omniomni/omni-engineering huggingface dataset:
Mercury(II) iodide. Mercury(II) iodide is a chemical compound with the molecular formula HgI2. It is typically produced synthetically but can also be found in nature as the extremely rare mineral coccinite. Unlike the related mercury(II) chloride it is hardly soluble in water (<100 ppm).

Production
Mercury(II) iodide is produced by adding an aqueous solution of potassium iodide to an aqueous solution of mercury(II) chloride with stirring; the precipitate is filtered off, washed and dried at 70 °C.

 HgCl2 + 2 KI → HgI2 + 2 KCl

Properties
Mercury(II) iodide displays thermochromism; when heated above 126 °C (400 K) it undergoes a phase transition, from the red alpha crystalline form to a pale yellow beta form. As the sample cools, it gradually reacquires its original colour. It has often been used for thermochromism demonstrations. A third form, which is orange, is also known; this can be formed by recrystallisation and is also metastable, eventually converting back to the red alpha form. The various forms can exist in a diverse range of crystal structures and as a result mercury(II) iodide possesses a surprisingly complex phase diagram.

Uses

Mercury(II) iodide is used for preparation of Nessler's reagent, used for detection of presence of ammonia.

Mercury(II) iodide is a semiconductor material, used in some x-ray and gamma ray detection and imaging devices operating at room temperatures.

In veterinary medicine, mercury(II) iodide is used in blister ointments in exostoses, bursal enlargement, etc. 

It can appear as a precipitate in many reactions.

See also

Mercury(I) iodide, Hg2I2

References

Iodides
Metal halides
Mercury(II) compounds
Semiconductor materials
Mercury(II) sulfate. Mercury(II) sulfate, commonly called mercuric sulfate, is the chemical compound HgSO4. It is an odorless salt that forms white granules or crystalline powder. In water, it separates into an insoluble basic sulfate with a yellow color and sulfuric acid.

Structure

The anhydrous compound features Hg2+ in a highly distorted tetrahedral HgO4 environment.  Two Hg-O distances are 2.22 Å and the others are 2.28 and 2.42 Å.  In the monohydrate,  Hg2+ adopts a linear coordination geometry with Hg-O (sulfate) and Hg-O (water) bond lengths of 2.179 and 2.228 Å, respectively.  Four weaker bonds are also observed with Hg---O distances >2.5 Å.

History
In 1932, the Japanese chemical company Chisso Corporation began using mercury sulfate as the catalyst for the production of acetaldehyde from acetylene and water. Though it was unknown at the time, methylmercury is formed as side product of this reaction. Exposure and consumption of the mercury waste products, including methylmercury, that were dumped into Minamata Bay by Chisso are believed to be the cause of Minamata disease in Minamata, Japan.

Production
Mercury sulfate can be produced
by treating mercury with hot concentrated sulfuric acid:

Alternatively yellow mercuric oxide reacts also with concentrated sulfuric acid.

Uses

Denigés' reagent
An acidic solution of mercury sulfate is known as Denigés' reagent. It was commonly used throughout the 20th century as a qualitative analysis reagent. If Denigés' reagent is added to a solution containing compounds that have tertiary alcohols, a yellow or red precipitate will form.

Hydration reactions 

Mercury sulfate, as well as other mercury(II) compounds, are commonly used as catalysts in oxymercuration-demercuration, a type of electrophilic addition reaction that results in hydration of an unsaturated compound. The hydration of an alkene gives an alcohol. The regioselectivity is that predicted by Markovnikov's rule. For an alkyne, the result is an enol, which tautomerizes to give the carbonyl. At one time, this chemistry was employed commercially for the preparation of acetaldehyde from acetylene:

A related and specialized example is the conversion of 2,5-dimethylhexyne-2,5-diol to 2,2,5,5-tetramethyltetrahydrofuran using aqueous mercury sulfate without the addition of acid.

Health issues
Inhalation of HgSO4 can result in acute poisoning: causing tightness in the chest, difficulties breathing, coughing and pain. Exposure of HgSO4 to the eyes can cause ulceration of conjunctiva and cornea. If mercury sulfate is exposed to the skin it may cause sensitization dermatitis. Lastly, ingestion of mercury sulfate will cause necrosis, pain, vomiting, and severe purging. Ingestion can result in death within a few hours due to peripheral vascular collapse.

It was used in the late 19th century to induce vomiting for medical reasons.

