Occupational Safety and Health Administration

From Wikipedia, the free encyclopedia

Occupational Safety and Health Administration

Agency overview
Formed	1971
Jurisdiction	Federal government of the United States
Headquarters	Frances Perkins Building Washington, D.C.
Employees	2,265 (2015)
Annual budget	$552 million (2015)
Agency executive	Loren Sweatt, Acting Assistant Secretary
Parent department	United States Department of Labor
Website	www.osha.gov

The Occupational Safety and Health Administration (OSHA) (/ˈoʊʃə/) is an agency of the United States Department of Labor. Congress established the agency under the Occupational Safety and Health Act (OSH Act), which President Richard M. Nixon signed into law on December 29, 1970. OSHA's mission is to "assure safe and healthy working conditions for working men and women by setting and enforcing standards and by providing training, outreach, education and assistance". The agency is also charged with enforcing a variety of whistleblower statutes and regulations. OSHA is currently headed by Acting Assistant Secretary of Labor Loren Sweatt. OSHA's workplace safety inspections have been shown to reduce injury rates and injury costs without adverse effects to employment, sales, credit ratings, or firm survival.

History

OSHA officially formed on April 28, 1971, the date that the OSH Act became effective. George Guenther was appointed as the agency's first director.

OSHA has a number of training, compliance assistance, and health and safety recognition programs throughout its history. The OSHA Training Institute, which trains government and private sector health and safety personnel, began in 1972. In 1978, the agency began a grantmaking program, now called the Susan Harwood Training Grant Program, to train workers and employers in reducing workplace hazards. OSHA started the Voluntary Protection Programs in 1982, which allow employers to apply as "model workplaces" to achieve special designation if they meet certain requirements.

OSH Act coverage

The OSH Act covers most private sector employers and their workers, in addition to some public sector employers and workers in the 50 states and certain territories and jurisdictions under federal authority. Those jurisdictions include the District of Columbia, Puerto Rico, the Virgin Islands, American Samoa, Guam, Northern Mariana Islands, Wake Island, Johnston Island, and the Outer Continental Shelf Lands as defined in the Outer Continental Shelf Lands Act.

Private sector employers

The OSH Act covers most private sector employers in all 50 states, the District of Columbia, and other U.S. jurisdictions—either directly through federal OSHA or through an OSHA approved state plan. 

State plans are OSHA-approved job safety and health programs operated by individual states instead of federal OSHA. Federal OSHA approves and monitors all state plans and provides as much as fifty percent of the funding for each program. State-run safety and health programs are required to be at least as effective as the federal OSHA program. 

The following 22 states or territories have OSHA-approved state programs: Alaska, Arizona, California, Hawaii, Indiana, Iowa, Kentucky, Maryland, Michigan, Minnesota, Nevada, New Mexico, North Carolina, Oregon, Puerto Rico, South Carolina, Tennessee, Utah, Vermont, Virginia, Washington, and Wyoming.

Federal OSHA provides coverage to certain workplaces specifically excluded from a state’s plan — for example, work in maritime industries or on military bases.

State and local governments

Workers at state and local government agencies are not covered by federal OSHA, but have OSH Act protections if they work in those states that have an OSHA-approved state program. OSH Act rules also permit states and territories to develop plans that cover only public sector (state and local government) workers. In these cases, private sector workers and employers remain under federal OSHA jurisdiction. Five additional states and one U.S. territory have OSHA approved state plans that cover public sector workers only: Connecticut, Illinois, Maine, New Jersey, New York, and the Virgin Islands.

Federal government agencies

OSHA’s protection applies to all federal agencies. Section 19 of the OSH Act makes federal agency heads responsible for providing safe and healthful working conditions for their workers. OSHA conducts inspections of federal facilities in response to workers’ reports of hazards and under programs that target high hazard federal workplaces.

Federal agencies must have a safety and health program that meets the same standards as private employers. OSHA issues “virtual fines” to federal agencies – following an inspection where violations are found, OSHA issues a press release stating the size the fine would be if the federal agency were a private sector employer. Under a 1998 amendment, the OSHA Act covers the U.S. Postal Service the same as any private sector employer.

Not covered under the OSH Act

The OSH Act does not cover the self-employed, immediate family members of farm employers, or workplace hazards regulated by another federal agency (for example, the Mine Safety and Health Administration, the Department of Energy, or Coast Guard).

Rights and responsibilities under OSH Act law

Employers have the responsibility to provide a safe workplace.

By law, employers must provide their workers with a workplace that does not have serious hazards and must follow all OSH Act safety and health standards. Employers must find and correct safety and health problems. The OSH Act further requires that employers must first try to eliminate or reduce hazards by making feasible changes in working conditions rather than relying on personal protective equipment such as masks, gloves, or earplugs. Switching to safer chemicals, enclosing processes to trap harmful fumes, or using ventilation systems to clean the air are examples of effective ways to eliminate or reduce risks. 

Employers must also:

• Inform workers about chemical hazards through training, labels, alarms, color-coded systems, chemical information sheets and other methods.
• Provide safety training to workers in a language and vocabulary they can understand.
• Keep accurate records of work-related injuries and illnesses.
• Perform tests in the workplace, such as air sampling, required by some OSH Act standards.
• Provide required personal protective equipment at no cost to workers. (Employers must pay for most types of required personal protective equipment.)
• Provide hearing exams or other medical tests when required by OSH Act standards.
• Post OSHA citations and annually post injury and illness summary data where workers can see them.
• Notify OSHA within eight hours of a workplace fatality. Notify OSHA within 24 hours of all work-related inpatient hospitalizations, all amputations, and all losses of an eye (1-800-321-OSHA [6742]).
• Prominently display the official OSHA Job Safety and Health – It’s the Law poster that describes rights and responsibilities under the OSH Act.
• Not retaliate or discriminate against workers for using their rights under the law, including their right to report a work-related injury or illness.

Workers have the right to:

• Working conditions that do not pose a risk of serious harm.
• File a confidential complaint with OSHA to have their workplace inspected.
• Receive information and training about hazards, methods to prevent harm, and the OSH Act standards that apply to their workplace. The training must be done in a language and vocabulary workers can understand.
• Receive copies of records of work-related injuries and illnesses that occur in their workplace.
• Receive copies of the results from tests and monitoring done to find and measure hazards in their workplace.
• Receive copies of their workplace medical records.
• Participate in an OSHA inspection and speak in private with the inspector.
• File a complaint with OSHA if they have been retaliated or discriminated against by their employer as the result of requesting an inspection or using any of their other rights under the OSH Act.
• File a complaint if punished or retaliated against for acting as a “whistleblower” under the 21 additional federal laws for which OSHA has jurisdiction.

