月: 2023年1月

What Is Rhodium?

Rhodium, with atomic number 45 and symbol Rh, is derived from the Greek “rhodeos,” meaning rose-colored, due to the rose hue of its salts. Classified as a “minor metal,” it is extremely rare in nature, ranking as the third-rarest stable isotope element with an abundance of 200 ppt in the Earth’s crust. It is valued in decorative, chemical catalytic, and industrial components for its hardness, electrical resistance, and corrosion resistance.

Uses for Rhodium

Used for plating decorative items, rhodium enhances color and strength. Its properties are also beneficial in computer reed switches, reflective mirrors, thermocouples, interference filters, and glass fiber production nozzles. Notably, rhodium serves as a critical three-way catalyst in automobiles, converting toxic nitrogen oxides into nitrogen, and catalyzing the production of oxo alcohols and acetic acid.

Properties of Rhodium

This silvery-white transition metal has a specific gravity of 12.5, a melting point of 1,966°C, and a boiling point of 3,960°C. Soft and ductile, it resists oxidation at room temperature but forms rhodium(III) oxide under intense heat. Rhodium withstands strong acids but is susceptible to oxidation by strong oxidizing agents at high temperatures.

Structure of Rhodium

At room temperature, rhodium exhibits a face-centered cubic structure, transitioning to a simple cubic lattice above 1,000°C. It can take oxidation states from -1 to +6 under high temperatures.

Other Information on Rhodium

1. Production of Rhodium

Discovered by William Hyde Wollaston in platinum ore, rhodium continues to be sourced as a platinum ore impurity.

2. Isotopes of Rhodium

The stable isotope ¹⁰³Rh is the most prevalent, with radioactive ¹⁰¹Rh having a half-life of 3.3 years. Other notable isotopes include ¹⁰²Rh and ^102mRh with significant half-lives, alongside a variety of radioisotopes and nuclear isomers like ^101mRh.

3. Isotopic Decay of Rhodium

Isotopes lighter than ¹⁰³Rh decay to ruthenium, while those heavier decay to palladium, showcasing the diverse radioactive pathways of rhodium isotopes.

What Is Resorcinol?

Resorcinol is a colorless crystalline powder, an organic compound with the formula C₆H₆O₂, a molecular weight of 110.11, and CAS number 108-46-3. It features two hydroxy groups attached at the meta positions of the benzene ring, classifying it as a benzenediol and a structural isomer of catechol and hydroquinone. Discovered by Austrian chemist Heinrich Frachevitz in 1864, resorcinol was originally called benzene-1,3-diol as recommended by the IUPAC in 1993.

Properties of Resorcinol

Resorcinol melts at 109~111°C and boils at 280°C, with a specific gravity of 1.27, appearing as a colorless solid at room temperature. It is insoluble in chloroform and carbon disulfide but highly soluble in water, alcohol, and ether, and absorbs moisture from the air. Resorcinol can oxidize under light and oxygen, turning pink, necessitating careful storage.

Uses of Resorcinol

Primarily, resorcinol is used in industrial adhesives for tire manufacturing and wood adhesives, offering excellent water, heat, and weather resistance. It also serves as a disinfectant due to its strong reducing power and is a precursor for fluorescent dyes among other applications.

Other Information on Resorcinol

1. Production Process of Resorcinol

Resorcinol synthesis involves dialkylation of propylene to 1,3-diisopropylbenzene, followed by oxidation and Hock rearrangement, yielding resorcinol and acetone. Industrially, it is produced via the oxidation of 1,3-diisopropylbenzene using Cumene‘s autoxidation process.

2. Reaction of Resorcinol

Resorcinol can be hydrogenated to dihydroresorcinol, or 1,3-cyclohexanedione, and reacts with potassium hydroxide to produce phloroglucinol, pyrocatechol, and diresorcinol. Trinitroresorcinol, an explosive, is made through nitration with concentrated nitric acid and sulfuric acid.

