# data[].attributes.body
2 V must be just above the decomposition voltage

# data[].attributes.attachment_link.description
In chemistry, a solution is a special type of homogeneous mixture composed of two or more substances. The term aqueous solution is when one of the solvents is water. In such a mixture, a solute is a substance dissolved in another substance, known as a solvent.

# data[].attributes.attachment_link.description
The voltage at which electrolysis is thermodynamically preferred is the difference of the electrode potentials as calculated using the Nernst equation. Often the actual voltage needed for electrolysis exceeds the thermodynamical value. The additional voltage is referred to as overpotential and is especially high in electrolysis reactions that involve gases, such as oxygen, hydrogen or chlorine.

# data[].attributes.attachment_link.description
In mathematics, extrapolation is the process of estimating, beyond the original observation range, the value of a variable on the basis of its relationship with another variable. It is similar to interpolation, which produces estimates between known observations, but extrapolation is subject to greater uncertainty and a higher risk of producing meaningless results.

# data[].attributes.attachment_link.description
The most common use of sulfuric acid (60% of total) is for fertilizer manufacture.[citation needed ] It is also a central substance in the chemical industry. Principal uses include fertilizer manufacturing (and other mineral processing), oil refining, wastewater processing, and chemical synthesis.

# data[].attributes.headline
Chemistry Investigate and compare the quantitative effects of changing(b) concentration of sulfuric acid on the rate of production of hydrogen gas during the electrolysis of acidified water.

# data[].attributes.subject
Chemistry Investigate and compare the quantitative effects of changing(b) concentration of sulfuric acid on the rate of production of hydrogen gas during the electrolysis of acidified water.

# data[].attributes.headline
Increasing linear graph through the origin

# data[].attributes.subject
Increasing linear graph through the origin

# data[].attributes.attachment_link.readable_text
In chemistry and manufacturing, electrolysis is a technique that uses direct electric current (DC) to drive an otherwise non-spontaneous chemical reaction. Electrolysis is commercially important as a stage in the separation of elements from naturally occurring sources such as ores using an electrolytic cell. The voltage that is needed for electrolysis to occur is called the decomposition potential. The word "lysis" means to separate or break, so in terms, electrolysis would mean "breakdown via electricity".

The word "electrolysis" was introduced by Michael Faraday in 1834, using the Greek words [ɛ̌ːlektron] "amber", which since the 17th century was associated with electrical phenomena, and meaning "dissolution". Nevertheless, electrolysis, as a tool to study chemical reactions and obtain pure elements, precedes the coinage of the term and formal description by Faraday.

In the early nineteenth century, William Nicholson and Anthony Carlisle sought to further Volta's experiments. They attached two wires to either side of a voltaic pile and placed the other ends in a tube filled with water. They noticed when the wires were brought together that each wire produced bubbles. One type was hydrogen, the other was oxygen.

In 1785 a Dutch scientist named Martin van Marum created an electrostatic generator that he used to reduce tin, zinc and antimony from their salts using a process later known as electrolysis. Though he unknowingly produced electrolysis, it was not until 1800 when William Nicholson and Anthony Carlisle discovered how electrolysis works.

In 1791 Luigi Galvani experimented with frog legs. He claimed that placing animal muscle between two dissimilar metal sheets resulted in electricity. Responding to these claims, Alessandro Volta conducted his own tests. This would give insight to Humphry Davy's ideas on electrolysis. During preliminary experiments, Humphry Davy hypothesized that when two elements combine together to form a compound, electrical energy is released. Humphry Davy would go on to create Decomposition Tables from his preliminary experiments on Electrolysis. The Decomposition Tables would give insight on the energies needed to break apart certain compounds.

During the later years of Humphry Davy's research, Michael Faraday became his assistant. While studying the process of electrolysis under Humphry Davy, Michael Faraday discovered two laws of electrolysis.

During the time of Maxwell and Faraday, concerns came about[vague ] for electropositive and electronegative activities.

In November 1875, Paul Émile Lecoq de Boisbaudran discovered gallium using electrolysis of gallium hydroxide, producing 3.4 mg of gallium. The following December, he presented his discovery of gallium to the Académie des sciences in Paris.

In 1886 Charles Martin Hall from America and Paul Héroult from France both filed for American patents for the electrolysis of aluminum, with Héroult submitting his in May, and Hall, in July. Hall was able to get his patent by proving through letters to his brother and family evidence that his method was discovered before the French patent was submitted. This became known as the Hall-Héroult process which benefited many industries because aluminum's price then dropped from four dollars to thirty cents per pound.

1785 - Martinus van Marum's electrostatic generator was used to reduce tin, zinc, and antimony from their salts using electrolysis.

1800 - William Nicholson and Anthony Carlisle (and also Johann Ritter), decomposed water into hydrogen and oxygen.

1808 - Potassium (1807), sodium (1807), barium, calcium and magnesium were discovered by Humphry Davy using electrolysis.

1821 - Lithium was discovered by the English chemist William Thomas Brande, who obtained it by electrolysis of lithium oxide.

1834 - Michael Faraday published his two laws of electrolysis, provided a mathematical explanation for them, and introduced terminology such as electrode, electrolyte, anode, cathode, anion, and cation.

1875 - Paul Émile Lecoq de Boisbaudran discovered gallium using electrolysis.

1886 - Fluorine was discovered by Henri Moissan using electrolysis.

1886 - Hall-Héroult process developed for making aluminium.

1890 - Castner-Kellner process developed for making sodium hydroxide.

Electrolysis is the passing of a direct electric current through an electrolyte producing chemical reactions at the electrodes and decomposition of the materials.

The main components required to achieve electrolysis are an electrolyte, electrodes, and an external power source. A partition (e.g. an ion-exchange membrane or a salt bridge) is optional to keep the products from diffusing to the vicinity of the opposite electrode.

The electrolyte is a chemical substance which contains free ions and carries electric current (e.g. an ion-conducting polymer, solution, or a ionic liquid compound). If the ions are not mobile, as in most solid salts, then electrolysis cannot occur. A liquid electrolyte is produced by:

The electrodes are immersed separated by a distance such that a current flows between them through the electrolyte and are connected to the power source which completes the electrical circuit. A direct current supplied by the power source drives the reaction causing ions in the electrolyte to be attracted toward the respective oppositely charged electrode.

Electrodes of metal, graphite and semiconductor material are widely used. Choice of suitable electrode depends on chemical reactivity between the electrode and electrolyte and manufacturing cost. Historically, when non-reactive anodes were desired for electrolysis, graphite (called plumbago in Faraday's time) or platinum were chosen. They were found to be some of the least reactive materials for anodes. Platinum erodes very slowly compared to other materials, and graphite crumbles and can produce carbon dioxide in aqueous solutions but otherwise does not participate in the reaction. Cathodes may be made of the same material, or they may be made from a more reactive one since anode wear is greater due to oxidation at the anode.

The key process of electrolysis is the interchange of atoms and ions by the removal or addition of electrons due to the applied current. The desired products of electrolysis are often in a different physical state from the electrolyte and can be removed by mechanical processes (e.g. by collecting gas above an electrode or precipitating a product out of the electrolyte).

The quantity of the products is proportional to the current, and when two or more electrolytic cells are connected in series to the same power source, the products produced in the cells are proportional to their equivalent weight. These are known as Faraday's laws of electrolysis.

Each electrode attracts ions that are of the opposite charge. Positively charged ions ( cations) move towards the electron-providing (negative) cathode. Negatively charged ions ( anions) move towards the electron-extracting (positive) anode. In this process electrons are effectively introduced at the cathode as a reactant and removed at the anode as a product. In chemistry, the loss of electrons is called oxidation, while electron gain is called reduction.