Temporary workers must be treated like permanent employees. Staffing agencies and host employers share a joint accountability over temporary workers. Both entities are therefore bound to comply with workplace health and safety requirements and to ensure worker safety and health. OSHA could hold both the host and temporary employers responsible for the violation of any condition.

Health and safety standards

The Occupational Safety and Health Act grant OSHA the authority to issue workplace health and safety regulations. These regulations include limits on hazardous chemical exposure, employee access to hazard information, requirements for the use of personal protective equipment, and requirements to prevent falls and hazards from operating dangerous equipment. 

The OSH Act's current Construction, General Industry, Maritime and Agriculture standards are designed to protect workers from a wide range of serious hazards. Examples of OSHA standards include requirements for employers to provide fall protection such as a safety harness/line or guardrails; prevent trenching cave-ins; prevent exposure to some infectious diseases; ensure the safety of workers who enter confined spaces; prevent exposure to harmful chemicals; put guards on dangerous machines; provide respirators or other safety equipment; and provide training for certain dangerous jobs in a language and vocabulary workers can understand. 

OSHA sets enforceable permissible exposure limits (PELs) to protect workers against the health effects of exposure to hazardous substances, including limits on the airborne concentrations of hazardous chemicals in the air. Most of OSHA’s PELs were issued shortly after adoption of the OSH Act in 1970. Attempts to issue more stringent PELs have been blocked by litigation from industry; thus, the vast majority of PELs have not been updated since 1971. The agency has issued non-binding, alternate occupational exposure limits that may better protect workers.

Employers must also comply with the General Duty Clause of the OSH Act. This clause requires employers to keep their workplaces free of serious recognized hazards and is generally cited when no specific OSHA standard applies to the hazard. 

In its first year of operation, OSHA was permitted to adopt regulations based on guidelines set by certain standards organizations, such as the American Conference of Governmental Industrial Hygienists, without going through all of the requirements of a typical rulemaking. OSHA is granted the authority to promulgate standards that prescribe the methods employers are legally required to follow to protect their workers from hazards. Before OSHA can issue a standard, it must go through a very extensive and lengthy process that includes substantial public engagement, notice and comment. The agency must show that a significant risk to workers exists and that there are feasible measures employers can take to protect their workers. 

In 2000, OSHA issued an ergonomics standard. In March 2001, Congress voted to repeal the standard through the Congressional Review Act. The repeal, one of the first major pieces of legislation signed by President George W. Bush, is the first instance that Congress has successfully used the Congressional Review Act to block regulation. 

Since 2001, OSHA has issued the following standards:

• 2002: Exit Routes, Emergency Action Plans, and Fire Prevention Plans
• 2004: Commercial Diving Operations
• 2004: Fire Protection in Shipyards
• 2006: Occupational Exposure to Hexavalent Chromium
• 2006: Assigned Protection Factors for Respiratory Protection Equipment
• 2007: Electrical Installation Standard
• 2007: Personal Protective Equipment Payment (Clarification)
• 2008: Vertical Tandem Lifts
• 2010: Cranes and Derricks in Construction
• 2010: General Working Conditions in Shipyards
• 2012: GHS Update to the Hazard Communication Standard
• 2014: New Recordkeeping and Reporting Requirements for Employers
• 2014: Revision to Electric Power Generation, Transmission, and Distribution; Electrical Protective Equipment
• 2016: Occupational Exposure to Respirable Crystalline Silica
• 2016: Update General Industry Walking-Working Surfaces and Fall Protection Standards 

Enforcement

OSHA is responsible for enforcing its standards on regulated entities. Compliance Safety and Health Officers carry out inspections and assess fines for regulatory violations. Inspections are planned for worksites in particularly hazardous industries. Inspections can also be triggered by a workplace fatality, multiple hospitalizations, worker complaints, or referrals. 

OSHA is a small agency, given the size of its mission: with its state partners, OSHA has approximately 2,400 inspectors covering more than 8 million workplaces where 130 million workers are employed. In Fiscal Year 2012 (ending Sept. 30), OSHA and its state partners conducted more than 83,000 inspections of workplaces across the United States — just a fraction of the nation’s worksites. According to a report by AFL–CIO, it would take OSHA 129 years to inspect all workplaces under its jurisdiction.

Enforcement plays an important part in OSHA’s efforts to reduce workplace injuries, illnesses, and fatalities. Inspections are initiated without advance notice, conducted using on-site or telephone and facsimile investigations, performed by trained compliance officers and scheduled based on the following priorities [highest to lowest]: imminent danger; catastrophes – fatalities or hospitalizations; worker complaints and referrals; targeted inspections – particular hazards, high injury rates; and follow-up inspections. 

Current workers or their representatives may file a complaint and ask OSHA to inspect their workplace if they believe that there is a serious hazard or that their employer is not following OSHA standards. Workers and their representatives have the right to ask for an inspection without OSHA telling their employer who filed the complaint. It is a violation of the OSH Act for an employer to fire, demote, transfer or in any way discriminate against a worker for filing a complaint or using other OSHA rights. 

When an inspector finds violations of OSHA standards or serious hazards, OSHA may issue citations and fines. A citation includes methods an employer may use to fix a problem and the date by which the corrective actions must be completed. 

OSHA’s fines are very low compared with other government agencies. They were raised for the first time since 1990 on Aug. 2, 2016 to comply with the 2015 Federal Civil Penalties Inflation Adjustment Act Improvements Act passed by Congress to advance the effectiveness of civil monetary penalties and to maintain their deterrent effect. The new law directs agencies to adjust their penalties for inflation each year. The maximum OSHA fine for a serious violation is $12,500 and the maximum fine for a repeat or willful violation is $125,000. In determining the amount of the proposed penalty, OSHA must take into account the gravity of the alleged violation and the employer’s size of the business, good faith and history of previous violations. Employers have the right to contest any part of the citation, including whether a violation actually exists. Workers only have the right to challenge the deadline by which a problem must be resolved. Appeals of citations are heard by the independent Occupational Safety and Health Review Commission (OSHRC).

OSHA carries out its enforcement activities through its 10 regional offices and 90 area offices. OSHA’s regional offices are located in Boston, New York City, Philadelphia, Atlanta, Chicago, Dallas, Kansas City metropolitan area, Denver, San Francisco, and Seattle.

