document_title: Geostationary orbit

question_text: the orbital speed of geo stationary satellite is

text: Geostationary orbit

text: History[edit]

text: The first appearance of a geostationary orbit in popular literature was in the first Venus Equilateral story by George O. Smith,[2] but Smith did not go into details. British science fiction author Arthur C. Clarke disseminated the idea widely, with more details on how it would work, in a 1945 paper entitled "Extra-Terrestrial Relays — Can Rocket Stations Give Worldwide Radio Coverage?", published in Wireless World magazine. Clarke acknowledged the connection in his introduction to The Complete Venus Equilateral.[3] The orbit, which Clarke first described as useful for broadcast and relay communications satellites,[4] is sometimes called the Clarke Orbit.[5] Similarly, the Clarke Belt is the part of space about 35,786 km (22,236 mi) above sea level, in the plane of the equator, where near-geostationary orbits may be implemented. The Clarke Orbit is about 265,000 km (165,000 mi) in circumference.

text: Practical uses[edit]

text: Most commercial communications satellites, broadcast satellites and SBAS satellites operate in geostationary orbits. A geostationary transfer orbit is used to move a satellite from low Earth orbit (LEO) into a geostationary orbit. The first satellite placed into a geostationary orbit was the Syncom-3, launched by a Delta D rocket in 1964.

text: A worldwide network of operational geostationary meteorological satellites is used to provide visible and infrared images of Earth's surface and atmosphere. These satellite systems include:

text:
the United States GOES
Meteosat, launched by the European Space Agency and operated by the European Weather Satellite Organization, EUMETSAT
the Japanese Himawari
Chinese Fengyun
India's INSAT series

text: A statite, a hypothetical satellite that uses a solar sail to modify its orbit, could theoretically hold itself in a geostationary "orbit" with different altitude and/or inclination from the "traditional" equatorial geostationary orbit.[6]

text: Communications[edit]

text: Satellites in geostationary orbits are far enough away from Earth that communication latency becomes significant — about a quarter of a second for a trip from one ground-based transmitter to the satellite and back to another ground-based transmitter; close to half a second for a round-trip communication from one Earth station to another and then back to the first.

text: For example, for ground stations at latitudes of φ = ±45° on the same meridian as the satellite, the time taken for a signal to pass from Earth to the satellite and back again can be computed using the cosine rule, given the geostationary orbital radius r (derived below), the Earth's radius R and the speed of light c, as

text: (Note that r is the orbital radius, the distance from the centre of the Earth, not the height above the equator.)

text: This delay presents problems for latency-sensitive applications such as voice communication.[7]

text: Geostationary satellites are directly overhead at the equator and appear lower in the sky to an observer nearer the poles. As the observer's latitude increases, communication becomes more difficult due to factors such as atmospheric refraction, Earth's thermal emission, line-of-sight obstructions, and signal reflections from the ground or nearby structures. At latitudes above about 81°, geostationary satellites are below the horizon and cannot be seen at all.[8] Because of this, some Russian communication satellites have used elliptical Molniya and Tundra orbits, which have excellent visibility at high latitudes.

text: Orbit allocation[edit]

text: Satellites in geostationary orbit must all occupy a single ring above the equator. The requirement to space these satellites apart to avoid harmful radio-frequency interference during operations means that there are a limited number of orbital "slots" available, and thus only a limited number of satellites can be operated in geostationary orbit. This has led to conflict between different countries wishing access to the same orbital slots (countries near the same longitude but differing latitudes) and radio frequencies. These disputes are addressed through the International Telecommunication Union's allocation mechanism.[9][10] In the 1976 Bogotá Declaration, eight countries located on the Earth's equator claimed sovereignty over the geostationary orbits above their territory, but the claims gained no international recognition.[11]

text: Orbital stability[edit]

text: A geostationary orbit can be achieved only at an altitude very close to 35,786 km (22,236 mi) and directly above the equator. This equates to an orbital velocity of 3.07 km/s (1.91 mi/s) and an orbital period of 1,436 minutes, which equates to almost exactly one sidereal day (23.934461223 hours). This ensures that the satellite will match the Earth's rotational period and has a stationary footprint on the ground. All geostationary satellites have to be located on this ring.

text: A combination of lunar gravity, solar gravity, and the flattening of the Earth at its poles causes a precession motion of the orbital plane of any geostationary object, with an orbital period of about 53 years and an initial inclination gradient of about 0.85° per year, achieving a maximal inclination of 15° after 26.5 years.[12] To correct for this orbital perturbation, regular orbital stationkeeping manoeuvres are necessary, amounting to a delta-v of approximately 50 m/s per year.

text: A second effect to be taken into account is the longitudinal drift, caused by the asymmetry of the Earth – the equator is slightly elliptical. There are two stable (at 75.3°E and 252°E) and two unstable (at 165.3°E and 14.7°W) equilibrium points. Any geostationary object placed between the equilibrium points would (without any action) be slowly accelerated towards the stable equilibrium position, causing a periodic longitude variation.[12] The correction of this effect requires station-keeping maneuvers with a maximal delta-v of about 2 m/s per year, depending on the desired longitude.

text: Solar wind and radiation pressure also exert small forces on satellites; over time, these cause them to slowly drift away from their prescribed orbits.

text: In the absence of servicing missions from the Earth or a renewable propulsion method, the consumption of thruster propellant for station keeping places a limitation on the lifetime of the satellite. Hall-effect thrusters, which are currently in use, have the potential to prolong the service life of a satellite by providing high-efficiency electric propulsion.

text: Limitations to usable life of geostationary satellites[edit]

text: When they run out of thruster fuel, the satellites are at the end of their service life, as they are no longer able to stay in their allocated orbital position. The transponders and other onboard systems generally outlive the thruster fuel and, by stopping N–S station keeping, some satellites can continue to be used in inclined orbits (where the orbital track appears to follow a figure-eight loop centred on the equator),[13][14] or else be elevated to a "graveyard" disposal orbit.

text: Derivation of geostationary altitude[edit]

text: In any circular orbit, the centripetal force required to maintain the orbit (Fc) is provided by the gravitational force on the satellite (Fg). To calculate the geostationary orbit altitude, one begins with this equivalence:

text: By Newton's second law of motion,[15] we can replace the forces F with the mass m of the object multiplied by the acceleration felt by the object due to that force:

text: We note that the mass of the satellite m appears on both sides — geostationary orbit is independent of the mass of the satellite.[c] Calculating the geostationary altitude, therefore, simplifies down to calculating the altitude where the magnitudes of the centripetal acceleration required for orbital motion and the gravitational acceleration provided by Earth's gravity are equal.

text: The centripetal acceleration's magnitude is:

text: where ω is the angular speed, and r is the orbital radius as measured from the Earth's center of mass.

text: The magnitude of the gravitational acceleration is:

text: where M is the mass of Earth, 5.9736 × 1024 kg, and G is the gravitational constant, (6.67428 ± 0.00067) × 10−11 m3 kg−1 s−2.

text: Equating the two accelerations gives:

text: The product GM is known with much greater precision than either factor alone; it is known as the geocentric gravitational constant μ = 398,600.4418 ± 0.0008 km3 s−2. Hence

text: The angular speed ω is found by dividing the angle travelled in one revolution (360° = 2π rad) by the orbital period (the time it takes to make one full revolution). In the case of a geostationary orbit, the orbital period is one sidereal day, or 7004861640905400000♠86164.09054 s).[16] This gives

text: The resulting orbital radius is 42,164 kilometres (26,199 mi). Subtracting the Earth's equatorial radius, 6,378 kilometres (3,963 mi), gives the altitude of 35,786 kilometres (22,236 mi).

text: Orbital speed is calculated by multiplying the angular speed by the orbital radius:

text: By the same formula, we can find the geostationary-type orbit of an object in relation to Mars (this type of orbit above is referred to as an areostationary orbit if it is above Mars). The geocentric gravitational constant GM (which is μ) for Mars has the value of 42,828 km3s−2, and the known rotational period (T) of Mars is 88,642.66 seconds. Since ω = 2π/T, using the formula above, the value of ω is found to be approx 7.088218×10−5 s−1. Thus r3 = 8.5243×1012 km3, whose cube root is 20,427 km (the orbital radius); subtracting the equatorial radius of Mars (3396.2 km) gives the orbital altitude of 17,031 km.

text: Orbital speed of a Mars geostationary orbit can be calculated as for Earth above:

text:
v
=
ω
r
≈
1.4479

km/s

≈
5

212

km/h

.

{\displaystyle v=\omega r\approx 1.4479~{\text{km/s}}\approx 5\,212~{\text{km/h}}.}


document_title: Climate of Chicago

question_text: when does the weather get warmer in chicago

root

text: Climate of Chicago

text: Official locations[edit]

text: The National Weather Service office in Chicago has one of the longest periods of official weather records, dating back to 1870, though all of the 1870 and 1871 weather records taken at 181 West Washington Street were lost in the Great Chicago Fire. As for the two major airports located in Chicago, Midway Airport began observations in 1928, and O'Hare Airport began observations in 1958. Both sites have served as official observation locations, with the latter being the current location where Chicago's official weather data is recorded. For Midway Airport, weather data prior to July 1, 1942 and after January 16, 1980 are not part of the official climate record of Chicago. All weather data taken at O'Hare from the beginning of observations in 1958 until January 17, 1980 are not part of the official climate record of Chicago.

tag: p

text: Here is a list of official weather locations for the Chicago office:[1]

text:
Dates
Location
Remarks

October 15, 1870 – October 8, 1871
181 W. Washington St.
All records lost due to Great Chicago Fire

October 15, 1871 – June 11, 1872
427 W. Randolph St.

June 11, 1872 – June 8, 1873
20 N. Wacker Dr.