When neutral atoms or molecules, such as those on the surface of an electrode, gain or lose electrons they become ions and may dissolve in the electrolyte and react with other ions.

When ions gain or lose electrons and become neutral, they will form compounds that separate from the electrolyte. Positive metal ions like Cu ²⁺ deposit onto the cathode in a layer. The terms for this are electroplating, electrowinning, and electrorefining.

When an ion gains or loses electrons without becoming neutral, its electronic charge is altered in the process.

For example, the electrolysis of brine produces hydrogen and chlorine gases which bubble from the electrolyte and are collected. The initial overall reaction is thus:

The reaction at the anode results in chlorine gas from chlorine ions:

The reaction at the cathode results in hydrogen gas and hydroxide ions:

Without a partition between the electrodes, the OH − ions produced at the cathode are free to diffuse throughout the electrolyte to the anode. As the electrolyte becomes more basic due to the production of OH −, less Cl ₂ emerges from the solution as it begins to react with the hydroxide producing hypochlorite (ClO −) at the anode:

The more opportunity the Cl ₂ has to interact with NaOH in the solution, the less Cl ₂ emerges at the surface of the solution and the faster the production of hypochlorite progresses. This depends on factors such as solution temperature, the amount of time the Cl ₂ molecule is in contact with the solution, and concentration of NaOH.

Likewise, as hypochlorite increases in concentration, chlorates are produced from them:

Other reactions occur, such as the self-ionization of water and the decomposition of hypochlorite at the cathode, the rate of the latter depends on factors such as diffusion and the surface area of the cathode in contact with the electrolyte.

Decomposition potential or decomposition voltage refers to the minimum voltage (difference in electrode potential) between anode and cathode of an electrolytic cell that is needed for electrolysis to occur.

The voltage at which electrolysis is thermodynamically preferred is the difference of the electrode potentials as calculated using the Nernst equation. Applying additional voltage, referred to as overpotential, can increase the rate of reaction and is often needed above the thermodynamic value. It is especially necessary for electrolysis reactions involving gases, such as oxygen, hydrogen or chlorine.

Oxidation of ions or neutral molecules occurs at the anode. For example, it is possible to oxidize ferrous ions to ferric ions at the anode:

Reduction of ions or neutral molecules occurs at the cathode. It is possible to reduce ferricyanide ions to ferrocyanide ions at the cathode:

Neutral molecules can also react at either of the electrodes. For example: p-benzoquinone can be reduced to hydroquinone at the cathode:

+ 2 e − + 2 H⁺ →

In the last example, H⁺ ions (hydrogen ions) also take part in the reaction and are provided by the acid in the solution, or by the solvent itself (water, methanol, etc.). Electrolysis reactions involving H⁺ ions are fairly common in acidic solutions. In aqueous alkaline solutions, reactions involving OH − (hydroxide ions) are common.

Sometimes the solvents themselves (usually water) are oxidized or reduced at the electrodes. It is even possible to have electrolysis involving gases, e.g. by using a gas diffusion electrode.

The amount of electrical energy that must be added equals the change in Gibbs free energy of the reaction plus the losses in the system. The losses can (in theory) be arbitrarily close to zero, so the maximum thermodynamic efficiency equals the enthalpy change divided by the free energy change of the reaction. In most cases, the electric input is larger than the enthalpy change of the reaction, so some energy is released in the form of heat. In some cases, for instance, in the electrolysis of steam into hydrogen and oxygen at high temperature, the opposite is true and heat energy is absorbed. This heat is absorbed from the surroundings, and the heating value of the produced hydrogen is higher than the electric input.

Galvanic cells and batteries use spontaneous, energy-releasing redox reactions to generate an electrical potential that provides useful power. When a secondary battery is charged, its redox reaction is run in reverse and the system can be considered as an electrolytic cell.

In manufacturing, electrolysis can be used for:

Electroplating, where a thin film of metal is deposited over a substrate material. Electroplating is used in many industries for either functional or decorative purposes, as in-vehicle bodies and nickel coins.

Electrochemical machining (ECM), where an electrolytic cathode is used as a shaped tool for removing material by anodic oxidation from a workpiece. ECM is often used as a technique for deburring or for etching metal surfaces like tools or knives with a permanent mark or logo.

Using a cell containing inert platinum electrodes, electrolysis of aqueous solutions of some salts leads to the reduction of the cations (such as metal deposition with, for example, zinc salts) and oxidation of the anions (such as the evolution of bromine with bromides). However, with salts of some metals (such as sodium) hydrogen is evolved at the cathode, and for salts containing some anions (such as sulfate ) oxygen is evolved at the anode. In both cases, this is due to water being reduced to form hydrogen or oxidized to form oxygen. In principle, the voltage required to electrolyze a salt solution can be derived from the standard electrode potential for the reactions at the anode and cathode. The standard electrode potential is directly related to the Gibbs free energy, Δ G, for the reactions at each electrode and refers to an electrode with no current flowing. An extract from the table of standard electrode potentials is shown below.

In terms of electrolysis, this table should be interpreted as follows:

Moving down the table, E° becomes more positive, and species on the left are more likely to be reduced: for example, zinc ions are more likely to be reduced to zinc metal than sodium ions are to be reduced to sodium metal.

Moving up the table, E° becomes more negative, and species on the right are more likely to be oxidized: for example, sodium metal is more likely to be oxidized to sodium ions than zinc metal is to be oxidized to zinc ions.

Using the Nernst equation the electrode potential can be calculated for a specific concentration of ions, temperature and the number of electrons involved. For pure water ( pH 7):

the electrode potential for the reduction producing hydrogen is −0.41 V,

the electrode potential for the oxidation producing oxygen is +0.82 V.

Comparable figures calculated in a similar way, for 1 M zinc bromide, ZnBr ₂, are −0.76 V for the reduction to Zn metal and +1.10 V for the oxidation producing bromine. The conclusion from these figures is that hydrogen should be produced at the cathode and oxygen at the anode from the electrolysis of water-which is at variance with the experimental observation that zinc metal is deposited and bromine is produced. The explanation is that these calculated potentials only indicate the thermodynamically preferred reaction. In practice, many other factors have to be taken into account such as the kinetics of some of the reaction steps involved. These factors together mean that a higher potential is required for the reduction and oxidation of water than predicted, and these are termed overpotentials. Experimentally it is known that overpotentials depend on the design of the cell and the nature of the electrodes.

For the electrolysis of a neutral (pH 7) sodium chloride solution, the reduction of sodium ion is thermodynamically very difficult and water is reduced evolving hydrogen leaving hydroxide ions in solution. At the anode the oxidation of chlorine is observed rather than the oxidation of water since the overpotential for the oxidation of chloride to chlorine is lower than the overpotential for the oxidation of water to oxygen. The hydroxide ions and dissolved chlorine gas react further to form hypochlorous acid. The aqueous solutions resulting from this process is called electrolyzed water and is used as a disinfectant and cleaning agent.

The electrochemical reduction or electrocatalytic conversion of can produce value-added chemicals such methane, ethylene, ethanol, etc. The electrolysis of carbon dioxide gives formate or carbon monoxide, but sometimes more elaborate organic compounds such as ethylene. This technology is under research as a carbon-neutral route to organic compounds.

Electrolysis of water produces hydrogen and oxygen in a ratio of 2 to 1 respectively.

The National Renewable Energy Laboratory estimated in 2006 that 1 kg of hydrogen (roughly equivalent to 3 kg, or 4 liters, of petroleum in energy terms) could be produced by wind powered electrolysis for between US$5.55 in the near term and US$2.27 in the longer term.