Record keeping requirements

Tracking and investigating workplace injuries and illnesses play an important role in preventing future injuries and illnesses. Under OSHA’s Recordkeeping regulation, certain covered employers in high hazard industries are required to prepare and maintain records of serious occupational injuries and illnesses. This information is important for employers, workers and OSHA in evaluating the safety of a workplace, understanding industry hazards, and implementing worker protections to reduce and eliminate hazards. 

Employers with more than ten employees and whose establishments are not classified as a partially exempt industry must record serious work-related injuries and illnesses using OSHA Forms 300, 300A and 301. Recordkeeping forms, requirements and exemption information are at OSHA’s website.

Whistleblower protection

OSHA enforces the whistleblower provisions of the Occupational Safety and Health Act and 21 other statutes protecting workers who report violations of various airline, commercial motor carrier, consumer product, environmental, financial reform, food safety, health care reform, nuclear, pipeline, public transportation agency, maritime and securities laws. Over the years, OSHA has been responsible for enforcing these laws that protect the rights of workers to speak up without fear of retaliation, regardless of the relationship of these laws to occupational safety and health matters.

Compliance assistance

OSHA has developed several training, compliance assistance, and health and safety recognition programs throughout its history. 

The OSHA Training Institute, which trains government and private sector health and safety personnel, began in 1972. In 1978, the agency began a grant making program, now called the Susan Harwood Training Grant Program, to train workers and employers in identifying and reducing workplace hazards.

The Voluntary Protection Program (VPP) recognize employers and workers in private industry and federal agencies who have implemented effective safety and health management programs and maintain injury and illness rates below the national average for their respective industries. In VPP, management, labor, and OSHA work cooperatively and proactively to prevent fatalities, injuries, and illnesses through a system focused on: hazard prevention and control, worksite analysis, training, and management commitment and worker involvement.

OSHA’s On-site Consultation Program offers free and confidential advice to small and medium-sized businesses in all states across the country, with priority given to high-hazard worksites. Each year, responding to requests from small employers looking to create or improve their safety and health management programs, OSHA’s On-site Consultation Program conducts over 29,000 visits to small business worksites covering over 1.5 million workers across the nation. On-site consultation services are separate from enforcement and do not result in penalties or citations. Consultants from state agencies or universities work with employers to identify workplace hazards, provide advice on compliance with OSHA standards, and assist in establishing safety and health management programs.

Under the consultation program, certain exemplary employers may request participation in OSHA’s Safety and Health Achievement Recognition Program (SHARP). Eligibility for participation includes, but is not limited to, receiving a full-service, comprehensive consultation visit, correcting all identified hazards and developing an effective safety and health management program. Worksites that receive SHARP recognition are exempt from programmed inspections during the period that the SHARP certification is valid.

OSHA also provides compliance assistance through its national and area offices. Through hundreds of publications in a variety of languages, website safety and health topics pages, and through compliance assistance staff OSHA provides information to employers and workers on specific hazards and OSHA rights and responsibilities.

Efficacy

A 2012 study in Science found that OSHA's random workplace safety inspections caused a "9.4% decline in injury rates" and a "26% reduction in injury cost" for the inspected firms. The study found "no evidence that these improvements came at the expense of employment, sales, credit ratings, or firm survival."

Much of the debate about OSHA regulations and enforcement policies revolves around the cost of regulations and enforcement, versus the actual benefit in reduced worker injury, illness and death. A 1995 study of several OSHA standards by the Office of Technology Assessment (OTA) found that OSHA relies "generally on methods that provide a credible basis for the determinations essential to rulemakings". Though it found that OSHA's finding and estimates are "subject to vigorous review and challenge", it stated that this is natural because "interested parties and experts involved in rulemakings have differing visions".

OSHA has come under considerable criticism for the ineffectiveness of its penalties, particularly its criminal penalties. The maximum penalty is a misdemeanor with a maximum of 6-months in jail. In response to the criticism, OSHA, in conjunction with the Department of Justice, has pursued several high-profile criminal prosecutions for violations under the Act, and has announced a joint enforcement initiative between OSHA and the United States Environmental Protection Agency (EPA) which has the ability to issue much higher fines than OSHA. Meanwhile, Congressional Democrats, labor unions and community safety and health advocates are attempting to revise the OSH Act to make it a felony with much higher penalties to commit a willful violation that results in the death of a worker. Some local prosecutors are charging company executives with manslaughter and other felonies when criminal negligence leads to the death of a worker.

A New York Times investigation in 2003 showed that over the 20-year period from 1982 to 2002, 2,197 workers died in 1,242 incidents in which OSHA investigators concluded that employers had willfully violated workplace safety laws. In 93% of these fatality cases arising from wilful violation, OSHA made no referral to the U.S. Department of Justice for criminal prosecution. The Times investigation found that OSHA had failed to pursue prosecution "even when employers had been cited before for the very same safety violation" and even in cases where multiple worker died. In interviews, current and former OSHA officials said that the low rates of criminal enforcement were the result of "a bureaucracy that works at every level to thwart criminal referrals. ... that fails to reward, and sometimes penalizes, those who push too hard for prosecution" and that " aggressive enforcement [was] suffocated by endless layers of review.

OSHA has also been criticized for taking too long to develop new regulations. For instance, speaking about OSHA under the George W. Bush presidency on the specific issue of combustible dust explosions, Chemical Safety Board appointee Carolyn Merritt said: "The basic disappointment has been this attitude of no new regulation. They don't want industry to be pestered. In some instances, industry has to be pestered in order to comply."