June 8, 1873 – January 1, 1887
Roanoke Building

January 1, 1887 – February 1, 1890
Chicago Opera House

February 1, 1890 – July 1, 1905
Auditorium Tower
Automatic rain gauge installed in 1897

July 1, 1905 – December 31, 1925
U.S. Court House
Supplemental observations through 1970

January 1, 1926 – June 30, 1942
University of Chicago
Supplemental observations through 1962

July 1, 1942 – January 16, 1980
Midway Airport
Continues to provide observations

January 17, 1980 – present
O'Hare Airport

tag: table

text: Note: Some of the addresses prior to 1909 are different than the post-1909 addresses

text: Data[edit]

text:
[show]Climate data for Chicago (Midway Airport), 1981–2010 normals,[a] extremes 1928–present

Month
Jan
Feb
Mar
Apr
May
Jun
Jul
Aug
Sep
Oct
Nov
Dec
Year

Record high °F (°C)
67
(19)
75
(24)
86
(30)
92
(33)
102
(39)
107
(42)
109
(43)
102
(39)
101
(38)
94
(34)
81
(27)
72
(22)
109
(43)

Mean maximum °F (°C)
52.3
(11.3)
57.2
(14)
73.0
(22.8)
82.2
(27.9)
88.2
(31.2)
94.1
(34.5)
96.5
(35.8)
94.1
(34.5)
90.5
(32.5)
82.2
(27.9)
68.7
(20.4)
55.7
(13.2)
97.7
(36.5)

Average high °F (°C)
31.5
(−0.3)
35.8
(2.1)
46.8
(8.2)
59.2
(15.1)
70.2
(21.2)
79.9
(26.6)
84.2
(29)
82.1
(27.8)
75.3
(24.1)
62.8
(17.1)
48.6
(9.2)
35.3
(1.8)
59.4
(15.2)

Daily mean °F (°C)
24.8
(−4)
28.7
(−1.8)
38.8
(3.8)
50.4
(10.2)
60.9
(16.1)
71.0
(21.7)
75.9
(24.4)
74.1
(23.4)
66.4
(19.1)
54.2
(12.3)
41.5
(5.3)
29.0
(−1.7)
51.4
(10.8)

Average low °F (°C)
18.2
(−7.7)
21.7
(−5.7)
30.9
(−0.6)
41.7
(5.4)
51.6
(10.9)
62.1
(16.7)
67.5
(19.7)
66.2
(19)
57.5
(14.2)
45.7
(7.6)
34.5
(1.4)
22.7
(−5.2)
43.5
(6.4)

Mean minimum °F (°C)
−3.3
(−19.6)
2.4
(−16.4)
14.3
(−9.8)
27.0
(−2.8)
38.4
(3.6)
48.6
(9.2)
56.8
(13.8)
56.1
(13.4)
43.1
(6.2)
31.1
(−0.5)
19.9
(−6.7)
2.2
(−16.6)
−8.7
(−22.6)

Record low °F (°C)
−25
(−32)
−20
(−29)
−7
(−22)
10
(−12)
28
(−2)
35
(2)
46
(8)
43
(6)
34
(1)
20
(−7)
−3
(−19)
−20
(−29)
−25
(−32)

Average precipitation inches (mm)
2.06
(52.3)
1.94
(49.3)
2.72
(69.1)
3.64
(92.5)
4.13
(104.9)
4.06
(103.1)
4.01
(101.9)
3.99
(101.3)
3.31
(84.1)
3.24
(82.3)
3.42
(86.9)
2.57
(65.3)
39.09
(992.9)

Average snowfall inches (cm)
11.5
(29.2)
9.1
(23.1)
5.4
(13.7)
1.0
(2.5)
trace
0
(0)
0
(0)
0
(0)
0
(0)
0.1
(0.3)
1.3
(3.3)
8.7
(22.1)
37.1
(94.2)

Average precipitation days (≥ 0.01 in)
10.7
8.8
11.2
11.1
11.4
10.3
9.9
9.0
8.2
10.2
11.2
11.1
123.1

Average snowy days (≥ 0.1 in)
8.1
5.5
3.8
0.7
0
0
0
0
0
0.1
1.8
6.7
26.7

Source: NOAA,[2][3][4] WRCC[5]

text:
[show]Climate data for Chicago (O'Hare Int'l Airport), 1981–2010 normals,[a] extremes 1871–present[b]

Month
Jan
Feb
Mar
Apr
May
Jun
Jul
Aug
Sep
Oct
Nov
Dec
Year

Record high °F (°C)
67
(19)
75
(24)
88
(31)
91
(33)
98
(37)
104
(40)
105
(41)
102
(39)
101
(38)
94
(34)
81
(27)
71
(22)
105
(41)

Mean maximum °F (°C)
51.4
(10.8)
56.5
(13.6)
72.4
(22.4)
81.7
(27.6)
87.2
(30.7)
93.6
(34.2)
95.5
(35.3)
93.4
(34.1)
89.7
(32.1)
81.1
(27.3)
67.6
(19.8)
55.0
(12.8)
96.8
(36)

Average high °F (°C)
31.0
(−0.6)
35.3
(1.8)
46.6
(8.1)
59.0
(15)
70.0
(21.1)
79.7
(26.5)
84.1
(28.9)
81.9
(27.7)
74.8
(23.8)
62.3
(16.8)
48.2
(9)
34.8
(1.6)
59.1
(15.1)

Daily mean °F (°C)
23.8
(−4.6)
27.7
(−2.4)
37.9
(3.3)
48.9
(9.4)
59.1
(15.1)
68.9
(20.5)
74.0
(23.3)
72.4
(22.4)
64.6
(18.1)
52.5
(11.4)
40.3
(4.6)
27.7
(−2.4)
49.9
(9.9)

Average low °F (°C)
16.5
(−8.6)
20.1
(−6.6)
29.2
(−1.6)
38.8
(3.8)
48.3
(9.1)
58.1
(14.5)
63.9
(17.7)
62.9
(17.2)
54.3
(12.4)
42.8
(6)
32.4
(0.2)
20.7
(−6.3)
40.8
(4.9)

Mean minimum °F (°C)
−5.6
(−20.9)
−0.3
(−17.9)
11.9
(−11.2)
23.5
(−4.7)
35.6
(2)
44.3
(6.8)
52.0
(11.1)
52.6
(11.4)
39.2
(4)
28.4
(−2)
16.9
(−8.4)
−0.1
(−17.8)
−10.9
(−23.8)

Record low °F (°C)
−27
(−33)
−21
(−29)
−12
(−24)
7
(−14)
27
(−3)
35
(2)
45
(7)
42
(6)
29
(−2)
14
(−10)
−2
(−19)
−25
(−32)
−27
(−33)

Average precipitation inches (mm)
1.73
(43.9)
1.79
(45.5)
2.50
(63.5)
3.38
(85.9)
3.68
(93.5)
3.45
(87.6)
3.70
(94)
4.90
(124.5)
3.21
(81.5)
3.15
(80)
3.15
(80)
2.25
(57.2)
36.89
(937)

Average snowfall inches (cm)
10.8
(27.4)
9.1
(23.1)
5.6
(14.2)
1.2
(3)
trace
0
(0)
0
(0)
0
(0)
0
(0)
0.2
(0.5)
1.2
(3)
8.2
(20.8)
36.3
(92.2)

Average precipitation days (≥ 0.01 in)
10.5
8.8
11.1
12.0
11.6
10.2
9.8
9.8
8.3
10.2
10.8
11.0
124.1

Average snowy days (≥ 0.1 in)
8.2
5.9
4.2
0.9
0
0
0
0
0
0.2
1.7
6.9
28.0

Average relative humidity (%)
72.2
71.6
69.7
64.9
64.1
65.6
68.5
70.7
71.1
68.6
72.5
75.5
69.6

Mean monthly sunshine hours
135.8
136.2
187.0
215.3
281.9
311.4
318.4
283.0
226.6
193.2
113.3
106.3
2,508.4

Percent possible sunshine
46
46
51
54
62
68
69
66
60
56
38
37
56

Source: NOAA (relative humidity and sun 1961–1990)[3][8][9]

text:
[hide]Climate data for Chicago Aurora Municipal Airport, Illinois 1981–2010 normals

Month
Jan
Feb
Mar
Apr
May
Jun
Jul
Aug
Sep
Oct
Nov
Dec
Year

Record high °F (°C)
67
(19)
72
(22)
83
(28)
93
(34)
105
(41)
106
(41)
111
(44)
105
(41)
103
(39)
90
(32)
81
(27)
69
(21)
111
(44)

Average high °F (°C)
29.1
(−1.6)
33.7
(0.9)
47.0
(8.3)
60.0
(15.6)
70.6
(21.4)
80.1
(26.7)
83.5
(28.6)
81.2
(27.3)
74.9
(23.8)
62.2
(16.8)
47.8
(8.8)
32.8
(0.4)
58.58
(14.75)

Daily mean °F (°C)
20.8
(−6.2)
24.9
(−3.9)
37.0
(2.8)
48.5
(9.2)
58.8
(14.9)
68.8
(20.4)
71.7
(22.1)
70.1
(21.2)
62.8
(17.1)
50.6
(10.3)
38.1
(3.4)
25.1
(−3.8)
48.1
(8.96)

Average low °F (°C)
12.6
(−10.8)
16.2
(−8.8)
26.9
(−2.8)
37.0
(2.8)
47.1
(8.4)
57.5
(14.2)
59.8
(15.4)
59.0
(15)
50.7
(10.4)
39.0
(3.9)
28.4
(−2)
17.5
(−8.1)
37.64
(3.13)

Record low °F (°C)
−33
(−36)
−25
(−32)
−15
(−26)
8
(−13)
21
(−6)
34
(1)
40
(4)
37
(3)
25
(−4)
11
(−12)
−11
(−24)
−25
(−32)
−33
(−36)

Average precipitation inches (mm)
1.47
(37.3)
1.62
(41.1)
2.23
(56.6)
3.38
(85.9)
4.26
(108.2)
4.01
(101.9)
4.19
(106.4)
4.12
(104.6)
3.44
(87.4)
3.12
(79.2)
3.14
(79.8)
1.96
(49.8)
36.86
(936.2)