About 4% of hydrogen gas produced worldwide is generated by electrolysis, and normally used onsite. Hydrogen is used for the creation of ammonia for fertilizer via the Haber process, and converting heavy petroleum sources to lighter fractions via hydrocracking. Recently, onsite electrolysis has been utilized to capture hydrogen for hydrogen fuel-cells in hydrogen vehicles.

Carbon/hydrocarbon assisted water electrolysis

Recently, to reduce the energy input, the utilization of carbon ( coal), alcohols (hydrocarbon solution), and organic solution ( glycerol, formic acid, ethylene glycol, etc.) with co-electrolysis of water has been proposed as a viable option. The carbon/hydrocarbon assisted water electrolysis (so-called CAWE) process for hydrogen generation would perform this operation in a single electrochemical reactor. This system energy balance can be required only around 40% electric input with 60% coming from the chemical energy of carbon or hydrocarbon. This process utilizes solid coal/carbon particles or powder as fuels dispersed in acid/alkaline electrolyte in the form of slurry and the carbon contained source co-assist in the electrolysis process as following theoretical overall reactions:

Carbon/Coal slurry (C + 2H ₂O) → CO ₂ + 2H ₂ E′ = 0.21 V (reversible voltage) / E ′ = 0.46 V (thermo-neutral voltage)

or

Carbon/Coal slurry (C + H ₂O) → CO + H ₂ E′ = 0.52 V (reversible voltage) / E ′ = 0.91 V (thermo-neutral voltage)

Thus, this CAWE approach is that the actual cell overpotential can be significantly reduced to below 1.0 V as compared to 1.5 V for conventional water electrolysis.

The current method of producing steel from iron ore is very carbon intensive, in part to the direct release of CO ₂ in the blast furnace. A study of steel making in Germany found that producing 1 ton of steel emitted 2.1 tons of with 22% of that being direct emissions from the blast furnace. As of 2022, steel production contributes from 7 to 9% of global emissions. Electrolysis of iron can eliminate direct emissions and further reduce emissions if the electricity is created from green energy.

The small-scale electrolysis of iron has been successfully reported by dissolving it in molten oxide salts and using a platinum anode. Oxygen anions form oxygen gas and electrons at the anode. Iron cations consume electrons and form iron metal at the cathode. This method was performed a temperature of 1550 °C which presents a significant challenge to maintaining the reaction. Particularly, anode corrosion is a concern at these temperatures.

Additionally, the low temperature reduction of iron oxide by dissolving it in alkaline water has been reported. The temperature is much lower than traditional iron production at 114 °C. The low temperatures also tend to correlate with higher current efficiencies, with an efficiency of 95% being reported. While these methods are promising, they struggle to be cost competitive because of the large economies of scale keeping the price of blast furnace iron low.

# data[].attributes.attachment_link.description
In chemistry and manufacturing, electrolysis is a technique that uses direct electric current (DC) to drive an otherwise non-spontaneous chemical reaction. Electrolysis is commercially important as a stage in the separation of elements from naturally occurring sources such as ores using an electrolytic cell.

# data[].attributes.attachment_link.description
This website uses cookies. They are harmless and used to f.x. anonymous data analysis. In order to use our website, please accept Cookies Terms of Use OK

# data[].attributes.attachment_link.description
The mole is the unit of measurement for amount of substance in the International System of Units (SI). The unit is defined as the amount or sample of a chemical substance that contains as many constitutive particles, e.g., atoms, molecules, ions, electrons, or photons, as there are atoms in 12 grams of carbon-12 ( 12C), the isotope of carbon with standard atomic weight 12 by definition.

# data[].attributes.attachment_link.readable_text
Pronunciation

​(PLAT-in-əm)

Appearance

silvery white

Standard atomic weight A r, std(Pt)

195.084(9) [1]

(Atomic number Z)

78

Group

group 10

Period
period 6

Block

d-block

Electron configuration

[ Xe] 4f ¹⁴ 5d ⁹ 6s ¹

Electrons per shell

2, 8, 18, 32, 17, 1

at Phase STP
solid

Melting point

2041.4 K ​(1768.3 °C, ​3214.9 °F)

(near Density r.t.)

21.45 g/cm ³

when liquid (at m.p.)

19.77 g/cm ³

Heat of fusion

22.17 kJ/mol

Vapor pressure

P (Pa)

1

10

100

1 k

10 k

100 k

at T (K)

2330

(2550)

2815

3143

3556

4094

Oxidation states

−3, −2, −1, 0, +1, +2, +3, +4, +5, +6 (a mildly basic oxide)

Ionization energies

1st: 870 kJ/mol

2nd: 1791 kJ/mol

Atomic radius

empirical: 139 pm

Spectral lines of platinum

Natural occurrence
primordial

Crystal structure

​ face-centered cubic (fcc)

Speed of sound thin rod

2800 m/s (at r.t.)

Magnetic ordering

paramagnetic

Molar magnetic susceptibility

+201.9 × 10 −6 cm ³/mol (290 K)[2]

Discovery

Antonio de Ulloa (1735)

Category: Platinum
| references

Platinum is a chemical element with the symbol Pt and atomic number 78. It is a dense, malleable, ductile, highly unreactive, precious, silverish-white transition metal. Its name is derived from the Spanish term platino, meaning "little silver".

Platinum is a member of the platinum group of elements and group 10 of the periodic table of elements. It has six naturally occurring isotopes. It is one of the rarer elements in Earth's crust, with an average abundance of approximately 5 μg/kg. It occurs in some nickel and copper ores along with some native deposits, mostly in South Africa, which accounts for 80% of the world production. Because of its scarcity in Earth's crust, only a few hundred tonnes are produced annually, and given its important uses, it is highly valuable and is a major precious metal commodity.

Platinum is one of the least reactive metals. It has remarkable resistance to corrosion, even at high temperatures, and is therefore considered a noble metal. Consequently, platinum is often found chemically uncombined as native platinum. Because it occurs naturally in the alluvial sands of various rivers, it was first used by pre-Columbian South American natives to produce artifacts. It was referenced in European writings as early as 16th century, but it was not until Antonio de Ulloa published a report on a new metal of Colombian origin in 1748 that it began to be investigated by scientists.

Platinum is used in catalytic converters, laboratory equipment, electrical contacts and electrodes, platinum resistance thermometers, dentistry equipment, and jewelry. Being a heavy metal, it leads to health problems upon exposure to its salts; but due to its corrosion resistance, metallic platinum has not been linked to adverse health effects. Compounds containing platinum, such as cisplatin, oxaliplatin and carboplatin, are applied in chemotherapy against certain types of cancer.

In early 2021, the value of platinum ranged from US$1,055 per troy ounce to US$1,320 per troy ounce.

Pure platinum is a lustrous, ductile, and malleable, silver-white metal. Platinum is more ductile than gold, silver or copper, thus being the most ductile of pure metals, but it is less malleable than gold. The metal has excellent resistance to corrosion, is stable at high temperatures and has stable electrical properties. Platinum does oxidize, forming , at 500 °C; this oxide can be easily removed thermally. It reacts vigorously with fluorine at 500 °C (932 °F) to form platinum tetrafluoride. It is also attacked by chlorine, bromine, iodine, and sulfur. Platinum is insoluble in hydrochloric and nitric acid, but dissolves in hot (a mixture of nitric and hydrochloric acids), to form chloroplatinic acid, .

Its physical characteristics and chemical stability make it useful for industrial applications. Its resistance to wear and tarnish is well suited to use in fine jewellery.

The most common oxidation states of platinum are +2 and +4. The +1 and +3 oxidation states are less common, and are often stabilized by metal bonding in bimetallic (or polymetallic) species. As is expected, tetracoordinate platinum(II) compounds tend to adopt 16-electron square planar geometries. Although elemental platinum is generally unreactive, it dissolves in hot aqua regia to give aqueous chloroplatinic acid ):

As a soft acid, platinum has a great affinity for sulfur, such as on dimethyl sulfoxide (DMSO); numerous DMSO complexes have been reported and care should be taken in the choice of reaction solvent.