3. Regulatory Information

Resorcinol is a notable hazardous substance and is designated as a deleterious substance. It necessitates careful handling.

4. Handling and Storage Precautions

• Store away from heat, sparks, and flames in a cool, dry, well-ventilated place.
• Avoid exposure to strong oxidizers, ammonia, and amino compounds, and do not inhale dust or vapors.
• Wear protective gear including gloves, eye protection, and masks, and ensure proper ventilation during use.
• Follow proper glove removal techniques and wash hands thoroughly after handling.

What Is Rubidium?

Rubidium is an alkali metal with an atomic number of 37 and an atomic weight of 85.4678. Discovered by Bunsen and Kirchhoff in Germany in 1861, it’s relatively abundant in the Earth’s crust, though rarely found in ore form. Rubidium is typically obtained as a byproduct of refining lithium and is designated as a hazardous material for its spontaneous combustibility.

Uses of Rubidium

Rubidium strengthens glass and enhances its electrical insulation properties, primarily mixed with glass as rubidium carbonate (Rb₂CO₃). Its applications include use in cathode ray tubes and atomic clocks, though its commercial and industrial uses are limited. Another application is in dating materials, utilizing the Rubidium87 to Strontium87 ratio to determine the materials’ age from their crystallization time.

Properties of Rubidium

A soft, silvery-white metal, rubidium exhibits a dark red color in flame reactions, similar to potassium. It has a low melting point of 39°C and a boiling point of 688°C, with a gaseous state that appears blue. Sharing properties with other alkali metals, Rubidium has the second lowest electronegativity among nonradioactive alkali elements and a low ionization energy of 406 kJ/mol. In compounds, its valence is +1, and Rubidium ions are readily taken up by plant and animal cells.

Structure of Rubidium

At room temperature, rubidium has a density of 1.532 g/cm³ and adopts a body-centered cubic structure. Its electron configuration is [Kr] 5s¹. Due to its large ionic radius, rubidium falls into the category of incompatible elements, which are not easily incorporated into the crystals of rock-forming minerals.

Other Information About Rubidium

1. Reactions of Rubidium

Rubidium is more reactive than potassium and sodium, capable of spontaneously combusting in air and quickly oxidizing to form peroxide (Rb₂O₂) and superoxide (RbO₂). It reacts violently with halogens and water, producing hydrogen gas that can ignite and explode. Rubidium can also form alloys with potassium, sodium, cesium, calcium, and gold, and amalgamate with mercury.

2. Natural Rubidium

The 23rd most abundant element in the Earth’s crust, rubidium is found in ores such as leucite, carnallite, pollucite, and zinnwaldite, with concentrations up to 1%. Lithia mica contains 0.3 to 3.5% rubidium, making it a commercial source. Seawater averages 125 μg/L of Rubidium, less than potassium but more than cesium. Major sources include rubidium microcline on the island of Elba, Italy, and porphyritic deposits at Burnick Lake, Canada.

3. Production of Rubidium

Annually, 2-4 tons of Rubidium compounds are produced. Separation from potassium allows for the production of pure rubidium alum through fractional crystallization from rubidium cesium alum.

4. Rubidium Isotopes

There are 24 known isotopes of Rubidium, including the stable isotope ⁸⁵Rb and the radioactive isotope ⁸⁷Rb, with natural abundances of 72.2% and 27.8%, respectively. ⁸⁷Rb, with a half-life of 4.88 x 10¹⁰ years, is used for dating as it decays to stable ⁸⁷Sr. ⁸²Rb, produced from ⁸²Sr decay, has a half-life of 1.273 minutes and is used in positron emission tomography of the heart.

What Is Ruthenium?

Ruthenium is a silvery-white, hard but brittle solid that easily transforms into powder. This element belongs to the platinum group, with the symbol Ru, atomic number 44, atomic weight 101.07, and CAS number 7440-18-8. Discovered by the German chemist Osan in 1828 from deposits in the Urals of Russia, its first pure isolation was achieved by the Russian chemist Krauss in 1844. Ruthenium naturally occurs in the earth’s crust alongside other platinum metals like iridium.