Average snowfall inches (cm)
10.0
(25.4)
7.0
(17.8)
3.7
(9.4)
0.8
(2)
0
(0)
0
(0)
0
(0)
0
(0)
0
(0)
0.1
(0.3)
1.4
(3.6)
7.5
(19.1)
30.5
(77.5)

Average precipitation days (≥ 0.01 in)
9.7
7.8
10.4
11.8
11.2
10.2
9.3
10.0
8.9
8.9
10.2
10.7
119.1

Average snowy days (≥ 0.1 in)
6.7
4.5
2.1
0.4
0
0
0
0
0
0.1
1.2
4.9
19.9

Source: NOAA (normals, 1981–2010)[10]

text: Seasons[edit]

text:
Chicago

Climate chart (explanation)

J
F
M
A
M
J
J
A
S
O
N
D

2.1

31
18

1.9

35
21

2.7

47
31

3.6

59
42

4.1

70
52

4.1

80
62

4

84
68

4

82
66

3.3

75
58

3.2

63
46

3.4

49
35

2.6

35
22

Average max. and min. temperatures in °F

Precipitation totals in inches

Source: NOAA

[show]Metric conversion

J
F
M
A
M
J
J
A
S
O
N
D

52

0
−8

49

2
−6

69

8
−1

92

15
5

105

21
11

103

27
17

102

29
20

101

28
19

84

24
14

82

17
8

87

9
1

65

2
−5

Average max. and min. temperatures in °C

Precipitation totals in mm

text: Winter[edit]

text: Winter in Chicago proves quite variable: Seasonal snowfall in the city has ranged from 9.8 inches (24.9 cm) (in 1920–21) up to 89.7 in (228 cm) (in 1978–79), and the average annual snowfall in Chicago is 36 inches (91 cm).[11] Most winters produce many snow falls during the season in light accumulations of around 2 in (5.1 cm). Cities on the other side of Lake Michigan usually receive more snow than Chicago because of the lake-effect snow that falls on these communities, even though northeasterly winds can sometimes bring lake-effect snow to Chicago area too. However, every three years or so during the winter Chicago experiences a heavier snowstorm that can produce over 10 in (25 cm) of snow over a 1- to 3-day period, a level of snowfall very often seen in cities on the "snowbelt" on other side of the lake such as Grand Rapids, Michigan, Kalamazoo, Michigan, and South Bend, Indiana.

text: Winter temperatures can vary tremendously within the span of one week. The daily average high temperature in January at O'Hare is 31.0 °F (−0.6 °C) with the average daily low of 16.5 °F (−8.6 °C) and the daily mean of 23.6 °F (−4.6 °C). Temperatures drop to or below 0 °F (−18 °C) on 5.5 nights annually at Midway and 8.2 nights at O'Hare and up to 10–14 nights in some far western and northern suburbs,[12] although subzero (°F) readings in the absence of snow cover are rare.[13] There have been streaks of multiple winter seasons without a single subzero reading, yet there have also been winters with 20 or more subzero readings.

text: The warming effect of Lake Michigan during the winter makes subzero temperatures somewhat less common on the lakefront than in the more inland parts of the city. Highs reach 50 °F (10 °C) an average of 8.8 days each winter from December to February at Midway.

text: Based on 30-year averages obtained from NOAA's National Climatic Data Center for the months of December, January and February, Weather Channel ranked Chicago the sixth-coldest major U.S. city as of 2014.[14]

text: Although it is extremely rare, temperatures during late winter (generally March 1–20) can reach up to and well over 80 °F (27 °C). In 2012, there were eight days in the month of March with temperatures 80 °F (27 °C)+ (with a ninth day occurring in many suburban areas) during the record-breaking March 2012 North American heat wave. The last couple of 80 °F (27 °C) days in this record-breaking stretch of warmth occurred after the vernal equinox.

text: Spring[edit]

text: Spring in Chicago is perhaps the city's wettest season: Winter conditions can persist well into April and even occasionally into May. Thunderstorms are especially prevalent in the spring time as the city's lakeside location makes it a center of conflicts between large volumes of warm and cold air, which can trigger a wide variety of severe weather.

text: On the other hand, large snowfalls can also occur in late March and in early April. For example, in 1970, over 10 in (25 cm) of snow fell in a storm that occurred on April 1–2.[15] Twelve years later, Opening Day for the Chicago White Sox was postponed due to another 9-inch (23 cm) snowfall that had occurred on April 5.[16] Even more extraordinary, over 18 in (46 cm) of snow fell on March 25–26, 1930, which remains one of the city's five biggest recorded snowstorms despite it occurring past the vernal equinox.[17] The average date for last measurable snowfall (≥0.1 in or 0.25 cm) is April 1.

text: Temperatures can vary tremendously in the springtime; at 100 °F (38 °C), March is the month with the greatest span between the record high and low. At O'Hare, temperatures as low as 7 °F (−14 °C) and 31 °F (−1 °C) have been recorded as late as April 7 and May 21, respectively. Conversely, in official records, the earliest triple-digit (100 °F (38 °C)) high occurred on June 1, 1934, when official readings were taken closer to Lake Michigan. Though rare, triple-digit heat has occurred in May at Midway Airport and in outlying suburban locations. Typically, the last freezing low of the season on average occurs on April 13 at Midway and ten days later at O'Hare.

text: During the springtime, the effects of Lake Michigan are most prevalent. During this season, the lake is still quite cold, as the effects of much warmer temperatures are slow to affect the large body of water of Lake Michigan. It is common for Lake Michigan shoreline and water temperatures to remain in the 40s even into May. If the winds blow from the east, or from Lake Michigan into the city, a wide discrepancy in temperatures in a matter of miles can be found, especially on particularly warm days. It is not uncommon for high temperatures to be officially recorded in the 80s and occasionally 90s at O'Hare, Midway, and in suburban locations but to have temperatures be only in the 40s or 50s along the immediate lakeshore.

text: Summer[edit]

text: On a typical summer day, humidity is usually moderately high, and temperatures ordinarily reach anywhere between 78 and 92 °F (26 and 33 °C). Overnight temperatures in summer usually drop to around 65–70 °F (18–21 °C), although even in July and August there can be several nights where the temperature drops below 60 °F (16 °C), particularly during the cooler summers. Conversely, on the other extreme, temperatures can on a rare basis remain above 80 °F (27 °C) overnight, though this level of overnight warmth is generally limited to the city proper with its urban heat island effects along with Lake Michigan nearby. On such warm nights, especially during strong heat waves, most suburban locations drop down to between 75 and 79 °F (24 and 26 °C) but quickly rebound in the early morning hours. During such strong heat waves, the outlying suburban areas can record temperatures more than 5 °F above city and lakeshore locations.

text: A perfect example of such an occurrence happening, although considered "unofficial", occurred during the Dust Bowl years. Midway Airport recorded a record eight consecutive 100 °F-plus days in July, 1936. In that heat wave, temperatures at the lake remained in the middle and upper 90s Fahrenheit (middle 30s Celsius), whereas Midway Airport recorded temperatures over 100 °F (38 °C) for nearly two weeks, peaking at 107 °F (42 °C) on July 11. The official record high for Chicago for July 11 is also from 1936—but is recorded as just 97 °F. Further west in what would today be the near and far suburbs (e.g. DuPage County and westward), temperatures reached a blistering 110 °F (43 °C) or still higher at points during this massive heat wave. These extreme temperatures are not shown in the official records because daily weather observations were taken at the University of Chicago until 1942. The University of Chicago is close to the lake, which can and does reduce temperatures in the immediate shoreline area in the summer.

text: Chicago's yearly precipitation comes in at an average of about 36 inches (910 mm), but during the summer, rain arises from short-lived hit-or-miss rain rather than actual prolonged rainfalls, and thunderstorms also occur with regularity at night. In a normal summer, temperatures exceed 90 °F (32 °C) on 23 days. Summer is both the rainiest and sunniest season here;[18] only the three months of June through August experience more than 65% of possible sunshine.[9]

text: In July 2012 Chicago reached and exceeded 100 °F (38 °C) for three consecutive days at O'Hare Airport with highs reaching 103 °F (39 °C) in the city and many suburban areas recording temperatures between 105 and 110 °F (41 and 43 °C). It was the first time in 65 years that Chicago had ever seen a triad of 100 °F days.[19] Chicago nearly recorded a fourth consecutive 100 °F day, but the temperature reached 98 degrees at O'Hare in late morning before a cold front came through the area and cooled temperatures off slightly in the city area, holding in the lower 90s. The effects of the cold front did not affect many suburban areas, as temperatures reached or exceeded 100 °F for a fourth consecutive day throughout much of the region.

text: During the summer, Lake Michigan continues to have an effect on Chicago weather, but it is not as common or as strong as it is during the spring months. On very hot days, temperatures can still be cooler along the immediate shoreline and slightly inland of the lake if winds blow from the east. Temperatures can be held in the 70s or 80s in these areas while outlying and suburban areas temperatures are rising well into the 90s. Temperatures can also reach extreme levels of heat on the immediate shoreline, such as when the air temperature reached 105 °F (41 °C) at Northerly Island during the aforementioned July 2012 heat wave.

text: Autumn[edit]

text: The extreme heat that Chicago is capable of experiencing during the height of the summer season can persist into the autumn season. Temperatures have reached 100 °F as late as September 7 (with 99 °F or 37 °C occurring as late as September 29), and temperatures have reached the lower-to-mid-90s Fahrenheit (low 30s Celsius) as late as October 6. Conversely, temperatures have dropped below freezing overnight as early as September 23, and subzero temperatures (below −18 °C) have arrived as early as November 23. Although extremely rare, temperatures at or above 70 °F (21 °C) have been recorded into early December, most recently in 2012 when 70 °F was recorded on December 3, some surrounding areas reaching temperatures as high as 72 °F (22 °C) to 75 °F (24 °C).