Further reading

References

External links

 National Pollutant Inventory – Mercury and compounds Fact Sheet
 NIOSH Pocket Guide to Chemical Hazards

Sulfates
Mercury(II) compounds
Mercury(II) acetate. Mercury(II) acetate, also known as mercuric acetate is a chemical compound, the mercury(II) salt of acetic acid, with the formula Hg(O2CCH3)2. Commonly abbreviated Hg(OAc)2, this compound is employed as a reagent to generate organomercury compounds from unsaturated organic precursors.  It is a white, water-soluble solid, but some samples can appear yellowish with time owing to decomposition.

Structure
Mercury(II) acetate is a crystalline solid consisting of isolated Hg(OAc)2 molecules with Hg-O distances of 2.07 Å. Three long, weak intermolecular Hg···O bonds of about 2.75 Å are also present, resulting in a slightly distorted square pyramidal coordination geometry at Hg.

Synthesis and reactions
Mercury(II) acetate can be produced by reaction of mercuric oxide with acetic acid.

HgO + 2 CH3COOH → Hg(CH3COO)2 + H2O

Inorganic reactions
Mercury(II) acetate in acetic acid solution reacts with H2S to rapidly precipitate the black (β) polymorph of HgS.  With gentle heating of the slurry, the black solid converts to the red form.  The mineral cinnabar is red HgS.  The precipitation of HgS as well as a few other sulfides, using hydrogen sulfide is a step in qualitative inorganic analysis.

Organic chemistry
Electron-rich arenes undergo "mercuration" upon treatment with Hg(OAc)2.  This behavior is illustrated with phenol:
C6H5OH + Hg(OAc)2 → C6H4(OH)-2-HgOAc + HOAc
The  acetate group (OAc) that remains on mercury can be displaced by chloride:
C6H4(OH)-2-HgOAc + NaCl → C6H4(OH)-2-HgCl + NaOAc

The Hg2+ center binds to alkenes, inducing the addition of hydroxide and alkoxide.  For example, treatment of methyl acrylate with mercuric acetate in methanol gives an α-mercuri ester:
Hg(OAc)2 + CH2=CHCO2CH3 + CH3OH  →  CH3OCH2CH(HgOAc)CO2CH3 + HOAc

Exploiting the high affinity of mercury(II) for sulfur ligands, Hg(OAc)2 can be used as a reagent to deprotect thiol groups in organic synthesis. Similarly Hg(OAc)2 has been used to convert thiocarbonate esters into dithiocarbonates:
(RS)2C=S + H2O + Hg(OAc)2 → (RS)2C=O + HgS + 2 HOAc

Mercury(II) acetate is used for oxymercuration reactions.

A famous use of Hg(OAc)2 was in the synthesis of idoxuridine.

Toxicity 

Mercuric acetate is a highly toxic compound, due to it being water-soluble and having mercury ions.
Symptoms of mercury poisoning include peripheral neuropathy, skin discoloration and desquamation (peeling and/or shedding of the skin). Chronic exposure may cause reduced intelligence and kidney failure.

References

Acetates
Mercury(II) compounds
Potassium acetate. Potassium acetate (also called potassium ethanoate), (CH3COOK) is the potassium salt of acetic acid. It is a hygroscopic solid at room temperature.

Preparation
It can be prepared by treating a potassium-containing base such as potassium hydroxide or potassium carbonate with acetic acid:

CH3COOH + KOH → CH3COOK + H2O
This sort of reaction is known as an acid-base neutralization reaction.

At saturation, the sesquihydrate in water solution (CH3COOK·1½H2O) begins to form semihydrate at 41.3 °C.

Applications

Deicing
Potassium acetate (as a substitute for calcium chloride or magnesium chloride) can be used as a deicer to remove ice or prevent its formation. It offers the advantage of being less aggressive on soils and much less corrosive: for this reason, it is preferred for airport runways although it is more expensive.

Fire extinguishing
Potassium acetate is the extinguishing agent used in Class K fire extinguishers because of its ability to cool and form a crust over burning oils.

Food additive
Potassium acetate is used in processed foods as a preservative and acidity regulator.  In the European Union, it is labeled by the E number E261; it is also approved for usage in the USA, Australia, and New Zealand. Potassium hydrogen diacetate (CAS #) with formula KH(OOCCH3)2 is a related food additive with the same E number as potassium acetate.

Medicine and biochemistry
In medicine, potassium acetate is used as part of electrolyte replacement protocols in the treatment of diabetic ketoacidosis because of its ability to break down to bicarbonate to help neutralize the acidotic state.

