Halogen Halogen

8/8/2019 Halogen Halogen

1/36

KIMIA

HALOGEN

KHAIRUNNISA NASUTIONLUPITA YESSICA

RISKI SUWISTIANISARIDHO RAYENDRA

XII IPA 4

8/8/2019 Halogen Halogen

2/36

Halogen

The halogens orhalogen elements are a series ofnonmetal elements from Group 17 IUPACStyle (formerly: VII, VIIA) of theperiodic table, comprising fluorine(F),chlorine (Cl),bromine

(Br), iodine (I), and astatine (At). The artificially created element 117, provisionally referred toby the systematic name ununseptium, may also be a halogen. The group of halogens is the only

periodic table group which contains elements in all three familiar states of matter at standardtemperature and pressure.

Abundance

Owing to their high reactivity, the halogens are found in the environment only in compounds or

as ions. Halide ions and oxoanionssuch as iodate (IO3) can be found in many minerals and in

seawater. Halogenated organic compounds can also be found as natural products in living

organisms. In their elemental forms, the halogens exist as diatomic molecules, but these only

have a fleeting existence in nature and are much more common in the laboratory and in industry.At room temperature and pressure, fluorine and chlorine are gases, bromine is a liquid and iodine

and astatine are solids; Group 17 is therefore the only periodic table group exhibiting all three

states of matterat room temperature.

Properties

Fluorine, (F); chlorine, (Cl); bromine, (Br); iodine, (I) at room temperature. The first two aregaseous, the third is liquid and the fourth is solid.

Like other groups, the candidates of this family show patterns in its electron configuration,

especially the outermost shells resulting in trends in chemical behavior:

Z ElementNo. of

electrons/shell

9 fluorine 2, 7

17 chlorine 2, 8, 7

35 bromine 2, 8, 18, 7

53 iodine 2, 8, 18, 18, 7

85 astatine 2, 8, 18, 32, 18, 7

8/8/2019 Halogen Halogen

3/36

The halogens show a series of trends when moving down the groupfor instance, decreasing

electronegativity and reactivity, and increasing melting andboiling point.

Halogen

Standard

Atomic

Weight(u)

Melting

Point

(K)

Boiling

Point

(K)

Electronegativity

(Pauling)

Fluorine 18.998 53.53 85.03 3.98

Chlorine 35.453 171.60 239.11 3.16

Bromine 79.904 265.80 332.00 2.96

Iodine 126.904 386.85 457.40 2.66

Astatine (210) 575 610 (?) 2.20

Diatomic halogen molecules

halogen molecule structure modeld(XX) / pm

(gas phase)

d(XX) / pm

(solid phase)

fluorine F2 143 149

chlorine Cl2 199 198

bromine Br2 228 227

iodine I2 266 272

The elements become less reactive and have higher melting points as the atomic number

increases.

8/8/2019 Halogen Halogen

4/36

Chemistry

Reactivity

Halogens are highly reactive, and as such can be harmful or lethal to biological organisms in

sufficient quantities. This high reactivity is due to the atoms being highly electronegativedue totheir high effective nuclear charge. They can gain an electron by reacting with atoms of other

elements. Fluorine is one of the most reactive elements in existence, attacking otherwise inert

materials such as glass, and forming compounds with the heaviernoble gases. It is a corrosive

and highly toxic gas. The reactivity of fluorine is such that if used or stored in laboratoryglassware, it can react with glass in the presence of small amounts of water to form silicon

tetrafluoride (SiF4). Thus fluorine must be handled with substances such as Teflon (which is

itself an organofluorine compound), extremely dry glass, or metals such as copper or steel whichform a protective layer of fluoride on their surface.

The high reactivity of fluorine means that once it does react with something, it bonds with it so

strongly that the resulting molecule is very inert and non-reactive to anything else. For example,Teflon is fluorine bonded with carbon.

Both chlorine and bromine are used as disinfectants for drinking water, swimming pools, fresh

wounds, spas, dishes, and surfaces. They kill bacteria and other potentially harmful

microorganisms through a process known as sterilization. Their reactivity is also put to use in

bleaching. Sodium hypochlorite, which is produced from chlorine, is the active ingredient ofmost fabricbleaches and chlorine-derived bleaches are used in the production of some paper

products.