text: Autumn, in some ways, is a calmer season than any of the other three in Chicago. This does not mean that wild weather can't happen. The first freeze of the season on average occurs on October 24 at Midway and 11 days earlier at O'Hare. In most years, a period of warm weather, known as an Indian summer, occurs well after the autumnal equinox has occurred. In these so-called Indian summers, unusually warm weather can persist for several days well into October, and in particularly warm autumns, into November. For example, during the 2005 American League Division Series, for both home games between the Chicago White Sox and the Boston Red Sox, temperatures soared to near 90 °F (32 °C), despite the fact that it was already well into October. On top of that, during the same warm stretch in October 2005, for two consecutive days the overnight temperature failed to drop below 70 °F (21 °C), a rare occurrence for Chicago in October.

text: Autumn can bring heavy rain showers, some of which are capable of producing serious flooding. As the winter solstice nears, the threat of a major winter or snowstorm grows, and there have been major winter storms around the Thanksgiving holiday, causing major delays at the city's two major airports. The first measurable (≥0.1 in or 0.25 cm) snow on average falls on November 19. However, in the 2012–13 winter season, the first measurable snow fell on December 20, 2012, eclipsing the previous record of December 17, 1899. The 2012-13 winter season would fail to produce a daily maximum temperature below freezing 32 °F (0 °C) until early January 2013, the first such time that has happened in Chicago weather records. The entire calendar year of 2012 did not record a temperature lower than 5 °F (−15 °C). The largest snowstorm before the winter solstice dropped 14.8 inches (380 mm) at Midway Airport in December 1929.

text: During the autumn, the effects of Lake Michigan are usually reversed from the spring or summer, particularly in the late autumn. Temperatures near the immediate lakeshore can be a few degrees warmer than in outlying areas, especially during nighttime due to the delayed effects of cooler temperatures on the large body of water. It is rare, though possible, during the Indian summer when unusually warm temperatures are occurring in the inland areas of the city and suburbs for temperatures to be somewhat cooler along the lake as often happens during the spring season.

text: Extremes[edit]

text: The highest temperature ever recorded in the Chicago city limits is an unofficial 109 °F (43 °C) on July 24, 1934, at Midway Airport. The official reading of 105 °F (41 °C) for that day was taken at the University of Chicago campus near the shoreline off Lake Michigan. The 105° high that day is the highest official temperature ever recorded in the city.[20] On July 29, 1916, the low temperature sank to only 85 °F (29 °C). Many suburban locations have all-time records that have surpassed 110 °F (43 °C).

text: Readings near the lake can be several degrees cooler than inland locations if lake breezes are present, which suggests that the higher unofficial reading may also be accurate. During the Chicago Heat Wave of 1995, official temperatures reached 104 °F (40 °C) at O'Hare Airport and 106 °F (41 °C) at Midway, but high humidity pushed the heat index to 15–20 °F (8–11 °C) above actual air temperatures, resulting in heat indexes well over 120 °F (49 °C). The coldest temperature ever recorded in Chicago city limits is −27 °F (−33 °C) at O'Hare on January 20, 1985, though unofficial temperatures as low as −33 °F (−36 °C) have been recorded at Chicago Aurora Airport in far western suburbs and in the rural areas to the west of Chicago.[21] On December 24, 1983, and January 18, 1994, the high temperature reached only −11 °F (−24 °C).

text: The greatest 24-hour precipitation was 6.86 in (174.2 mm) at O'Hare on July 23, 2011, while the heaviest 24-hour snowfall was 18.6 in (47.2 cm), again at O'Hare, on January 2, 1999. Daily snow depth has amassed as high as 29 in (74 cm) on January 14, 1979; as of 2018, January 1979 alone holds 8 of the top 10 daily snow-depth measurements.

text: Windy City[edit]

text: Chicago is known as the Windy City. The "Windy City" moniker did not originally refer to Chicago's climate. It is believed to have been created by a New York newspaper writer deriding Chicagoans' bluster as they promoted their city as the site of the 1893 Columbian Exposition. In terms of climate, Chicago is slightly windier than the average American city. Average wind speeds range from 8 miles per hour (13 km/h) in late summer to 12 miles per hour (19 km/h) in spring months.

text: Lake breeze[edit]

text: Chicago can be cooler and moister than other parts of Illinois because of its proximity to the relatively cooler waters of Lake Michigan, effects which are most pronounced during spring and early summer. A frequent lakeshore breeze pushes much cooler, moister air into Chicago than the usual hot air of the Plains States (usually a moist air mass depending on upper level circulation), but the effect can be so localized that only the immediate waterfront neighborhoods (both north and southside lake adjacent communities) are cooler than inland parts of the city. South, west and southwestern suburbs can be more than 20 °F (11 °C) warmer than the lakefront at some times of year.[citation needed] The lake breeze also has other effects, including dense fog spilling into the city. Because of the closed-loop circulation pattern with a lake breeze that moves back and forth across the city, it is thought to significantly increase low-level ozone counts.[22] Differing wind direction on either side of the thermal dividing line allows for sharp updrafts under certain conditions, favorable to thunderstorm development. Offshore or land breezes shutdown the lake breeze and can have the opposite effect, with over lake convection occurring, usually in mid- to late summer and autumn. As a general rule, the opposite trend occurs during the winter—winter temperatures are warmer along the lakeshore and downtown than inland.

text: Notes[edit]

text:
^ Jump up to: a b Mean monthly maxima and minima (i.e. the expected highest and lowest temperature readings at any point during the year or given month) calculated based on data at said location from 1981 to 2010.
Jump up ^ Official records for Chicago were kept at various locations in downtown from January 1871 to 31 December 1925, University of Chicago from 1 January 1926 to 30 June 1942, Midway Airport from 1 July 1942 to 16 January 1980, and at O'Hare Airport since 17 January 1980.[6][7]

tag: ol


document_title: The Young Offenders (TV series)

question_text: the young offenders (tv series) how many episodes

question_tokens: the young offenders ( tv series ) how many episodes

text: The Young Offenders (TV series)

text: Plot[edit]

text: The series follows the adventurous life of Cork teenagers Conor MacSweeney (Alex Murphy) and Jock O'Keeffe (Chris Walley). Their criminal activities often brings themselves into trouble with the police, and their positions at their school is repeatedly put in jeopardy. Their continuous feud with their principal, Barry Walsh (P.J. Gallagher), is well-depicted throughout the series' narrative, and their feelings for Principal Walsh's daughters Siobhan and Linda is also shown through the show.

text: Cast[edit]

text:
Alex Murphy as Conor MacSweeney
Chris Walley as Jock O'Keeffe
Hilary Rose as Mairead MacSweeney
Dominic MacHale as Sergeant Healy
P. J. Gallagher as Principal Barry Walsh
Jennifer Barry as Siobhan Walsh
Demi Isaac Oviawe as Linda Walsh
Shane Casey as Billy Murphy
Chris Kent as Conor's Dad
Orla Fitzgerald as Orla Walsh
Cora Fenton as Jock's Mum
Michael Sands as Jock's Dad

tag: ul

text: Production[edit]

text: The film, The Young Offenders, was released in 2016, which was generally well-reviewed, and won an Irish Film & Television Award. With the popularity of the film, on 9 May 2017, it was confirmed that a six-episode television programme, based on the film with the same name, had been ordered by RTÉ, to be broadcast in 2018.[2] Creator, Peter Foott commented, "It's incredibly exciting to be able to work with the BBC and with RTÉ to bring these characters back to the screen, they have been so supportive of the project and will be a wonderful home for it. The public response to the film, and specifically the characters themselves, was just so overwhelming that we really felt there were a lot more stories to tell".[2]

text: The first series began broadcasting on 1 February 2018, broadcast by RTÉ Two in Ireland and airing through the online television service BBC Three.[3] With a popular and high reception, the series was recommissioned for a second series, the projected release date was not disclosed.[4]

text: Episodes[edit]

text:
No.
Title
Directed by
Written by
Original air date
Overnight
viewership

1
"Episode 1"[5]
Peter Foott
Peter Foott
1 February 2018
277,000[6]

Conor MacSweeney and Jock O'Keeffe are two young students with a passion for criminal activity, including robbery, under-age smoking, threat and other acts. The duo go to the roof of a building and steal lead to sell on. The landlord reports them to the police, and they are chased on bikes. Wearing masks to hide their identity, they ditch their bikes and their school uniform. The police retrieve Conor's jumper, and it is revealed that it belongs to him, due to his name being labelled in it. As they are about to be arrested for robbery, two other school students, wearing the masks that Conor and Jock had worn during the robbery, vandalise the school roof, which clears their names. It is then revealed that Conor and Jock had persuaded the two boys to do so.

2
"Episode 2"[7]
Peter Foott
Peter Foott
8 February 2018
N/A

Conor and Jock are back to their criminal ways, but Conor's true feelings are exposed when he falls for Principal Walsh's daughter Linda. Linda invites the duo over for their barbecue dinner, much to the dismay of Principal Walsh. When Conor tells Linda about his feelings, her father immediately jumps in, and tells Linda to help her mother in the kitchen – who did not need any help at all, Linda tells Connor to wait. A wrestling match pursuits between Conor and Principal Walsh, and the Principal insults him, leading to Conor to retaliate and give his principal a black eye. This leads to an argument between Linda and Conor. Meanwhile, Jock tries to cope with the grief of rejection for the first time, when Linda's sister Siobhan does so.

3
"Episode 3"[8]
Peter Foott
Peter Foott
15 February 2018
N/A

Conor convinces Linda and Siobhan to play truant off school so that he and Jock can convince them to give them a second chance. Conor and Jock steal four bikes for the quartet to ride to a river bridge on, and Conor and Jock strip to their underwear to jump from the bridge into the water, which the two girls are unsatisfied with, and refuse to wear the inappropriate swimsuits that their supposed lovers attempt to give them. Conor packs a picnic, which the girls refuse to eat. Jock confronts Siobhan for her ungratefulness towards his love attempts, and Conor reveals to her that it is the anniversary of Jock's mother's death. Siobhan tells Jock to take them to the graveyard, which he does. Upon their arrival, they see Jock's alcohol-addicted grieving father over his late wife's grave. When he leaves, the four sing a song by her grave, which makes the girls give their lovers another chance. Conor and Jock get the tattoos that they wanted; Conor's mother pre-despises her son's ideas, but her opinions are changed afterwards, when Conor shows her his new tattoo which says "I love me mam", which breaks mother and son into tears. Principal Walsh approves of his daughters' relationships, whilst Jock decides to show gratitude towards his grieving father, having seriously disliked him previously due to his alcoholism.

