Listen to this article

세슘

Caesium
세륨에 혼란스러워 하기로 했다.
세슘, 55Cs
Some pale gold metal, with a liquid-like texture and lustre, sealed in a glass ampoule
세슘
발음/ˈsiːziəm/ (SEE-zee-əm)
대체 이름세슘(미국)
외모창백한 금
표준 원자량Ar°(C)
  • 132.90545196±0.00000006
  • 132.91±0.01(요약한)[1]
세슘은 주기율 표 상의.
수소 헬륨
리튬 베릴륨 붕소 카본 질소 산소 불소 네온
나트륨 마그네슘 알루미늄 실리콘 인 유황 염소 아르곤
칼륨 칼슘 스칸듐 티타늄 바나듐 크롬 망간 철 코발트 니켈 구리 아연 갈륨 게르마늄 비소 셀레늄 브롬 크립톤
루비듐 스트론튬 이트륨 지르코늄 니오브 몰리브덴 테크네튬 루테늄 로듐 팔라듐 실버 카드뮴 인듐 주석 안티몬 텔루루 요오드 제논
세슘 바륨 란타넘 세륨 프라세오디뮴 네오디뮴 프로메튬 사마리움 유로피움 가돌리늄 터비움 디스프로슘 홀뮴 엘비움 툴륨 이터비움 루테튬 하프늄 탄탈룸 텅스텐 레늄 오스뮴 이리듐 플래티넘 골드 수은(원소) 탈륨 이끌다 비스무트 폴로늄 아스타틴 라돈
프랑슘 라듐 악티늄 토륨 프로탁티늄 우라늄 넵투늄 플루토늄 아메리슘 퀴륨 베르켈륨 칼리포늄 아인스타이늄 페르미움 멘델레비움 노벨륨 로렌슘 러더포디움 두브늄 시보르기움 보리움 하시움 마이트네리움 다름슈타디움 뢴트제늄 코페르니슘 니혼리움 플레로비움 모스코비움 리버모리움 테네신 오가네손
Rb
↑
Cs
↓
프루
크세논 ← 세슘 → 바륨
원자 번호 (Z)55
그룹.그룹 1:수소와 알칼리 금속이다.
기간6교시
블록 s블록
전자 구성[Xe] 6s1
셸당 전자 수2, 8, 18, 18, 8, 1
물리 속성
단계 STP에서단단한
녹는점301.7K(28.5°C, 83.3°F)
비등점944 K (671 °C, 1240 °F)
밀도 (근처)1.93g/cm3
액상일 때(로)1.843g/cm3
임계점1938 K, 9.4 MPa[2]
융해열2.09 kJ/mol
기화열63.9 kJ/mol
몰 열용량32.210 J/(mol·K)
증기 압력
P (Pa) 1 10 100 1k 10k 100k
(K)에서 418 469 534 623 750 940
원자 특성
산화 상태-1, +1[3](강염기성 산화물)
전기 음성도폴링 스케일: 0.79
이온화 에너지
  • 첫 번째: 375.7 kJ/mol
  • 두 번째: 2234.3 kJ/mol
  • 3차: 3400kJ/mol
원자 반지름경험적: 265 pm
공유 반지름244±11 pm
반데르발스 반지름오후 343시
Color lines in a spectral range
세슘 스펙트럼 라인
기타 속성
자연발생원시적인
결정 구조 체심입방체(BCC)
Bodycentredcubic crystal structure for caesium
열팽창97 µm/(mkK) (25 °C에서)
열전도율35.9 W/(mµK)
전기 저항률205 NΩm (20 °C에서)
자기 순서상사성[4]
영률1.7 GPa
벌크 계수1.6 GPa
모스 경도0.2
브리넬 경도0.14 MPa
CAS 번호7440-46-2
역사
명명하늘색이라는 라틴어 세시우스에서 유래한 것
검출로버트 분젠과 구스타프 키르히호프(1860)
첫 번째 분리칼 세터버그(1882)
세슘의 주요 동위원소
이소토페 아부노댄스 반감기 (t1/2) 붕괴 모드 프로덕트
133Cs 100% 안정적인.
134Cs 동기 2.0648년 ε 134Xe
β− 134바
135Cs 추적하다 2.3×106 y β− 135바
137Cs 동기 30.17년[5] β− 137바
카테고리 : 세슘
레퍼런스

세슘(IUPAC[6] 철자)은 화학 원소로 [note 1]기호는 Cs이고 원자 번호는 55입니다.은빛으로 빛나는 부드러운 알칼리 금속으로, 녹는점이 28.5°C(83.3°F)로, [note 2]상온 또는 그 부근에서 액체 상태의 5가지 원소 금속 중 하나입니다.세슘은 루비듐이나 칼륨과 비슷한 물리적, 화학적 성질을 가지고 있다.모든 금속 중 가장 반응성이 높은 금속으로 -116°C(-177°F)에서도 물과 반응합니다.폴링 척도에서 0.79의 값을 갖는 가장 작은 전기음성 원소입니다.그것은 오직 하나의 안정 동위원소인 세슘-133을 가지고 있다.세슘은 대부분 꽃가루 알갱이에서 채굴된다.이 원소는 40개의 알려진 동위원소를 가지고 있으며, 바륨, 수은과 함께 가장 많은 [11]동위원소를 가진 원소 중 하나이다.핵분열 생성물인 세슘-137은 [why?]원자로에서 생성된 폐기물에서 추출된다.

독일 화학자 로버트 분젠과 물리학자 구스타프 키르히호프는 1860년 새롭게 개발된 화염분광법에 의해 세슘을 발견했다.세슘의 첫 번째 소규모 용도는 진공관과 광전지의 "게터"로서 사용되었습니다.1967년, 빛의 속도가 우주에서 가장 일정한 차원이라는 아인슈타인의 증거에 따라, 국제 단위계는 세슘-133의 방출 스펙트럼에서 두 개의 특정한 파장을 사용하여 두 번째와 미터를 정의했습니다.그 이후로 세슘은 고정밀 원자시계에 널리 사용되어 왔다.

1990년대 이후, 이 원소의 가장 큰 용도는 시추액용 포름산 세슘이었지만, 전기 생산, 전자제품, 화학 분야에서 다양하게 사용되고 있다.방사성 동위원소 세슘-137의 반감기는 약 30년이며 의료용, 산업용 게이지, 수문학 등에 사용된다.비방사성 세슘 화합물은 약간의 독성이 있을 뿐이지만 순수한 금속이 물과 폭발적으로 반응하는 경향은 세슘이 유해 물질로 간주되고 방사성 동위원소가 환경에 상당한 건강과 생태학적 위험을 초래한다는 것을 의미한다.

특성.

물리 속성

Y-shaped yellowish crystal in glass ampoule, looking like the branch of a pine tree
아르곤에 저장된 고순도 세슘-133.

실온에서 고체인 모든 원소 중에서 세슘이 가장 부드럽습니다. 세슘의 경도는 0.2Mohs입니다.그것은 매우 연성이 있고 창백한 금속으로, 미량의 [12][13][14]산소가 있으면 어두워진다.미네랄 오일이 있으면(운반 중에 가장 잘 보관되는 곳) 금속 광택이 없어지고 더 칙칙한 회색빛을 띠게 됩니다.28.5°C(83.3°F)의 녹는점을 가지고 있어 상온에 가까운 액체 상태의 몇 안 되는 원소 금속 중 하나입니다.수은은 [note 3][16]세슘보다 낮은 녹는점을 가진 유일한 안정된 원소 금속이다.또한, 금속은 [17]수은을 제외한 모든 금속 중 가장 낮은 641°C(1,186°F)로 다소 낮은 끓는점을 가지고 있습니다.그것의 화합물은 파란색이나[18][19] 보라색[19] 빛깔로 타오른다.

세슘 결정(금색)과 루비듐 결정(은색) 비교

세슘은 다른 알칼리 금속, 금, 그리고 수은과 합금을 형성합니다.650°C(1,202°F) 미만의 온도에서는 코발트, 철, 몰리브덴, 니켈, 백금, 탄탈 또는 텅스텐과 합금되지 않습니다.그것은 안티몬, 갈륨, 인듐, 토륨과 함께 잘 정의된 금속간 화합물을 형성하는데,[12] 이들은 감광성이 있다.이는 다른 모든 알칼리 금속(리튬 제외)과 혼합됩니다. 몰 분포가 41% 세슘, 47% 칼륨, 12% 나트륨인 합금은 -78°C(-108°[16][20]F)로 알려진 금속 합금 중 가장 낮은 녹는점을 가집니다.몇 가지 아말감이 연구되었습니다.CsHg는
2 검은색에 보라색 금속 광택이 나는 반면, CsHg는 황금색이며 금속 [21]광택도 있습니다.

세슘의 황금색은 그룹이 하강하면서 알칼리 금속의 전자를 자극하는 데 필요한 빛의 빈도가 감소하기 때문입니다.루비듐을 통과하는 리튬의 경우 이 주파수는 자외선에 있지만 세슘의 경우 스펙트럼의 청자색 끝에 도달한다. 즉, 알칼리 금속의 플라스몬 주파수는 리튬에서 세슘으로 낮아진다.따라서 세슘은 제비꽃빛을 우선적으로 투과하고 부분적으로 흡수하는 반면 다른 색(주파수가 낮은 색)은 반사되기 때문에 [22]노란색으로 보입니다.