In 2007, Gerhard Ertl won the Nobel Prize in Chemistry for determining the detailed molecular mechanisms of the catalytic oxidation of carbon monoxide over platinum ( catalytic converter).

Platinum has six naturally occurring isotopes: ¹⁹⁰Pt, ¹⁹²Pt, ¹⁹⁴Pt, ¹⁹⁵Pt, ¹⁹⁶Pt, and ¹⁹⁸Pt. The most abundant of these is ¹⁹⁵Pt, comprising 33.83% of all platinum. It is the only stable isotope with a non-zero spin. The spin of ¹/ ₂ and other favourable magnetic properties of the nucleus are utilised in . Due to its spin and large abundance, ¹⁹⁵Pt satellite peaks are also often observed in ¹H and ³¹P NMR spectroscopy ( e.g., for Pt-phosphine and Pt-alkyl complexes). ¹⁹⁰Pt is the least abundant at only 0.01%. Of the naturally occurring isotopes, only ¹⁹⁰Pt is unstable, though it decays with a half-life of 6.5×10 years, causing an activity of 15 Bq/kg of natural platinum. ¹⁹⁸Pt can undergo alpha decay, but its decay has never been observed (the half-life is known to be longer than 3.2×10 years); therefore, it is considered stable. Platinum also has 34 synthetic isotopes ranging in atomic mass from 165 to 204, making the total number of known isotopes 40. The least stable of these are ¹⁶⁵Pt and ¹⁶⁶Pt, with half-lives of 260 µs, whereas the most stable is ¹⁹³Pt with a half-life of 50 years. Most platinum isotopes decay by some combination of beta decay and alpha decay. ¹⁸⁸Pt, ¹⁹¹Pt, and ¹⁹³Pt decay primarily by electron capture. ¹⁹⁰Pt and ¹⁹⁸Pt are predicted to have energetically favorable double beta decay paths.

Platinum is an extremely rare metal, occurring at a concentration of only 0.005 ppm in Earth's crust. It is sometimes mistaken for silver. Platinum is often found chemically uncombined as native platinum and as alloy with the other platinum-group metals and iron mostly. Most often the native platinum is found in secondary deposits in alluvial deposits. The alluvial deposits used by pre-Columbian people in the Chocó Department, Colombia are still a source for platinum-group metals. Another large alluvial deposit is in the Ural Mountains, Russia, and it is still mined.

In nickel and copper deposits, platinum-group metals occur as sulfides (e.g., (Pt,Pd)S), tellurides (e.g., PtBiTe), antimonides (PdSb), and arsenides (e.g. ), and as end alloys with nickel or copper. Platinum arsenide, sperrylite (), is a major source of platinum associated with nickel ores in the Sudbury Basin deposit in Ontario, Canada. At Platinum, Alaska, about 17,000 kg (550,000 ozt) was mined between 1927 and 1975. The mine ceased operations in 1990. The rare sulfide mineral cooperite, (Pt,Pd,Ni)S, contains platinum along with palladium and nickel. Cooperite occurs in the Merensky Reef within the Bushveld complex, Gauteng, South Africa.

In 1865, chromites were identified in the Bushveld region of South Africa, followed by the discovery of platinum in 1906. In 1924, the geologist Hans Merensky discovered a large supply of platinum in the Bushveld Igneous Complex in South Africa. The specific layer he found, named the Merensky Reef, contains around 75% of the world's known platinum. The large copper-nickel deposits near Norilsk in Russia, and the Sudbury Basin, Canada, are the two other large deposits. In the Sudbury Basin, the huge quantities of nickel ore processed make up for the fact platinum is present as only 0.5 ppm in the ore. Smaller reserves can be found in the United States, for example in the Absaroka Range in Montana. In 2010, South Africa was the top producer of platinum, with an almost 77% share, followed by Russia at 13%; world production in 2010 was 192,000 kg (423,000 lb).

Large platinum deposits are present in the state of Tamil Nadu, India.

Hexachloroplatinic acid mentioned above is probably the most important platinum compound, as it serves as the precursor for many other platinum compounds. By itself, it has various applications in photography, zinc etchings, indelible ink, plating, mirrors, porcelain coloring, and as a catalyst.

Treatment of hexachloroplatinic acid with an ammonium salt, such as ammonium chloride, gives ammonium hexachloroplatinate, which is relatively insoluble in ammonium solutions. Heating this ammonium salt in the presence of hydrogen reduces it to elemental platinum. Potassium hexachloroplatinate is similarly insoluble, and hexachloroplatinic acid has been used in the determination of potassium ions by gravimetry.

When hexachloroplatinic acid is heated, it decomposes through platinum(IV) chloride and platinum(II) chloride to elemental platinum, although the reactions do not occur stepwise:

All three reactions are reversible. Platinum(II) and platinum(IV) bromides are known as well. Platinum hexafluoride is a strong oxidizer capable of oxidizing oxygen.

Platinum(IV) oxide, , also known as ' Adams' catalyst', is a black powder that is soluble in potassium hydroxide (KOH) solutions and concentrated acids. and the less common PtO both decompose upon heating. Platinum(II,IV) oxide, , is formed in the following reaction:

Unlike palladium acetate, platinum(II) acetate is not commercially available. Where a base is desired, the halides have been used in conjunction with sodium acetate. The use of platinum(II) acetylacetonate has also been reported.

Several barium platinides have been synthesized in which platinum exhibits negative oxidation states ranging from −1 to −2. These include BaPt, , and . Caesium platinide, , a dark-red transparent crystalline compound has been shown to contain Pt anions. Platinum also exhibits negative oxidation states at surfaces reduced electrochemically. The negative oxidation states exhibited by platinum are unusual for metallic elements, and they are attributed to the relativistic stabilization of the 6s orbitals.

Zeise's salt, containing an ethylene ligand, was one of the first organometallic compounds discovered. Dichloro(cycloocta-1,5-diene)platinum(II) is a commercially available olefin complex, which contains easily displaceable cod ligands ("cod" being an abbreviation of 1,5-cyclooctadiene). The cod complex and the halides are convenient starting points to platinum chemistry.

Cisplatin, or cis-diamminedichloroplatinum(II) is the first of a series of square planar platinum(II)-containing chemotherapy drugs. Others include carboplatin and oxaliplatin. These compounds are capable of crosslinking DNA, and kill cells by similar pathways to alkylating chemotherapeutic agents. (Side effects of cisplatin include nausea and vomiting, hair loss, tinnitus, hearing loss, and nephrotoxicity.)

Organoplatinum compounds such as the above antitumour agents, as well as soluble inorganic platinum complexes, are routinely characterised using ¹⁹⁵Pt nuclear magnetic resonance spectroscopy.

Archaeologists have discovered traces of platinum in the gold used in ancient Egyptian burials as early as 1200 BC. For example, a small box from burial of Shepenupet II was found to be decorated with gold-platinum hieroglyphics. However, the extent of early Egyptians' knowledge of the metal is unclear. It is quite possible they did not recognize there was platinum in their gold.

The metal was used by pre-Columbian Americans near modern-day Esmeraldas, Ecuador to produce artifacts of a white gold-platinum alloy. Archeologists usually associate the tradition of platinum-working in South America with the La Tolita Culture (circa 600 BC - AD 200), but precise dates and location is difficult, as most platinum artifacts from the area were bought secondhand through the antiquities trade rather than obtained by direct archeological excavation. To work the metal, they would combine gold and platinum powders by sintering. The resulting gold-platinum alloy would then be soft enough to shape with tools. The platinum used in such objects was not the pure element, but rather a naturally occurring mixture of the platinum group metals, with small amounts of palladium, rhodium, and iridium.