In molecular biology, potassium acetate is used to precipitate Sodium dodecyl sulfate (SDS) and SDS-bound proteins to allow their removal from DNA. 

Potassium acetate is used in mixtures applied for tissue preservation, fixation, and mummification. Most museums today use a formaldehyde-based method recommended by Kaiserling in 1897 which contains potassium acetate. This process was used to soak Lenin's corpse.

Use in executions

Potassium acetate was incorrectly used in place of potassium chloride when putting a prisoner to death in Oklahoma in January 2015. Charles Frederick Warner was executed on January 15, 2015 with potassium acetate; this was not public knowledge until the scheduled execution of Richard Glossip was called off. In August 2017, the U.S. state of Florida executed Mark James Asay using a combination of etomidate, rocuronium bromide, and potassium acetate. The drug was also used in the February 2023 execution of Donald Dillbeck, once again in combination with etomidate and rocuronium bromide.

Industry
Potassium acetate is used as a catalyst in the production of polyurethanes.

Historical
It is used as a diuretic and urinary alkalizer. Before modern chemistry, it was variously called terra foliata tartari, sal Sennerti, tartarus regeneratus, arcanum tartari and sal diureticus. In 1760 it was used in the preparation of Cadet's fuming liquid ((CH3)2As)2O, the first organometallic compound ever produced.

References

External links

Acetates
Potassium compounds
Lethal injection components
E-number additives
Lithium metaborate. Lithium metaborate is a chemical compound of lithium, boron, and oxygen with elemental formula .  It is often encountered as a hydrate, , where n is usually 2 or 4.  However, these formulas do not describe the actual structure of the solids.

Lithium metaborate is one of the borates, a large family of salts (ionic compounds) with anions consisting of boron, oxygen, and hydrogen.

Structure

Lithium metaborate has several crystal forms.

The α form consists of infinite chains of trigonal planar metaborate anions .

The γ form is stable at 15 kbar and 950 °C. It has a polymeric cation consisting of a tridimensional regular array of  tetrahedra sharing oxygen vertices, alernating with lithium cations, each also surrounded by four oxygen atoms. The B-O distances are 148.3 pm, the Li-O distances are 196 pm.

Lithium metaborate forms glass relatively easily, and consists of approximately 40% tetrahedral borate anions, and 60% trigonal planar boron. The ratio of tetrahedral to trigonal boron has been shown to be strongly temperature dependent in the liquid and supercooled liquid state.

Applications

Laboratory

Molten lithium metaborate, often mixed with lithium tetraborate , is used to dissolve oxide samples for analysis by XRF, AAS, ICP-OES, ICP-AES, and ICP-MS, modern versions of classical bead test. The process may be used also to facilitate the dissolution of oxides in acids for wet analysis. Small amounts of lithium bromide  or lithium iodide  may be added as mold and crucible release agents.

Lithium metaborate dissolves acidic oxides  with x < y, such as SiO2, , , , , , , , and Fe2O3.  Lithium tetraborate, on the other hand, dissolves basic oxides with x > y, such as CaO, MgO and other oxides of the alkali metals and alkaline earth metals.  Most oxides are best dissolved in a mixture of the two lithium borate salts, for spectrochemical analysis.

References 

Borates
Lithium salts
FuturePlay. The International Academic Conference on the Future of Game Design and Technology, also known as FuturePlay, is an international academic conference that focuses on the future of video game design and technology. The most recent conference under that title was held in 2010.

Overview
Future Play evolved from the former Computer Game Technology (CGT) Conference that had been initiated and run by Algoma University since 2002. The conference was held annually in Toronto, Ontario, Canada. Future Play 2008 was hosted by Algoma and UOIT, each year bringing together key players from academics, industry, and government to advance game design and technology.

The conference focused on three themes: Future game development, future game impacts and application, future game talent. These issues were addressed through keynotes from leaders in academia and industry, peer-reviewed paper sessions, panel sessions, workshops, and exhibitions of posters, games, and the latest game technologies and supports from industry vendors.

Academic competitions

Games competition
The Future Play Game Exhibition and Competition allows academic, independent, experimental, and/or student developers to gain recognition for their game-related work in a juried competition. Developers can submit to the exhibition in the following categories:
 Future Games: this category is for games that experiment with new game technologies or design concepts.
 Future Game Impacts and Applications: this category is for games that are built not necessarily for pure entertainment purposes (i.e. serious games), such as games that are for learning purposes.
 Future Game Talent: this category is primarily for students who have built an original, entertainment game title who wish to show off their talents, earn publicity for their school, gain personal recognition, and perhaps even find a publisher for the game.
Games competition winners are generally recognized in the trade press.