Hydrogen halides

The halogens all form binary compounds with hydrogen known as the hydrogen halides (HF,

HCl, HBr, HI, and HAt), a series of particularly strong acids. When in aqueous solution, the

hydrogen halides are known ashydrohalic acids. HAt, or "hydroastatic acid", should also qualify,but it is not typically included in discussions of hydrohalic acid due to astatine's extreme

instability toward alpha decay.

Interhalogen compounds

The halogens react with each other to form interhalogen compounds. Diatomic interhalogen

compounds such as BrF, ICl, and ClF bear resemblance to the pure halogens in some respects.

The properties and behaviour of a diatomic interhalogen compound tend to be intermediatebetween those of its parent halogens. Some properties, however, are found in neither parent

halogen. For example, Cl2 and I2 are soluble in CCl4, but ICl is not since it is apolar molecule

due to the relatively largeelectronegativitydifference between I and Cl.

8/8/2019 Halogen Halogen

5/36

Organohalogen compounds

Many synthetic organic compounds such as plasticpolymers, and a few natural ones, contain

halogen atoms; these are known as halogenatedcompounds ororganic halides. Chlorine is by farthe most abundant of the halogens, and the only one needed in relatively large amounts (as

chloride ions) by humans. For example, chloride ions play a key role in brain function bymediating the action of the inhibitory transmitterGABA and are also used by the body toproduce stomach acid. Iodine is needed in trace amounts for the production ofthyroidhormones

such as thyroxine. On the other hand, neither fluorine nor bromine are believed to be essential for

humans.

Polyhalogenated compounds

Polyhalogenated compounds are industrially created compounds substituted with multiple

halogens. Many of them are very toxic and bioaccumulate in humans, and have a very wide

application range. They include the much maligned PCB's, PBDE's, and PFC's as well asnumerous other compounds.

Reactivity with water

Fluorine reacts vigorously with water to produce oxygen(O2) and hydrogen fluoride (HF):[3]

2 F2(g) + 2 H2O(l) O2(g) + 4 HF(aq)

Chlorine has minimal solubility of 0.7g Cl2 per kg of water at ambient temperature (21oC).[4]

Dissolved chlorine reacts to formhydrochloric acid (HCl) and hypochlorous acid, a solution that

can be used as a disinfectantorbleach:

Cl2(g) + H2O(l) HCl(aq) + HClO(aq)

Bromine has a solubility of 3.41 g per 100 g of water, [5] but it slowly reacts to form hydrogen

bromide (HBr) and hypobromous acid (HBrO):

Br2(g) + H2O(l) HBr(aq) + HBrO(aq)

Iodine, however, is minimally soluble in water (0.03 g/100 g water @ 20 C) and does not react

with it.[6]However, iodine will form an aqueous solution in the presence of iodide ion, such as by

addition ofpotassium iodide (KI), because the triiodide ion is formed.

8/8/2019 Halogen Halogen

6/36

Fluorine

oxygen fluorine neon

-

F

Cl

9FPeriodic table

Appearance

Tan or Yellow gas

General properties

Name,symbol, number fluorine, F, 9

Pronunciation /f lrn/

Element category halogenGroup, period,block 17,2,p

Standard atomic weight 18.9984032(5) gmol1

Electron configuration 1s2 2s2 2p5

Electrons per shell 2, 7 (Image)

Physical properties

Phase gas

Density(0 C, 101.325 kPa)

1.7 g/L

Liquid density at b.p. 1.108 gcm3

8/8/2019 Halogen Halogen

7/36

Melting point 53.53 K219.62 ,C363.32 ,F

Boiling point 85.03 K188.12 ,C306.62 ,F

Critical point 144.13 K, 5.172 MPa

Heat of fusion (F2) 0.510kJmol1

Heat of vaporization (F2) 6.62 kJmol1

Specific heat capacity(25 C) (F2)31.304 Jmol1K1

Vapor pressure

P (Pa) 1 10 100 1 k 10 k 100 k

at T (K) 38 44 50 58 69 85

Atomic properties

Oxidation states1

(Weaklyacidic oxide)

Electronegativity 3.98 (Pauling scale)

Ionization energies

(more)

1st: 1681.0 kJmol1

2nd: 3374.2 kJmol1

3rd: 6050.4 kJmol1

Covalent radius 573 pm(seecovalent radius of fluorine)