The first European reference to platinum appears in 1557 in the writings of the Italian humanist Julius Caesar Scaliger as a description of an unknown noble metal found between Darién and Mexico, "which no fire nor any Spanish artifice has yet been able to liquefy". From their first encounters with platinum, the Spanish generally saw the metal as a kind of impurity in gold, and it was treated as such. It was often simply thrown away, and there was an official decree forbidding the adulteration of gold with platinum impurities.

In 1735, Antonio de Ulloa and Jorge Juan y Santacilia saw Native Americans mining platinum while the Spaniards were travelling through Colombia and Peru for eight years. Ulloa and Juan found mines with the whitish metal nuggets and took them home to Spain. Antonio de Ulloa returned to Spain and established the first mineralogy lab in Spain and was the first to systematically study platinum, which was in 1748. His historical account of the expedition included a description of platinum as being neither separable nor calcinable. Ulloa also anticipated the discovery of platinum mines. After publishing the report in 1748, Ulloa did not continue to investigate the new metal. In 1758, he was sent to superintend mercury mining operations in Huancavelica.

In 1741, Charles Wood, a British metallurgist, found various samples of Colombian platinum in Jamaica, which he sent to William Brownrigg for further investigation.

In 1750, after studying the platinum sent to him by Wood, Brownrigg presented a detailed account of the metal to the Royal Society, stating that he had seen no mention of it in any previous accounts of known minerals. Brownrigg also made note of platinum's extremely high melting point and refractoriness toward borax.[clarification needed ] Other chemists across Europe soon began studying platinum, including Andreas Sigismund Marggraf, Torbern Bergman, Jöns Jakob Berzelius, William Lewis, and Pierre Macquer. In 1752, Henrik Scheffer published a detailed scientific description of the metal, which he referred to as "white gold", including an account of how he succeeded in fusing platinum ore with the aid of arsenic. Scheffer described platinum as being less pliable than gold, but with similar resistance to corrosion.

Carl von Sickingen researched platinum extensively in 1772. He succeeded in making malleable platinum by alloying it with gold, dissolving the alloy in hot , precipitating the platinum with ammonium chloride, igniting the ammonium chloroplatinate, and hammering the resulting finely divided platinum to make it cohere. Franz Karl Achard made the first platinum crucible in 1784. He worked with the platinum by fusing it with arsenic, then later volatilizing the arsenic.

Platinum, along with the rest of the platinum-group metals, is obtained commercially as a by-product from nickel and copper mining and processing. During electrorefining of copper, noble metals such as silver, gold and the platinum-group metals as well as selenium and tellurium settle to the bottom of the cell as "anode mud", which forms the starting point for the extraction of the platinum-group metals.

One suitable method for purification for the raw platinum, which contains platinum, gold, and the other platinum-group metals, is to process it with aqua regia, in which palladium, gold and platinum are dissolved, whereas osmium, iridium, ruthenium and rhodium stay unreacted. The gold is precipitated by the addition of iron(II) chloride and after filtering off the gold, the platinum is precipitated as ammonium chloroplatinate by the addition of ammonium chloride. Ammonium chloroplatinate can be converted to platinum by heating. Unprecipitated hexachloroplatinate(IV) may be reduced with elemental zinc, and a similar method is suitable for small scale recovery of platinum from laboratory residues. Mining and refining platinum has environmental impacts.

The most common use of platinum is as a catalyst in chemical reactions, often as platinum black. It has been employed as a catalyst since the early 19th century, when platinum powder was used to catalyze the ignition of hydrogen. Its most important application is in automobiles as a catalytic converter, which allows the complete combustion of low concentrations of unburned hydrocarbons from the exhaust into carbon dioxide and water vapor. Platinum is also used in the petroleum industry as a catalyst in a number of separate processes, but especially in catalytic reforming of straight-run naphthas into higher-octane gasoline that becomes rich in aromatic compounds. , also known as Adams' catalyst, is used as a hydrogenation catalyst, specifically for vegetable oils. Platinum also strongly catalyzes the decomposition of hydrogen peroxide into water and oxygen and it is used in fuel cells as a catalyst for the reduction of oxygen.

From 1889 to 1960, the meter was defined as the length of a platinum-iridium (90:10) alloy bar, known as the international prototype of the meter. The previous bar was made of platinum in 1799. Until May 2019, the kilogram was defined as the mass of the international prototype of the kilogram, a cylinder of the same platinum-iridium alloy made in 1879.

The Standard Platinum Resistance Thermometer (SPRT) is one of the four types of thermometers used to define the International Temperature Scale of 1990 (ITS-90), the international calibration standard for temperature measurements. The resistance wire in the thermometer is made of pure platinum (e.g. The NIST manufactured the wires from platinum bar stock with a chemical purity of 99.999% by weight). In addition to laboratory uses, Platinum Resistance Thermometry (PRT) also has many industrial applications, industrial standards include ASTM E1137 and IEC 60751.

The standard hydrogen electrode also uses a platinized platinum electrode due to its corrosion resistance, and other attributes.

Platinum is a precious metal commodity; its bullion has the ISO currency code of XPT. Coins, bars, and ingots are traded or collected. Platinum finds use in jewellery, usually as a 90-95% alloy, due to its inertness. It is used for this purpose for its prestige and inherent bullion value. Jewellery trade publications advise jewellers to present minute surface scratches (which they term patina) as a desirable feature in attempt to enhance value of platinum products.

In watchmaking, Vacheron Constantin, Patek Philippe, Rolex, Breitling, and other companies use platinum for producing their limited edition watch series. Watchmakers appreciate the unique properties of platinum, as it neither tarnishes nor wears out (the latter quality relative to gold).

During periods of sustained economic stability and growth, the price of platinum tends to be as much as twice the price of gold, whereas during periods of economic uncertainty, the price of platinum tends to decrease due to reduced industrial demand, falling below the price of gold. Gold prices are more stable in slow economic times, as gold is considered a safe haven. Although gold is also used in industrial applications, especially in electronics due to its use as a conductor, its demand is not so driven by industrial uses. In the 18th century, platinum's rarity made King Louis XV of France declare it the only metal fit for a king.

1,000 cubic centimeters of 99.9% pure platinum, worth about US$696,000 at 29 Jun 2016 prices

In the laboratory, platinum wire is used for electrodes; platinum pans and supports are used in thermogravimetric analysis because of the stringent requirements of chemical inertness upon heating to high temperatures (~1000 °C). Platinum is used as an alloying agent for various metal products, including fine wires, noncorrosive laboratory containers, medical instruments, dental prostheses, electrical contacts, and thermocouples. Platinum-cobalt, an alloy of roughly three parts platinum and one part cobalt, is used to make relatively strong permanent magnets. Platinum-based anodes are used in ships, pipelines, and steel piers. Platinum drugs are used to treat a wide variety of cancers, including testicular and ovarian carcinomas, melanoma, small-cell and non-small-cell lung cancer, myelomas and lymphomas.

Platinum's rarity as a metal has caused advertisers to associate it with exclusivity and wealth. "Platinum" debit and credit cards have greater privileges than " gold" cards. " Platinum awards" are the second highest possible, ranking above "gold", " silver" and " bronze", but below diamond. For example, in the United States, a musical album that has sold more than 1 million copies will be credited as "platinum", whereas an album that has sold more than 10 million copies will be certified as "diamond". Some products, such as blenders and vehicles, with a silvery-white color are identified as "platinum". Platinum is considered a precious metal, although its use is not as common as the use of gold or silver. The frame of the Crown of Queen Elizabeth The Queen Mother, manufactured for her coronation as Consort of King George VI, is made of platinum. It was the first British crown to be made of this particular metal.