화학적 성질

찬물에 소량의 세슘을 첨가하면 폭발성이 있다.

세슘 금속은 매우 반응성이 높고 매우 열성이 강하다.공기 중에서 자연 발화하며, 다른 알칼리 금속(주기율표의 [12]첫 번째 그룹)보다 낮은 온도에서도 물과 폭발적으로 반응합니다.-116°C(-177°F)[16]의 낮은 온도에서 얼음과 반응합니다.이 높은 반응성 때문에 세슘 금속은 유해 물질로 분류된다.미네랄 오일과 같은 건조하고 포화 상태의 탄화수소로 저장 및 배송됩니다.아르곤과 같은 불활성 가스에서만 취급할 수 있습니다.하지만, 세슘-물 폭발은 종종 비슷한 양의 나트륨이 있는 나트륨-물 폭발보다 덜 강력합니다.물과 접촉하면 세슘이 순식간에 폭발해 수소가 [23]쌓일 시간이 거의 남지 않기 때문이다.세슘은 진공 밀폐된 붕규산염 유리 앰플에 저장할 수 있습니다.세슘은 약 100그램(3.5온스) 이상의 양으로 밀폐된 스테인리스 스틸 [12]용기에 담겨 출고됩니다.

세슘의 화학은 다른 알칼리 금속, 특히 주기율표의 [24]세슘 위에 있는 원소인 루비듐과 유사합니다.알칼리 금속의 예상대로, 유일한 공통 산화 상태는 +1이다.[note 4]다른 (비방사성) 알칼리 [26]금속보다 원자 질량이 높고 전기 양성이 강하다는 사실에서 약간의 차이가 발생합니다.세슘은 가장 전기 양성이 강한 화학 [note 5][16]원소이다.세슘 이온은 또한 가벼운 알칼리 금속보다 크고 덜 단단합니다.

컴파운드

27 small grey spheres in 3 evenly spaced layers of nine. 8 spheres form a regular cube and 8 of those cubes form a larger cube. The grey spheres represent the caesium atoms. The center of each small cube is occupied by a small green sphere representing a chlorine atom. Thus, every chlorine is in the middle of a cube formed by caesium atoms and every caesium is in the middle of a cube formed by chlorine.
CsCl에서 Cs와 Cl의 입방정위 볼 앤 스틱 모델

대부분의 세슘 화합물은 다양한 음이온에 이온 결합하는 양이온+
 Cs로서 이 원소를 포함합니다.주목할 만한 예외 중 하나는 카이사이드 음이온([3]Cs)이며−
, 다른 것은 몇 가지 아산화물입니다(아래 산화물에 대한 섹션 참조).보다 최근에는 세슘이 p-블록 원소로 작용하여 높은 산화 상태로 높은 불소(즉n,[28] n > 1)를 고압 하에서 형성할 수 있을 것으로 예측된다.이 예측은 추가 [29]실험을 통해 검증될 필요가 있다.

Cs의+ 소금은 음이온 자체가 착색되지 않는 한 보통 무색이다.단순염의 대부분은 흡습성이지만 가벼운 알칼리 금속의 해당 소금보다는 덜하다.인산염,[30] 아세트산염, 탄산염, 할로겐화물, 산화물, 질산염, 황산염은 수용성이다.이중염은 용해성이 낮은 경우가 많고, 세슘 알루미늄 황산염의 낮은 용해성은 광석에서 Cs를 정제하는 데 이용된다.안티몬(CsSbCl
4), 비스무트, 카드뮴, 구리, 철, 납을 포함한 이중염도 용해성이 [12]낮다.

Caesium hydroxide (CsOH) is hygroscopic and strongly basic.[24] It rapidly etches the surface of semiconductors such as silicon.[31] CsOH has been previously regarded by chemists as the "strongest base", reflecting the relatively weak attraction between the large Cs+ ion and OH−;[18] it is indeed the strongest Arrhenius base; however, a number of compounds such as n-butyllithium, sodium amide, sodium hydride, caesium hydride, etc., which cannot be dissolved in water as reacting violently with it but rather only used in some anhydrous polar aprotic solvents, are far more basic on the basis of the Brønsted–Lowry acid–base theory.[24]

A stoichiometric mixture of caesium and gold will react to form yellow caesium auride (Cs+Au−) upon heating. The auride anion here behaves as a pseudohalogen. The compound reacts violently with water, yielding caesium hydroxide, metallic gold, and hydrogen gas; in liquid ammonia it can be reacted with a caesium-specific ion exchange resin to produce tetramethylammonium auride. The analogous platinum compound, red caesium platinide (Cs2Pt), contains the platinide ion that behaves as a pseudochalcogen.[32]

Complexes

Like all metal cations, Cs+ forms complexes with Lewis bases in solution. Because of its large size, Cs+ usually adopts coordination numbers greater than 6, the number typical for the smaller alkali metal cations. This difference is apparent in the 8-coordination of CsCl. This high coordination number and softness (tendency to form covalent bonds) are properties exploited in separating Cs+ from other cations in the remediation of nuclear wastes, where 137Cs+ must be separated from large amounts of nonradioactive K+.[33]

Halides

Monatomic caesium halide wires grown inside double-wall carbon nanotubes (TEM image).[34]

Caesium fluoride (CsF) is a hygroscopic white solid that is widely used in organofluorine chemistry as a source of fluoride anions.[35] Caesium fluoride has the halite structure, which means that the Cs+ and F− pack in a cubic closest packed array as do Na+ and Cl− in sodium chloride.[24] Notably, caesium and fluorine have the lowest and highest electronegativities, respectively, among all the known elements.

Caesium chloride (CsCl) crystallizes in the simple cubic crystal system. Also called the "caesium chloride structure",[26] this structural motif is composed of a primitive cubic lattice with a two-atom basis, each with an eightfold coordination; the chloride atoms lie upon the lattice points at the edges of the cube, while the caesium atoms lie in the holes in the centre of the cubes. This structure is shared with CsBr and CsI, and many other compounds that do not contain Cs. In contrast, most other alkaline halides have the sodium chloride (NaCl) structure.[26] The CsCl structure is preferred because Cs+ has an ionic radius of 174 pm and Cl−
 181 pm.[36]

Oxides

The stick and ball diagram shows three regular octahedra, which are connected to the next one by one surface and the last one shares one surface with the first. All three have one edge in common. All eleven vertices are purple spheres representing caesium, and at the center of each octahedron is a small red sphere representing oxygen.
Cs
11O
3 cluster

More so than the other alkali metals, caesium forms numerous binary compounds with oxygen. When caesium burns in air, the superoxide CsO
2 is the main product.[37] The "normal" caesium oxide (Cs
2O) forms yellow-orange hexagonal crystals,[38] and is the only oxide of the anti-CdCl
2 type.[39] It vaporizes at 250 °C (482 °F), and decomposes to caesium metal and the peroxide Cs
2O
2 at temperatures above 400 °C (752 °F). In addition to the superoxide and the ozonide CsO
3,[40][41] several brightly coloured suboxides have also been studied.[42] These include Cs
7O, Cs
4O, Cs
11O
3, Cs
3O (dark-green[43]), CsO, Cs
3O
2,[44] as well as Cs
7O
2.[45][46] The latter may be heated in a vacuum to generate Cs
2O.[39] Binary compounds with sulfur, selenium, and tellurium also exist.[12]

Isotopes

Main article: Isotopes of caesium

Caesium has 40 known isotopes, ranging in mass number (i.e. number of nucleons in the nucleus) from 112 to 151. Several of these are synthesized from lighter elements by the slow neutron capture process (S-process) inside old stars[47] and by the R-process in supernova explosions.[48] The only stable caesium isotope is 133Cs, with 78 neutrons. Although it has a large nuclear spin (7/2+), nuclear magnetic resonance studies can use this isotope at a resonating frequency of 11.7 MHz.[49]

A graph showing the energetics of caesium-137 (nuclear spin: I=7/2+, half-life of about 30 years) decay. With a 94.6% probability, it decays by a 512 keV beta emission into barium-137m (I=11/2-, t=2.55min); this further decays by a 662 keV gamma emission with an 85.1% probability into barium-137 (I=3/2+). Alternatively, caesium-137 may decay directly into barium-137 by a 0.4% probability beta emission.
Decay of caesium-137

The radioactive 135Cs has a very long half-life of about 2.3 million years, the longest of all radioactive isotopes of caesium. 137Cs and 134Cs have half-lives of 30 and two years, respectively. 137Cs decomposes to a short-lived 137mBa by beta decay, and then to nonradioactive barium, while 134Cs transforms into 134Ba directly. The isotopes with mass numbers of 129, 131, 132 and 136, have half-lives between a day and two weeks, while most of the other isotopes have half-lives from a few seconds to fractions of a second. At least 21 metastable nuclear isomers exist. Other than 134mCs (with a half-life of just under 3 hours), all are very unstable and decay with half-lives of a few minutes or less.[50][51]