Van der Waals radius 147 pm

Miscellanea

Crystal structure cubic

Magnetic ordering nonmagnetic

Thermal conductivity (300 K) 27.7 m Wm1K1

CAS registry number 7782-41-4

Most stable isotopes

Main article:Isotopes of fluorine

iso NA half-life DM DE(MeV) DP

18F syn 109.77 min+ (97%) 0.64 18O

(3%) 1.656 18O19F 100% 19F is stable with 10 neutronsvde

Fluorine is the chemical element withatomic number9, represented by the symbol F. Fluorine

forms a single bond with itself in elemental form, resulting in the diatomic F2 molecule. F2(fluorine) is a supremely reactive, poisonous, pale, yellowish brown gas. Elemental fluorine is

the most chemically reactive and electronegativeof all the elements. For example, it will readily"burn" hydrocarbons at room temperature, in contrast to the combustion of hydrocarbons by

oxygen, which requires an input of energy with a spark. Therefore, molecular fluorine is highly

dangerous, more so than otherhalogens such as the poisonous chlorinegas.

Fluorine's highest electronegativity and small atomic radius give unique properties to many of its

compounds. For example, the enrichment of 235U, the principal nuclear fuel, relies on the

volatility of UF6. Also, the carbonfluorine bond is one of the strongest bonds in organic

chemistry. This contributes to the stability and persistence offluoroalkane based organofluorinecompounds, such as PTFE/(Teflon) and PFOS. The carbonfluorine bond's inductive effects

result in the strength of many fluorinated acids, such as triflic acidandtrifluoroacetic acid. Drugs

8/8/2019 Halogen Halogen

8/36

are often fluorinated at biologically reactive positions, to prevent their metabolism and prolong

their half-lives.

Characteristics

Physical characteristics

F2 is a corrosive pale yellow or brown[1]gas that is a powerful oxidizing agent. It is the most

reactive and most electronegative of all the elements on the classic Pauling scale (4.0), and

readily forms compounds with most other elements. It is found in the -1 oxidation state, exceptwhen bonded to another fluorine in F2 which gives it an oxidation number of 0. Fluorine

combines with the noble gasesargon, krypton,xenon, andradon. Even in dark, cool conditions,

fluorine reacts explosively with hydrogen. The reaction with hydrogen can occur at extremelylow temperatures, using liquid hydrogen and solid fluorine. It is so reactive that metals, water, aswell as most othersubstances, burn with a bright flame in a jet of fluorine gas. In moist air, it

reacts with water to form the also dangerous hydrofluoric acid.

Chemical characteristics

Fluorine as a freely reacting oxidant gives the strongest oxidants known.

Occurrence

Fluorite (CaF2) crystals

Fluorine is obtained from two sources, mainly. The production of phosphate fertilizers from

apatite generates a byproduct of hydrogen fluoride that is collected. The other main source is themineral fluorite, CaF2, which is widespread.

8/8/2019 Halogen Halogen

9/36

Organofluorine compounds are naturally rare compounds. Though F2 is too reactive to have any

natural biological role, fluorine is incorporated into compounds with biological activity.

Naturally occurring organofluorine compounds are rare, the most notable example isfluoroacetate, which functions as a plant defence against herbivores in at least 40 plants in

Australia, Brazil, and Africa.[2] The enzyme adenosyl-fluoride synthase catalyzes the formation

of 5'-deoxy-5'-fluoroadenosine. Fluorine is not an essential nutrient, but its importance inpreventing tooth decay is well-recognized.[3] The effect is predominantly topical, although prior

to 1981 it was considered primarily systemic (occurring through ingestion).[4]

From the perspective of cosmology, fluorine is relatively rare because the solar temperatures

needed to make it also enable it to quickly fuse with hydrogen to form oxygen and helium , orwith helium to become neon. Most fluorine is created either during a supernova when a neutrino

hits an atom of neon, or when a blue Wolf-Rayet star massing over 40 solar masses has a stellar

wind blowing the fluorine out of the star before hydrogen or helium can destroy it. [5]

Isotopes

Although fluorine has multiple isotopes, only one of these isotopes (19F) is stable, and the othershave short half-lives and are not found in nature. Fluorine is thus a mononuclidic element.

The nuclide 18F is the radionuclide of fluorine with the longest half life (about 110 minutes), and

commercially is an important source ofpositrons, finding its major use in positron emission

tomography scanning.

Chemistry of fluorine compounds

Fluorine forms a variety of very different compounds, owing to its small atomic size and

covalent behavior. Elemental fluorine is a dangerously powerful oxidant, reflecting the extremeelectronegativity of fluorine. Hydrofluoric acid is extremely dangerous, whereas, in syntheticdrugs incorporating an aromatic ring (e.g.,flumazenil), fluorine is used to help prevent toxication

or to delay metabolism[citation needed].