Bromine

From Wikipedia, the free encyclopedia

Bromine,  ₃₅Br

Bromine
Pronunciation	/ˈbroʊmiːn, -mɪn, -maɪn/ ​(BROH-meen, -⁠min, -⁠myn)
Appearance	reddish-brown
Standard atomic weight Ar, std(Br)	[79.901, 79.907] conventional: 79.904
Bromine in the periodic table
Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson Cl ↑ Br ↓  I  selenium ← bromine → krypton
Atomic number (Z)	35
Group	group 17 (halogens)
Period	period 4
Block	p-block
Element category	reactive nonmetal
Electron configuration	[Ar] 3d10 4s2 4p5
Electrons per shell	2, 8, 18, 7
Physical properties
Phase at STP	liquid
Melting point	265.8 K ​(−7.2 °C, ​19 °F)
Boiling point	332.0 K ​(58.8 °C, ​137.8 °F)
Density (near r.t.)	Br2, liquid: 3.1028 g/cm3
Triple point	265.90 K, ​5.8 kPa
Critical point	588 K, 10.34 MPa
Heat of fusion	(Br2) 10.571 kJ/mol
Heat of vaporisation	(Br2) 29.96 kJ/mol
Molar heat capacity	(Br2) 75.69 J/(mol·K)
Vapour pressure P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 185 201 220 244 276 332
Atomic properties
Oxidation states	−1, +1, +3, +4, +5, +7 (a strongly acidic oxide)
Electronegativity	Pauling scale: 2.96
Ionisation energies	1st: 1139.9 kJ/mol 2nd: 2103 kJ/mol 3rd: 3470 kJ/mol
Atomic radius	empirical: 120 pm
Covalent radius	120±3 pm
Van der Waals radius	185 pm
Spectral lines of bromine
Other properties
Natural occurrence	primordial
Crystal structure	​orthorhombic
Speed of sound	206 m/s (at 20 °C)
Thermal conductivity	0.122 W/(m·K)
Electrical resistivity	7.8×1010 Ω·m (at 20 °C)
Magnetic ordering	diamagnetic
Magnetic susceptibility	−56.4·10−6 cm3/mol
CAS Number	7726-95-6
History
Discovery and first isolation	Antoine Jérôme Balard and Carl Jacob Löwig (1825)
Main isotopes of bromine
Iso­tope Abun­dance Half-life (t1/2) Decay mode Pro­duct 79Br 51% stable 81Br 49% stable

Bromine is a chemical element with the symbol Br and atomic number 35. It is the third-lightest halogen, and is a fuming liquid with a deep red color. At room temperature, Bromine evaporates readily to form a red to amber coloured gas. Bromine's properties are intermediate between those of chlorine and iodine. Isolated independently by two chemists, Carl Jacob Löwig (in 1825) and Antoine Jérôme Balard (in 1826), its name was derived from the Ancient Greek βρῶμος ("stench"), referencing its sharp and disagreeable smell.

Elemental bromine is very reactive and thus does not occur free in nature, but in colourless soluble crystalline mineral halide salts, analogous to table salt. While it is rather rare in the Earth's crust, the high solubility of the bromide ion (Br⁻) has caused its accumulation in the oceans. Commercially the element is easily extracted from brine pools, mostly in the United States, Israel and China. The mass of bromine in the oceans is about one three-hundredth that of chlorine.

At high temperatures, organobromine compounds readily dissociate to yield free bromine atoms, a process that stops free radical chemical chain reactions. This effect makes organobromine compounds useful as fire retardants, and more than half the bromine produced worldwide each year is put to this purpose. The same property causes ultraviolet sunlight to dissociate volatile organobromine compounds in the atmosphere to yield free bromine atoms, causing ozone depletion. As a result, many organobromide compounds—such as the pesticide methyl bromide—are no longer used. Bromine compounds are still used in well drilling fluids, in photographic film, and as an intermediate in the manufacture of organic chemicals.

Large amounts of bromide salts are toxic from the action of soluble bromide ion, causing bromism. However, a clear biological role for bromide ion and hypobromous acid has recently been elucidated, and it now appears that bromine is an essential trace element in humans. The role of biological organobromine compounds in sea life such as algae has been known for much longer. As a pharmaceutical, the simple bromide ion (Br⁻) has inhibitory effects on the central nervous system, and bromide salts were once a major medical sedative, before replacement by shorter-acting drugs. They retain niche uses as antiepileptics.

History

Antoine Balard, one of the discoverers of bromine

 

Bromine was discovered independently by two chemists, Carl Jacob Löwig and Antoine Balard, in 1825 and 1826, respectively.

Löwig isolated bromine from a mineral water spring from his hometown Bad Kreuznach in 1825. Löwig used a solution of the mineral salt saturated with chlorine and extracted the bromine with diethyl ether. After evaporation of the ether a brown liquid remained. With this liquid as a sample of his work he applied for a position in the laboratory of Leopold Gmelin in Heidelberg. The publication of the results was delayed and Balard published his results first.

Balard found bromine chemicals in the ash of seaweed from the salt marshes of Montpellier. The seaweed was used to produce iodine, but also contained bromine. Balard distilled the bromine from a solution of seaweed ash saturated with chlorine. The properties of the resulting substance were intermediate between those of chlorine and iodine; thus he tried to prove that the substance was iodine monochloride (ICl), but after failing to do so he was sure that he had found a new element, and named it muride, derived from the Latin word muria for brine.

After the French chemists Louis Nicolas Vauquelin, Louis Jacques Thénard, and Joseph-Louis Gay-Lussac approved the experiments of the young pharmacist Balard, the results were presented at a lecture of the Académie des Sciences and published in Annales de Chimie et Physique. In his publication, Balard states that he changed the name from muride to brôme on the proposal of M. Anglada. Brôme (bromine) derives from the Greek βρωμος (stench). Other sources claim that the French chemist and physicist Joseph-Louis Gay-Lussac suggested the name brôme for the characteristic smell of the vapors. Bromine was not produced in large quantities until 1858, when the discovery of salt deposits in Stassfurt enabled its production as a by-product of potash.

Apart from some minor medical applications, the first commercial use was the daguerreotype. In 1840, bromine was discovered to have some advantages over the previously used iodine vapor to create the light sensitive silver halide layer in daguerreotypy.

Potassium bromide and sodium bromide were used as anticonvulsants and sedatives in the late 19th and early 20th centuries, but were gradually superseded by chloral hydrate and then by the barbiturates. In the early years of the First World War, bromine compounds such as xylyl bromide were used as poison gas.

Properties

Illustrative and secure bromine sample for teaching

 

Bromine is the third halogen, being a nonmetal in group 17 of the periodic table. Its properties are thus similar to those of fluorine, chlorine, and iodine, and tend to be intermediate between those of the two neighbouring halogens, chlorine and iodine. Bromine has the electron configuration [Ar]3d¹⁰4s²4p⁵, with the seven electrons in the fourth and outermost shell acting as its valence electrons. Like all halogens, it is thus one electron short of a full octet, and is hence a strong oxidising agent, reacting with many elements in order to complete its outer shell. Corresponding to periodic trends, it is intermediate in electronegativity between chlorine and iodine (F: 3.98, Cl: 3.16, Br: 2.96, I: 2.66), and is less reactive than chlorine and more reactive than iodine. It is also a weaker oxidising agent than chlorine, but a stronger one than iodine. Conversely, the bromide ion is a weaker reducing agent than iodide, but a stronger one than chloride. These similarities led to chlorine, bromine, and iodine together being classified as one of the original triads of Johann Wolfgang Döbereiner, whose work foreshadowed the periodic law for chemical elements. It is intermediate in atomic radius between chlorine and iodine, and this leads to many of its atomic properties being similarly intermediate in value between chlorine and iodine, such as first ionisation energy, electron affinity, enthalpy of dissociation of the X₂ molecule (X = Cl, Br, I), ionic radius, and X–X bond length. The volatility of bromine accentuates its very penetrating, choking, and unpleasant odour.