The isotope 135Cs is one of the long-lived fission products of uranium produced in nuclear reactors.[52] However, this fission product yield is reduced in most reactors because the predecessor, 135Xe, is a potent neutron poison and frequently transmutes to stable 136Xe before it can decay to 135Cs.[53][54]

The beta decay from 137Cs to 137mBa is a strong emission of gamma radiation.[55] 137Cs and 90Sr are the principal medium-lived products of nuclear fission, and the prime sources of radioactivity from spent nuclear fuel after several years of cooling, lasting several hundred years.[56] Those two isotopes are the largest source of residual radioactivity in the area of the Chernobyl disaster.[57] Because of the low capture rate, disposing of 137Cs through neutron capture is not feasible and the only current solution is to allow it to decay over time.[58]

Almost all caesium produced from nuclear fission comes from the beta decay of originally more neutron-rich fission products, passing through various isotopes of iodine and xenon.[59] Because iodine and xenon are volatile and can diffuse through nuclear fuel or air, radioactive caesium is often created far from the original site of fission.[60] With nuclear weapons testing in the 1950s through the 1980s, 137Cs was released into the atmosphere and returned to the surface of the earth as a component of radioactive fallout. It is a ready marker of the movement of soil and sediment from those times.[12]

Occurrence

A white mineral, from which white and pale pink crystals protrude
Pollucite, a caesium mineral
See also: Caesium minerals

Caesium is a relatively rare element, estimated to average 3 parts per million in the Earth's crust.[61] It is the 45th most abundant element and the 36th among the metals. Nevertheless, it is more abundant than such elements as antimony, cadmium, tin, and tungsten, and two orders of magnitude more abundant than mercury and silver; it is 3.3% as abundant as rubidium, with which it is closely associated, chemically.[12]

Due to its large ionic radius, caesium is one of the "incompatible elements".[62] During magma crystallization, caesium is concentrated in the liquid phase and crystallizes last. Therefore, the largest deposits of caesium are zone pegmatite ore bodies formed by this enrichment process. Because caesium does not substitute for potassium as readily as rubidium does, the alkali evaporite minerals sylvite (KCl) and carnallite (KMgCl
3·6H
2O) may contain only 0.002% caesium. Consequently, caesium is found in few minerals. Percentage amounts of caesium may be found in beryl (Be
3Al
2(SiO
3)
6) and avogadrite ((K,Cs)BF
4), up to 15 wt% Cs2O in the closely related mineral pezzottaite (Cs(Be
2Li)Al
2Si
6O
18), up to 8.4 wt% Cs2O in the rare mineral londonite ((Cs,K)Al
4Be
4(B,Be)
12O
28), and less in the more widespread rhodizite.[12] The only economically important ore for caesium is pollucite Cs(AlSi
2O
6), which is found in a few places around the world in zoned pegmatites, associated with the more commercially important lithium minerals, lepidolite and petalite. Within the pegmatites, the large grain size and the strong separation of the minerals results in high-grade ore for mining.[63]

The world's most significant and richest known source of caesium is the Tanco Mine at Bernic Lake in Manitoba, Canada, estimated to contain 350,000 metric tons of pollucite ore, representing more than two-thirds of the world's reserve base.[63][64] Although the stoichiometric content of caesium in pollucite is 42.6%, pure pollucite samples from this deposit contain only about 34% caesium, while the average content is 24 wt%.[64] Commercial pollucite contains more than 19% caesium.[65] The Bikita pegmatite deposit in Zimbabwe is mined for its petalite, but it also contains a significant amount of pollucite. Another notable source of pollucite is in the Karibib Desert, Namibia.[64] At the present rate of world mine production of 5 to 10 metric tons per year, reserves will last for thousands of years.[12]

Production

Mining and refining pollucite ore is a selective process and is conducted on a smaller scale than for most other metals. The ore is crushed, hand-sorted, but not usually concentrated, and then ground. Caesium is then extracted from pollucite primarily by three methods: acid digestion, alkaline decomposition, and direct reduction.[12][66]

In the acid digestion, the silicate pollucite rock is dissolved with strong acids, such as hydrochloric (HCl), sulfuric (H
2SO
4), hydrobromic (HBr), or hydrofluoric (HF) acids. With hydrochloric acid, a mixture of soluble chlorides is produced, and the insoluble chloride double salts of caesium are precipitated as caesium antimony chloride (Cs
4SbCl
7), caesium iodine chloride (Cs
2ICl), or caesium hexachlorocerate (Cs
2(CeCl
6)). After separation, the pure precipitated double salt is decomposed, and pure CsCl is precipitated by evaporating the water.

The sulfuric acid method yields the insoluble double salt directly as caesium alum (CsAl(SO
4)
2·12H
2O). The aluminium sulfate component is converted to insoluble aluminium oxide by roasting the alum with carbon, and the resulting product is leached with water to yield a Cs
2SO
4 solution.[12]

Roasting pollucite with calcium carbonate and calcium chloride yields insoluble calcium silicates and soluble caesium chloride. Leaching with water or dilute ammonia (NH
4OH) yields a dilute chloride (CsCl) solution. This solution can be evaporated to produce caesium chloride or transformed into caesium alum or caesium carbonate. Though not commercially feasible, the ore can be directly reduced with potassium, sodium, or calcium in vacuum to produce caesium metal directly.[12]

Most of the mined caesium (as salts) is directly converted into caesium formate (HCOO−Cs+) for applications such as oil drilling. To supply the developing market, Cabot Corporation built a production plant in 1997 at the Tanco mine near Bernic Lake in Manitoba, with a capacity of 12,000 barrels (1,900 m3) per year of caesium formate solution.[67] The primary smaller-scale commercial compounds of caesium are caesium chloride and nitrate.[68]

Alternatively, caesium metal may be obtained from the purified compounds derived from the ore. Caesium chloride and the other caesium halides can be reduced at 700 to 800 °C (1,292 to 1,472 °F) with calcium or barium, and caesium metal distilled from the result. In the same way, the aluminate, carbonate, or hydroxide may be reduced by magnesium.[12]

The metal can also be isolated by electrolysis of fused caesium cyanide (CsCN). Exceptionally pure and gas-free caesium can be produced by 390 °C (734 °F) thermal decomposition of caesium azide CsN
3, which can be produced from aqueous caesium sulfate and barium azide.[66] In vacuum applications, caesium dichromate can be reacted with zirconium to produce pure caesium metal without other gaseous products.[68]

Cs
2Cr
2O
7 + 2 Zr → 2 Cs + 2 ZrO
2+ Cr
2O
3

The price of 99.8% pure caesium (metal basis) in 2009 was about $10 per gram ($280/oz), but the compounds are significantly cheaper.[64]

History

Three middle-aged men, with the one in the middle sitting down. All wear long jackets, and the shorter man on the left has a beard.
Gustav Kirchhoff (left) and Robert Bunsen (centre) discovered caesium with their newly invented spectroscope.

In 1860, Robert Bunsen and Gustav Kirchhoff discovered caesium in the mineral water from Dürkheim, Germany. Because of the bright blue lines in the emission spectrum, they derived the name from the Latin word caesius, meaning sky-blue.[note 6][69][70][71] Caesium was the first element to be discovered with a spectroscope, which had been invented by Bunsen and Kirchhoff only a year previously.[16]

To obtain a pure sample of caesium, 44,000 litres (9,700 imp gal; 12,000 US gal) of mineral water had to be evaporated to yield 240 kilograms (530 lb) of concentrated salt solution. The alkaline earth metals were precipitated either as sulfates or oxalates, leaving the alkali metal in the solution. After conversion to the nitrates and extraction with ethanol, a sodium-free mixture was obtained. From this mixture, the lithium was precipitated by ammonium carbonate. Potassium, rubidium, and caesium form insoluble salts with chloroplatinic acid, but these salts show a slight difference in solubility in hot water, and the less-soluble caesium and rubidium hexachloroplatinate ((Cs,Rb)2PtCl6) were obtained by fractional crystallization. After reduction of the hexachloroplatinate with hydrogen, caesium and rubidium were separated by the difference in solubility of their carbonates in alcohol. The process yielded 9.2 grams (0.32 oz) of rubidium chloride and 7.3 grams (0.26 oz) of caesium chloride from the initial 44,000 litres of mineral water.[70]

From the caesium chloride, the two scientists estimated the atomic weight of the new element at 123.35 (compared to the currently accepted one of 132.9).[70] They tried to generate elemental caesium by electrolysis of molten caesium chloride, but instead of a metal, they obtained a blue homogeneous substance which "neither under the naked eye nor under the microscope showed the slightest trace of metallic substance"; as a result, they assigned it as a subchloride (Cs
2Cl). In reality, the product was probably a colloidal mixture of the metal and caesium chloride.[72] The electrolysis of the aqueous solution of chloride with a mercury cathode produced a caesium amalgam which readily decomposed under the aqueous conditions.[70] The pure metal was eventually isolated by the German chemist Carl Setterberg while working on his doctorate with Kekulé and Bunsen.[71] In 1882, he produced caesium metal by electrolysing caesium cyanide, avoiding the problems with the chloride.[73]

Historically, the most important use for caesium has been in research and development, primarily in chemical and electrical fields. Very few applications existed for caesium until the 1920s, when it came into use in radio vacuum tubes, where it had two functions; as a getter, it removed excess oxygen after manufacture, and as a coating on the heated cathode, it increased the electrical conductivity. Caesium was not recognized as a high-performance industrial metal until the 1950s.[74] Applications for nonradioactive caesium included photoelectric cells, photomultiplier tubes, optical components of infrared spectrophotometers, catalysts for several organic reactions, crystals for scintillation counters, and in magnetohydrodynamic power generators.[12] Caesium is also used as a source of positive ions in secondary ion mass spectrometry (SIMS).