Fluorides are compounds that combine fluorine with some positively charged counterpart. They

often consist of crystalline ionic salts. Fluorine compounds with metals are among the moststable of salts.

Hydrogen fluoride is a weak acid when dissolved in water, but is still very corrosive and attacks

glass. Consequently, fluorides of alkali metalsproduce basic solutions. For example, a 1 M

solution of NaF in water has a pH of 8.59 compared to a 1 M solution of NaOH, a strong base,which has a pH of 14.00.[6]

The fluoride ion is basic, therefore hydrofluoric acid is a weak acid in water solution. However,

water is not an inert solvent in this case: When less basic solvents such as anhydrous acetic acid

are used, hydrofluoric acid is the strongest of the hydrohalogenic acids. Also, owing to thebasicity of the fluoride ion, soluble fluorides give basic water solutions. The fluoride ion is a

Lewis base, and has a high affinity to certain elements such as calcium and silicon. For example,

deprotection of silicon protecting groups is achieved with a fluoride. The fluoride ion ispoisonous.

8/8/2019 Halogen Halogen

10/36

Noble gas compounds

The reactivity of fluorine toward the noble gas xenon was first reported by Neil Bartlett in 1962.

Fluorides ofkryptonand radon have also been prepared. Argon fluorohydridehas been observed

at cryogenic temperatures.

Organofluorine compounds

The carbon-fluoride bond is covalent and very stable. The use of a fluorocarbon polymer,

poly(tetrafluoroethene) or Teflon, is an example: It is thermostable and waterproof enough to beused in frying pans. Organofluorines may be safely used in applications such as drugs, without

the risk of release of toxic fluoride. In synthetic drugs,toxication can be prevented. For example,

anaromatic ring is useful but presents a safety problem: enzymes in the body metabolize some ofthem into poisonous epoxides. When thepara position is substituted with fluorine, the aromatic

ring is protected and epoxide is no longer produced.

The substitution of fluorine forhydrogen in organic compounds offers a very large number ofcompounds. These compounds often feature single C-F units, but -CF3 and -OCF3 group providefurther variation, and more recently the -SF5 group.

[7]

Chlorine

sulfur chlorine argon

F

Cl

Br

17ClPeriodic table

8/8/2019 Halogen Halogen

11/36

Appearance

pale yellow-green gas

General properties

Name,symbol, number chlorine, Cl, 17

Pronunciation /k l r in /KLOR-een

Element category Halogen

Group, period,block 17,3,p

Standard atomic weight 35.453(2) gmol1

Electron configuration [Ne] 3s2 3p5

Electrons per shell 2, 8, 7 (Image)

Physical properties

Phase gas

Density(0 C, 101.325 kPa)

3.2 g/L

Liquid density at b.p. 1.5625[1] gcm3

Melting point 171.6 K-101.5 ,C-150.7 ,F

Boiling point 239.11 K-34.04 ,C-29.27 ,F

Critical point 416.9 K, 7.991 MPa

Heat of fusion (Cl2) 6.406kJmol

1

Heat of vaporization (Cl2) 20.41kJmol

1

Specific heat capacity(25 C) (Cl2)

33.949 Jmol1K1

Vapor pressure

P (Pa) 1 10 100 1 k 10 k 100 k

at T (K) 128 139 153 170 197 239

Atomic properties

Oxidation states7, 6, 5, 4, 3, 2, 1, -1(strongly acidicoxide)

Electronegativity 3.16 (Pauling scale)

Ionization energies

(more)

1st: 1251.2 kJmol1

2nd: 2298 kJmol1

3rd: 3822 kJmol1

Covalent radius 1024pm

Van der Waals radius 175pm

Miscellanea

Crystal structure orthorhombic

Magnetic ordering diamagnetic[2]

Electrical resistivity (20 C) > 10 m

8/8/2019 Halogen Halogen

12/36

Thermal conductivity (300 K) 8.9x10-3 Wm1K1

Speed of sound (gas, 0 C) 206 m/s

CAS registry number 7782-50-5

Most stable isotopes

Main article:Isotopes of chlorine

iso NA half-life DM DE(MeV) DP35Cl 75.77% 35Cl isstable with 18 neutrons

36Cl trace 3.01105 y 0.709 36Ar

- 36S37Cl 24.23% 37Cl isstable with 20 neutronsvde

Chlorine is the chemical element with atomic number17 and symbol Cl. It is a halogen, found

in theperiodic table ingroup 17 (formerly VII, VIIa, or VIIb). As the chlorideion, which is part

ofcommon salt and other compounds, it is abundant in nature and necessary to most forms of

life, including humans. In its elemental form (Cl2 or "dichlorine") under standard conditions,chlorine is a powerfuloxidant and is used inbleaching and disinfectants, as well as an essential

reagent in the chemical industry. As a common disinfectant, chlorine compounds are used inswimming pools to keep them clean and sanitary. In theupper atmosphere, chlorine-containing

molecules such as chlorofluorocarbons have been implicated in the destruction of the ozone

layer.