According to the Centers for Disease Control and Prevention, short-term exposure to platinum salts may cause irritation of the eyes, nose, and throat, and long-term exposure may cause both respiratory and skin allergies. The current OSHA standard is 2 micrograms per cubic meter of air averaged over an 8-hour work shift. The National Institute for Occupational Safety and Health has set a recommended exposure limit (REL) for platinum as 1 mg/m ³ over an 8-hour workday.

As platinum is a catalyst in the manufacture of the silicone rubber and gel components of several types of medical implants (breast implants, joint replacement prosthetics, artificial lumbar discs, vascular access ports, etc.), the possibility that platinum could enter the body and cause adverse effects has merited study. The Food and Drug Administration and other institutions have reviewed the issue and found no evidence to suggest toxicity in vivo. Chemically unbounded platinum has been identified by the FDA as a "fake cancer 'cure'". The misunderstanding is created by healthcare workers who are using inappropriately the name of the metal as a slang term for platinum-based chemotherapy medications like cisplatin.[citation needed ] They are platinum compounds not the metal itself.

1. ^ Meija, Juris; et al. (2016). "Atomic weights of the elements 2013 (IUPAC Technical Report)". Pure and Applied Chemistry. 88 (3): 265-91. doi:10.1515/pac-2015-0305.

2. ^ "Chapter 6.11 Platinum" (PDF), Air Quality Guidelines (Second ed.), WHO Regional Office for Europe, Copenhagen, Denmark, 2000, archived (PDF) from the original on 18 October 2012

3. ^ Wheate, N. J.; Walker, S.; Craig, G. E.; Oun, R. (2010). "The status of platinum anticancer drugs in the clinic and in clinical trials" (PDF). Dalton Transactions. 39 (35): 8113-27. doi:10.1039/C0DT00292E. hdl:. PMID 20593091.

4. Lagowski, J. J., ed. (2004). Chemistry Foundations and Applications. 3. Thomson Gale. pp. 267-268. ISBN 978-0-02-865724-0.

5. ^ CRC press encyclopedia of materials and finishes, 2nd edition, Mel Schwartz, 2002

6. ^ Materials handbook, fifteenth edition, McGraw-Hill, by John Vaccari, 2002

7. ^ CRC contributors (2007-2008). "Platinum". In Lide, David R. (ed.). CRC Handbook of Chemistry and Physics. 4. New York: CRC Press. p. 26. ISBN 978-0-8493-0488-0.

8. Kauffman, George B.; Thurner, Joseph J.; Zatko, David A. (1967). Ammonium Hexachloroplatinate(IV). Inorganic Syntheses. 9. pp. 182-185. doi:10.1002/9780470132401.ch51. ISBN 978-0-470-13240-1.

9. ^ Han, Y.; Huynh, H. V.; Tan, G. K. (2007). "Mono- vs Bis(carbene) Complexes: A Detailed Study on Platinum(II)−Benzimidazolin-2-ylidenes". Organometallics. 26 (18): 4612-4617. doi:10.1021/om700543p.

10. ^ Audi, Georges; Bersillon, Olivier; Blachot, Jean; Wapstra, Aaldert Hendrik (2003), "The NUBASE evaluation of nuclear and decay properties", Nuclear Physics A, 729: 3-128, Bibcode:2003NuPhA.729....3A, doi:10.1016/j.nuclphysa.2003.11.001

11. ^ "The History of Platinum". Alaska Community Database Online. ExploreNorth. Archived from the original on 22 December 2010. Platinum is located on the Bering Sea coast, below Red Mountain on the south spit of Goodnews Bay.

12. ^ R. Grant Cawthorn (1999). "Seventy-fifth Anniversary of the Discovery of the Platiniferous Merensky Reef". Platinum Metals Review.

13. Seymour, R. J.; O'Farrelly, J. I. (2001). "Platinum-group metals". Kirk Othmer Encyclopedia of Chemical Technology. Wiley. doi:10.1002/0471238961.1612012019052513.a01.pub2. ISBN 978-0471238966.

14. ^ "Mining Platinum in Montana". New York Times. 13 August 1998. Archived from the original on 3 February 2008.

15. ^ Krebs, Robert E. (1998). "Platinum". The History and Use of our Earth's Chemical Elements. Greenwood Press. pp. 124-127. ISBN 978-0-313-30123-0.

16. ^ Smith, G. F.; Gring, J. L. (1933). "The Separation and Determination of the Alkali Metals Using Perchloric Acid. V. Perchloric Acid and Chloroplatinic Acid in the Determination of Small Amounts of Potassium in the Presence of Large Amounts of Sodium". Journal of the American Chemical Society. 55 (10): 3957-3961. doi:10.1021/ja01337a007.

17. ^ Schweizer, A. E.; Kerr, G. T. (1978). "Thermal Decomposition of Hexachloroplatinic Acid". Inorganic Chemistry. 17 (8): 2326-2327. doi:10.1021/ic50186a067.

18. Jansen, Martin (2005). "Effects of relativistic motion of electrons on the chemistry of gold and platinum". Solid State Sciences. 7 (12): 1464-1474. Bibcode:2005SSSci...7.1464J. doi:10.1016/j.solidstatesciences.2005.06.015.

19. ^ Riddell, Imogen A.; Lippard, Stephen J. (2018). "Cisplatin and Oxaliplatin:Our Current Understanding of Their Actions". In Sigel, Astrid; Sigel, Helmut; Freisinger, Eva; Sigel, Roland K. O. (eds.). Metallo-Drugs:Development and Action of Anticancer Agents. Metal Ions in Life Sciences. 18. pp. 1-42. doi:10.1515/9783110470734-007. ISBN 978-3-11-046984-4. PMID 29394020.

20. ^ Richards, A. D.; Rodger, A. (2007). "Synthetic metallomolecules as agents for the control of DNA structure" (PDF). Chemical Society Reviews. 36 (3): 471-483. doi:10.1039/b609495c. PMID 17325786.

21. ^ Carinder, James A.; Morrison, Pilar M.; Morrison, David G.; Jack E. Saux III (7 July 2014). Practical Oncology Protocols. Mill City Press, Incorporated. p. 22. ISBN 978-1-62652-816-1. Archived from the original on 9 November 2017.

22. ^ Taguchi, Takashi; Nazneen, Arifa; Abid, M. Ruhul; Razzaque, Mohammed S. (2005). Cisplatin-Associated Nephrotoxicity and Pathological Events. Contributions to Nephrology. 148. pp. 107-121. doi:10.1159/000086055. ISBN 978-3-8055-7858-5. PMID 15912030. S2CID 24509477.

23. ^ Berthelot, M. (1901). "Sur les métaux égyptiens: Présence du platine parmi les caractères d'inscriptions hiéroglyphiques, confié à mon examn" [On Egyptian metals: Presence of platinum among the characters of hieroglyphic inscriptions, entrusted to my examination]. Comptes rendus de l'Académie des Sciences (in French). 132: 729.

24. ^ Ogden, Jack M. (1976). "The So-Called 'Platinum' Inclusions in Egyptian Goldwork". The Journal of Egyptian Archaeology. SAGE Publications. 62 (1): 138-144. doi:10.1177/030751337606200116. ISSN 0307-5133. S2CID 192364303.

25. ^ David A. Scott and Warwick Bray (1980). "Ancient Platinum Technology in South America: Its use by the Indians in Pre-Hispanic Times". Platinum Metals Review.

26. ^ Bergsøe, Paul (1936). "Metallurgy of Gold and Platinum among the Pre-Columbian Indians". Nature. Springer Science and Business Media LLC. 137 (3453): 29. doi:10.1038/137029a0. ISSN 0028-0836. S2CID 4100269.