Since 1967, the International System of Measurements has based the primary unit of time, the second, on the properties of caesium. The International System of Units (SI) defines the second as the duration of 9,192,631,770 cycles at the microwave frequency of the spectral line corresponding to the transition between two hyperfine energy levels of the ground state of caesium-133.[75] The 13th General Conference on Weights and Measures of 1967 defined a second as: "the duration of 9,192,631,770 cycles of microwave light absorbed or emitted by the hyperfine transition of caesium-133 atoms in their ground state undisturbed by external fields".

Applications

Petroleum exploration

The largest present-day use of nonradioactive caesium is in caesium formate drilling fluids for the extractive oil industry.[12] Aqueous solutions of caesium formate (HCOO−Cs+)—made by reacting caesium hydroxide with formic acid—were developed in the mid-1990s for use as oil well drilling and completion fluids. The function of a drilling fluid is to lubricate drill bits, to bring rock cuttings to the surface, and to maintain pressure on the formation during drilling of the well. Completion fluids assist the emplacement of control hardware after drilling but prior to production by maintaining the pressure.[12]

The high density of the caesium formate brine (up to 2.3 g/cm3, or 19.2 pounds per gallon),[76] coupled with the relatively benign nature of most caesium compounds, reduces the requirement for toxic high-density suspended solids in the drilling fluid—a significant technological, engineering and environmental advantage. Unlike the components of many other heavy liquids, caesium formate is relatively environment-friendly.[76] Caesium formate brine can be blended with potassium and sodium formates to decrease the density of the fluids to that of water (1.0 g/cm3, or 8.3 pounds per gallon). Furthermore, it is biodegradable and may be recycled, which is important in view of its high cost (about $4,000 per barrel in 2001).[77] Alkali formates are safe to handle and do not damage the producing formation or downhole metals as corrosive alternative, high-density brines (such as zinc bromide ZnBr
2 solutions) sometimes do; they also require less cleanup and reduce disposal costs.[12]

Atomic clocks

A room with a black box in the foreground and six control cabinets with space for five to six racks each. Most, but not all, of the cabinets are filled with white boxes.
Atomic clock ensemble at the U.S. Naval Observatory
A laboratory table with some optical devices on it.
FOCS-1, a continuous cold caesium fountain atomic clock in Switzerland, started operating in 2004 at an uncertainty of one second in 30 million years

Caesium-based atomic clocks use the electromagnetic transitions in the hyperfine structure of caesium-133 atoms as a reference point. The first accurate caesium clock was built by Louis Essen in 1955 at the National Physical Laboratory in the UK.[78] Caesium clocks have improved over the past half-century and are regarded as "the most accurate realization of a unit that mankind has yet achieved."[75] These clocks measure frequency with an error of 2 to 3 parts in 1014, which corresponds to an accuracy of 2 nanoseconds per day, or one second in 1.4 million years. The latest versions are more accurate than 1 part in 1015, about 1 second in 20 million years.[12] The caesium standard is the primary standard for standards-compliant time and frequency measurements.[79] Caesium clocks regulate the timing of cell phone networks and the Internet.[80]

Definition of the second

The second, symbol s, is the SI unit of time. It is defined by taking the fixed numerical value of the caesium frequency ΔνCs, the unperturbed ground-state hyperfine transition frequency of the caesium-133 atom, to be 9192631770 when expressed in the unit Hz, which is equal to s−1.

Electric power and electronics

Caesium vapour thermionic generators are low-power devices that convert heat energy to electrical energy. In the two-electrode vacuum tube converter, caesium neutralizes the space charge near the cathode and enhances the current flow.[81]

Caesium is also important for its photoemissive properties, converting light to electron flow. It is used in photoelectric cells because caesium-based cathodes, such as the intermetallic compound K
2CsSb, have a low threshold voltage for emission of electrons.[82] The range of photoemissive devices using caesium include optical character recognition devices, photomultiplier tubes, and video camera tubes.[83][84] Nevertheless, germanium, rubidium, selenium, silicon, tellurium, and several other elements can be substituted for caesium in photosensitive materials.[12]

Caesium iodide (CsI), bromide (CsBr) and caesium fluoride (CsF) crystals are employed for scintillators in scintillation counters widely used in mineral exploration and particle physics research to detect gamma and X-ray radiation. Being a heavy element, caesium provides good stopping power with better detection. Caesium compounds may provide a faster response (CsF) and be less hygroscopic (CsI).

Caesium vapour is used in many common magnetometers.[85]

The element is used as an internal standard in spectrophotometry.[86] Like other alkali metals, caesium has a great affinity for oxygen and is used as a "getter" in vacuum tubes.[87] Other uses of the metal include high-energy lasers, vapour glow lamps, and vapour rectifiers.[12]

Centrifugation fluids

The high density of the caesium ion makes solutions of caesium chloride, caesium sulfate, and caesium trifluoroacetate (Cs(O
2CCF
3)) useful in molecular biology for density gradient ultracentrifugation.[88] This technology is used primarily in the isolation of viral particles, subcellular organelles and fractions, and nucleic acids from biological samples.[89]

Chemical and medical use

Some fine white powder on a laboratory watch glass
Caesium chloride powder

Relatively few chemical applications use caesium.[90] Doping with caesium compounds enhances the effectiveness of several metal-ion catalysts for chemical synthesis, such as acrylic acid, anthraquinone, ethylene oxide, methanol, phthalic anhydride, styrene, methyl methacrylate monomers, and various olefins. It is also used in the catalytic conversion of sulfur dioxide into sulfur trioxide in the production of sulfuric acid.[12]

Caesium fluoride enjoys a niche use in organic chemistry as a base[24] and as an anhydrous source of fluoride ion.[91] Caesium salts sometimes replace potassium or sodium salts in organic synthesis, such as cyclization, esterification, and polymerization. Caesium has also been used in thermoluminescent radiation dosimetry (TLD): When exposed to radiation, it acquires crystal defects that, when heated, revert with emission of light proportionate to the received dose. Thus, measuring the light pulse with a photomultiplier tube can allow the accumulated radiation dose to be quantified.

Nuclear and isotope applications

Caesium-137 is a radioisotope commonly used as a gamma-emitter in industrial applications. Its advantages include a half-life of roughly 30 years, its availability from the nuclear fuel cycle, and having 137Ba as a stable end product. The high water solubility is a disadvantage which makes it incompatible with large pool irradiators for food and medical supplies.[92] It has been used in agriculture, cancer treatment, and the sterilization of food, sewage sludge, and surgical equipment.[12][93] Radioactive isotopes of caesium in radiation devices were used in the medical field to treat certain types of cancer,[94] but emergence of better alternatives and the use of water-soluble caesium chloride in the sources, which could create wide-ranging contamination, gradually put some of these caesium sources out of use.[95][96] Caesium-137 has been employed in a variety of industrial measurement gauges, including moisture, density, levelling, and thickness gauges.[97] It has also been used in well logging devices for measuring the electron density of the rock formations, which is analogous to the bulk density of the formations.[98]

Caesium-137 has been used in hydrologic studies analogous to those with tritium. As a daughter product of fission bomb testing from the 1950s through the mid-1980s, caesium-137 was released into the atmosphere, where it was absorbed readily into solution. Known year-to-year variation within that period allows correlation with soil and sediment layers. Caesium-134, and to a lesser extent caesium-135, have also been used in hydrology to measure the caesium output by the nuclear power industry. While they are less prevalent than either caesium-133 or caesium-137, these bellwether isotopes are produced solely from anthropogenic sources.[99]

Other uses

Electrons beamed from an electron gun hit and ionize neutral fuel atoms; in a chamber surrounded by magnets, the positive ions are directed toward a negative grid that accelerates them. The force of the engine is created by expelling the ions from the rear at high velocity. On exiting, the positive ions are neutralized from another electron gun, ensuring that neither the ship nor the exhaust is electrically charged and are not attracted.
Schematics of an electrostatic ion thruster developed for use with caesium or mercury fuel

Caesium and mercury were used as a propellant in early ion engines designed for spacecraft propulsion on very long interplanetary or extraplanetary missions. The fuel was ionized by contact with a charged tungsten electrode. But corrosion by caesium on spacecraft components has pushed development in the direction of inert gas propellants, such as xenon, which are easier to handle in ground-based tests and do less potential damage to the spacecraft.[12] Xenon was used in the experimental spacecraft Deep Space 1 launched in 1998.[100][101] Nevertheless, field-emission electric propulsion thrusters that accelerate liquid metal ions such as caesium have been built.[102]