Uses of Ruthenium

Ruthenium’s applications are diverse, spanning various fields:

1. Addition to Metals

Alloyed with platinum, palladium, titanium, and molybdenum, ruthenium enhances resistance to corrosion, oxidation, and wear. Ruthenium-palladium alloys find use in jewelry, ornaments, and dentistry, while ruthenium-platinum alloys serve in electrical contacts and decorations.

2. Medical Materials

The beta-decay isotope Ru-106 is utilized in radiotherapy for ocular tumors, especially malignant melanoma of the uvea. Ruthenium complexes are being explored for their potential anti-cancer properties, offering greater hydrolysis resistance and expected tumor selectivity compared to platinum complexes.

3. Catalysts

Ruthenium compounds are valuable as catalysts in various chemical reactions. Homogeneous catalyst applications include olefin metathesis, catalyzed by ruthenium trichloride and Grubbs catalysts, and enantioselective hydrogenation of ketones. Heterogeneous catalyst uses involve Fischer-Tropsch synthesis.

4. Others

Additional uses of ruthenium include cathodic protection of structures and as components of composite metal oxide anodes in chlorine production from salt water via electrolysis.

Properties of Ruthenium

With a melting point of 2,450°C and a boiling point of 3,700°C, Ruthenium is a solid at room temperature, featuring a density of 12.43 g/cm³. It shows remarkable chemical stability, being insoluble in acids, including aqua regia, and only soluble in molten alkalis, forming ruthenates (RuO₄^2-). Ruthenium significantly enhances hardness and strength when alloyed in small quantities with platinum, palladium, or titanium.

Other Information on Ruthenium

1. Ruthenium Manufacturing Process

Ruthenium is obtained commercially as a by-product from nickel, copper, and platinum ore processing. During the electrolytic refining of copper and nickel, it precipitates as anodic mud along with other precious metals, from which ruthenium can be extracted. It is also sourced from spent nuclear fuel as a fission or neutron absorption product.

2. Ruthenium Reactions

Ruthenium can be oxidized to ruthenium dioxide (RuO₂) and further to ruthenium tetroxide (RuO₄), a potent oxidant, using sodium meta-periodate. Ruthenium tetroxide serves mainly as an intermediate in ruthenium purification from ores and radioactive waste.

3. Regulatory Information

Ruthenium is not regulated under any safety laws.

4. Handling and Storage Precautions

Recommended precautions for handling and storage include sealing the container tightly, storing it in a dry, cool, and dark place, using it in well-ventilated areas or outdoors, avoiding contact with strong oxidizers, and wearing protective gear. After handling, thorough hand washing is advised, along with immediate rinsing of skin or eyes upon contact.

What Is Lithium?

Lithium is the smallest alkali metal element with an atomic number of 3 and an atomic weight of 6.941. It is among the softest metals, has a low melting point, and is highly reactive, readily reacting with oxygen, water, and nitrogen in the air and easily oxidizing.

Uses of Lithium

Lithium finds extensive use in the ceramic industry, particularly in glass and ceramics, where its compound, lithium carbonate, is used as a modifier in glazes. It is also employed as an additive in heat-resistant glass and optical glass.

Primarily, lithium is vital in battery materials, including primary lithium batteries and secondary lithium-ion batteries, which are rechargeable and used repeatedly. Lithium-ion batteries power various modern electronic devices, such as smartphones, tablets, notebook computers, and electric vehicles. They are prized for their high energy density, long service life, lightweight, high charge efficiency, and low self-discharge rate, making lithium-ion batteries indispensable in modern society.

Beyond batteries, lithium serves in nuclear reactor materials, as a polymerization catalyst for organic synthesis, and as an alloying element with magnesium and aluminum. Its vivid red flame reaction is employed in fireworks and as a dehumidifying agent leveraging its reactivity with moisture.

