Sodium-ion battery

Battery-grade salts of sodium are cheap and abundant, much more so than those of lithium. This makes them a cost-effective alternative especially for applications where weight and energy density are of minor importance. These cells can be completely drained (to zero charge) without damaging the active materials. They can be stored and shipped safely. Moreover, sodium-ion batteries have excellent electrochemical features in terms of charge-discharge, reversibility, coulombic efficiency and high specific discharge capacity.

In November of 2017 French Network on Electrochemical Energy Storage (RS2E) announced the intention to produce a 18650 format battery by 2020. The battery will be 3.5V, 90Wh/Kg, perform more than 2,000 charge and discharge cycles without significant loss of performance, and life expectancy of more than 10 years in continuous use.

SIBs store energy in chemical bonds of the anode. Charging the battery forces Na⁺ ions to de-intercalate from the cathode and migrate towards the anode. Charge balancing electrons pass from the cathode through the external circuit containing the charger and into the anode. During discharge the process reverses. Once a circuit is completed electrons pass back from the anode to the cathode and the Na⁺ ions travel back to the cathode.

Anode

Aquion originally used a mix of activated carbon and titanium phosphate NaTi₂(PO₄)₃ that relied mostly on pseudocapacitance to store charge, resulting in a low energy density and a tilted voltage-charge slope. In many ways, titanium phosphate is similar to iron phosphate used in some other batteries, but with a low (anodic) electrode potential.The initial electrolyte was an aqueous sodium sulphate solution. Later a more soluble <5M NaClO₄ was used

Cellulose

In one study, tin-coated wood anodes replaced stiff anode bases. The wood fibers proved withstood more than 400 charging cycles. After hundreds of cycles, the wood ended up wrinkled but intact. Computer models indicated that the wrinkles effectively reduce stress during charging and recharging. Na ions move via the fibrous cell walls and diffuse at the tin film surface.
Another study used MoS₂/graphene composite paper as an electrode, yielding 230 Ah/kg with Coulombic efficiency reaching approximately 99%

Cathode

Tests of Na₂FePO₄F and Li₂FePO₄F cathode materials indicated that a sodium iron phosphate cathode can replace a lithium iron phosphate cathode in a Li cell. The lithium-ion and sodium-ion combination would lower manufacturing costs.
P2-Na_2/3[Fe_1/2Mn_1/2]O₂ delivered 190  Ah/kg of reversible capacity in sodium cells using electrochemically active Fe³⁺/Fe⁴⁺ redox at room temperature. Triclinic Na₂FeP₂O₇ was examined as rechargeable sodium ion batteries by a glass-ceramics method. The precursor glass, also made of Na₂FeP₂O₇, was prepared by melt-quenching. Na₂FeP₂O₇ and exhibited 2.9 V, 88 Ah/kg.
Separately, chromium cathodes employed the reaction:

NaF + (1−x)VPO₄ + xCrPO₄ → NaV_1−xCr_xPO₄F

The effects of Cr doping on cathode performance materials was analyzed in terms of crystal structure, charge/discharge curves and cycle performance and indicated that the Cr-doped materials expressed better cycle stability. The initial reversible capacity was 83.3 Ah/kg and the first charge/discharge efficiency was about 90.3%. The reversible capacity retention of the material was 91.4% after the 20th cycle.

Michael Faraday

Michael Faraday, who came from a very poor family, became one of the greatest scientists in history. His achievement was remarkable in a time when science was usually the preserve of people born into wealthy families. The unit of electrical capacitance is named the farad in his honor, with the symbol F.

Michael Faraday was born on September 22, 1791 in London, England, UK. He was the third child of James and Margaret Faraday. His father was a blacksmith who suffered poor health. Before marriage, his mother had been a servant. The family lived in a degree of poverty.

Michael Faraday attended a local school until he was 13, where he received a basic education. To earn money for the family he started working as a delivery boy for a bookshop. He worked hard and impressed his employer. After a year, he was promoted to become an apprentice bookbinder.

Sir Humphry Davy was one of the most famous scientists in the world. Faraday jumped at the chance and attended four lectures about one of the newest problems in chemistry – defining acidity. He watched Davy perform experiments at the lectures.

This was the world he wanted to live in, he told himself. He took notes and then made so many additions to the notes that he produced a 300 page handwritten book, which he bound and sent to Davy as a tribute.