Caesium nitrate is used as an oxidizer and pyrotechnic colorant to burn silicon in infrared flares,[103] such as the LUU-19 flare,[104] because it emits much of its light in the near infrared spectrum.[105] Caesium compounds may have been used as fuel additives to reduce the radar signature of exhaust plumes in the Lockheed A-12 CIA reconnaissance aircraft.[106] Caesium and rubidium have been added as a carbonate to glass because they reduce electrical conductivity and improve stability and durability of fibre optics and night vision devices. Caesium fluoride or caesium aluminium fluoride are used in fluxes formulated for brazing aluminium alloys that contain magnesium.[12]

Magnetohydrodynamic (MHD) power-generating systems were researched, but failed to gain widespread acceptance.[107] Caesium metal has also been considered as the working fluid in high-temperature Rankine cycle turboelectric generators.[108]

Caesium salts have been evaluated as antishock reagents following the administration of arsenical drugs. Because of their effect on heart rhythms, however, they are less likely to be used than potassium or rubidium salts. They have also been used to treat epilepsy.[12]

Caesium-133 can be laser cooled and used to probe fundamental and technological problems in quantum physics. It has a particularly convenient Feshbach spectrum to enable studies of ultracold atoms requiring tunable interactions.[109]

Health and safety hazards

Caesium
Hazards
GHS labelling:[110]
GHS02: Flammable GHS05: Corrosive
Danger
H260, H314
P223, P231+P232, P280, P305+P351+P338, P370+P378, P422
NFPA 704 (fire diamond)
3
4
3
Graph of percentage of the radioactive output by each nuclide that form after a nuclear fallout vs. logarithm of time after the incident. In curves of various colours, the predominant source of radiation are depicted in order: Te-132/I-132 for the first five or so days; I-131 for the next five; Ba-140/La-140 briefly; Zr-95/Nb-95 from day 10 until about day 200; and finally Cs-137. Other nuclides producing radioactivity, but not peaking as a major component are Ru, peaking at about 50 days, and Cs-134 at around 600 days.
The portion of the total radiation dose (in air) contributed by each isotope plotted against time after the Chernobyl disaster. Caesium-137 became the primary source of radiation about 200 days after the accident.[111]

Nonradioactive caesium compounds are only mildly toxic, and nonradioactive caesium is not a significant environmental hazard. Because biochemical processes can confuse and substitute caesium with potassium, excess caesium can lead to hypokalemia, arrhythmia, and acute cardiac arrest, but such amounts would not ordinarily be encountered in natural sources.[112][113]

The median lethal dose (LD50) for caesium chloride in mice is 2.3 g per kilogram, which is comparable to the LD50 values of potassium chloride and sodium chloride.[114] The principal use of nonradioactive caesium is as caesium formate in petroleum drilling fluids because it is much less toxic than alternatives, though it is more costly.[76]

Caesium metal is one of the most reactive elements and is highly explosive in the presence of water. The hydrogen gas produced by the reaction is heated by the thermal energy released at the same time, causing ignition and a violent explosion. This can occur with other alkali metals, but caesium is so potent that this explosive reaction can be triggered even by cold water.[12]

It is highly pyrophoric: the autoignition temperature of caesium is −116 °C (−177 °F), and it ignites explosively in air to form caesium hydroxide and various oxides. Caesium hydroxide is a very strong base, and will rapidly corrode glass.[17]

The isotopes 134 and 137 are present in the biosphere in small amounts from human activities, differing by location. Radiocaesium does not accumulate in the body as readily as other fission products (such as radioiodine and radiostrontium). About 10% of absorbed radiocaesium washes out of the body relatively quickly in sweat and urine. The remaining 90% has a biological half-life between 50 and 150 days.[115] Radiocaesium follows potassium and tends to accumulate in plant tissues, including fruits and vegetables.[116][117][118] Plants vary widely in the absorption of caesium, sometimes displaying great resistance to it. It is also well-documented that mushrooms from contaminated forests accumulate radiocaesium (caesium-137) in the fungal sporocarps.[119] Accumulation of caesium-137 in lakes has been a great concern after the Chernobyl disaster.[120][121] Experiments with dogs showed that a single dose of 3.8 millicuries (140 MBq, 4.1 μg of caesium-137) per kilogram is lethal within three weeks;[122] smaller amounts may cause infertility and cancer.[123] The International Atomic Energy Agency and other sources have warned that radioactive materials, such as caesium-137, could be used in radiological dispersion devices, or "dirty bombs".[124]

See also

Notes

  1. ^ Caesium is the spelling recommended by the International Union of Pure and Applied Chemistry (IUPAC).[7] The American Chemical Society (ACS) has used the spelling cesium since 1921,[8][9] following Webster's New International Dictionary. The element was named after the Latin word caesius, meaning "bluish grey".[10] In medieval and early modern writings caesius was spelled with the ligature æ as cæsius; hence, an alternative but now old-fashioned orthography is cæsium. More spelling explanation at ae/oe vs e.
  2. ^ Along with rubidium (39 °C [102 °F]), francium (estimated at 27 °C [81 °F]), mercury (−39 °C [−38 °F]), and gallium (30 °C [86 °F]); bromine is also liquid at room temperature (melting at −7.2 °C [19.0 °F]), but it is a halogen and not a metal. Preliminary work with copernicium and flerovium suggests that they are gaseous metals at room temperature.
  3. ^ The radioactive element francium may also have a lower melting point, but its radioactivity prevents enough of it from being isolated for direct testing.[15] Copernicium and flerovium may also have lower melting points.
  4. ^ It differs from this value in caesides, which contain the Cs− anion and thus have caesium in the −1 oxidation state.[3] Additionally, 2013 calculations by Mao-sheng Miao indicate that under conditions of extreme pressure (greater than 30 GPa), the inner 5p electrons could form chemical bonds, where caesium would behave as the seventh 5p element. This discovery indicates that higher caesium fluorides with caesium in oxidation states from +2 to +6 could exist under such conditions.[25]
  5. ^ Francium's electropositivity has not been experimentally measured due to its high radioactivity. Measurements of the first ionization energy of francium suggest that its relativistic effects may lower its reactivity and raise its electronegativity above that expected from periodic trends.[27]
  6. ^ Bunsen quotes Aulus Gellius Noctes Atticae II, 26 by Nigidius Figulus: Nostris autem veteribus caesia dicts est quae Graecis, ut Nigidus ait, de colore coeli quasi coelia.