Characteristics

Physical characteristics

At standard temperature and pressure, two chlorine atoms form the diatomic moleculeCl2. Thisis a pale yellow-green gas that has its distinctive strong smell, the smell of bleach. The bonding

between the two atoms is relatively weak (only 242.580 0.004 kJ/mol), which makes the Cl2molecule highly reactive.

Chemical characteristics

Along with fluorine, bromine, iodine, andastatine, chlorine is a member of the halogenseriesthat forms the group 17 of the periodic table. Chlorine forms compounds with almost all of the

elements to give compounds that are usually called chlorides. Chlorine gas reacts with most

organic compounds, and will even sluggishly support the combustion ofhydrocarbons.[3]

Hydrolysis

At 10 C and atmospheric pressure, one liter of water dissolves 3.10 L of gaseous chlorine, [4]

Solutions of chlorine in water contain chlorine (Cl2), hydrochloric acid, and hypochlorous acid:

Cl2 + H2O HCl + HClO

8/8/2019 Halogen Halogen

13/36

This conversion to the right is called disproportionation, because the ingredient chlorine both

increases and decreases in formal oxidation state. The solubility of chlorine in water is increased

if the water contains dissolved alkali hydroxide, and in this way, chlorine bleach is produced.

Cl2 + 2 OH- ClO- + Cl- + H2O

Chlorine gas only exists in a neutral or acidic solution.

Compounds

Chlorine exists in all odd numbered oxidation states from 1 to +7, as well as the elemental state

of zero (see Table). Progressing through the states, hydrochloric acid can be oxidized usingmanganese dioxide, orhydrogen chloride gas oxidized catalytically by air to form elemental

chlorine gas.

Oxidation

stateName Formula Illustrative compounds

1 chlorides Cl ionic chlorides, organic chlorides, hydrochloric acid

0 chlorine Cl2 elemental chlorine

+1 hypochlorites ClO sodium hypochlorite, calcium hypochlorite

+3 chlorites ClO2 sodium chlorite

+4chlorine

dioxideClO2

+5 chlorates ClO3 sodium chlorate,potassium chlorate,chloric acid

+7 perchlorates ClO4 perchloric acid, perchlorate salts such as magnesiumperchlorate, dichlorine heptoxide

Interhalogen compounds

Chlorine oxidizes bromide and iodide salts to bromine and iodine, respectively. But it cannot

oxidize fluoride to fluorine. It makes are variety of "interhalogen compounds" such as thechlorine fluorides, chlorine monofluoride (ClF), chlorine trifluoride (ClF3), chlorine

pentafluoride (ClF5). Chlorides of bromine and iodine are also known.

Organo chlorine compounds

Chlorine is used extensively in organic in substitution and addition reactions. Chlorine often

imparts many desired properties to an organic compound, in part due to its electronegativity.Organochlorine compounds are also serious pollutants, either as side products of industrial

processes or as persistent pesticidess.

Many important industrial products are produced via organochlorine intermediates. Examplesinclude polycarbonates, polyurethanes, silicones, polytetrafluoroethylene, carboxymethyl

cellulose, and propylene oxide. Like the other halogens, chlorine participates in free-radical

substitution reactions with hydrogen-containing organic compounds. When applied to organic

substrates, reaction is oftenbut not invariablynon-regioselective, and, hence, may result in amixture of isomeric products. It is often difficult to control the degree of substitution as well, so

multiple substitutions are common. If the different reaction products are easily separated, e.g., by

8/8/2019 Halogen Halogen

14/36

distillation, substitutive free-radical chlorination (in some cases accompanied by concurrent

thermal dehydrochlorination) may be a useful synthetic route. Industrial examples of this are the

production ofmethyl chloride, methylene chloride, chloroform, and carbon tetrachloride frommethane, allyl chloridefrom propylene, and trichloroethylene, and tetrachloroethylene from 1,2-

dichloroethane.