And then there was a fortunate (for Faraday) accident. Sir Humphry Davy was hurt in an explosion when an experiment went wrong: this temporarily affected his ability to write. Faraday managed to get work for a few days taking notes for Davy, who had been impressed by the book Faraday had sent him. There were some advantages to being a bookbinder after all!

When his short time as Davy’s note-taker ended, Faraday sent a note to Davy, asking if he might be employed as his assistant. Soon after this, one of Davy’s laboratory assistants was fired for misconduct, and Davy sent a message to Faraday asking him if he would like the job of chemical assistant.

Faraday began work at the Royal Institution of Great Britain at the age of 21 on March 1, 1813.

After just seven months at the Royal Institution, Davy took Faraday as his secretary on a tour of Europe that lasted 18 months.

During this time Faraday met great scientists such as André-Marie Ampère in Paris and Alessandro Voltain Milan. In some ways, the tour acted like a university education, and Faraday learned a lot from it.

In 1816, aged 24, Faraday gave his first ever lecture, on the properties of matter, to the City Philosophical Society. And he published his first ever academic paper, discussing his analysis of calcium hydroxide, in the Quarterly Journal of Science.

In 1821, aged 29, he was promoted to be Superintendent of House and Laboratory of the Royal Institution. He also married Sarah Barnard. He and his bride lived in rooms in the Royal Institution for most of the next 46 years: no longer in attic rooms, they lived in a comfortable suite Humphry Davy himself had once lived in.

In 1824, aged 32, he was elected to the Royal Society. This was recognition that he had become a notable scientist in his own right.

In 1825, aged 33, he became Director of the Royal Institution’s Laboratory.

In 1833, aged 41, he became Fullerian Professor of Chemistry at the Royal Institution of Great Britain. He held this position for the rest of his life.

In 1848, aged 54, and again in 1858 he was offered the Presidency of the Royal Society, but he turned it down

Michael Faraday, (born September 22, 1791, Newington, Surrey, England—died August 25, 1867, Hampton Court, Surrey), English physicist and chemist whose many experiments contributed greatly to the understanding of electromagnetism.

Faraday, who became one of the greatest scientists of the 19th century, began his career as a chemist. He wrote a manual of practical chemistry that reveals his mastery of the technical aspects of his art, discovered a number of new organic compounds, among them benzene, and was the first to liquefy a “permanent” gas (i.e., one that was believed to be incapable of liquefaction). His major contribution, however, was in the field of electricity and magnetism. He was the first to produce an electric current from a magnetic field, invented the first electric motor and dynamo, demonstrated the relation between electricity and chemical bonding, discovered the effect of magnetism on light, and discovered and named diamagnetism, the peculiar behaviour of certain substances in strong magnetic fields. He provided the experimental, and a good deal of the theoretical, foundation upon which James Clerk Maxwellerected classical electromagnetic field theory.

Electrolysis

In 1800, Sir Humphry Davy, inventor of the miner’s safety lamp, began testing the chemical effects of electricity and found out that decomposition occurred when passing electrical current through substances. This process was later called electrolysis.

He made new discoveries by installing the world’s largest and most powerful electric battery in the vaults of the Royal Institution of London, connecting the battery to charcoal electrodes produced the first electric light.

Voltaic cells use a spontaneous chemical reaction to drive an electric current through an external circuit. These cells are important because they are the basis for the batteries that fuel modern society. But they aren't the only kind of electrochemical cell. It is also possible to construct a cell that does work on a chemical system by driving an electric current through the system. These cells are called electrolytic cells. Electrolysis is used to drive an oxidation-reduction reaction in a direction in which it does not occur spontaneously. Voltaic cells use the energy given off in a spontaneous reaction to do electrical work. Electrolytic cells use electrical work as source of energy to drive the reaction in the opposite direction.

A Galvanic cell converts chemical energy into electrical energy.

Here, the redox reaction is spontaneous and is responsible for the production of electrical energy.

The two half-cells are set up in different containers, being connected through the salt bridge or porous partition.

Here the anode is negative and cathode is the positive electrode. The reaction at the anode is oxidation and that at the cathode is reduction.

An electrolytic cell converts electrical energy into chemical energy.

The redox reaction is not spontaneous and electrical energy has to be supplied to initiate the reaction.

Both the electrodes are placed in a same container in the solution of molten electrolyte.

Here, the anode is positive and cathode is the negative electrode. The reaction at the anode is oxidation and that at the cathode is reduction.

The external battery supplies the electrons. They enter through the cathode and come out through the anode.

