An Examination Of The Chemical Processes In Batteries Biology Essay

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Recent developments in electronics have revolutionized portable devices-you can now take radios, telephones, video recorders, and computers anywhere. These gadgets have created a need for portable electric power, most of which is supplied by batteries. Batteries store energy in chemical form and deliver it when needed as electricity. Batteries are remarkable devices in their own right. They have no visible moving parts yet they manage to push electric currents through circuits. Batteries continue to improve from year to year, with designers always trying to increase energy capacity, reliability, and reusability while decreasing size, weight, and cost. Nonetheless, batteries remain the limiting factor in such technologies as electric vehicles and portable computers.

History of the battery

The first battery was invented in 1800 by Alessando Volta. In1800 Alessando Volta of Italy built the voltaic plié and discovered the first practical method of generating electricity. Constructed of alternating discs of zinc and copper, with pieces of cardboard socked in brine between the metals, the voltic pile produced electrical current. The metallic conducting arc was used to carry the electricity over a greater distance. Alessando Volta's voltaic plié was the first battery that produces a reliable, steady current of electricity


The next major advance came in 1866 when Georges Leclanché developed a much improved battery. Leclanché assembled his cell in a porous pot. The cathode consisted of crushed manganese dioxide with a little carbon added. The anode was a zinc rod. The cathode was packed into the pot with a carbon rod inserted as a current collector. The anode and the pot were then immersed in an ammonium chloride solution, which acted as the electrolyte. The liquid seeped through the porous cup and made contact with the cathode material. Even though it was a heavy wet cell prone to breaking, Leclanch's invention represented an advance over previous batteries and it became an immediate success, gaining wide use in telegraph systems within two years of its development

Further improvements came in the 1880's when Carl Gassner, a German scientist, invented the first dry cell. Gassner used zinc as the container to house the cell's other components; at the same time, he used the sealed zinc container as the anode. The cathode surrounded a carbon rod. Gassner also added zinc chloride to the electrolyte, which markedly reduced corrosion of the zinc when the cell was idling, adding considerably to its shelf life.

Classification of battery

Primary cell

Secondary cell

Primary batteries


Secondary batteries


Type of battery

Lead -acid battery

Wet and dry cells

Carbon-zinc cell

Alkaline cell

Nickel-cadmium cell

Mercury cell

Fuel cell

Lead- acid batter

A lead -acid battery is an electricity storage device that uses a reversible chemical reaction to store energy. It uses a combination of lead plates or girds and an electrolyte consisting.


In the charged state, each cell contains electrodes of elemental lead (Pb) and lead(IV) oxide (PbO2) in an electrolyte of approximately 33.5% v/v (4.2 Molar) sulphuric acid (H2SO4).

In the discharged state both electrodes turn into lead (II) sulphate (PbSO4) and the electrolyte loses its dissolved sulphuric acid and becomes primarily water. Due to the freezing-point depression of water, as the battery discharges and the concentration of sulphuric acid decreases, the electrolyte is more likely to freeze during winter weather.

The chemical reactions are (discharged to charged ):

Anode (oxidation):

Cathode (reduction):

Because of the open cells with liquid electrolyte in most lead-acid batteries, overcharging with high charging voltages generates oxygen and hydrogen gas by electrolysis of water, forming an explosive mix. The acid electrolyte is also corrosive.

Practical cells are usually not made with pure lead but have small amounts of antimony, tin, calcium or selenium alloyed in the plate material to add strength and simplify manufacture

We and dry cells

Wet cell



Dry cell




Carbon-zinc cell

A zinc-carbon dry cell or battery is packaged in a zinc can that serves as both a container and negative terminal. It was developed from the wet Leclanché cell. The positive terminal is a carbon rod surrounded by a mixture of manganese dioxide and carbon powder. The electrolyte used is a paste of zinc chloride and ammonium chloride dissolved in water. Zinc chloride cells are an improved version from the original ammonium chloride variety. Zinc-carbon batteries are the least expensive primary batteries and thus a popular choice by manufacturers when devices are sold with batteries included. They are commonly labeled as "General Purpose" batteries. They can be used in remote controls, flashlights, clocks, or transistor radios, since the power drain is not too heavy.