4
"Episode 4"
Peter Foott
Peter Foott
22 February 2018
N/A

The ancient refrigerator of the MacSweeney household officially dies, and Conor and his mother prepare for a lovely family road trip to collect a new fridge from a farmer for a cheap price. Much to Mairead's dismay, Jock arrives, and asks to go with them. After previously saying no, she is persuaded to do so by Conor and Jock's short upset about this. Jock causes dismay in the car journey, and writes a song, which results in the trio singing it and Mairead crashing the car, causing the front left tyre to pop. Conor has an idea to put a block of wood below the car to remove the rock and the tyre, after none of them can find the jack. Jock then finds it in the front where he was looking shortly afterwards, which causes more frustration to Mairead's head. When they arrive, Mairead purchases the new fridge, and the farmer wants to kill a duck who is causing dismay to his farm animals. Conor and Jock volunteer to do so, and record a video, much to the farmer's joy. However, on the way home, it is revealed that they saved the duck from murder, having faked it's death with the help of Conor and Jock. Jock volunteers to take the duck to his home for hospitality. Mairead drops Jock off home, and tells him to stay away from Conor, and Jock tells Mairead that she is not perfect and that he does not criticise her. Jock leaves the duck in back car seat. As he enters his house, Mairead realises that Jock's drunken father is abusing his son, and storms into the house to save Jock from further abuse. She tells Jock to pack his bags, and takes Jock to live with her and Conor, much to the best friends' delights. The trio, plus the duck, truly bond, and they soon become a happy "family".

5
"Episode 5"
Peter Foott
Peter Foott
1 March 2018
N/A

Money becomes tight in the MacSweeney household when Jock moves in, and the new family trio need to work out a way to keep food last and bring in extra cash. Conor and Jock decide to take their work experience at the fish market where Mairead works, and Jock informs Conor that if they steal a rare blue fin tuna fish for a dodgy businessman, they will receive thousands of pounds, which they would intend to give to Mairead to pay the bills and get the money for her. Conor agrees, and the troublesome duo steal the fish from their place of work at night. After a long attempt to get the tuna into the trolley, they set off to deliver the fish to the businessman. After trying to protect their fish from their rival gang, which results in the tuna falling from the trolley, Jock comes up with a new idea to deliver the fish. He takes a life boat ring and ties the end of the rope to the tuna and puts it in the water to make it easier. When meeting up with Siobhan and Linda, Conor discovers that Jock had toileted in the hand-drier at school which resulted in Conor's hands being covered in faeces. A small argument between the best friends results in Conor dropping the ring into the water, for a swimming-in-training Jock going to retrieve it. Jock agrees to show Siobhan the fish, but when they surface it, it is revealed that the other half of the tuna has been eaten and that it is the wrong fish. The quartet ditch, and Conor and Jock go home to find Sergeant Healy sitting in the kitchen. He shows them a video, which sees Conor and Jock's numerous attempts to get the fish into the trolley. When Mairead arrives home, she shows the official commence of her new relationship with the guard by kissing him, much to the displeasure of Conor and Jock, and she surprisingly insults her son and his best friend, as they could not even steal the correct fish. Meanwhile, Siobhan discovers that she is pregnant with Jock's child.

6
"Episode 6"
Peter Foott
Peter Foott
8 March 2018
N/A

Siobhan prepares to tell Jock about her pregnancy. She tells her mother, as she believes that her mother will be supportive, but the complete opposite happens as her mother is angry with shock. Jock and Conor decide to take Siobhan and Linda to the cinema, in an attempt to get Conor and Linda to "shift". Jock puts on his "fake Billy" mask and steals alcohol from a local shop. The security guards sees this and chases him, but Jock hides and loses the guard. The guard asks Billy if he has seen "fake Billy", which he has not, though sees Jock, wearing his grey tracksuit and hooded to protect seeing his face, and decides to follow him. Jock and Conor wait on the second tier of a double-decker bus for Siobhan and Linda, and Jock drops his mask prior to going up the stairs to the second tier, and Billy gets on bus to search for "fake Billy". Billy realises that Jock is wearing a grey tracksuit, but Jock convinces Billy that it was not him. Siobhan and Linda get on the bus and Billy keeps them hostage after Sergeant Healy sees him in possession of a knife. He orders the bus driver to keep on driving until they lose Healy, but Healy keeps on following them, with the car being driven by Mairead. Healy calls Conor, and Billy answers, and Jock convinces Billy to order Healy to get refreshments, which would tempt Billy to give himself up. Billy feeds the entire bus passengers individual food, whilst Jock serves them with alcohol. Siobhan informs Jock that she is pregnant and shocks him, and he does not speak for a while. The passengers sing a song to pass the time, led by Billy. The bus is blocked off front and back by police and they threaten to kill Billy if he does not give himself up. Conor, feeling sorry for Billy, comes up with a plan. All bus passengers wear the "fake Billy" mask, and they all say "I am Billy" as they come off the bus, Billy escapes by wandering into the woods as the guards question Healy if he was sure he saw Billy get on the bus. Jock informs Mairead that he is becoming a dad, much to her anger, and Siobhan informs her father of her pregnancy. Linda and Conor finally kiss, though inaccurately, before Linda kisses Conor properly before she leaves.

text: Reception[edit]

text: Following the broadcasting of the pilot episode, Aoife Kelly of the Irish Independent released a positive review of the programme, commenting, "It takes the best of the movie and builds on it".[9] Bruce Dessau stated that the programme was a brilliant show to watch, and that he believed that it was a "male version" of Channel 4 series Derry Girls, with other journalists believing that a potential crossover could be produced in the future.[10]


document_title: Uranium

question_text: where did the uranium on earth come from

question_tokens: where did the uranium on earth come from

text: Uranium

text: Characteristics

text: When refined, uranium is a silvery white, weakly radioactive metal. It has a Mohs hardness of 6, sufficient to scratch glass and approximately equal to that of titanium, rhodium, manganese and niobium. It is malleable, ductile, slightly paramagnetic, strongly electropositive and a poor electrical conductor.[8][9] Uranium metal has a very high density of 19.1 g/cm3,[10] denser than lead (11.3 g/cm3),[11] but slightly less dense than tungsten and gold (19.3 g/cm3).[12][13]

text: Uranium metal reacts with almost all non-metal elements (with the exception of the noble gases) and their compounds, with reactivity increasing with temperature.[14] Hydrochloric and nitric acids dissolve uranium, but non-oxidizing acids other than hydrochloric acid attack the element very slowly.[8] When finely divided, it can react with cold water; in air, uranium metal becomes coated with a dark layer of uranium oxide.[9] Uranium in ores is extracted chemically and converted into uranium dioxide or other chemical forms usable in industry.

text: Uranium-235 was the first isotope that was found to be fissile. Other naturally occurring isotopes are fissionable, but not fissile. On bombardment with slow neutrons, its uranium-235 isotope will most of the time divide into two smaller nuclei, releasing nuclear binding energy and more neutrons. If too many of these neutrons are absorbed by other uranium-235 nuclei, a nuclear chain reaction occurs that results in a burst of heat or (in special circumstances) an explosion. In a nuclear reactor, such a chain reaction is slowed and controlled by a neutron poison, absorbing some of the free neutrons. Such neutron absorbent materials are often part of reactor control rods (see nuclear reactor physics for a description of this process of reactor control).

text: As little as 15 lb (7 kg) of uranium-235 can be used to make an atomic bomb.[15] The first nuclear bomb used in war, Little Boy, relied on uranium fission, but the very first nuclear explosive (the Gadget used at Trinity) and the bomb that destroyed Nagasaki (Fat Man) were both plutonium bombs.

text: Uranium metal has three allotropic forms:[16]

text:
α (orthorhombic) stable up to 668 °C. Orthorhombic, space group No. 63, Cmcm, lattice parameters a = 285.4 pm, b = 587 pm, c = 495.5 pm.[17]
β (tetragonal) stable from 668 °C to 775 °C. Tetragonal, space group P42/mnm, P42nm, or P4n2, lattice parameters a = 565.6 pm, b = c = 1075.9 pm.[17]
γ (body-centered cubic) from 775 °C to melting point—this is the most malleable and ductile state. Body-centered cubic, lattice parameter a = 352.4 pm.[17]

text: Applications

text: Military

text: The major application of uranium in the military sector is in high-density penetrators. This ammunition consists of depleted uranium (DU) alloyed with 1–2% other elements, such as titanium or molybdenum.[18] At high impact speed, the density, hardness, and pyrophoricity of the projectile enable the destruction of heavily armored targets. Tank armor and other removable vehicle armor can also be hardened with depleted uranium plates. The use of depleted uranium became politically and environmentally contentious after the use of such munitions by the US, UK and other countries during wars in the Persian Gulf and the Balkans raised questions concerning uranium compounds left in the soil (see Gulf War Syndrome).[15]

text: Depleted uranium is also used as a shielding material in some containers used to store and transport radioactive materials. While the metal itself is radioactive, its high density makes it more effective than lead in halting radiation from strong sources such as radium.[8] Other uses of depleted uranium include counterweights for aircraft control surfaces, as ballast for missile re-entry vehicles and as a shielding material.[9] Due to its high density, this material is found in inertial guidance systems and in gyroscopic compasses.[9] Depleted uranium is preferred over similarly dense metals due to its ability to be easily machined and cast as well as its relatively low cost.[19] The main risk of exposure to depleted uranium is chemical poisoning by uranium oxide rather than radioactivity (uranium being only a weak alpha emitter).