All four stable halogens experience intermolecular van der Waals forces of attraction, and their strength increases together with number of electrons among all homonuclear diatomic halogen molecules. Thus, the melting and boiling points of bromine are intermediate between those of chlorine and iodine. As a result of the increasing molecular weight of the halogens down the group, the density and heats of fusion and vaporisation of bromine are again intermediate between those of chlorine and iodine, although all their heats of vaporisation are fairly low (leading to high volatility) thanks to their diatomic molecular structure. The halogens darken in colour as the group is descended: fluorine is a very pale yellow gas, chlorine is greenish-yellow, and bromine is a reddish-brown volatile liquid that melts at −7.2 °C and boils at 58.8 °C. (Iodine is a shiny black solid.) This trend occurs because the wavelengths of visible light absorbed by the halogens increase down the group. Specifically, the colour of a halogen, such as bromine, results from the electron transition between the highest occupied antibonding π_g molecular orbital and the lowest vacant antibonding σ_u molecular orbital. The colour fades at low temperatures, so that solid bromine at −195 °C is pale yellow.

Like solid chlorine and iodine, solid bromine crystallises in the orthorhombic crystal system, in a layered lattice of Br₂ molecules. The Br–Br distance is 227 pm (close to the gaseous Br–Br distance of 228 pm) and the Br···Br distance between molecules is 331 pm within a layer and 399 pm between layers (compare the van der Waals radius of bromine, 195 pm). This structure means that bromine is a very poor conductor of electricity, with a conductivity of around 5 × 10⁻¹³ Ω⁻¹ cm⁻¹ just below the melting point, although this is better than the essentially undetectable conductivity of chlorine.

At a pressure of 55 GPa (roughly 540,000 times atmospheric pressure) bromine undergoes an insulator-to-metal transition. At 75 GPa it changes to a face-centered orthorhombic structure. At 100 GPa it changes to a body centered orthorhombic monatomic form.

Isotopes

Bromine has two stable isotopes, ⁷⁹Br and ⁸¹Br. These are its only two natural isotopes, with ⁷⁹Br making up 51% of natural bromine and ⁸¹Br making up the remaining 49%. Both have nuclear spin 3/2− and thus may be used for nuclear magnetic resonance, although ⁸¹Br is more favourable. The relatively 1:1 distribution of the two isotopes in nature is helpful in identification of bromine containing compounds using mass spectroscopy. Other bromine isotopes are all radioactive, with half-lives too short to occur in nature. Of these, the most important are ⁸⁰Br (t_1/2 = 17.7 min), ^80mBr (t_1/2 = 4.421 h), and ⁸²Br (t_1/2 = 35.28 h), which may be produced from the neutron activation of natural bromine. The most stable bromine radioisotope is ⁷⁷Br (t_1/2 = 57.04 h). The primary decay mode of isotopes lighter than ⁷⁹Br is electron capture to isotopes of selenium; that of isotopes heavier than ⁸¹Br is beta decay to isotopes of krypton; and ⁸⁰Br may decay by either mode to stable ⁸⁰Se or ⁸⁰Kr.

Chemistry and compounds

Bromine is intermediate in reactivity between chlorine and iodine, and is one of the most reactive elements. Bond energies to bromine tend to be lower than those to chlorine but higher than those to iodine, and bromine is a weaker oxidising agent than chlorine but a stronger one than iodine. This can be seen from the standard electrode potentials of the X₂/X⁻ couples (F, +2.866 V; Cl, +1.395 V; Br, +1.087 V; I, +0.615 V; At, approximately +0.3 V). Bromination often leads to higher oxidation states than iodination but lower or equal oxidation states to chlorination. Bromine tends to react with compounds including M–M, M–H, or M–C bonds to form M–Br bonds.

Hydrogen bromide

The simplest compound of bromine is hydrogen bromide, HBr. It is mainly used in the production of inorganic bromides and alkyl bromides, and as a catalyst for many reactions in organic chemistry. Industrially, it is mainly produced by the reaction of hydrogen gas with bromine gas at 200–400 °C with a platinum catalyst. However, reduction of bromine with red phosphorus is a more practical way to produce hydrogen bromide in the laboratory:

2 P + 6 H₂O + 3 Br₂ → 6 HBr + 2 H₃PO₃

H₃PO₃ + H₂O + Br₂ → 2 HBr + H₃PO₄

At room temperature, hydrogen bromide is a colourless gas, like all the hydrogen halides apart from hydrogen fluoride, since hydrogen cannot form strong hydrogen bonds to the large and only mildly electronegative bromine atom; however, weak hydrogen bonding is present in solid crystalline hydrogen bromide at low temperatures, similar to the hydrogen fluoride structure, before disorder begins to prevail as the temperature is raised. Aqueous hydrogen bromide is known as hydrobromic acid, which is a strong acid (pK_a = −9) because the hydrogen bonds to bromine are too weak to inhibit dissociation. The HBr/H₂O system also involves many hydrates HBr·nH₂O for n = 1, 2, 3, 4, and 6, which are essentially salts of bromine anions and hydronium cations. Hydrobromic acid forms an azeotrope with boiling point 124.3 °C at 47.63 g HBr per 100 g solution; thus hydrobromic acid cannot be concentrated beyond this point by distillation.

Unlike hydrogen fluoride, anhydrous liquid hydrogen bromide is difficult to work with as a solvent, because its boiling point is low, it has a small liquid range, its dielectric constant is low and it does not dissociate appreciably into H₂Br⁺ and HBr⁻
₂ ions – the latter, in any case, are much less stable than the bifluoride ions (HF⁻
₂) due to the very weak hydrogen bonding between hydrogen and bromine, though its salts with very large and weakly polarising cations such as Cs⁺ and NR⁺
₄ (R = Me, Et, Buⁿ) may still be isolated. Anhydrous hydrogen bromide is a poor solvent, only able to dissolve small molecular compounds such as nitrosyl chloride and phenol, or salts with very low lattice energies such as tetraalkylammonium halides.

Other binary bromides

Silver bromide (AgBr)

 

Nearly all elements in the periodic table form binary bromides. The exceptions are decidedly in the minority and stem in each case from one of three causes: extreme inertness and reluctance to participate in chemical reactions (the noble gases, with the exception of xenon in the very unstable XeBr₂); extreme nuclear instability hampering chemical investigation before decay and transmutation (many of the heaviest elements beyond bismuth); and having an electronegativity higher than bromine's (oxygen, nitrogen, fluorine, and chlorine), so that the resultant binary compounds are formally not bromides but rather oxides, nitrides, fluorides, or chlorides of bromine. (Nonetheless, nitrogen tribromide is named as a bromide as it is analogous to the other nitrogen trihalides.)