Properties of Lithium

Lithium, a soft, silvery-white metal, is the lightest of all metallic elements. It is widely distributed in ores and rocks, such as lithia mica, spodumene, and lepidolite, and remains stable in dry air, showing minimal oxidation. However, it reacts with moisture, even at room temperature, producing nitrides when heated and oxides. Lithium reacts violently with water, producing hydrogen gas that ignites, though less violently than potassium or sodium, its fellow alkali metals. It boasts the highest melting point of 180°C and boiling point of 1,330°C among the alkali metals.

Other Information on Lithium

1. Distribution of Lithium

Lithium is broadly distributed across the earth but exists mainly in the form of compounds due to its high reactivity. It comprises 0.004% of the earth’s crust and is extracted from salt lake brines and ores such as amblygonite (2LiF・Al₂O₃・P₂O₅), spodumene (Li₂O・Al₂O₃・4SiO₂), petalite (Li₂O・Al₂O₃・8SiO₂), and lepidolite (K(Li,Al)₃(Al,Si,Rb)₄O₁₀(F,OH)₂).

2. Production Method of Lithium

Lithium is extracted from ores and brine, which are then converted into lithium carbonate before undergoing electrolysis to produce solid lithium. Initially, lithium carbonate is extracted from ores by roasting and crushing, followed by treatment with sulfuric acid. Impurities are removed by adding sodium carbonate and calcium hydroxide. Similarly, lithium carbonate from brine is processed by sun-drying lithium chloride brine, adding sodium carbonate to precipitate lithium carbonate. Finally, metallic lithium is produced by electrolyzing lithium chloride, obtained from lithium carbonate, with hydrochloric acid and potassium chloride, resulting in lithium precipitation at the cathode and chlorine at the anode.

What Is Lanthanum?

Lanthanum, with the symbol La, was discovered by Swedish chemist Carl Gustaf Mosander in 1839 as an impurity in cerium nitrate. It is part of group 3 of the periodic table and lends its name to the lanthanide series, which consists of 15 elements from lanthanum to lutetium (Lu).

Though classified as a rare earth element, lanthanides, including lanthanum, are more prevalent in the Earth’s crust than lead, occurring in minerals such as phosphates, carbonates, and silicates.

Uses of Lanthanum

Lanthanum oxide (La₂O₃), employed in the manufacture of ceramic capacitors and optical lenses, and LaNi₅, an alloy with hydrogen storage capabilities, is crucial in the negative electrodes of nickel-hydrogen rechargeable batteries.

The isolation and purification of lanthanum from other lanthanides are challenging and costly. As a result, it is commonly used in a mischmetal form, comprising more than 50% lanthanum, rather than in a pure state.

Furthermore, lanthanum carbonate serves as a phosphorus absorption inhibitor in patients suffering from renal failure by forming non-absorbable phosphorus compounds in the intestinal tract.

Properties of Lanthanum

Lanthanum is a soft, silvery-white metal that tarnishes slowly in the air. It is ductile enough to be cut with a knife. It reacts with halogens to produce trihalides and forms binary compounds with nonmetals like boron, carbon, nitrogen, and phosphorus upon heating. Lanthanum hydroxide, La(OH)₃, results from its reaction with water, while dilute sulfuric acid yields a colorless solution of the hydrated tripositive ion [La(H₂O)₉]³⁺, due to the absence of f-electrons in La³⁺.

It is the least volatile among lanthanides, with a melting point of 920°C and a boiling point of 3,464°C.

Structure of Lanthanum

Lanthanum has an atomic weight of 138.91, a density of 6.162 g/cm³ at room temperature, and a liquid density of 5.94 g/cm³ at its melting point. It boasts the largest atomic radius among the lanthanides. Its electron configuration is [Xe] 5d¹ 6s², and it usually exhibits a +3 oxidation state in chemical reactions.

At room temperature, lanthanum adopts a hexagonal structure, transitioning to face-centered cubic at 310°C, and becoming body-centered cubic at 865°C.