text: During the later stages of World War II, the entire Cold War, and to a lesser extent afterwards, uranium-235 has been used as the fissile explosive material to produce nuclear weapons. Initially, two major types of fission bombs were built: a relatively simple device that uses uranium-235 and a more complicated mechanism that uses plutonium-239 derived from uranium-238. Later, a much more complicated and far more powerful type of fission/fusion bomb (thermonuclear weapon) was built, that uses a plutonium-based device to cause a mixture of tritium and deuterium to undergo nuclear fusion. Such bombs are jacketed in a non-fissile (unenriched) uranium case, and they derive more than half their power from the fission of this material by fast neutrons from the nuclear fusion process.[20]

text: Civilian

text: The main use of uranium in the civilian sector is to fuel nuclear power plants. One kilogram of uranium-235 can theoretically produce about 20 terajoules of energy (2×1013 joules), assuming complete fission; as much energy as 1500 tonnes of coal.[6]

text: Commercial nuclear power plants use fuel that is typically enriched to around 3% uranium-235.[6] The CANDU and Magnox designs are the only commercial reactors capable of using unenriched uranium fuel. Fuel used for United States Navy reactors is typically highly enriched in uranium-235 (the exact values are classified). In a breeder reactor, uranium-238 can also be converted into plutonium through the following reaction:[9]

text: Before (and, occasionally, after) the discovery of radioactivity, uranium was primarily used in small amounts for yellow glass and pottery glazes, such as uranium glass and in Fiestaware.[21]

text: The discovery and isolation of radium in uranium ore (pitchblende) by Marie Curie sparked the development of uranium mining to extract the radium, which was used to make glow-in-the-dark paints for clock and aircraft dials.[22] This left a prodigious quantity of uranium as a waste product, since it takes three tonnes of uranium to extract one gram of radium. This waste product was diverted to the glazing industry, making uranium glazes very inexpensive and abundant. Besides the pottery glazes, uranium tile glazes accounted for the bulk of the use, including common bathroom and kitchen tiles which can be produced in green, yellow, mauve, black, blue, red and other colors.

text: Uranium was also used in photographic chemicals (especially uranium nitrate as a toner),[9] in lamp filaments for stage lighting bulbs,[23] to improve the appearance of dentures,[24] and in the leather and wood industries for stains and dyes. Uranium salts are mordants of silk or wool. Uranyl acetate and uranyl formate are used as electron-dense "stains" in transmission electron microscopy, to increase the contrast of biological specimens in ultrathin sections and in negative staining of viruses, isolated cell organelles and macromolecules.

text: The discovery of the radioactivity of uranium ushered in additional scientific and practical uses of the element. The long half-life of the isotope uranium-238 (4.51×109 years) makes it well-suited for use in estimating the age of the earliest igneous rocks and for other types of radiometric dating, including uranium–thorium dating, uranium–lead dating and uranium–uranium dating. Uranium metal is used for X-ray targets in the making of high-energy X-rays.[9]

text: History

text: Pre-discovery use

text: The use of uranium in its natural oxide form dates back to at least the year 79 CE, when it was used to add a yellow color to ceramic glazes.[9] Yellow glass with 1% uranium oxide was found in a Roman villa on Cape Posillipo in the Bay of Naples, Italy, by R. T. Gunther of the University of Oxford in 1912.[25] Starting in the late Middle Ages, pitchblende was extracted from the Habsburg silver mines in Joachimsthal, Bohemia (now Jáchymov in the Czech Republic), and was used as a coloring agent in the local glassmaking industry.[26] In the early 19th century, the world's only known sources of uranium ore were these mines.

text: Discovery

text: The discovery of the element is credited to the German chemist Martin Heinrich Klaproth. While he was working in his experimental laboratory in Berlin in 1789, Klaproth was able to precipitate a yellow compound (likely sodium diuranate) by dissolving pitchblende in nitric acid and neutralizing the solution with sodium hydroxide.[26] Klaproth assumed the yellow substance was the oxide of a yet-undiscovered element and heated it with charcoal to obtain a black powder, which he thought was the newly discovered metal itself (in fact, that powder was an oxide of uranium).[26][27] He named the newly discovered element after the planet Uranus (named after the primordial Greek god of the sky), which had been discovered eight years earlier by William Herschel.[28]

text: In 1841, Eugène-Melchior Péligot, Professor of Analytical Chemistry at the Conservatoire National des Arts et Métiers (Central School of Arts and Manufactures) in Paris, isolated the first sample of uranium metal by heating uranium tetrachloride with potassium.[26][29]

text: Henri Becquerel discovered radioactivity by using uranium in 1896.[14] Becquerel made the discovery in Paris by leaving a sample of a uranium salt, K2UO2(SO4)2 (potassium uranyl sulfate), on top of an unexposed photographic plate in a drawer and noting that the plate had become "fogged".[30] He determined that a form of invisible light or rays emitted by uranium had exposed the plate.

text: Fission research

text: A team led by Enrico Fermi in 1934 observed that bombarding uranium with neutrons produces the emission of beta rays (electrons or positrons from the elements produced; see beta particle).[31] The fission products were at first mistaken for new elements with atomic numbers 93 and 94, which the Dean of the Faculty of Rome, Orso Mario Corbino, christened ausonium and hesperium, respectively.[32][33][34][35] The experiments leading to the discovery of uranium's ability to fission (break apart) into lighter elements and release binding energy were conducted by Otto Hahn and Fritz Strassmann[31] in Hahn's laboratory in Berlin. Lise Meitner and her nephew, the physicist Otto Robert Frisch, published the physical explanation in February 1939 and named the process "nuclear fission".[36] Soon after, Fermi hypothesized that the fission of uranium might release enough neutrons to sustain a fission reaction. Confirmation of this hypothesis came in 1939, and later work found that on average about 2.5 neutrons are released by each fission of the rare uranium isotope uranium-235.[31] Fermi urged Alfred O.C. Nier to separate uranium isotopes for determination of the fissile component, and on February 29, 1940, Nier used an instrument he built at the University of Minnesota to separate the world's first uranium-235 sample in the Tate Laboratory. After mailed to Columbia University's cyclotron, John Dunning confirmed the sample to be the isolated fissile material on March 1.[37] Further work found that the far more common uranium-238 isotope can be transmuted into plutonium, which, like uranium-235, is also fissile by thermal neutrons. These discoveries led numerous countries to begin working on the development of nuclear weapons and nuclear power.

text: On 2 December 1942, as part of the Manhattan Project, another team led by Enrico Fermi was able to initiate the first artificial self-sustained nuclear chain reaction, Chicago Pile-1. An initial plan using enriched uranium-235 was abandoned as it was as yet unavailable in sufficient quantities.[38] Working in a lab below the stands of Stagg Field at the University of Chicago, the team created the conditions needed for such a reaction by piling together 400 short tons (360 metric tons) of graphite, 58 short tons (53 metric tons) of uranium oxide, and six short tons (5.5 metric tons) of uranium metal, a majority of which was supplied by Westinghouse Lamp Plant in a makeshift production process.[31][39]

text: Nuclear weaponry

text: Two major types of atomic bombs were developed by the United States during World War II: a uranium-based device (codenamed "Little Boy") whose fissile material was highly enriched uranium, and a plutonium-based device (see Trinity test and "Fat Man") whose plutonium was derived from uranium-238. The uranium-based Little Boy device became the first nuclear weapon used in war when it was detonated over the Japanese city of Hiroshima on 6 August 1945. Exploding with a yield equivalent to 12,500 tonnes of TNT, the blast and thermal wave of the bomb destroyed nearly 50,000 buildings and killed approximately 75,000 people (see Atomic bombings of Hiroshima and Nagasaki).[30] Initially it was believed that uranium was relatively rare, and that nuclear proliferation could be avoided by simply buying up all known uranium stocks, but within a decade large deposits of it were discovered in many places around the world.[40]

text: Reactors

text: The X-10 Graphite Reactor at Oak Ridge National Laboratory (ORNL) in Oak Ridge, Tennessee, formerly known as the Clinton Pile and X-10 Pile, was the world's second artificial nuclear reactor (after Enrico Fermi's Chicago Pile) and was the first reactor designed and built for continuous operation. Argonne National Laboratory's Experimental Breeder Reactor I, located at the Atomic Energy Commission's National Reactor Testing Station near Arco, Idaho, became the first nuclear reactor to create electricity on 20 December 1951.[41] Initially, four 150-watt light bulbs were lit by the reactor, but improvements eventually enabled it to power the whole facility (later, the town of Arco became the first in the world to have all its electricity come from nuclear power generated by BORAX-III, another reactor designed and operated by Argonne National Laboratory).[42][43] The world's first commercial scale nuclear power station, Obninsk in the Soviet Union, began generation with its reactor AM-1 on 27 June 1954. Other early nuclear power plants were Calder Hall in England, which began generation on 17 October 1956,[44] and the Shippingport Atomic Power Station in Pennsylvania, which began on 26 May 1958. Nuclear power was used for the first time for propulsion by a submarine, the USS Nautilus, in 1954.[31][45]

text: Prehistoric naturally occurring fission

text: In 1972, the French physicist Francis Perrin discovered fifteen ancient and no longer active natural nuclear fission reactors in three separate ore deposits at the Oklo mine in Gabon, West Africa, collectively known as the Oklo Fossil Reactors. The ore deposit is 1.7 billion years old; then, uranium-235 constituted about 3% of the total uranium on Earth.[46] This is high enough to permit a sustained nuclear fission chain reaction to occur, provided other supporting conditions exist. The capacity of the surrounding sediment to contain the nuclear waste products has been cited by the U.S. federal government as supporting evidence for the feasibility to store spent nuclear fuel at the Yucca Mountain nuclear waste repository.[46]

text: Contamination and the Cold War legacy

text: Above-ground nuclear tests by the Soviet Union and the United States in the 1950s and early 1960s and by France into the 1970s and 1980s[19] spread a significant amount of fallout from uranium daughter isotopes around the world.[47] Additional fallout and pollution occurred from several nuclear accidents.[48]

text: Uranium miners have a higher incidence of cancer. An excess risk of lung cancer among Navajo uranium miners, for example, has been documented and linked to their occupation.[49] The Radiation Exposure Compensation Act, a 1990 law in the USA, required $100,000 in "compassion payments" to uranium miners diagnosed with cancer or other respiratory ailments.[50]