27. ^ Watson, Wm; Brownrigg, William (1749). "Several Papers concerning a New Semi-Metal, Called Platina; Communicated to the Royal Society by Mr. Wm. Watson F. R. S" (PDF). Philosophical Transactions. 46 (491-496): 584-596. Bibcode:1749RSPT...46..584W. doi:10.1098/rstl.1749.0110. S2CID 186213277. Archived from the original (PDF) on 21 October 2013.

28. ^ Hunt, L. B.; Lever, F. M. (1969). "Platinum Metals: A Survey of Productive Resources to industrial Uses" (PDF). Platinum Metals Review. 13 (4): 126-138. Archived (PDF) from the original on 29 October 2008.

29. ^ Kauffman, George B.; Teter, Larry A. & Rhoda, Richard N. (1963). Recovery of Platinum from Laboratory Residues. Inorg. Synth. Inorganic Syntheses. 7. pp. 232-236. doi:10.1002/9780470132388.ch61. ISBN 978-0-470-13238-8.

30. ^ "Guide tothe Realization of the ITS-90 - Platinum Resistance Thermometry" (PDF). International Committee for Weights and Measures.

31. ^ "Standard Reference Material 1750:Standard Platinum Resistance Thermometers,13.8033 K to 429.7485 K" (PDF). NIST.

32. ^ "Platinum versus Gold". The Speculative Invertor. 14 April 2002. Archived from the original on 26 October 2008.

33. ^ Apps, Michael G; Choi, Eugene H Y; Wheate, Nial J (August 2015). "The state-of-play and future of platinum drugs". Endocrine-Related Cancer. Society for Endocrinology. 22 (4): R219-R233. doi:10.1530/ERC-15-0237. PMID 26113607.

34. ^ "CDC - NIOSH Pocket Guide to Chemical Hazards - Platinum". www.cdc.gov. Archived from the original on 21 November 2015.

35. ^ "187 Fake Cancer 'Cures' Consumers Should Avoid". U.S. Food and Drug Administration. Archived from the original on 2 May 2017.

# data[].attributes.attachment_link.description
Pronunciation ​(PLAT-in-əm ) Appearance silvery white Standard atomic weight A r, std(Pt) 195.084(9) [1] (Atomic number Z) 78 Group group 10 Period period 6 Block d-block Electron configuration [ Xe] 4f 14 5d 9 6s 1 Electrons per shell 2, 8, 18, 32, 17, 1 at Phase STP solid Melting point 2041.4 K ​(1768.3 °C, ​3214.9 °F) (near Density r.t.)

# data[].attributes.attachment_link.title
Electrolysis of water experiment using pencils, h2o electrolysis, electrolysis water

# data[].attributes.attachment_link.title
electrolysis of water acidified dilute sulfuric acid sulphuric acid products electrode equations anode cathode apparatus diagram cell electrolyte sodium hydroxide gcse chemistry KS4 science igcse O level revision notes

# data[].attributes.attachment_link.description
introduction to electrolysis explaining which liquids conduct and why can some liquids conduct electricity and others cannot? - electrolytes, electrodes, what happens on electrode surfaces, the products of electrolysis. There are tables of electrode reactions, descriptions of experimental methods of electrolysis and a summary table of the electrolysis products from many common melts or aqueous solutions that undergo electrolysis when a d.c.

# data[].attributes.attachment_link.readable_text
In galvanic cells, chemical energy is converted into electrical energy. The opposite is true for electrolytic cells. In electrolytic cells, electrical energy causes nonspontaneous reactions to occur in a process known as electrolysis. The charging electric car pictured at the beginning of this module shows one such process. Electrical energy is converted into the chemical energy in the battery as it is charged. Once charged, the battery can be used to power the automobile.

The same principles are involved in electrolytic cells as in galvanic cells. We will look at three electrolytic cells and the quantitative aspects of electrolysis.

The Electrolysis of Molten Sodium Chloride

In molten sodium chloride, the ions are free to migrate to the electrodes of an electrolytic cell. A simplified diagram of the cell commercially used to produce sodium metal and chlorine gas is shown in Figure 1. Sodium is a strong reducing agent and chlorine is used to purify water, and is used in antiseptics and in paper production. The reactions are

The power supply (battery) must supply a minimum of 4 V, but, in practice, the applied voltages are typically higher because of inefficiencies in the process itself.

The Electrolysis of Water

It is possible to split water into hydrogen and oxygen gas by electrolysis. Acids are typically added to increase the concentration of hydrogen ion in solution (Figure 2). The reactions are

Note that the sulfuric acid is not consumed and that the volume of hydrogen gas produced is twice the volume of oxygen gas produced. The minimum applied voltage is 1.229 V.

The Electrolysis of Aqueous Sodium Chloride

The electrolysis of aqueous sodium chloride is the more common example of electrolysis because more than one species can be oxidized and reduced. Considering the anode first, the possible reactions are

These values suggest that water should be oxidized at the anode because a smaller potential would be needed-using reaction (ii) for the oxidation would give a less-negative cell potential. When the experiment is run, it turns out chlorine, not oxygen, is produced at the anode. The unexpected process is so common in electrochemistry that it has been given the name overpotential. The overpotential is the difference between the theoretical cell voltage and the actual voltage that is necessary to cause electrolysis. It turns out that the overpotential for oxygen is rather high and effectively makes the reduction potential more positive. As a result, under normal conditions, chlorine gas is what actually forms at the anode.

Now consider the cathode. Three reductions could occur:

Reaction (v) is ruled out because it has such a negative reduction potential. Under standard state conditions, reaction (iii) would be preferred to reaction (iv). However, the pH of a sodium chloride solution is 7, so the concentration of hydrogen ions is only 1× 10−7 M. At such low concentrations, reaction (iii) is unlikely and reaction (iv) occurs. The overall reaction is then

As the reaction proceeds, hydroxide ions replace chloride ions in solution. Thus, sodium hydroxide can be obtained by evaporating the water after the electrolysis is complete. Sodium hydroxide is valuable in its own right and is used for things like oven cleaner, drain opener, and in the production of paper, fabrics, and soap.

Electroplating

An important use for electrolytic cells is in electroplating. Electroplating results in a thin coating of one metal on top of a conducting surface. Reasons for electroplating include making the object more corrosion resistant, strengthening the surface, producing a more attractive finish, or for purifying metal. The metals commonly used in electroplating include cadmium, chromium, copper, gold, nickel, silver, and tin. Common consumer products include silver-plated or gold-plated tableware, chrome-plated automobile parts, and jewelry. We can get an idea of how this works by investigating how silver-plated tableware is produced (Figure 3).

In the figure, the anode consists of a silver electrode, shown on the left. The cathode is located on the right and is the spoon, which is made from inexpensive metal. Both electrodes are immersed in a solution of silver nitrate. As the potential is increased, current flows. Silver metal is lost at the anode as it goes into solution.

The mass of the cathode increases as silver ions from the solution are deposited onto the spoon

The net result is the transfer of silver metal from the anode to the cathode. The quality of the object is usually determined by the thickness of the deposited silver and the rate of deposition.

Quantitative Aspects of Electrolysis

The amount of current that is allowed to flow in an electrolytic cell is related to the number of moles of electrons. The number of moles of electrons can be related to the reactants and products using stoichiometry. Recall that the SI unit for current ( I) is the ampere (A), which is the equivalent of 1 coulomb per second ([latex]1\text{ A}=1\text{ }\frac{\text{C}}{\text{s}}[/latex]). The total charge ( Q, in coulombs) is given by

Where t is the time in seconds, n the number of moles of electrons, and F is the Faraday constant.