Other Information on Lanthanum

1. Production of Lanthanum

As the third most abundant lanthanide in the Earth’s crust, lanthanum can be extracted by heating lanthanum oxide with hydrofluoric acid and ammonium chloride or fluoride, then reduced to metallic form with alkali or alkaline earth metals in vacuum or argon atmosphere. Pure lanthanum is also obtainable through the electrolysis of a molten mixture of LaCl₃, NaCl, or KCl at elevated temperatures.

2. Lanthanum Isotopes

Lanthanum’s isotopes range in atomic weight from 117 to 155, with naturally occurring stable isotopes being ¹³⁹La and ¹³⁸La. ¹³⁹La, the most prevalent at 99.91%, and ¹³⁸La, a radioactive isotope with a half-life of 10⁵ x 10⁹ years. In total, thirty-eight radioactive isotopes and three nuclear isomers of lanthanum have been identified.

3. Hazards of Lanthanum

Lanthanum, while low to moderately toxic, requires careful handling. Its injection can lead to adverse effects such as hypotension, hyperglycemia, liver alteration, and spleen degeneration. It also influences human metabolism, potentially lowering cholesterol levels, appetite, blood pressure, and blood clotting risk, and acts as a painkiller when injected into the brain.

What Is Iodine?

Iodine, a diatomic molecule with the molecular formula I₂, is a significant element in the halogen group with atomic number 53. Its CAS registration number is 7553-56-2. Essential for human health, iodine is primarily found in the thyroid gland where it plays a crucial role in hormone production.

Often referred to as the “mineral of the sea,” iodine is plentiful in seaweed and seafood, including kelp and wakame.

Uses of Iodine

Iodine finds extensive applications across various fields, particularly in chemistry and medicine. Key uses include:

• Organic synthesis intermediates and catalysts, stabilizers for organic compounds, rare metals smelting, and analytical reagents
• Pharmaceuticals, dietary supplements, and disinfectants
• Diagnostic and therapeutic radiopharmaceuticals (radioactive iodine-131)
• Livestock feed additives, dyes, photoengraving processes, and agrochemicals
• Thin film thickness measurement, water pipe inspection, and oil field exploration

1. Analytical Chemistry

The addition of starch to an iodine solution triggers an iodo-starch reaction, turning the solution blue-purple. This reaction is a cornerstone of iodometric titration.

2. Medical Field

Iodine is widely utilized as a disinfectant. Common formulations include iodine tincture (an alcohol solution of iodine), Lugol’s solution (a glycerol solution of iodine and potassium iodide), and povidone-iodine (a complex of iodine and polyvinylpyrrolidone).

It is also a key component in contrast media for X-ray diagnostics and various gargles.

3. Other Industrial Applications

As a catalyst in industry and agriculture, iodine serves multiple purposes, including as a raw material for antifungal agents. It’s also finding uses in advanced technologies such as polarizing films for LCD panels and solar cells.

Properties of Iodine

Figure 1. Basic Information on Iodine

Iodine appears as blackish-purple, shiny crystals at room temperature, with a pungent odor and the capability to sublime. It has a molecular weight of 253.808, a melting point of 113.7°C, and a boiling point of 184.4°C. In liquid form, iodine is reddish-brown, and its vapor is purple. The density of iodine is 4.933 g/mL. It is soluble in ethanol and diethyl ether but has limited solubility in water (0.3 g/L).

Types of Iodine

Iodine is available primarily as reagents for research and development and as industrial chemicals.

1. Reagent Products for Research and Development

Used in R&D for pharmaceuticals, reagent preparation, and organic synthesis, iodine is offered in quantities suited for laboratory use, such as 5g, 25g, 100g, and 500g. Due to its sublimation properties, it is typically shipped and stored under refrigeration.

2. Industrial Chemicals

For industrial applications, iodine is packaged in 20 kg and 50 kg fiber drums, facilitating easy handling in factory settings. It is intended for use in a variety of industries, including X-ray contrast media, fungicides, liquid crystals, and more.