text: During the Cold War between the Soviet Union and the United States, huge stockpiles of uranium were amassed and tens of thousands of nuclear weapons were created using enriched uranium and plutonium made from uranium. Since the break-up of the Soviet Union in 1991, an estimated 600 short tons (540 metric tons) of highly enriched weapons grade uranium (enough to make 40,000 nuclear warheads) have been stored in often inadequately guarded facilities in the Russian Federation and several other former Soviet states.[15] Police in Asia, Europe, and South America on at least 16 occasions from 1993 to 2005 have intercepted shipments of smuggled bomb-grade uranium or plutonium, most of which was from ex-Soviet sources.[15] From 1993 to 2005 the Material Protection, Control, and Accounting Program, operated by the federal government of the United States, spent approximately US $550 million to help safeguard uranium and plutonium stockpiles in Russia.[15] This money was used for improvements and security enhancements at research and storage facilities. Scientific American reported in February 2006 that in some of the facilities security consisted of chain link fences which were in severe states of disrepair. According to an interview from the article, one facility had been storing samples of enriched (weapons grade) uranium in a broom closet before the improvement project; another had been keeping track of its stock of nuclear warheads using index cards kept in a shoe box.[51]

text: Occurrence

text: Origin

text: Along with all elements having atomic weights higher than that of iron, uranium is only naturally formed in supernovae.[52] Primordial thorium and uranium are only produced in the r-process (rapid neutron capture), because the s-process (slow neutron capture) is too slow and cannot pass the gap of instability after bismuth.[53][54] Besides the two extant primordial uranium isotopes, 235U and 238U, the r-process also produced significant quantities of 236U, which has a shorter half-life and has long since decayed completely to 232Th, which was itself enriched by the decay of 244Pu, accounting for the observed higher-than-expected abundance of thorium and lower-than-expected abundance of uranium.[55] While the natural abundance of uranium has been supplemented by the decay of extinct 242Pu (half-life 0.375 million years) and 247Cm (half-life 16 million years), producing 238U and 235U respectively, this occurred to an almost negligible extent due to the shorter half-lives of these parents and their lower production than 236U and 244Pu, the parents of thorium: the 247Cm:235U ratio at the formation of the Solar System was (7.0 ± 1.6) × 10−5.[56]

text: Biotic and abiotic

text: Uranium is a naturally occurring element that can be found in low levels within all rock, soil, and water. Uranium is the 51st element in order of abundance in the Earth's crust. Uranium is also the highest-numbered element to be found naturally in significant quantities on Earth and is almost always found combined with other elements.[9] The decay of uranium, thorium, and potassium-40 in the Earth's mantle is thought to be the main source of heat[57][58] that keeps the outer core liquid and drives mantle convection, which in turn drives plate tectonics.

text: Uranium's average concentration in the Earth's crust is (depending on the reference) 2 to 4 parts per million,[8][19] or about 40 times as abundant as silver.[14] The Earth's crust from the surface to 25 km (15 mi) down is calculated to contain 1017 kg (2×1017 lb) of uranium while the oceans may contain 1013 kg (2×1013 lb).[8] The concentration of uranium in soil ranges from 0.7 to 11 parts per million (up to 15 parts per million in farmland soil due to use of phosphate fertilizers), and its concentration in sea water is 3 parts per billion.[19]

text: Uranium is more plentiful than antimony, tin, cadmium, mercury, or silver, and it is about as abundant as arsenic or molybdenum.[9][19] Uranium is found in hundreds of minerals, including uraninite (the most common uranium ore), carnotite, autunite, uranophane, torbernite, and coffinite.[9] Significant concentrations of uranium occur in some substances such as phosphate rock deposits, and minerals such as lignite, and monazite sands in uranium-rich ores[9] (it is recovered commercially from sources with as little as 0.1% uranium[14]).

text: Some bacteria, such as Shewanella putrefaciens, Geobacter metallireducens and some strains of Burkholderia fungorum, use uranium for their growth and convert U(VI) to U(IV).[59][60]

text: Some organisms, such as the lichen Trapelia involuta or microorganisms such as the bacterium Citrobacter, can absorb concentrations of uranium that are up to 300 times the level of their environment.[61] Citrobacter species absorb uranyl ions when given glycerol phosphate (or other similar organic phosphates). After one day, one gram of bacteria can encrust themselves with nine grams of uranyl phosphate crystals; this creates the possibility that these organisms could be used in bioremediation to decontaminate uranium-polluted water.[26][62] The proteobacterium Geobacter has also been shown to bioremediate uranium in ground water.[63] The mycorrhizal fungus Glomus intraradices increases uranium content in the roots of its symbiotic plant.[64]

text: In nature, uranium(VI) forms highly soluble carbonate complexes at alkaline pH. This leads to an increase in mobility and availability of uranium to groundwater and soil from nuclear wastes which leads to health hazards. However, it is difficult to precipitate uranium as phosphate in the presence of excess carbonate at alkaline pH. A Sphingomonas sp. strain BSAR-1 has been found to express a high activity alkaline phosphatase (PhoK) that has been applied for bioprecipitation of uranium as uranyl phosphate species from alkaline solutions. The precipitation ability was enhanced by overexpressing PhoK protein in E. coli.[65]

text: Plants absorb some uranium from soil. Dry weight concentrations of uranium in plants range from 5 to 60 parts per billion, and ash from burnt wood can have concentrations up to 4 parts per million.[26] Dry weight concentrations of uranium in food plants are typically lower with one to two micrograms per day ingested through the food people eat.[26]

text: Production and mining

text: Worldwide production of U3O8 (yellowcake) in 2013 amounted to 70,015 tonnes, of which 22,451 t (32%) was mined in Kazakhstan. Other important uranium mining countries are Canada (9,331 t), Australia (6,350 t), Niger (4,518 t), Namibia (4,323 t) and Russia (3,135 t).[66]

text: Uranium ore is mined in several ways: by open pit, underground, in-situ leaching, and borehole mining (see uranium mining).[6] Low-grade uranium ore mined typically contains 0.01 to 0.25% uranium oxides. Extensive measures must be employed to extract the metal from its ore.[67] High-grade ores found in Athabasca Basin deposits in Saskatchewan, Canada can contain up to 23% uranium oxides on average.[68] Uranium ore is crushed and rendered into a fine powder and then leached with either an acid or alkali. The leachate is subjected to one of several sequences of precipitation, solvent extraction, and ion exchange. The resulting mixture, called yellowcake, contains at least 75% uranium oxides U3O8. Yellowcake is then calcined to remove impurities from the milling process before refining and conversion.[69]

text:
Commercial-grade uranium can be produced through the reduction of uranium halides with alkali or alkaline earth metals.[9] Uranium metal can also be prepared through electrolysis of KUF
5 or UF
4, dissolved in molten calcium chloride (CaCl
2) and sodium chloride (NaCl) solution.[9] Very pure uranium is produced through the thermal decomposition of uranium halides on a hot filament.[9]

text: Resources and reserves

text: It is estimated that 5.5 million tonnes of uranium exists in ore reserves that are economically viable at US$59 per lb of uranium,[70] while 35 million tonnes are classed as mineral resources (reasonable prospects for eventual economic extraction).[71] Prices went from about $10/lb in May 2003 to $138/lb in July 2007. This has caused a big increase in spending on exploration,[70] with US$200 million being spent worldwide in 2005, a 54% increase on the previous year.[71] This trend continued through 2006, when expenditure on exploration rocketed to over $774 million, an increase of over 250% compared to 2004. The OECD Nuclear Energy Agency said exploration figures for 2007 would likely match those for 2006.[70]

text: Australia has 31% of the world's known uranium ore reserves[72] and the world's largest single uranium deposit, located at the Olympic Dam Mine in South Australia.[73] There is a significant reserve of uranium in Bakouma a sub-prefecture in the prefecture of Mbomou in Central African Republic.

text: Some nuclear fuel comes from nuclear weapons being dismantled,[74] such as from the Megatons to Megawatts Program.

text: An additional 4.6 billion tonnes of uranium are estimated to be in sea water (Japanese scientists in the 1980s showed that extraction of uranium from sea water using ion exchangers was technically feasible).[75][76] There have been experiments to extract uranium from sea water,[77] but the yield has been low due to the carbonate present in the water. In 2012, ORNL researchers announced the successful development of a new absorbent material dubbed HiCap which performs surface retention of solid or gas molecules, atoms or ions and also effectively removes toxic metals from water, according to results verified by researchers at Pacific Northwest National Laboratory.[78][79]

text: Supplies

text: This section needs to be updated. Please update this article to reflect recent events or newly available information. (September 2016)

text: In 2005, seventeen countries produced concentrated uranium oxides: Canada (27.9% of world production), Australia (22.8%), Kazakhstan (10.5%), Russia (8.0%), Namibia (7.5%), Niger (7.4%), Uzbekistan (5.5%), the United States (2.5%), Argentina (2.1%), Ukraine (1.9%) and China (1.7%).[81] Kazakhstan continues to increase production and may have become the world's largest producer of uranium by 2009 with an expected production of 12,826 tonnes, compared to Canada with 11,100 t and Australia with 9,430 t.[82][83] In the late 1960s, UN geologists also discovered major uranium deposits and other rare mineral reserves in Somalia. The find was the largest of its kind, with industry experts estimating the deposits at over 25% of the world's then known uranium reserves of 800,000 tons.[84]

text: The ultimate available supply is believed to be sufficient for at least the next 85 years,[71] although some studies indicate underinvestment in the late twentieth century may produce supply problems in the 21st century.[85] Uranium deposits seem to be log-normal distributed. There is a 300-fold increase in the amount of uranium recoverable for each tenfold decrease in ore grade.[86] In other words, there is little high grade ore and proportionately much more low grade ore available.