Bromination of metals with Br₂ tends to yield lower oxidation states than chlorination with Cl₂ when a variety of oxidation states is available. Bromides can be made by reaction of an element or its oxide, hydroxide, or carbonate with hydrobromic acid, and then dehydrated by mildly high temperatures combined with either low pressure or anhydrous hydrogen bromide gas. These methods work best when the bromide product is stable to hydrolysis; otherwise, the possibilities include high-temperature oxidative bromination of the element with bromine or hydrogen bromide, high-temperature bromination of a metal oxide or other halide by bromine, a volatile metal bromide, carbon tetrabromide, or an organic bromide. For example, niobium(V) oxide reacts with carbon tetrabromide at 370 °C to form niobium(V) bromide. Another method is halogen exchange in the presence of excess "halogenating reagent", for example:

FeCl₃ + BBr₃ (excess) → FeBr₃ + BCl₃

When a lower bromide is wanted, either a higher halide may be reduced using hydrogen or a metal as a reducing agent, or thermal decomposition or disproportionation may be used, as follows:

3 WBr₅ + Al thermal gradient→475°C → 240°C 3 WBr₄ + AlBr₃

EuBr₃ + 1/2 H₂ → EuBr₂ + HBr

2 TaBr₄ 500°C→  TaBr₃ + TaBr₅

Most of the bromides of the pre-transition metals (groups 1, 2, and 3, along with the lanthanides and actinides in the +2 and +3 oxidation states) are mostly ionic, while nonmetals tend to form covalent molecular bromides, as do metals in high oxidation states from +3 and above. Silver bromide is very insoluble in water and is thus often used as a qualitative test for bromine.

Bromine halides

The halogens form many binary, diamagnetic interhalogen compounds with stoichiometries XY, XY₃, XY₅, and XY₇ (where X is heavier than Y), and bromine is no exception. Bromine forms a monofluoride and monochloride, as well as a trifluoride and pentafluoride. Some cationic and anionic derivatives are also characterised, such as BrF⁻
₂, BrCl⁻
₂, BrF⁺
₂, BrF⁺
₄, and BrF⁺
₆. Apart from these, some pseudohalides are also known, such as cyanogen bromide (BrCN), bromine thiocyanate (BrSCN), and bromine azide (BrN₃).

The pale-brown bromine monofluoride (BrF) is unstable at room temperature, disproportionating quickly and irreversibly into bromine, bromine trifluoride, and bromine pentafluoride. It thus cannot be obtained pure. It may be synthesised by the direct reaction of the elements, or by the comproportionation of bromine and bromine trifluoride at high temperatures. Bromine monochloride (BrCl), a red-brown gas, quite readily dissociates reversibly into bromine and chlorine at room temperature and thus also cannot be obtained pure, though it can be made by the reversible direct reaction of its elements in the gas phase or in carbon tetrachloride. Bromine monofluoride in ethanol readily leads to the monobromination of the aromatic compounds PhX (para-bromination occurs for X = Me, Bu^t, OMe, Br; meta-bromination occurs for the deactivating X = –CO₂Et, –CHO, –NO₂); this is due to heterolytic fission of the Br–F bond, leading to rapid electrophilic bromination by Br⁺.

At room temperature, bromine trifluoride (BrF₃) is a straw-coloured liquid. It may be formed by directly fluorinating bromine at room temperature and is purified through distillation. It reacts explosively with water and hydrocarbons, but is a less violent fluorinating reagent than chlorine trifluoride. It reacts vigorously with boron, carbon, silicon, arsenic, antimony, iodine, and sulfur to give fluorides, and also reacts with most metals and their oxides: as such, it is used to oxidise uranium to uranium hexafluoride in the nuclear industry. Refractory oxides tend to be only partially fluorinated, but here the derivatives KBrF₄ and BrF₂SbF₆ remain reactive. Bromine trifluoride is a useful nonaqueous ionising solvent, since it readily dissociates to form BrF⁺
₂ and BrF⁻
₄ and thus conducts electricity.

Bromine pentafluoride (BrF₅) was first synthesised in 1930. It is produced on a large scale by direct reaction of bromine with excess fluorine at temperatures higher than 150 °C, and on a small scale by the fluorination of potassium bromide at 25 °C. It is a very vigorous fluorinating agent, although chlorine trifluoride is still more violent. Bromine pentafluoride explodes on reaction with water and fluorinates silicates at 450 °C.

Polybromine compounds

Although dibromine is a strong oxidising agent with a high first ionisation energy, very strong oxidisers such as peroxydisulfuryl fluoride (S₂O₆F₂) can oxidise it to form the cherry-red Br⁺
₂ cation. A few other bromine cations are known, namely the brown Br⁺
₃ and dark brown Br⁺
₅. The tribromide anion, Br⁻
₃, has also been characterised; it is analogous to triiodide.

Bromine oxides and oxoacids

Bromine oxides are not as well-characterised as chlorine oxides or iodine oxides, as they are all fairly unstable: it was once thought that they could not exist at all. Dibromine monoxide is a dark-brown solid which, while reasonably stable at −60 °C, decomposes at its melting point of −17.5 °C; it is useful in bromination reactions and may be made from the low-temperature decomposition of bromine dioxide in a vacuum. It oxidises iodine to iodine pentoxide and benzene to 1,4-benzoquinone; in alkaline solutions, it gives the hypobromite anion.

So-called "bromine dioxide", a pale yellow crystalline solid, may be better formulated as bromine perbromate, BrOBrO₃. It is thermally unstable above −40 °C, violently decomposing to its elements at 0 °C. Dibromine trioxide, syn-BrOBrO₂, is also known; it is the anhydride of hypobromous acid and bromic acid. It is an orange crystalline solid which decomposes above −40 °C; if heated too rapidly, it explodes around 0 °C. A few other unstable radical oxides are also known, as are some poorly characterised oxides, such as dibromine pentoxide, tribromine octoxide, and bromine trioxide.

