte, which allows for ionic movement, and the two electrodes are the constituent components of a lithium-ion battery cell. Electrochemistry.The participants in the electrochemical reactions in a lithium-ion battery are the negative and positive electrodes with the electrolyte providing a conductive medium for lithium ions to move between the electrodes. The cathode (marked +) half-reaction is: Li_1-xCO₂ + xLi⁺ + xe^– Li_1-хCoO₂ +xLi⁺ +xe^– LiCoO₂Li_1-хCoO₂ +xLi⁺ +xe^– LiCoO₂ LiCO2. The anode (marked -) half reaction is: xLiC₆ xLi⁺ + xe^– Li_1-хCoO₂ +xLi⁺ +xe^– LiCoO₂Li_1-хCoO₂ +xLi⁺ +xe^– LiCoO₂ xC₆.xLiC₆ xLi⁺+xe^– +xC₆ The overall reaction has its limits. Overdischarge supersaturates lithi-um cobalt oxide, leading to the production of lithium oxide, possibly by the following irreversible reaction: Li⁺ + xe^– + LiCO₂ Li⁺ + e^– + LiCoO₂ = Li₂O + CO.Li₂O + CoO Overcharge up to 5.2 volts leads to the synthesis of cobalt(IV) oxide, as evidenced by x-ray diffraction: LiCO₂ Li⁺ + e^– + LiCoO₂ = Li⁺ + CO₂ + xe^-. LiCoO₂ = Li⁺ + CoO₂ + e^– The cell's energy is equal to the voltage times the charge. Each gram of lithium represents Faraday's constant/6.941 or 13,901 coulombs. At 3 V, this gives 41.7 kJ per gram of lithium, or 11.6 kWh per kg. This is a bit more than the heat of combustion of gasoline, but does not consider the other materials that go into a lithium battery and that make lithium batteries many times heavier per unit of energy. Elec-trolytes. The cell voltages given in the Electrochemistry section are larger than the potential at which aqueous solutions will electrolyze. Liquid electrolytes in lithium-ion batteries consist of lithium salts, such as LiPF₆, LiBF₆, or LiClO₄ in an organic solvent, such as ethy-lene carbonate, dimethyl carbonate, and diethyl carbonate. A liquid electrolyte acts as a conductive pathway for the movement of cations passing from the negative to the positive electrodes during discharge. Typical conductivities of liquid electrolyte at room temperature (20 °C (68 °F)) are in the range of 10 mS/cm, increasing by approximately 30–40% at 40 °C (104 °F) and decreasing slightly at 0 °C (32 °F). The combination of linear and cyclic carbonates (e.g., ethylene carbonate (EC) and dimethyl carbonate (DMC)) offers high conductivity and SEI-forming ability. A mixture of a high ionic conductivity and low viscosity carbonate solvents is needed, because the two properties are mutually exclusive in a single material. Organic solvents easily de-compose on the negative electrodes during charge. When appropriate organic solvents are used as the electrolyte, the solvent decomposes on initial charging and forms a solid layer called the solid electrolyte interphase (SEI), which is electrically insulating yet provides signi-ficant ionic conductivity. The interphase prevents further decompo-sition of the electrolyte after the second charge.