Academic papers
Paper submissions for the conference were original, unpublished research. Papers often included late-breaking advances and work-in-progress reports from ongoing research. All submissions were peer-reviewed and evaluated based on originality, technical and/or research content/depth, correctness, relevance to conference, contributions and readability.

2008 program
Future Play 2008 was hosted by Algoma University and the University of Ontario Institute of Technology. It was hosted in Toronto, Ontario, Canada. Speakers for the conference were Ken Perlin and Katie Salen. The theme of the conference was "Research. Play. Share."

References

External links
 FuturePlay website 
 FuturePlay blog

Recurring events established in 2002
International conferences in Canada
Video game conferences
2002 establishments in Ontario
Sodium persulfate. Sodium persulfate is the inorganic compound with the formula Na2S2O8.  It is the sodium salt of peroxydisulfuric acid, H2S2O8, an oxidizing agent. It is a white solid that dissolves in water.  It is almost non-hygroscopic and has good shelf-life.

Production
The salt is prepared by the electrolytic oxidation of sodium bisulfate:

Oxidation is conducted at a platinum anode. In this way about 165,000 tons were produced in 2005. 

The standard redox potential of sodium persulfate into hydrogen sulfate is 2.1 V, which is higher than that of hydrogen peroxide (1.8 V) but lower than ozone (2.2 V). The sulfate radical formed in situ has a standard electrode potential of 2.7 V.

However, there are a few drawbacks in utilizing platinum anodes to produce the salts; the manufacturing process is inefficient due to oxygen evolution and the product could contain contaminants coming from platinum corrosion (mainly due to extremely oxidizing nature of the sulfate radical). Thus, boron-doped diamond electrodes have been proposed as alternatives to the conventional platinum electrodes.

Structure
The sodium and potassium salts adopt very similar structures in the solid state, according to X-ray crystallography. In the sodium salt, the O-O distance is 1.476Å. The sulfate groups are tetrahedral, with three short S-O distances near 1.44 and one long S-O bond at 1.64Å.

Applications
It is mainly used as a  radical initiator for emulsion polymerization reactions for styrene based polymers such as Acrylonitrile butadiene styrene.  Also applicable for accelerated curing of low formaldehyde adhesives.

Other uses
It is a bleach, both standalone (particularly in hair cosmetics) and as a detergent component. It is a replacement for ammonium persulfate in etching mixtures for zinc and printed circuit boards, and is used for pickling of copper and some other metals.

It is also used as a soil conditioner and for soil and groundwater remediation and in manufacture of dyestuffs, modification of starch, bleach activator, desizing agent for oxidative desizing, etc.

Organic chemistry
Sodium persulfate is a specialized oxidizing agent in chemistry, classically in the Elbs persulfate oxidation and the Boyland–Sims oxidation reactions. It is also used in radical reactions; for example in a synthesis of diapocynin from apocynin where iron(II) sulfate is the radical initiator.

Safety
The salt is an oxidizer and forms combustible mixtures with organic materials such as paper.

References

Persulfates
Sodium compounds
Oxidizing agents
Joe Celko. Joe Celko is an American relational database expert from Austin, Texas. He has participated on the ANSI X3H2 Database Standards Committee, and helped write the SQL-89 and SQL-92 standards. He is the author of a Morgan-Kaufmann series of books on SQL, and over 1200 published articles on SQL and other database topics. He had been a full-time statistician for several years.

He is credited with coining the term lasagna code and popularizing Michael J. Kamfonas' nested set model for trees in SQL, a taxonomy of data encoding schemes, and several other design patterns in SQL DDL and DML.

Early life and education 
Celko achieved his undergraduate and first master's degrees in math at Georgia State University. He also has a second master's degree in computer science from Georgia Tech.

Career 
Celko has a very wide range of industries and applications starting in 1965 with a National Science Fair prize job at the Pittman-Dunn Research labs at the Frankford Arsenal. His public sector consulting work has been for prison systems, medical organizations, NASA and defense contractors. He helped set up programming and Software Engineering standards for the US Army at AIRMICS and wrote regular columns on it in the trade press on Software Engineering.