References

  1. ^ "Standard Atomic Weights: Caesium". CIAAW. 2013.
  2. ^ Haynes, William M., ed. (2011). CRC Handbook of Chemistry and Physics (92nd ed.). Boca Raton, FL: CRC Press. p. 4.121. ISBN 1-4398-5511-0.
  3. ^ a b c Dye, J. L. (1979). "Compounds of Alkali Metal Anions". Angewandte Chemie International Edition. 18 (8): 587–598. doi:10.1002/anie.197905871.
  4. ^ "Magnetic susceptibility of the elements and inorganic compounds". Handbook of Chemistry and Physics (PDF) (87th ed.). CRC press. ISBN 0-8493-0487-3. Retrieved 2010-09-26.
  5. ^ "NIST Radionuclide Half-Life Measurements". NIST. Retrieved 2011-03-13.
  6. ^ "IUPAC Periodic Table of Elements". International Union of Pure and Applied Chemistry.
  7. ^ International Union of Pure and Applied Chemistry (2005). Nomenclature of Inorganic Chemistry (IUPAC Recommendations 2005). Cambridge (UK): RSC–IUPAC. ISBN 0-85404-438-8. pp. 248–49. Electronic version..
  8. ^ Coghill, Anne M.; Garson, Lorrin R., eds. (2006). The ACS Style Guide: Effective Communication of Scientific Information (3rd ed.). Washington, D.C.: American Chemical Society. p. 127. ISBN 978-0-8412-3999-9.
  9. ^ Coplen, T. B.; Peiser, H. S. (1998). "History of the recommended atomic-weight values from 1882 to 1997: a comparison of differences from current values to the estimated uncertainties of earlier values" (PDF). Pure Appl. Chem. 70 (1): 237–257. doi:10.1351/pac199870010237. S2CID 96729044.
  10. ^ OED entry for "caesium". Second edition, 1989; online version June 2012. Retrieved 07 September 2012. Earlier version first published in New English Dictionary, 1888.
  11. ^ "Isotopes". Ptable.
  12. ^ a b c d e f g h i j k l m n o p q r s t u v w x y z aa Butterman, William C.; Brooks, William E.; Reese Jr., Robert G. (2004). "Mineral Commodity Profile: Cesium" (PDF). United States Geological Survey. Archived from the original (PDF) on February 7, 2007. Retrieved 2009-12-27.
  13. ^ Heiserman, David L. (1992). Exploring Chemical Elements and their Compounds. McGraw-Hill. pp. 201–203. ISBN 978-0-8306-3015-8.
  14. ^ Addison, C. C. (1984). The Chemistry of the Liquid Alkali Metals. Wiley. ISBN 978-0-471-90508-0. Retrieved 2012-09-28.
  15. ^ "Francium". Periodic.lanl.gov. Retrieved 2010-02-23.
  16. ^ a b c d e Kaner, Richard (2003). "C&EN: It's Elemental: The Periodic Table – Cesium". American Chemical Society. Retrieved 2010-02-25.
  17. ^ a b "Chemical Data – Caesium – Cs". Royal Society of Chemistry. Retrieved 2010-09-27.
  18. ^ a b Lynch, Charles T. (1974). CRC Handbook of Materials Science. CRC Press. p. 13. ISBN 978-0-8493-2321-8.
  19. ^ a b Clark, Jim (2005). "Flame Tests". chemguide. Retrieved 2012-01-29.
  20. ^ Taova, T. M.; et al. (June 22, 2003). "Density of melts of alkali metals and their Na-K-Cs and Na-K-Rb ternary systems" (PDF). Fifteenth symposium on thermophysical properties, Boulder, Colorado, United States. Archived from the original (PDF) on October 9, 2006. Retrieved 2010-09-26.
  21. ^ Deiseroth, H. J. (1997). "Alkali metal amalgams, a group of unusual alloys". Progress in Solid State Chemistry. 25 (1–2): 73–123. doi:10.1016/S0079-6786(97)81004-7.
  22. ^ Addison, C. C. (1984). The chemistry of the liquid alkali metals. Wiley. p. 7. ISBN 9780471905080.
  23. ^ Gray, Theodore (2012) The Elements, Black Dog & Leventhal Publishers, p. 131, ISBN 1-57912-895-5.
  24. ^ a b c d e Greenwood, N. N.; Earnshaw, A. (1984). Chemistry of the Elements. Oxford, UK: Pergamon Press. ISBN 978-0-08-022057-4.
  25. ^ Moskowitz, Clara. "A Basic Rule of Chemistry Can Be Broken, Calculations Show". Scientific American. Retrieved 2013-11-22.
  26. ^ a b c Holleman, Arnold F.; Wiberg, Egon; Wiberg, Nils (1985). "Vergleichende Übersicht über die Gruppe der Alkalimetalle". Lehrbuch der Anorganischen Chemie (in German) (91–100 ed.). Walter de Gruyter. pp. 953–955. ISBN 978-3-11-007511-3.
  27. ^ Andreev, S. V.; Letokhov, V. S.; Mishin, V. I. (1987). "Laser resonance photoionization spectroscopy of Rydberg levels in Fr". Physical Review Letters. 59 (12): 1274–76. Bibcode:1987PhRvL..59.1274A. doi:10.1103/PhysRevLett.59.1274. PMID 10035190.
  28. ^ Miao, Mao-sheng (2013). "Caesium in high oxidation states and as a p-block element". Nature Chemistry. 5 (10): 846–852. doi:10.1038/nchem.1754. ISSN 1755-4349.
  29. ^ Sneed, D.; Pravica, M.; Kim, E.; Chen, N.; Park, C.; White, M. (2017-10-01). "Forcing Cesium into Higher Oxidation States Using Useful hard x-ray Induced Chemistry under High Pressure". Journal of Physics. Conference Series. 950 (11, 2017). doi:10.1088/1742-6596/950/4/042055. ISSN 1742-6588.
  30. ^ Hogan, C. M. (2011)."Phosphate". Archived from the original on 2012-10-25. Retrieved 2012-06-17. in Encyclopedia of Earth. Jorgensen, A. and Cleveland, C.J. (eds.). National Council for Science and the Environment. Washington DC
  31. ^ Köhler, Michael J. (1999). Etching in microsystem technology. Wiley-VCH. p. 90. ISBN 978-3-527-29561-6.
  32. ^ Jansen, Martin (2005-11-30). "Effects of relativistic motion of electrons on the chemistry of gold and platinum". Solid State Sciences. 7 (12): 1464–1474. Bibcode:2005SSSci...7.1464J. doi:10.1016/j.solidstatesciences.2005.06.015.
  33. ^ Moyer, Bruce A.; Birdwell, Joseph F.; Bonnesen, Peter V.; Delmau, Laetitia H. (2005). Use of Macrocycles in Nuclear-Waste Cleanup: A Realworld Application of a Calixcrown in Cesium Separation Technology. Macrocyclic Chemistry. pp. 383–405. doi:10.1007/1-4020-3687-6_24. ISBN 978-1-4020-3364-3..
  34. ^ Senga, Ryosuke; Suenaga, Kazu (2015). "Single-atom electron energy loss spectroscopy of light elements". Nature Communications. 6: 7943. Bibcode:2015NatCo...6.7943S. doi:10.1038/ncomms8943. PMC 4532884. PMID 26228378.
  35. ^ Evans, F. W.; Litt, M. H.; Weidler-Kubanek, A. M.; Avonda, F. P. (1968). "Reactions Catalyzed by Potassium Fluoride. 111. The Knoevenagel Reaction". Journal of Organic Chemistry. 33 (5): 1837–1839. doi:10.1021/jo01269a028.
  36. ^ Wells, A. F. (1984). Structural Inorganic Chemistry (5th ed.). Oxford Science Publications. ISBN 978-0-19-855370-0.
  37. ^ Cotton, F. Albert; Wilkinson, G. (1962). Advanced Inorganic Chemistry. John Wiley & Sons, Inc. p. 318. ISBN 978-0-471-84997-1.
  38. ^ Lide, David R., ed. (2006). CRC Handbook of Chemistry and Physics (87th ed.). Boca Raton, FL: CRC Press. pp. 451, 514. ISBN 0-8493-0487-3.
  39. ^ a b Tsai, Khi-Ruey; Harris, P. M.; Lassettre, E. N. (1956). "The Crystal Structure of Cesium Monoxide". Journal of Physical Chemistry. 60 (3): 338–344. doi:10.1021/j150537a022. Archived from the original on September 24, 2017.
  40. ^ Vol'nov, I. I.; Matveev, V. V. (1963). "Synthesis of cesium ozonide through cesium superoxide". Bulletin of the Academy of Sciences, USSR Division of Chemical Science. 12 (6): 1040–1043. doi:10.1007/BF00845494.
  41. ^ Tokareva, S. A. (1971). "Alkali and Alkaline Earth Metal Ozonides". Russian Chemical Reviews. 40 (2): 165–174. Bibcode:1971RuCRv..40..165T. doi:10.1070/RC1971v040n02ABEH001903. S2CID 250883291.
  42. ^ Simon, A. (1997). "Group 1 and 2 Suboxides and Subnitrides — Metals with Atomic Size Holes and Tunnels". Coordination Chemistry Reviews. 163: 253–270. doi:10.1016/S0010-8545(97)00013-1.
  43. ^ Tsai, Khi-Ruey; Harris, P. M.; Lassettre, E. N. (1956). "The Crystal Structure of Tricesium Monoxide". Journal of Physical Chemistry. 60 (3): 345–347. doi:10.1021/j150537a023.
  44. ^ Okamoto, H. (2009). "Cs-O (Cesium-Oxygen)". Journal of Phase Equilibria and Diffusion. 31: 86–87. doi:10.1007/s11669-009-9636-5. S2CID 96084147.
  45. ^ Band, A.; Albu-Yaron, A.; Livneh, T.; Cohen, H.; Feldman, Y.; Shimon, L.; Popovitz-Biro, R.; Lyahovitskaya, V.; Tenne, R. (2004). "Characterization of Oxides of Cesium". The Journal of Physical Chemistry B. 108 (33): 12360–12367. doi:10.1021/jp036432o.
  46. ^ Brauer, G. (1947). "Untersuchungen ber das System Csium-Sauerstoff". Zeitschrift für Anorganische Chemie. 255 (1–3): 101–124. doi:10.1002/zaac.19472550110.