Like the other halides, chlorine undergoes electrophilic additions reactions, the most notable one

being the chlorination of alkenes and aromatic compounds with a Lewis acid catalyst. Organicchlorine compounds tend to be less reactive in nucleophilic substitution reactions than the

corresponding bromine or iodine derivatives, but they tend to be cheaper. They may be activated

for reaction by substituting with a tosylate group, or by the use of a catalytic amount of sodiumiodide.

Chlorides

Chlorine combines with almost all elements to give chlorides. Compounds with oxygen,

nitrogen, xenon, and krypton are known, but do not form by direct reaction of the elements.[5]

Chloride is one of the most common anions in nature. Hydrogen chloride and its aqueous

solution, hydrochloric acid, are produced on megaton scale annually both as valued intermediates

but sometimes as undesirable pollutants.

Chlorine oxides

Chlorine forms a variety of oxides: chlorine dioxide (ClO2), dichlorine monoxide (Cl2O),

dichlorine heptoxide (Cl2O7). The anionic derivatives of these same oxides are also well known

including chlorate (ClO3), chlorite (ClO2),hypochlorite(ClO

), andperchlorate (ClO4), andchloramine (NH2Cl).

[6] The acid derivatives of these anions are hypochlorous acid (HOCl),

chloric acid(HClO3) andperchloric acid (HClO4).

In hot concentrated alkali solution hypochlorite disproportionates:

2 ClO Cl + ClO2

ClO + ClO2 Cl + ClO3

Sodium chlorate andpotassium chloratecan be crystallized from solutions formed by the above

reactions. If their crystals are heated, they undergo a further, final disproportionation:

4 ClO3 Cl + 3 ClO4

This same progression from chloride to perchlorate can be accomplished by electrolysis. The

anode reaction progression is:[7]

ReactionElectrode

potential

Cl + 2 OH ClO + H2O + 2 e +0.89 volts

ClO + 2 OH ClO2 + H2O + 2 e +0.67 volts

8/8/2019 Halogen Halogen

15/36

ClO2 + 2 OH ClO3 + H2O + 2 e +0.33 volts

ClO3 + 2 OH ClO4 + H2O + 2 e +0.35 volts

Each step is accompanied at the cathode by

2 H2O + 2 e

2 OH

+ H2 (0.83 volts)

Occurrence

In nature, chlorine is found primarily as the chloride ion, a component of the salt that is

deposited in the earth or dissolved in the oceans about 1.9% of the mass of seawater is

chloride ions. Even higher concentrations of chloride are found in the Dead Sea and inunderground brine deposits. Most chloride salts are soluble in water, thus, chloride-containing

minerals are usually only found in abundance in dry climates or deep underground. Common

chloride minerals include halite (sodium chloride), sylvite (potassium chloride), and carnallite

(potassium magnesium chloride hexahydrate). Over 2000 naturally occurring organic chlorinecompounds are known.[8]

In the interstellar medium, chlorine is produced insupernovaevia ther-process.[9]

Isotopes

Chlorine has a wide range ofisotopes, the two principal stableisotopes being 35Cl (75.77%) and37Cl (24.23%); they give chlorine atoms an apparent atomic weight of 35.4527 g/mol.

Trace amounts ofradioactive36Cl exist in the environment, in a ratio of about 7x1013 to 1 with

stable isotopes. 36Cl is produced in the atmosphere by spallation of 36Ar by interactions withcosmic rayprotons. In the subsurface environment, 36Cl is generated primarily as a result of

neutron captureby 35Cl ormuon capture by 40Ca. 36Cl decays to 36S and to 36Ar, with a combined

half-life of 308,000 years. The half-life of this hydrophilic nonreactive isotope makes it suitable

forgeologic dating in the range of 60,000 to 1 million years. Additionally, large amounts of 36Clwere produced by irradiation of seawater during atmospheric detonations of nuclear weapons

between 1952 and 1958. The residence time of36Cl in the atmosphere is about 1 week. Thus, asan event marker of 1950s water in soil and ground water, 36Cl is also useful for dating waters less

than 50 years before the present. 36Cl has seen use in other areas of the geological sciences,

including dating ice and sediments.

Laboratory methods

Small amounts of chlorine gas can be made in the laboratory by combining hydrochloric acid

and manganese dioxide. Alternatively a strong acid such as sulfuric acid or hydrochloric acid