His commercial consulting work has been varied—auto loans, entertainment, aerospace, education, shipping (freight lines and ocean shipping), petro-chemical, software companies, and odd start-ups.

He has taught training classes in South Africa, Turkey, the UK, Norway, Sweden, Finland, The Netherlands, Brazil, Belgium and Switzerland. He has also created and taught on-line classes for MySQL and SQLUniversity.org. He was part of the founding faculty of Neumont University in Salt Lake City, a for-profit university for software developers.

Bibliography 

 Joe Celko's Analytics and OLAP in SQL (Morgan-Kaufmann, 2006) 
 Joe Celko's Data and Databases (Morgan Kaufmann, 1999) 
 Joe Celko's SQL for Smarties (5th Edition, Morgan-Kaufmann, 2014) 
 Joe Celko's SQL Programming Style (Morgan-Kaufmann, 2005) 
 Joe Celko's SQL Puzzles and Answers (2nd edition, Morgan-Kaufmann 2006) 
 Joe Celko's Trees and Hierarchies in SQL for Smarties, 2nd Edition (Morgan-Kaufmann, 2012)  
 Joe Celko’s Thinking in Sets: Auxiliary, Temporal, and Virtual Tables in SQL. (Morgan-Kaufmann, 2008) 
 Joe Celko’s Data, Measurements and Standards in SQL (Morgan-Kaufmann, 2010) 
 Joe Celko's Complete Guide to NoSQL (Morgan-Kaufmann, 2014)

References

External links 

 Personal Bio

Technical writers
American software engineers
Living people
Georgia State University alumni
Georgia Tech alumni
Year of birth missing (living people)
Lead(II) chromate. Lead(II) chromate is an inorganic compound with the chemical formula . It is a bright yellow solid that is very poorly soluble in water.  It occurs also as the mineral crocoite.  It is used as a pigment.

Structure

Two polymorphs of lead chromate are known, orthorhombic and the more stable monoclinic form. Monoclinic lead chromate is used in paints under the name chrome yellow, and many other names. Lead chromate adopts the monazite structure, meaning that the connectivity of the atoms is very similar to other compounds of the type . Pb(II) has a distorted coordination sphere being surrounded by eight oxides with Pb-O distances ranging from 2.53 to 2.80 Å. The chromate anion is tetrahedral, as usual.  Unstable polymorphs of lead chromate are the greenish yellow orthorhombic form and a red-orange tetragonal form.

Applications 

Approximately 37,000 tons were produced in 1996. The main applications are as a pigment in paints, under the name chrome yellow.

Preparation
Lead(II) chromate can be produced by treating sodium chromate with lead salts such as lead(II) nitrate or by combining lead(II) oxide with chromic acid.

Related lead sulfochromate pigments are produced by the replacement of some chromate by sulfate, resulting in a mixed lead-chromate-sulfate compositions . This replacement is possible because sulfate and chromate are isostructural. Since sulfate is colorless, sulfochromates with high values of x are less intensely colored than lead chromate. In some cases, chromate is replaced by molybdate.

Reactions 
Heating in hydroxide solution produces chrome red, a red or orange powder made by PbO and . Also, in hydroxide solution lead chromate slowly dissolves forming plumbite complex.

Safety hazards 
Despite containing both lead and hexavalent chromium, lead chromate is not acutely lethal because of its very low solubility. The LD50 for rats is only 5,000 mg/kg. Lead chromate must be treated with great care in its manufacture, the main concerns being dust of the chromate precursor.  Lead chromate is highly regulated in advanced countries. As one of the greatest threats comes from inhalation of particles, so much effort has been devoted to production of low-dust forms of the pigment.

In the 1800s, the product was used to impart a bright yellow color to some types of candy. It is used (illegally) to enhance the color of certain spices, particularly turmeric, particularly in Bangladesh.

Unlike other lead-based paint pigments, lead chromate is still widely used, especially in road marking paint.

In 2023 and 2024, consumption of adulterated cinnamon led to at least 136 cases of lead toxicity in children in the United States as reported by the US Centers for Disease Control and Prevention. The affected products were recalled. The US Food and Drug Administration determined that the ratio of lead to chromium in the cinammon indicated that lead chromate had been added to the cinnamon.

See also
Crocoite
Iranite

References 
Safety (MSDS) data for lead chromate
Chemical Profile for lead chromate
J.T. Baker MSDS

Chromates
Lead(II) compounds
Pyrotechnic oxidizers
Environmental impact of paint
Oxidizing agents