Other Information on Iodine

1. Dissolution in Aqueous Potassium Iodide Solution

Figure 2. Reaction of Iodine With Iodide Ions

Although iodine has limited solubility in water, it dissolves well in aqueous potassium iodide (KI) solutions, forming triiodide ions through a reaction with iodide ions.

2. Characteristics of Iodine Crystals

Iodine crystals exhibit an orthorhombic structure and are non-magnetic. They have an electrical resistivity of 1.3×10⁷Ω⋅m at 0°C and a thermal conductivity of 0.449 W/(m⋅K) at 300K. Iodine’s oxidizing power is lower than that of other halogens such as fluorine, chlorine, and bromine.

3. Toxicity of Iodine

Figure 3. Hazardous Properties of Iodine

Due to its toxic nature, iodine handling requires proper ventilation and protective equipment. It is classified by GHS as follows:

• Acute toxicity (oral): Category 4
• Acute toxicity (inhalation: vapor): Category 1
• Skin corrosion/irritation: Category 2
• Eye damage/irritation: Category 2
• Skin sensitization: Category 1
• Specific target organ toxicity (single exposure): Category 3 (respiratory irritation)
• Specific target organ toxicity (repeated exposure): Category 1 (thyroid gland)

What Is Methyl Ethyl Ketone?

Methyl ethyl ketone, a clear and colorless liquid, is recognized by various names such as ethyl methyl ketone, methyl acetone, and 2-butanone, abbreviated as MEK. As a chemical governed by strict regulations, it’s classified under safety laws due to its toxic and hazardous nature.

Uses of Methyl Ethyl Ketone

Methyl ethyl ketone serves as a versatile solvent in numerous industries, facilitating processes in resin processing, paint formulation, and the production of adhesives and printing inks. It’s also utilized in petroleum refining, manufacturing of polyurethane and other resins, and as a key ingredient in lacquers. Moreover, its applications extend to industrial cleaning, PVC treatments, film processing, and as a vulcanization accelerator.

Properties of Methyl Ethyl Ketone

With its distinct odor and a molecular weight of 72.10, MEK is characterized by its solubility in water and miscibility with various organic solvents. Its physical attributes include a low melting point and a boiling point at 80°C, alongside its high volatility and flammability, emphasizing the need for careful handling.

Other Information on Methyl Ethyl Ketone

1. How Methyl Ethyl Ketone Is Produced

The synthesis of methyl ethyl ketone typically involves the dehydrogenation of 2-butanol, a process that can be initiated through indirect or direct hydration of n-butene. The indirect method involves sulfuric acid esterification followed by hydrolysis, while the direct method uses a catalyzed hydration process. Dehydrogenation of 2-butanol, catalyzed by Cu-Zn, ultimately produces methyl ethyl ketone.

2. Precautions for Handling Methyl Ethyl Ketone

Given its potential health risks, including respiratory and skin irritation, as well as its classification as a possible carcinogen, proper safety measures are crucial when handling methyl ethyl ketone. This entails using protective gear, ensuring good ventilation, and strict adherence to storage guidelines to mitigate its flammable and explosive nature.

What Is Ethyl Ethynyl Ether?

Ethyl ethynyl ether is an ether compound that exists as a gas at room temperature. Its IUPAC name is “ethyl ethynyl ether,” and due to its low boiling point of 10.8 °C, it is transported and stored in gas cylinders at room temperature. It is highly flammable and is classified as a “flammable gas” under various safety regulations.

At concentrations between 2% and 10.1% in air, ethyl ethynyl ether is easily ignited by electric sparks or heating elements. Therefore, it should be handled in well-ventilated areas to prevent accumulation within the flammable range.

Uses of Ethyl Ethynyl Ether

1. Current Main Uses

Ethyl ethynyl ether is utilized in several applications, including as an active pharmaceutical ingredient, a synthetic raw material for chemical materials, and an intermediate. It is sometimes used in aerosol sprays due to its gaseous state at room temperature.