  47. ^ Busso, M.; Gallino, R.; Wasserburg, G. J. (1999). "Nucleosynthesis in Asymptotic Giant Branch Stars: Relevance for Galactic Enrichment and Solar System Formation" (PDF). Annual Review of Astronomy and Astrophysics. 37: 239–309. Bibcode:1999ARA&A..37..239B. doi:10.1146/annurev.astro.37.1.239. Retrieved 2010-02-20.
  48. ^ Arnett, David (1996). Supernovae and Nucleosynthesis: An Investigation of the History of Matter, from the Big Bang to the Present. Princeton University Press. p. 527. ISBN 978-0-691-01147-9.
  49. ^ Goff, C.; Matchette, Michael A.; Shabestary, Nahid; Khazaeli, Sadegh (1996). "Complexation of caesium and rubidium cations with crown ethers in N,N-dimethylformamide". Polyhedron. 15 (21): 3897–3903. doi:10.1016/0277-5387(96)00018-6.
  50. ^ Brown, F.; Hall, G. R.; Walter, A. J. (1955). "The half-life of Cs137". Journal of Inorganic and Nuclear Chemistry. 1 (4–5): 241–247. Bibcode:1955PhRv...99..188W. doi:10.1016/0022-1902(55)80027-9.
  51. ^ Sonzogni, Alejandro. "Interactive Chart of Nuclides". National Nuclear Data Center: Brookhaven National Laboratory. Archived from the original on 2008-05-22. Retrieved 2008-06-06.
  52. ^ Ohki, Shigeo; Takaki, Naoyuki (14–16 October 2002). Transmutation of Cesium-135 with Fast Reactors (PDF). Seventh Information Exchange Meeting on Actinide and Fission Product Partitioning and Transmutation. Jeju, Korea. Archived from the original (PDF) on 2011-09-28. Retrieved 2010-09-26.
  53. ^ "20 Xenon: A Fission Product Poison" (PDF). CANDU Fundamentals (Report). CANDU Owners Group Inc. Archived from the original (PDF) on July 23, 2011. Retrieved 2010-09-15.
  54. ^ Taylor, V. F.; Evans, R. D.; Cornett, R. J. (2008). "Preliminary evaluation of 135Cs/137Cs as a forensic tool for identifying source of radioactive contamination". Journal of Environmental Radioactivity. 99 (1): 109–118. doi:10.1016/j.jenvrad.2007.07.006. PMID 17869392.
  55. ^ "Cesium Radiation Protection". U.S. Environmental Protection Agency. 2006-06-28. Archived from the original on March 15, 2011. Retrieved 2010-02-15.
  56. ^ Zerriffi, Hisham (2000-05-24). IEER Report: Transmutation – Nuclear Alchemy Gamble (Report). Institute for Energy and Environmental Research. Retrieved 2010-02-15.
  57. ^ Chernobyl's Legacy: Health, Environmental and Socia-Economic Impacts and Recommendations to the Governments of Belarus, Russian Federation and Ukraine (PDF) (Report). International Atomic Energy Agency. Archived from the original (PDF) on 2010-02-15. Retrieved 2010-02-18.
  58. ^ Kase, Takeshi; Konashi, Kenji; Takahashi, Hiroshi; Hirao, Yasuo (1993). "Transmutation of Cesium-137 Using Proton Accelerator". Journal of Nuclear Science and Technology. 30 (9): 911–918. doi:10.3327/jnst.30.911.
  59. ^ Knief, Ronald Allen (1992). "Fission Fragments". Nuclear engineering: theory and technology of commercial nuclear power. Taylor & Francis. p. 42. ISBN 978-1-56032-088-3.
  60. ^ Ishiwatari, N.; Nagai, H. "Release of xenon-137 and iodine-137 from UO2 pellet by pulse neutron irradiation at NSRR". Nippon Genshiryoku Gakkaishi. 23 (11): 843–850. OSTI 5714707.
  61. ^ Turekian, K. K.; Wedepohl, K. H. (1961). "Distribution of the elements in some major units of the Earth's crust". Geological Society of America Bulletin. 72 (2): 175–192. Bibcode:1961GSAB...72..175T. doi:10.1130/0016-7606(1961)72[175:DOTEIS]2.0.CO;2. ISSN 0016-7606.
  62. ^ Rowland, Simon (1998-07-04). "Cesium as a Raw Material: Occurrence and Uses". Artemis Society International. Retrieved 2010-02-15.
  63. ^ a b Černý, Petr; Simpson, F. M. (1978). "The Tanco Pegmatite at Bernic Lake, Manitoba: X. Pollucite" (PDF). Canadian Mineralogist. 16: 325–333. Retrieved 2010-09-26.
  64. ^ a b c d Polyak, Désirée E. "Cesium" (PDF). U.S. Geological Survey. Retrieved 2009-10-17.
  65. ^ Norton, J. J. (1973). "Lithium, cesium, and rubidium—The rare alkali metals". In Brobst, D. A.; Pratt, W. P. (eds.). United States mineral resources. Vol. Paper 820. U.S. Geological Survey Professional. pp. 365–378. Archived from the original on 2010-07-21. Retrieved 2010-09-26.
  66. ^ a b Burt, R. O. (1993). "Caesium and cesium compounds". Kirk-Othmer encyclopedia of chemical technology. Vol. 5 (4th ed.). New York: John Wiley & Sons, Inc. pp. 749–764. ISBN 978-0-471-48494-3.
  67. ^ Benton, William; Turner, Jim (2000). "Cesium formate fluid succeeds in North Sea HPHT field trials" (PDF). Drilling Contractor (May/June): 38–41. Retrieved 2010-09-26.
  68. ^ a b Eagleson, Mary, ed. (1994). Concise encyclopedia chemistry. Eagleson, Mary. Berlin: de Gruyter. p. 198. ISBN 978-3-11-011451-5.
  69. ^ Oxford English Dictionary, 2nd Edition
  70. ^ a b c d Kirchhoff, G.; Bunsen, R. (1861). "Chemische Analyse durch Spectralbeobachtungen" (PDF). Annalen der Physik und Chemie. 189 (7): 337–381. Bibcode:1861AnP...189..337K. doi:10.1002/andp.18611890702. hdl:2027/hvd.32044080591324.
  71. ^ a b Weeks, Mary Elvira (1932). "The discovery of the elements. XIII. Some spectroscopic discoveries". Journal of Chemical Education. 9 (8): 1413–1434. Bibcode:1932JChEd...9.1413W. doi:10.1021/ed009p1413.
  72. ^ Zsigmondy, Richard (2007). Colloids and the Ultra Microscope. Read books. p. 69. ISBN 978-1-4067-5938-9.
  73. ^ Setterberg, Carl (1882). "Ueber die Darstellung von Rubidium- und Cäsiumverbindungen und über die Gewinnung der Metalle selbst". Justus Liebig's Annalen der Chemie. 211: 100–116. doi:10.1002/jlac.18822110105.
  74. ^ Strod, A. J. (1957). "Cesium—A new industrial metal". American Ceramic Bulletin. 36 (6): 212–213.
  75. ^ a b "Cesium Atoms at Work". Time Service Department—U.S. Naval Observatory—Department of the Navy. Archived from the original on February 23, 2015. Retrieved 2009-12-20.
  76. ^ a b c Downs, J. D.; Blaszczynski, M.; Turner, J.; Harris, M. (February 2006). Drilling and Completing Difficult HP/HT Wells With the Aid of Cesium Formate Brines-A Performance Review. IADC/SPE Drilling Conference. Miami, Florida, USASociety of Petroleum Engineers. doi:10.2118/99068-MS. Archived from the original on 2007-10-12.
  77. ^ Flatern, Rick (2001). "Keeping cool in the HPHT environment". Offshore Engineer (February): 33–37.
  78. ^ Essen, L.; Parry, J. V. L. (1955). "An Atomic Standard of Frequency and Time Interval: A Caesium Resonator". Nature. 176 (4476): 280–282. Bibcode:1955Natur.176..280E. doi:10.1038/176280a0. S2CID 4191481.
  79. ^ Markowitz, W.; Hall, R.; Essen, L.; Parry, J. (1958). "Frequency of Cesium in Terms of Ephemeris Time". Physical Review Letters. 1 (3): 105–107. Bibcode:1958PhRvL...1..105M. doi:10.1103/PhysRevLett.1.105.
  80. ^ Reel, Monte (2003-07-22). "Where timing truly is everything". The Washington Post. p. B1. Archived from the original on 2013-04-29. Retrieved 2010-01-26.
  81. ^ Rasor, Ned S.; Warner, Charles (September 1964). "Correlation of Emission Processes for Adsorbed Alkali Films on Metal Surfaces". Journal of Applied Physics. 35 (9): 2589–2600. Bibcode:1964JAP....35.2589R. doi:10.1063/1.1713806.
  82. ^ "Cesium Supplier & Technical Information". American Elements. Retrieved 2010-01-25.
  83. ^ Smedley, John; Rao, Triveni; Wang, Erdong (2009). "K2CsSb Cathode Development". AIP Conference Proceedings. 1149 (1): 1062–1066. Bibcode:2009AIPC.1149.1062S. doi:10.1063/1.3215593.
  84. ^ Görlich, P. (1936). "Über zusammengesetzte, durchsichtige Photokathoden". Zeitschrift für Physik. 101 (5–6): 335–342. Bibcode:1936ZPhy..101..335G. doi:10.1007/BF01342330. S2CID 121613539.
  85. ^ Groeger, S.; Pazgalev, A. S.; Weis, A. (2005). "Comparison of discharge lamp and laser pumped cesium magnetometers". Applied Physics B. 80 (6): 645–654. arXiv:physics/0412011. Bibcode:2005ApPhB..80..645G. doi:10.1007/s00340-005-1773-x. S2CID 36065775.
  86. ^ Haven, Mary C.; Tetrault, Gregory A.; Schenken, Jerald R. (1994). "Internal Standards". Laboratory instrumentation. New York: John Wiley and Sons. p. 108. ISBN 978-0-471-28572-4.