The four oxoacids, hypobromous acid (HOBr), bromous acid (HOBrO), bromic acid (HOBrO₂), and perbromic acid (HOBrO₃), are better studied due to their greater stability, though they are only so in aqueous solution. When bromine dissolves in aqueous solution, the following reactions occur:

Br2 + H2O	⇌ HOBr + H+ + Br−	Kac = 7.2 × 10−9 mol2 l−2
Br2 + 2 OH−	⇌ OBr− + H2O + Br−	Kalk = 2 × 108 mol−1 l

Hypobromous acid is unstable to disproportionation. The hypobromite ions thus formed disproportionate readily to give bromide and bromate:

3 BrO− ⇌ 2 Br− + BrO− 3

K = 1015

Bromous acids and bromites are very unstable, although the strontium and barium bromites are known. More important are the bromates, which are prepared on a small scale by oxidation of bromide by aqueous hypochlorite, and are strong oxidising agents. Unlike chlorates, which very slowly disproportionate to chloride and perchlorate, the bromate anion is stable to disproportionation in both acidic and aqueous solutions. Bromic acid is a strong acid. Bromides and bromates may comproportionate to bromine as follows:

BrO⁻
₃ + 5 Br⁻ + 6 H⁺ → 3 Br₂ + 3 H₂O

There were many failed attempts to obtain perbromates and perbromic acid, leading to some rationalisations as to why they should not exist, until 1968 when the anion was first synthesised from the radioactive beta decay of unstable ⁸³SeO²⁻
₄. Today, perbromates are produced by the oxidation of alkaline bromate solutions by fluorine gas. Excess bromate and fluoride are precipitated as silver bromate and calcium fluoride, and the perbromic acid solution may be purified. The perbromate ion is fairly inert at room temperature but is thermodynamically extremely oxidising, with extremely strong oxidising agents needed to produce it, such as fluorine or xenon difluoride. The Br–O bond in BrO⁻
₄ is fairly weak, which corresponds to the general reluctance of the 4p elements (especially arsenic, selenium, and bromine) to attain their maximum possible oxidation state, as they come after the scandide contraction characterised by the poor shielding afforded by the radial-nodeless 3d orbitals.

Organobromine compounds

Structure of N-bromosuccinimide, a common brominating reagent in organic chemistry

 

Like the other carbon–halogen bonds, the C–Br bond is a common functional group that forms part of core organic chemistry. Formally, compounds with this functional group may be considered organic derivatives of the bromide anion. Due to the difference of electronegativity between bromine (2.96) and carbon (2.55), the carbon in a C–Br bond is electron-deficient and thus electrophilic. The reactivity of organobromine compounds resembles but is intermediate between the reactivity of organochlorine and organoiodine compounds. For many applications, organobromides represent a compromise of reactivity and cost.

Organobromides are typically produced by additive or substitutive bromination of other organic precursors. Bromine itself can be used, but due to its toxicity and volatility safer brominating reagents are normally used, such as N-bromosuccinimide. The principal reactions for organobromides include dehydrobromination, Grignard reactions, reductive coupling, and nucleophilic substitution.

Organobromides are the most common organohalides in nature, even though the concentration of bromide is only 0.3% of that for chloride in sea water, because of the easy oxidation of bromide to the equivalent of Br⁺, a potent electrophile. The enzyme bromoperoxidase catalyzes this reaction. The oceans are estimated to release 1–2 million tons of bromoform and 56,000 tons of bromomethane annually.

Bromine addition to alkene reaction mechanism

 

An old qualitative test for the presence of the alkene functional group is that alkenes turn brown aqueous bromine solutions colourless, forming a bromohydrin with some of the dibromoalkane also produced. The reaction passes through a short-lived strongly electrophilic bromonium intermediate. This is an example of a halogen addition reaction.

Occurrence and production

View of salt evaporation pans on the Dead Sea, where Jordan (right) and Israel (left) produce salt and bromine 31°9′0″N 35°27′0″E

 

Bromine is significantly less abundant in the crust than fluorine or chlorine, comprising only 2.5 parts per million of the Earth's crustal rocks, and then only as bromide salts. It is the forty-sixth most abundant element in Earth's crust. It is significantly more abundant in the oceans, resulting from long-term leaching. There, it makes up 65 parts per million, corresponding to a ratio of about one bromine atom for every 660 chlorine atoms. Salt lakes and brine wells may have higher bromine concentrations: for example, the Dead Sea contains 0.4% bromide ions. It is from these sources that bromine extraction is mostly economically feasible.

The main sources of bromine are in the United States and Israel. The element is liberated by halogen exchange, using chlorine gas to oxidise Br⁻ to Br₂. This is then removed with a blast of steam or air, and is then condensed and purified. Today, bromine is transported in large-capacity metal drums or lead-lined tanks that can hold hundreds of kilograms or even tonnes of bromine. The bromine industry is about one-hundredth the size of the chlorine industry. Laboratory production is unnecessary because bromine is commercially available and has a long shelf life.

Applications

A wide variety of organobromine compounds are used in industry. Some are prepared from bromine and others are prepared from hydrogen bromide, which is obtained by burning hydrogen in bromine.

Flame retardants

Tetrabromobisphenol A

 

Brominated flame retardants represent a commodity of growing importance, and make up the largest commercial use of bromine. When the brominated material burns, the flame retardant produces hydrobromic acid which interferes in the radical chain reaction of the oxidation reaction of the fire. The mechanism is that the highly reactive hydrogen radicals, oxygen radicals, and hydroxy radicals react with hydrobromic acid to form less reactive bromine radicals (i.e., free bromine atoms). Bromine atoms may also react directly with other radicals to help terminate the free radical chain-reactions that characterise combustion.

To make brominated polymers and plastics, bromine-containing compounds can be incorporated into the polymer during polymerisation. One method is to include a relatively small amount of brominated monomer during the polymerisation process. For example, vinyl bromide can be used in the production of polyethylene, polyvinyl chloride or polypropylene. Specific highly brominated molecules can also be added that participate in the polymerisation process For example, tetrabromobisphenol A can be added to polyesters or epoxy resins, where it becomes part of the polymer. Epoxys used in printed circuit boards are normally made from such flame retardant resins, indicated by the FR in the abbreviation of the products (FR-4 and FR-2). In some cases the bromine containing compound may be added after polymerisation. For example, decabromodiphenyl ether can be added to the final polymers.

A number of gaseous or highly volatile brominated halomethane compounds are non-toxic and make superior fire suppressant agents by this same mechanism, and are particular effective in enclosed spaces such as submarines, airplanes, and spacecraft. However, they are expensive and their production and use has been greatly curtailed due to their effect as ozone-depleting agents. They are no longer used in routine fire extinguishers, but retain niche uses in aerospace and military automatic fire-suppression applications. They include bromochloromethane (Halon 1011, CH₂BrCl), bromochlorodifluoromethane (Halon 1211, CBrClF₂), and bromotrifluoromethane (Halon 1301, CBrF₃).