Michael Faraday defined the cathode of a cell as the electrode to which cations (positively charged ions, like silver ions Ag⁺  ) flow within the cell, to be reduced by reacting with electrons (negatively charged) from that electrode.

Likewise he defined the anode as the electrode to which anions (negatively charged ions, like chloride ions Cl⁻) flow within the cell, to be oxidized by depositing electrons on the electrode.

An electrolytic cell converts electrical energy into chemical energy and not the other way round.

The ensuing redox reaction in the process is not a spontaneous one and for the reaction to start, electric energy has to be introduced in the apparatus.

 

 

Galvanic cell

The Daniell cell, invented in 1836 by British chemist John Frederic Daniell, was the first practical source of electricity, becoming an industry standard and seeing widespread adoption as a power source for electrical telegraph networks. It consisted of a copper pot filled with a copper sulfate solution, in which was immersed an unglazed earthenware container filled with sulfuric acid and a zinc electrode.

In 1866, Georges Leclanché invented a battery that consisted of a zinc anode and a manganese dioxide cathode wrapped in a porous material, dipped in a jar of ammonium chloride solution. The manganese dioxide cathode had a little carbon mixed into it as well, which improved conductivity and absorption. It provided a voltage of 1.4 volts. This cell achieved very quick success in telegraphy, signalling and electric bell work.

Dry Cell

·         The container of the dry cell is made of zinc which also serves as one of the electrodes.

·         The other electrode is a carbon rod in the centre of the cell.

·         The zinc container is lined with a porous paper.

·         A moist mixture of ammonium chloride, man­ganese dioxide, zinc chloride and a porous inert filler occupy the space between the paper lined zinc container and the carbon rod.

·         The cell is sealed with a material like wax.

·         As the cell operates, the zinc is oxidised to Zn²⁺

·         Anode reaction: Zn → Zn²⁺ + 2e^-   

·         The electrons are utilized at carbon rod (cathode) as the ammonium ions are reduced.

·         Cathode Reaction: 2NH₄⁺+2e →2NH₃ + H₂ 

·         The cell reaction is:  Zn+ 2 NH₄⁺ → Zn²⁺ + 2NH₃ + H₂

·         Hydrogen is oxidized by MnO₂ in the cell:  2MnO₂ + H₂ →2MnO(OH)

·         Ammonia produced at cathode combines with zinc ions to form complex ion.  Zn²⁺ + 4NH₃ →[Zn(NH₃)₄]₂₊ 

·         E_cell is 1.6 volt

The battery History

1800 Voltaic Pile—Alessandro Volta invented the Voltaic Pile and discovered the first practical method of generating electricity. Constructed of alternating discs of zinc and copper with pieces of cardboard soaked in brine between the metals, the Voltaic Pile produced electrical current. The metallic conducting arc was used to carry the electricity over a greater distance. Alessandro Volta's voltaic pile was the first "wet cell battery" that produced a reliable, steady current of electricity.

1836 Daniell Cell—The Voltaic Pile could not deliver an electrical current for a long period of time. Englishman, John F. Daniell invented the Daniell Cell that used two electrolytes: copper sulfate and zinc sulfate. The Daniel Cell lasted longer than the Volta cell or pile. This battery, which produced about 1.1 volts, was used to power objects such as telegraphs, telephones, and doorbells, remained popular in homes for over 100 years

1859 Rechargeable—French inventor, Gaston Plante developed the first practical storage lead-acid battery that could be recharged (secondary battery). This type of battery is primarily used in cars today.

1881—J.A. Thiebaut patented the first battery with both the negative electrode and porous pot placed in a zinc cup.

1881—Carl Gassner invented the first commercially successful dry cell battery (zinc-carbon cell).

1899—Waldmar Jungner invented the first nickel-cadmium rechargeable battery.

1901 Alkaline Storage—Thomas Alva Edison invented the alkaline storage battery. Thomas Edison's alkaline cell had iron as the anode material (-) and nickelic oxide as the cathode material (+)

1949 Alkaline-Manganese Battery—Lew Urry developed the small alkaline battery in 1949. The inventor was working for the Eveready Battery Co. at their research laboratory in Parma, Ohio. Alkaline batteries last five to eight times as long as zinc-carbon cells, their predecessors

Most research activities today revolve around improving lithium-based systems, first commercialized by Sony in 1991. Besides powering cellular phones, laptops, digital cameras, power tools and medical devices, Li-ion is also used for electric vehicles and satellites. The battery has a number of benefits, most notably its high specific energy, simple charging, low maintenance and being environmentally benign.