Moles of electrons can be used in stoichiometry problems. The time required to deposit a specified amount of metal might also be requested, as in the second of the following examples.

Example 1

Converting Current to Moles of Electrons

In one process used for electroplating silver, a current of 10.23 A was passed through an electrolytic cell for exactly 1 hour. How many moles of electrons passed through the cell? What mass of silver was deposited at the cathode from the silver nitrate solution?

Solution

Faraday's constant can be used to convert the charge ( Q) into moles of electrons ( n). The charge is the current ( I) multiplied by the time

[latex]n=\frac{Q}{F}=\frac{\frac{10.23\text{ C}}{\text{s}}\times\text{1 hr}\times\frac{60\text{ min}}{\text{hr}}\times\frac{60\text{ s}}{\text{min}}}{\text{96,485 C/mol e}^-}=\frac{\text{36,830 C}}{\text{96,485 C/mol e}^-}=0.3817\text{ mol e}^-[/latex]

From the problem, the solution contains AgNO ₃, so the reaction at the cathode involves 1 mole of electrons for each mole of silver

The atomic mass of silver is 107.9 g/mol, so

mass Ag = [latex]0.3817\text{ mol e}^-\times\frac{1\text{ mol Ag}}{1\text{ mol e}^-}\times\frac{107.9\text{g Ag}}{1\text{ mol Ag}}=41.19\text{g Ag}[/latex]

Check your answer: From the stoichiometry, 1 mole of electrons would produce 1 mole of silver. Less than one-half a mole of electrons was involved and less than one-half a mole of silver was produced.

Check Your Learning

Aluminum metal can be made from aluminum ions by electrolysis. What is the half-reaction at the cathode? What mass of aluminum metal would be recovered if a current of 2.50 × 10 ³ A passed through the solution for 15.0 minutes? Assume the yield is 100%.

mol Al = 210.0 g Al.

Example 2

Time Required for Deposition

In one application, a 0.010-mm layer of chromium must be deposited on a part with a total surface area of 3.3 m ² from a solution of containing chromium(III) ions. How long would it take to deposit the layer of chromium if the current was 33.46 A? The density of chromium (metal) is 7.19 g/cm ³.

Solution

This problem brings in a number of topics covered earlier. An outline of what needs to be done is:

If the total charge can be determined, the time required is just the charge divided by the current

The total charge can be obtained from the amount of Cr needed and the stoichiometry

The amount of Cr can be obtained using the density and the volume Cr required

The volume Cr required is the thickness times the area

Solving in steps, and taking care with the units, the volume of Cr required is

[latex]\text{volume}=\left(\text{0.010 mm}\times\frac{\text{1 cm}}{\text{10 mm}}\right)\times\left(\text{3.3 m}^{2}\times\left(\frac{\text{10,000 cm}^{2}}{\text{1 m}^2}\right)\right)=\text{33 cm}^3[/latex]

Cubic centimeters were used because they match the volume unit used for the density. The amount of Cr is then

mass = volume × density = [latex]33\cancel{\text{ cm}^3}\times\frac{7.19\text{ g}}{\cancel{\text{cm}^3}}=237\text{g Cr}[/latex]
mol Cr =[latex]237\text{g Cr}\times\frac{1\text{ mol Cr}}{52.00\text{g Cr}}=4.56\text{ mol Cr}[/latex]

Since the solution contains chromium(III) ions, 3 moles of electrons are required per mole of Cr. The total charge is then

[latex]Q=4.56\text{ mol Cr}\times\frac{3\text{ mol e}^-}{1\text{ mol Cr}}\times\frac{96485\text{ C}}{\text{mol e}^-}=1.32\times{10^{6}}\text{C}[/latex]

The time required is then

[latex]t=\frac{Q}{I}=\frac{1.32\times10^6\text{ C}}{33.46\text{ C/s}}=3.95\times10^4\text{ s}=11.0\text{hr}[/latex]

Check your answer: In a long problem like this, a single check is probably not enough. Each of the steps gives a reasonable number, so things are probably correct. Pay careful attention to unit conversions and the stoichiometry.

Check Your Learning

What mass of zinc is required to galvanize the top of a 3.00 m × 5.50 m sheet of iron to a thickness of 0.100 mm of zinc? If the zinc comes from a solution of Zn(NO ₃) ₂ and the current is 25.5 A, how long will it take to galvanize the top of the iron? The density of zinc is 7.140 g/cm ³.

Answer: 231 g Zn required 446 minutes.

Key Concepts and Summary

Using electricity to force a nonspontaneous process to occur is electrolysis. Electrolytic cells are electrochemical cells with negative cell potentials (meaning a positive Gibbs free energy), and so are nonspontaneous. Electrolysis can occur in electrolytic cells by introducing a power supply, which supplies the energy to force the electrons to flow in the nonspontaneous direction. Electrolysis is done in solutions, which contain enough ions so current can flow. If the solution contains only one material, like the electrolysis of molten sodium chloride, it is a simple matter to determine what is oxidized and what is reduced. In more complicated systems, like the electrolysis of aqueous sodium chloride, more than one species can be oxidized or reduced and the standard reduction potentials are used to determine the most likely oxidation (the half-reaction with the largest [most positive] standard reduction potential) and reduction (the half-reaction with the smallest [least positive] standard reduction potential). Sometimes unexpected half-reactions occur because of overpotential. Overpotential is the difference between the theoretical half-reaction reduction potential and the actual voltage required. When present, the applied potential must be increased, making it possible for a different reaction to occur in the electrolytic cell. The total charge, Q, that passes through an electrolytic cell can be expressed as the current ( I) multiplied by time ( Q = It) or as the moles of electrons ( n) multiplied by Faraday's constant ( Q = nF). These relationships can be used to determine things like the amount of material used or generated during electrolysis, how long the reaction must proceed, or what value of the current is required.

Chemistry End of Chapter Exercises

1. Identify the reaction at the anode, reaction at the cathode, the overall reaction, and the approximate potential required for the electrolysis of the following molten salts. Assume standard states and that the standard reduction potentials in Standard Electrode (Half-Cell) Potentials are the same as those at each of the melting points. Assume the efficiency is 100%.

2. What mass of each product is produced in each of the electrolytic cells of the previous problem if a total charge of 3.33 × 105 C passes through each cell? Assume the voltage is sufficient to perform the reduction.

3. How long would it take to reduce 1 mole of each of the following ions using the current indicated? Assume the voltage is sufficient to perform the reduction.

4. A current of 2.345 A passes through the cell shown in Figure for 45 minutes. What is the volume of the hydrogen collected at room temperature if the pressure is exactly 1 atm? Assume the voltage is sufficient to perform the reduction. (Hint: Is hydrogen the only gas present above the water?)

5. An irregularly shaped metal part made from a particular alloy was galvanized with zinc using a Zn(NO₃)₂ solution. When a current of 2.599 A was used, it took exactly 1 hour to deposit a 0.01123-mm layer of zinc on the part. What was the total surface area of the part? The density of zinc is 7.140 g/cm³. Assume the efficiency is 100%.

Selected Answers

4. 0.79 L

Glossary

electrolysis
process using electrical energy to cause a nonspontaneous process to occur

electrolytic cell
electrochemical cell in which electrolysis is used; electrochemical cell with negative cell potentials

electroplating
depositing a thin layer of one metal on top of a conducting surface

overpotential
difference between the theoretical potential and actual potential in an electrolytic cell; the "extra" voltage required to make some nonspontaneous electrochemical reaction to occur

# data[].attributes.attachment_link.description
In galvanic cells, chemical energy is converted into electrical energy. The opposite is true for electrolytic cells. In electrolytic cells, electrical energy causes nonspontaneous reactions to occur in a process known as electrolysis. The charging electric car pictured at the beginning of this module shows one such process.