Other uses

Baltimore's Emerson Bromo-Seltzer Tower, originally part of the headquarters of Emerson Drug Company, which made Bromo-Seltzer

 

Silver bromide is used, either alone or in combination with silver chloride and silver iodide, as the light sensitive constituent of photographic emulsions.

Ethylene bromide was an additive in gasolines containing lead anti-engine knocking agents. It scavenges lead by forming volatile lead bromide, which is exhausted from the engine. This application accounted for 77% of the bromine use in 1966 in the US. This application has declined since the 1970s due to environmental regulations.

Poisonous bromomethane was widely used as pesticide to fumigate soil and to fumigate housing, by the tenting method. Ethylene bromide was similarly used. These volatile organobromine compounds are all now regulated as ozone depletion agents. The Montreal Protocol on Substances that Deplete the Ozone Layer scheduled the phase out for the ozone depleting chemical by 2005, and organobromide pesticides are no longer used (in housing fumigation they have been replaced by such compounds as sulfuryl fluoride, which contain neither the chlorine or bromine organics which harm ozone). Before the Montreal protocol in 1991 (for example) an estimated 35,000 tonnes of the chemical were used to control nematodes, fungi, weeds and other soil-borne diseases.

In pharmacology, inorganic bromide compounds, especially potassium bromide, were frequently used as general sedatives in the 19th and early 20th century. Bromides in the form of simple salts are still used as anticonvulsants in both veterinary and human medicine, although the latter use varies from country to country. For example, the U.S. Food and Drug Administration (FDA) does not approve bromide for the treatment of any disease, and it was removed from over-the-counter sedative products like Bromo-Seltzer, in 1975. Commercially available organobromine pharmaceuticals include the vasodilator nicergoline, the sedative brotizolam, the anticancer agent pipobroman, and the antiseptic merbromin. Otherwise, organobromine compounds are rarely pharmaceutically useful, in contrast to the situation for organofluorine compounds. Several drugs are produced as the bromide (or equivalents, hydrobromide) salts, but in such cases bromide serves as an innocuous counterion of no biological significance.

Other uses of organobromine compounds include high-density drilling fluids, dyes (such as Tyrian purple and the indicator bromothymol blue), and pharmaceuticals. Bromine itself, as well as some of its compounds, are used in water treatment, and is the precursor of a variety of inorganic compounds with an enormous number of applications (e.g. silver bromide for photography). Zinc–bromine batteries are hybrid flow batteries used for stationary electrical power backup and storage; from household scale to industrial scale.

Biological role and toxicity

Bromine

Hazards
GHS pictograms
GHS signal word	Danger
GHS hazard statements	H314, H330, H400
GHS precautionary statements	P260, P273, P280, P284, P305+351+338, P310
NFPA 704	0 3 0

2-Octyl 4-bromo-3-oxobutanoate, an organobromine compound found in mammalian cerebrospinal fluid

 

A 2014 study suggests that bromine (in the form of bromide ion) is a necessary cofactor in the biosynthesis of collagen IV, making the element essential to basement membrane architecture and tissue development in animals. Nevertheless, no clear deprivation symptoms or syndromes have been documented. In other biological functions, bromine may be non-essential but still beneficial when it takes the place of chlorine. For example, in the presence of hydrogen peroxide, H₂O₂, formed by the eosinophil, and either chloride or bromide ions, eosinophil peroxidase provides a potent mechanism by which eosinophils kill multicellular parasites (such as, for example, the nematode worms involved in filariasis) and some bacteria (such as tuberculosis bacteria). Eosinophil peroxidase is a haloperoxidase that preferentially uses bromide over chloride for this purpose, generating hypobromite (hypobromous acid), although the use of chloride is possible.

Although α-haloesters are generally thought of as highly reactive, and therefore, toxic intermediates in organic synthesis, mammals, including humans, cats, and rats, appear to biosynthesize traces of an α-bromoester, 2-octyl 4-bromo-3-oxobutanoate, which is found in their cerebrospinal fluid and appears to play a yet unclarified role in inducing REM sleep. Neutrophil myeloperoxidase can use H₂O₂ and Br^- to brominate deoxycytidine, which could result in DNA mutations. Marine organisms are the main source of organobromine compounds, and it is in these organisms that the essentiality of bromine is on much firmer ground. More than 1600 such organobromine compounds were identified by 1999. The most abundant is methyl bromide (CH₃Br), of which an estimated 56,000 tonnes is produced by marine algae each year. The essential oil of the Hawaiian alga Asparagopsis taxiformis consists of 80% bromoform. Most of such organobromine compounds in the sea are made by the action of a unique algal enzyme, vanadium bromoperoxidase.

The bromide anion is not very toxic: a normal daily intake is 2 to 8 milligrams. However, high levels of bromide chronically impair the membrane of neurons, which progressively impairs neuronal transmission, leading to toxicity, known as bromism. Bromide has an elimination half-life of 9 to 12 days, which can lead to excessive accumulation. Doses of 0.5 to 1 gram per day of bromide can lead to bromism. Historically, the therapeutic dose of bromide is about 3 to 5 grams of bromide, thus explaining why chronic toxicity (bromism) was once so common. While significant and sometimes serious disturbances occur to neurologic, psychiatric, dermatological, and gastrointestinal functions, death from bromism is rare. Bromism is caused by a neurotoxic effect on the brain which results in somnolence, psychosis, seizures and delirium.

Elemental bromine is toxic and causes chemical burns on human flesh. Inhaling bromine gas results in similar irritation of the respiratory tract, causing coughing, choking, and shortness of breath, and death if inhaled in large enough amounts. Chronic exposure may lead to frequent bronchial infections and a general deterioration of health. As a strong oxidising agent, bromine is incompatible with most organic and inorganic compounds. Caution is required when transporting bromine; it is commonly carried in steel tanks lined with lead, supported by strong metal frames. The Occupational Safety and Health Administration (OSHA) of the United States has set a permissible exposure limit (PEL) for bromine at a time-weighted average (TWA) of 0.1 ppm. The National Institute for Occupational Safety and Health (NIOSH) has set a recommended exposure limit (REL) of TWA 0.1 ppm and a short-term limit of 0.3 ppm. The exposure to bromine immediately dangerous to life and health (IDLH) is 3 ppm. Bromine is classified as an extremely hazardous substance in the United States as defined in Section 302 of the U.S. Emergency Planning and Community Right-to-Know Act (42 U.S.C. 11002), and is subject to strict reporting requirements by facilities which produce, store, or use it in significant quantities.