Electro chemistry

In a zinc-carbon dry cell, the outer zinc container is the negative terminal. The zinc is oxidised according to the following half-equation

Zn(s) â†’ Zn2+ (aq) + 2 e-

A graphite rod surrounded by a powder containing manganese (IV) oxide is the positive terminal. The manganese dioxide is mixed with carbon powder to increase the electrical conductivity. The reaction is as follows:

2MnO2(s) + H2 (g) â†’ Mn2O3(s) + H2O (l)

The H2 comes from the NH4+ (aq):

2NH4+ (aq) + 2 e- â†’ H2 (g) + 2NH3 (aq)

And the NH3 combines with the Zn2+.

In this half-reaction, the manganese is reduced from an oxidation state of (+4) to (+3).

There are other possible side-reactions, but the overall reaction in a zinc-carbon cell can be represented as:

Zn(s) + 2MnO2(s) + 2NH4+ (aq) â†’ Mn2O3(s) + Zn (NH3)22+ (aq) + H2O (l)

Alkaline battery

Alkaline batteries and alkaline cells are a type of disposable battery or rechargeable battery dependent upon the reaction between zinc and manganese dioxide (Zn/MnO2).

Compared with zinc-carbon batteries of the Leclanché or zinc chloride types, while all produce approximately 1.5 volts per cell, alkaline batteries have a higher energy density and longer shelf-life. Compared with silver-oxide batteries, which alkalines commonly compete against in button cells, they have lower energy density and shorter lifetimes but lower cost.

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 which are offered in the same nominal voltages and physical size. Other battery systems also use alkaline electrolytes, but they use different active materials for the electrodes.

The alkaline battery was invented by Canadian engineer Lewis Urry in the 1950s while working for the Eveready Battery Company

Electro chemistry

In an alkaline battery, the anode (negative terminal) is made of zinc powder (which allows more surface area for increased rate of reaction therefore increased electron flow) and the cathode (positive terminal) is composed of manganese dioxide. Alkaline batteries are comparable to zinc-carbon batteries, but the difference is that alkaline batteries use potassium hydroxide (KOH) as an electrolyte rather than ammonium chloride or zinc chloride

The half-reactions are:

Zn (s) + 2OHâˆ’ (aq) â†’ ZnO (s) + H2O (l) + 2eâˆ’

2MnO2 (s) + H2O (l) + 2eâˆ’ â†’Mn2O3 (s) + 2OHâˆ’ (aq)


Nickel-cadmium battery



There are two types of NiCd batteries: sealed and vented. This article mainly deals with sealed cells


The chemical reactions in a NiCd battery during discharge are:

At the cadmium electrode, and

At the nickel electrode. The net reaction during discharge is

Mercury battery

A mercury battery (also called mercuric oxide battery or mercury cell) is a non-rechargeable electrochemical battery, a primary cell. Due to the content of mercury, and the resulting environmental concerns, the sale of mercury batteries is banned in many countries. Mercury batteries were made in button types for watches, hearing aids, and calculators, and in larger forms for other applications.

Overall reaction for Mercury Batter

The diagram and the overall reaction for a mercury battery is shown above. The anode is a zinc inner case (like the dry cell), surrounded by stainless steel. The anode is then put in contact with an electrolyte containing zinc oxide and mercury (II) oxide. One of its main differences with the dry cell is that it provides a more constant voltage and a considerably longer life. This makes the mercury battery ideal for medicinal and electronic industries

Fuel Cells

In principle, a fuel cell operates like a battery. Unlike a battery, a fuel cell does not run down or require recharging. It will produce energy in the form of electricity and heat as long as fuel is supplied.


Hydrogen fuel is fed into the "anode" of the fuel cell. Oxygen (or air) enters the fuel cell through the cathode.

Encouraged by a catalyst, the hydrogen atom splits into a proton and an electron, which take different paths to the cathode. .The proton passes through the electrolyte. The electrons create a separate current that can be utilized before they return to the cathode, to be reunited with the hydrogen and oxygen in a molecule of water.

A fuel cell system which includes a "fuel reformer" can utilize the hydrogen from any hydrocarbon fuel - from natural gas to methanol, and even gasoline. Since the fuel cell relies on chemistry and not combustion, emissions from this type of a system would still be much smaller than emissions from the cleanest fuel combustion processes.

Fuel cells use combustion reactions (which are a type of redox reactions) to generate electricity. The overall reaction for one type of fuel cell is shown below:

How a Battery Produces Electricity

A battery uses chemical potential energy to pump electrons from its positive terminal to its negative terminal. Since electrostatic forces push the electrons the other direction, the battery must do work on the electrons as it moves them. Each time the battery transfers an electron, it uses up a small portion of its chemical potential energy. After transferring a certain number of electrons, the battery runs out of chemical potential energy and must be recharged or discarded.