When cooled to a liquid, ethyl ethynyl ether can also serve as a solvent for solid-liquid extraction. It quickly volatilizes at room temperature, providing a rapid means of obtaining solid products. However, diethyl ether is often preferred due to its lower cost and ease of handling.

2. Future Application Expansion

Future applications may include its use as a substitute for liquefied petroleum gas (LPG). Research is ongoing to explore the use of ethers like ethyl ethynyl ether as civilian gas alternatives to LPG, due to their flammability and ability to be stored and transported as liquids when cooled.

Dimethyl ether is currently the most promising candidate for this application, but ethyl ethynyl ether, with similar properties, may also be considered.

Properties of Ethyl Ethynyl Ether

1. Basic Structure and Properties

Ethyl ethynyl ether, with the molecular formula C3H8O, is an ether consisting of an ethyl group and a methyl group bonded together. It is a structural isomer of propanol. Key properties include a molecular weight of 60.1, a specific gravity of 0.725 (as a liquid at 0 °C), and solubility in water, acetone, and chloroform. It is also miscible with ethanol.

2. Industrial Synthesis

Symmetric ethers like diethyl ether are commonly synthesized through the dehydration-condensation reaction of alcohols. However, asymmetric ethers like ethyl ethynyl ether are typically synthesized using the Williamson ether synthesis method, which involves the SN2 reaction of metal alkoxides and haloalkanes. For ethyl ethynyl ether, metal ethoxide, and methyl bromide are used as raw materials.

Other Information on Ethyl Ethynyl Ether

1. Symmetric and Asymmetric Ethers

Ethyl ethynyl ether is an example of an asymmetric ether. Symmetric ethers have two identical alkyl or aryl groups, while asymmetric ethers have two different groups. Examples of symmetric ethers include dimethyl ether and diethyl ether, while asymmetric ethers include ethyl ethynyl ether and anisole.

2. Challenges with Ethyl Ethynyl Ether

The primary challenge with ethyl ethynyl ether lies in its production cost. The Williamson ether synthesis process used for its production is more complex and costly compared to the methods used for producing diethyl ether or dimethyl ether.

What Is Methyl Isobutyl Ketone?

Methyl isobutyl ketone, a colorless organic solvent with the formula C6H12O, is also known by its IUPAC name, 4-methyl-2-pentanone. Synthesized from acetone, it undergoes a three-step process, including an aldol reaction to form diacetone alcohol, dehydration to methyl oxide, and final hydrogenation.

Uses of Methyl Isobutyl Ketone

Primarily used as a solvent in the manufacturing of synthetic resins, paints, adhesives, and inks, methyl isobutyl ketone’s solvency properties extend to various other applications including pharmaceutical extractions, and as a component in antifreeze solutions, among others. It also finds use in magnetic tape production and in extracting penicillins.

Properties of Methyl Isobutyl Ketone

With a melting point of -80°C and a boiling range of around 116.5°C, methyl isobutyl ketone’s physical and chemical properties make it a versatile solvent. Soluble in water, ethanol, ether, and acetone, it is commercially available mainly as a liquid solvent due to its specific sweet odor and flammability.

Other Information on Methyl Isobutyl Ketone

1. Safety

Recognized as a flammable liquid with various health risks, methyl isobutyl ketone is subject to strict safety classifications and handling guidelines under global safety standards.

2. Handling Instructions

Proper safety gear, including protective gloves and glasses, should be worn when handling Methyl Isobutyl Ketone. Work areas should be well-ventilated and equipped with safety installations to mitigate its flammable nature and potential health hazards.

3. Storage

For safety, store methyl isobutyl ketone in specified containers away from incompatible substances and fireproof, well-ventilated areas to prevent accidental ignition or health risks.

4. Regulatory Compliance

Methyl isobutyl ketone is classified as a specified chemical substance, requiring businesses to adhere to stringent environmental and health safety measures, including reporting and management practices aimed at reducing environmental impact and ensuring workplace safety.