  87. ^ McGee, James D. (1969). Photo-electronic image devices: proceedings of the fourth symposium held at Imperial College, London, September 16–20, 1968. Vol. 1. Academic Press. p. 391. ISBN 978-0-12-014528-7.
  88. ^ Manfred Bick, Horst Prinz, "Cesium and Cesium Compounds" in Ullmann's Encyclopedia of Industrial Chemistry 2005, Wiley-VCH, Weinheim. doi:10.1002/14356007.a06_153.
  89. ^ Desai, Mohamed A., ed. (2000). "Gradient Materials". Downstream processing methods. Totowa, N.J.: Humana Press. pp. 61–62. ISBN 978-0-89603-564-5.
  90. ^ Burt, R. O. (1993). "Cesium and cesium compounds". Kirk-Othmer encyclopedia of chemical technology. Vol. 5 (4th ed.). New York: John Wiley & Sons. p. 759. ISBN 978-0-471-15158-6.
  91. ^ Friestad, Gregory K.; Branchaud, Bruce P.; Navarrini, Walter and Sansotera, Maurizio (2007) "Cesium Fluoride" in Encyclopedia of Reagents for Organic Synthesis, John Wiley & Sons. doi:10.1002/047084289X.rc050.pub2
  92. ^ Okumura, Takeshi (2003-10-21). "The material flow of radioactive cesium-137 in the U.S. 2000" (PDF). United States Environmental Protection Agency. Archived from the original (PDF) on July 20, 2011. Retrieved 2009-12-20.
  93. ^ Jensen, N. L. (1985). "Cesium". Mineral facts and problems. Vol. Bulletin 675. U.S. Bureau of Mines. pp. 133–138.
  94. ^ "IsoRay's Cesium-131 Medical Isotope Used In Milestone Procedure Treating Eye Cancers At Tufts-New England Medical Center". Medical News Today. 2007-12-17. Retrieved 2010-02-15.
  95. ^ Bentel, Gunilla Carleson (1996). "Caesium-137 Machines". Radiation therapy planning. McGraw-Hill Professional. pp. 22–23. ISBN 978-0-07-005115-7. Retrieved 2010-09-26.
  96. ^ National Research Council (U.S.). Committee on Radiation Source Use and Replacement (2008). Radiation source use and replacement: abbreviated version. National Academies Press. ISBN 978-0-309-11014-3.
  97. ^ Loxton, R.; Pope, P., eds. (1995). "Level and density measurement using non-contact nuclear gauges". Instrumentation : A Reader. London: Chapman & Hall. pp. 82–85. ISBN 978-0-412-53400-3.
  98. ^ Timur, A.; Toksoz, M. N. (1985). "Downhole Geophysical Logging". Annual Review of Earth and Planetary Sciences. 13: 315–344. Bibcode:1985AREPS..13..315T. doi:10.1146/annurev.ea.13.050185.001531.
  99. ^ Kendall, Carol. "Isotope Tracers Project – Resources on Isotopes – Cesium". National Research Program – U.S. Geological Survey. Retrieved 2010-01-25.
  100. ^ Marcucci, M. G.; Polk, J. E. (2000). "NSTAR Xenon Ion Thruster on Deep Space 1: Ground and flight tests (invited)". Review of Scientific Instruments. 71 (3): 1389–1400. Bibcode:2000RScI...71.1389M. doi:10.1063/1.1150468.
  101. ^ Sovey, James S.; Rawlin, Vincent K.; Patterson, Michael J. "A Synopsis of Ion Propulsion Development Projects in the United States: SERT I to Deep Space I" (PDF). NASA. Archived from the original (PDF) on June 29, 2009. Retrieved 2009-12-12.
  102. ^ Marrese, C.; Polk, J.; Mueller, J.; Owens, A.; Tajmar, M.; Fink, R. & Spindt, C. (October 2001). In-FEEP Thruster Ion Beam Neutralization with Thermionic and Field Emission Cathodes. 27th International Electric Propulsion Conference. Pasadena, California. pp. 1–15. Archived from the original (PDF) on 2010-05-27. Retrieved 2010-01-25.
  103. ^ "Infrared illumination compositions and articles containing the same". United States Patent 6230628. Freepatentsonline.com. Retrieved 2010-01-25.
  104. ^ "LUU-19 Flare". Federation of American Scientists. 2000-04-23. Archived from the original on 2010-08-06. Retrieved 2009-12-12.
  105. ^ Charrier, E.; Charsley, E. L.; Laye, P. G.; Markham, H. M.; Berger, B.; Griffiths, T. T. (2006). "Determination of the temperature and enthalpy of the solid–solid phase transition of caesium nitrate by differential scanning calorimetry". Thermochimica Acta. 445: 36–39. doi:10.1016/j.tca.2006.04.002.
  106. ^ Crickmore, Paul F. (2000). Lockheed SR-71: the secret missions exposed. Osprey. p. 47. ISBN 978-1-84176-098-8.
  107. ^ National Research Council (U.S.) (2001). Energy research at DOE—Was it worth it?. National Academy Press. pp. 190–194. doi:10.17226/10165. ISBN 978-0-309-07448-3. Retrieved 2010-09-26.
  108. ^ Roskill Information Services (1984). Economics of Caesium and Rubidium (Reports on Metals & Minerals). London, United Kingdom: Roskill Information Services. p. 51. ISBN 978-0-86214-250-6.
  109. ^ Chin, Cheng; Grimm, Rudolf; Julienne, Paul; Tiesinga, Eite (2010-04-29). "Feshbach resonances in ultracold gases". Reviews of Modern Physics. 82 (2): 1225–1286. arXiv:0812.1496. Bibcode:2010RvMP...82.1225C. doi:10.1103/RevModPhys.82.1225. S2CID 118340314.
  110. ^ "Cesium 239240". Sigma-Aldrich. 2021-09-17. Retrieved 2021-12-21.
  111. ^ Data from The Radiochemical Manual and Wilson, B. J. (1966) The Radiochemical Manual (2nd ed.).
  112. ^ Melnikov, P.; Zanoni, L. Z. (June 2010). "Clinical effects of cesium intake". Biological Trace Element Research. 135 (1–3): 1–9. doi:10.1007/s12011-009-8486-7. PMID 19655100. S2CID 19186683.
  113. ^ Pinsky, Carl; Bose, Ranjan; Taylor, J. R.; McKee, Jasper; Lapointe, Claude; Birchall, James (1981). "Cesium in mammals: Acute toxicity, organ changes and tissue accumulation". Journal of Environmental Science and Health, Part A. 16 (5): 549–567. doi:10.1080/10934528109375003.
  114. ^ Johnson, Garland T.; Lewis, Trent R.; Wagner, D. Wagner (1975). "Acute toxicity of cesium and rubidium compounds". Toxicology and Applied Pharmacology. 32 (2): 239–245. doi:10.1016/0041-008X(75)90216-1. PMID 1154391.
  115. ^ Rundo, J. (1964). "A Survey of the Metabolism of Caesium in Man". British Journal of Radiology. 37 (434): 108–114. doi:10.1259/0007-1285-37-434-108. PMID 14120787.
  116. ^ Nishita, H.; Dixon, D.; Larson, K. H. (1962). "Accumulation of Cs and K and growth of bean plants in nutrient solution and soils". Plant and Soil. 17 (2): 221–242. doi:10.1007/BF01376226. S2CID 10293954.
  117. ^ Avery, S. (1996). "Fate of caesium in the environment: Distribution between the abiotic and biotic components of aquatic and terrestrial ecosystems". Journal of Environmental Radioactivity. 30 (2): 139–171. doi:10.1016/0265-931X(96)89276-9.
  118. ^ Salbu, Brit; Østby, Georg; Garmo, Torstein H.; Hove, Knut (1992). "Availability of caesium isotopes in vegetation estimated from incubation and extraction experiments". Analyst. 117 (3): 487–491. Bibcode:1992Ana...117..487S. doi:10.1039/AN9921700487. PMID 1580386.
  119. ^ Vinichuk, M. (2010). "Accumulation of potassium, rubidium and caesium (133Cs and 137Cs) in various fractions of soil and fungi in a Swedish forest". Science of the Total Environment. 408 (12): 2543–2548. Bibcode:2010ScTEn.408.2543V. doi:10.1016/j.scitotenv.2010.02.024. PMID 20334900.
  120. ^ Smith, Jim T.; Beresford, Nicholas A. (2005). Chernobyl: Catastrophe and Consequences. Berlin: Springer. ISBN 978-3-540-23866-9.
  121. ^ Eremeev, V. N.; Chudinovskikh, T. V.; Batrakov, G. F.; Ivanova, T. M. (1991). "Radioactive isotopes of caesium in the waters and near-water atmospheric layer of the Black Sea". Physical Oceanography. 2 (1): 57–64. doi:10.1007/BF02197418. S2CID 127482742.
  122. ^ Redman, H. C.; McClellan, R. O.; Jones, R. K.; Boecker, B. B.; Chiffelle, T. L.; Pickrell, J. A.; Rypka, E. W. (1972). "Toxicity of 137-CsCl in the Beagle. Early Biological Effects". Radiation Research. 50 (3): 629–648. Bibcode:1972RadR...50..629R. doi:10.2307/3573559. JSTOR 3573559. PMID 5030090.
  123. ^ "Chinese 'find' radioactive ball". BBC News. 2009-03-27. Retrieved 2010-01-25.
  124. ^ Charbonneau, Louis (2003-03-12). "IAEA director warns of 'dirty bomb' risk". The Washington Post. Reuters. p. A15. Archived from the original on 2008-12-05. Retrieved 2010-04-28.

External links

Listen to this article (3 hours and 27 minutes)
Spoken Wikipedia icon
This audio file was created from a revision of this article dated 29 November 2010 (2010-11-29), and does not reflect subsequent edits.
(Audio help · More spoken articles)