But a battery sitting on the shelf stops transferring electrons long before it runs out of chemical potential energy. With each transfer, the negative terminal becomes more negatively charged and the positive terminal becomes more positively charged. The amount of separated charge on the terminals increases and so does the voltage rise across the battery-the battery must do more and more work to transfer each additional electron. Eventually, the electrostatic forces become so strong that the battery can't transfer any more electrons. The battery just sits on the shelf with negative charge on its negative terminal and positive charge on its positive terminal and it remains that way almost indefinitely .However, when you install the battery in a flashlight and turn the flashlight on electric circuit connects the two terminals to one another (Fig. 17.3.1)

Electrons flow from the negative terminal, through the light bulb, to the positive terminal and the amount of separated charge on the battery's terminals decreases.

The battery begins to pump electrons again. The battery pumps electrons onto the negative terminal and the flashlight returns those electrons to the positive terminal. The electrons flow around and around this circuit, receiving energy from the battery and delivering that energy to the light bulb, until the battery's chemical potential energy is exhausted or you turn the flashlight off.

But how does a battery use its chemical potential energy to pump electrons from its positive terminal to its negative terminal? Many batteries are based on electron transfers from atoms of one element to those of another. These different atoms have different affinities for their outermost or valence electrons and many of the transfers result in releases of energy. When an atom that binds its valence electrons relatively strongly is missing some of them, it may extract electrons from another atom that binds them relatively weakly. Overall, the electrons move from one atom to the other and some potential energy is released. This process is the principal source of a battery's energy.

Choosing which atoms to use in a battery is by examining their properties.

These properties depend in an orderly fashion on the numbers of protons and electrons the atoms have. One way to see this order is to arrange the atoms in a periodic table (Fig. 17.3.2). In this table, the atoms are arranged in horizontal rows according to their atomic numbers-the numbers of protons they contain.

Since atoms are normally electrically neutral, their atomic numbers also indicate the numbers of electrons they contain. The atom with atomic number 1 is hydrogen

(H), with atomic number 2 is helium (He), and so on.

The peculiar structure of the table comes from the way in which electrons fill the atomic orbitals surrounding the nucleus of each atom. Because of the Pauli Exclusion Principle, all electrons of a particular spin-either spin-up or spin down- must be in different orbitals. The electrons fill the orbitals from the lowest energy orbitals on up until the atom has the right number of electrons. The electrons in the last few orbitals filled determine most of the chemical properties of the atom, particularly the atom's behaviour in a battery. This filling process is quite complicated, but there are a few simple observations we can make.

Some atoms have just enough electrons to completely fill a major electronic shell. These atoms are extremely stable, unwilling to give up any electrons and uninterested in any additional electrons. These atoms are the noble gases, helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe), and radon (Rn), found in the rightmost vertical column of Fig. 17.3.2.

Some atoms have just one or two electrons more than are required to fill a major electronic shell and are relatively willing to give those electrons up. These atoms are the alkali metals, lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr), and the alkaline earths, beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra), found in the leftmost and second to leftmost vertical columns of Fig. 17.3.2.

Still other atoms have almost enough electrons to complete a major electronic shell and are relatively aggressive at attracting more. These atoms are found just to the left of the noble gases on the right side of Fig. 17.3.2. They include nitrogen (N), oxygen (O), sulphur (S), and the halogens, fluorine (F), chlorine (Cl), bromine (B), iodine (I), and astatine (At).

The remaining atoms fall in between. While their major electronic shells are only partly complete, they tend to exchange electrons in order to complete minor electronic shells. The atoms in the long horizontal stretch between scandium (Sc) and zinc (Zn) are called the transition metals and differ from one another by how much of one minor shell they have completed. The atoms shown at the bottom of Fig 17.3.2 are called the rare earths and differ by how much of another minor shell they have completed.

Many of these atoms are important for batteries. The main issue for batteries is just how strongly the atoms attract electrons. This tendency to attract electrons is called electronegativity and is measured in various ways. One scheme developed by American chemist Linus Pauling (1901-1994) is called Pauling electronegativity. The more strongly an atom attracts electrons, the higher its Pauling electronegativity. Values range from 0.7 for cesium (Cs) atoms, which easily give up electrons, to 4.0 for fluorine (F) atoms, which attracts electrons aggressively. Pauling electronegativities for the other atoms Batteries generally work by transferring electrons from atoms with low Pauling electronegativities to ones with high Pauling electronegativities.