text: Compounds

text: Oxidation states and oxides

text: Oxides

text:
Calcined uranium yellowcake, as produced in many large mills, contains a distribution of uranium oxidation species in various forms ranging from most oxidized to least oxidized. Particles with short residence times in a calciner will generally be less oxidized than those with long retention times or particles recovered in the stack scrubber. Uranium content is usually referenced to U
3O
8, which dates to the days of the Manhattan Project when U
3O
8 was used as an analytical chemistry reporting standard.

text:
Phase relationships in the uranium-oxygen system are complex. The most important oxidation states of uranium are uranium(IV) and uranium(VI), and their two corresponding oxides are, respectively, uranium dioxide (UO
2) and uranium trioxide (UO
3).[87] Other uranium oxides such as uranium monoxide (UO), diuranium pentoxide (U
2O
5), and uranium peroxide (UO
4·2H
2O) also exist.

text:
The most common forms of uranium oxide are triuranium octoxide (U
3O
8) and UO
2.[88] Both oxide forms are solids that have low solubility in water and are relatively stable over a wide range of environmental conditions. Triuranium octoxide is (depending on conditions) the most stable compound of uranium and is the form most commonly found in nature. Uranium dioxide is the form in which uranium is most commonly used as a nuclear reactor fuel.[88] At ambient temperatures, UO
2 will gradually convert to U
3O
8. Because of their stability, uranium oxides are generally considered the preferred chemical form for storage or disposal.[88]

text: Aqueous chemistry

text:
Salts of many oxidation states of uranium are water-soluble and may be studied in aqueous solutions. The most common ionic forms are U3+ (brown-red), U4+ (green), UO+
2 (unstable), and UO2+
2 (yellow), for U(III), U(IV), U(V), and U(VI), respectively.[89] A few solid and semi-metallic compounds such as UO and US exist for the formal oxidation state uranium(II), but no simple ions are known to exist in solution for that state. Ions of U3+ liberate hydrogen from water and are therefore considered to be highly unstable. The UO2+
2 ion represents the uranium(VI) state and is known to form compounds such as uranyl carbonate, uranyl chloride and uranyl sulfate. UO2+
2 also forms complexes with various organic chelating agents, the most commonly encountered of which is uranyl acetate.[89]

text: Unlike the uranyl salts of uranium and polyatomic ion uranium-oxide cationic forms, the uranates, salts containing a polyatomic uranium-oxide anion, are generally not water-soluble.

text: Carbonates

text: The interactions of carbonate anions with uranium(VI) cause the Pourbaix diagram to change greatly when the medium is changed from water to a carbonate containing solution. While the vast majority of carbonates are insoluble in water (students are often taught that all carbonates other than those of alkali metals are insoluble in water), uranium carbonates are often soluble in water. This is because a U(VI) cation is able to bind two terminal oxides and three or more carbonates to form anionic complexes.

text:
Pourbaix diagrams[90]

Uranium in a non-complexing aqueous medium (e.g. perchloric acid/sodium hydroxide).[90]
Uranium in carbonate solution

Relative concentrations of the different chemical forms of uranium in a non-complexing aqueous medium (e.g. perchloric acid/sodium hydroxide).[90]
Relative concentrations of the different chemical forms of uranium in an aqueous carbonate solution.[90]

text: Effects of pH

text: The uranium fraction diagrams in the presence of carbonate illustrate this further: when the pH of a uranium(VI) solution increases, the uranium is converted to a hydrated uranium oxide hydroxide and at high pHs it becomes an anionic hydroxide complex.

text: When carbonate is added, uranium is converted to a series of carbonate complexes if the pH is increased. One effect of these reactions is increased solubility of uranium in the pH range 6 to 8, a fact that has a direct bearing on the long term stability of spent uranium dioxide nuclear fuels.

text: Hydrides, carbides and nitrides

text: Uranium metal heated to 250 to 300 °C (482 to 572 °F) reacts with hydrogen to form uranium hydride. Even higher temperatures will reversibly remove the hydrogen. This property makes uranium hydrides convenient starting materials to create reactive uranium powder along with various uranium carbide, nitride, and halide compounds.[91] Two crystal modifications of uranium hydride exist: an α form that is obtained at low temperatures and a β form that is created when the formation temperature is above 250 °C.[91]

text:
Uranium carbides and uranium nitrides are both relatively inert semimetallic compounds that are minimally soluble in acids, react with water, and can ignite in air to form U
3O
8.[91] Carbides of uranium include uranium monocarbide (UC), uranium dicarbide (UC
2), and diuranium tricarbide (U
2C
3). Both UC and UC
2 are formed by adding carbon to molten uranium or by exposing the metal to carbon monoxide at high temperatures. Stable below 1800 °C, U
2C
3 is prepared by subjecting a heated mixture of UC and UC
2 to mechanical stress.[92] Uranium nitrides obtained by direct exposure of the metal to nitrogen include uranium mononitride (UN), uranium dinitride (UN
2), and diuranium trinitride (U
2N
3).[92]

text: Halides

text:
All uranium fluorides are created using uranium tetrafluoride (UF
4); UF
4 itself is prepared by hydrofluorination of uranium dioxide.[91] Reduction of UF
4 with hydrogen at 1000 °C produces uranium trifluoride (UF
3). Under the right conditions of temperature and pressure, the reaction of solid UF
4 with gaseous uranium hexafluoride (UF
6) can form the intermediate fluorides of U
2F
9, U
4F
17, and UF
5.[91]

text:
At room temperatures, UF
6 has a high vapor pressure, making it useful in the gaseous diffusion process to separate the rare uranium-235 from the common uranium-238 isotope. This compound can be prepared from uranium dioxide and uranium hydride by the following process:[91]

text:
The resulting UF
6, a white solid, is highly reactive (by fluorination), easily sublimes (emitting a vapor that behaves as a nearly ideal gas), and is the most volatile compound of uranium known to exist.[91]

text:
One method of preparing uranium tetrachloride (UCl
4) is to directly combine chlorine with either uranium metal or uranium hydride. The reduction of UCl
4 by hydrogen produces uranium trichloride (UCl
3) while the higher chlorides of uranium are prepared by reaction with additional chlorine.[91] All uranium chlorides react with water and air.

text:
Bromides and iodides of uranium are formed by direct reaction of, respectively, bromine and iodine with uranium or by adding UH
3 to those element's acids.[91] Known examples include: UBr
3, UBr
4, UI
3, and UI
4. Uranium oxyhalides are water-soluble and include UO
2F
2, UOCl
2, UO
2Cl
2, and UO
2Br
2. Stability of the oxyhalides decrease as the atomic weight of the component halide increases.[91]

text: Isotopes

text: Natural concentrations

text: Natural uranium consists of three major isotopes: uranium-238 (99.28% natural abundance), uranium-235 (0.71%), and uranium-234 (0.0054%). All three are radioactive, emitting alpha particles, with the exception that all three of these isotopes have small probabilities of undergoing spontaneous fission, rather than alpha emission. There are also five other trace isotopes: uranium-239, which is formed when 238U undergoes spontaneous fission, releasing neutrons that are captured by another 238U atom; uranium-237, which is formed when 238U captures a neutron but emits two more, which then decays to neptunium-237; and finally, uranium-233, which is formed in the decay chain of that neptunium-237. It is also expected that thorium-232 should be able to undergo double beta decay, which would produce uranium-232, but this has not yet been observed experimentally.[93]

text: Uranium-238 is the most stable isotope of uranium, with a half-life of about 4.468×109 years, roughly the age of the Earth. Uranium-235 has a half-life of about 7.13×108 years, and uranium-234 has a half-life of about 2.48×105 years.[94] For natural uranium, about 49% of its alpha rays are emitted by each of 238U atom, and also 49% by 234U (since the latter is formed from the former) and about 2.0% of them by the 235U. When the Earth was young, probably about one-fifth of its uranium was uranium-235, but the percentage of 234U was probably much lower than this.

text: Uranium-238 is usually an α emitter (occasionally, it undergoes spontaneous fission), decaying through the "Uranium Series" of nuclear decay, which has 18 members, into lead-206, by a variety of different decay paths.[14]

text: The decay series of 235U, which is called the actinium series, has 15 members and eventually decays into lead-207.[14] The constant rates of decay in these decay series makes the comparison of the ratios of parent to daughter elements useful in radiometric dating.

text: Uranium-234, which is a member of the uranium series (the decay chain of uranium-238), decays to lead-206 through a series of relatively short-lived isotopes.

text: Uranium-233 is made from thorium-232 by neutron bombardment, usually in a nuclear reactor, and 233U is also fissile.[9] Its decay series ends at bismuth-209 and thallium-205.

text: Uranium-235 is important for both nuclear reactors and nuclear weapons, because it is the only uranium isotope existing in nature on Earth in any significant amount that is fissile. This means that it can be split into two or three fragments (fission products) by thermal neutrons.[14]

text: Uranium-238 is not fissile, but is a fertile isotope, because after neutron activation it can produce plutonium-239, another fissile isotope. Indeed, the 238U nucleus can absorb one neutron to produce the radioactive isotope uranium-239. 239U decays by beta emission to neptunium-239, also a beta-emitter, that decays in its turn, within a few days into plutonium-239. 239Pu was used as fissile material in the first atomic bomb detonated in the "Trinity test" on 15 July 1945 in New Mexico.[31]

text: Enrichment

text: In nature, uranium is found as uranium-238 (99.2742%) and uranium-235 (0.7204%). Isotope separation concentrates (enriches) the fissionable uranium-235 for nuclear weapons and most nuclear power plants, except for gas cooled reactors and pressurised heavy water reactors. Most neutrons released by a fissioning atom of uranium-235 must impact other uranium-235 atoms to sustain the nuclear chain reaction. The concentration and amount of uranium-235 needed to achieve this is called a 'critical mass'.

text: To be considered 'enriched', the uranium-235 fraction should be between 3% and 5%.[95] This process produces huge quantities of uranium that is depleted of uranium-235 and with a correspondingly increased fraction of uranium-238, called depleted uranium or 'DU'. To be considered 'depleted', the uranium-235 isotope concentration should be no more than 0.3%.[96] The price of uranium has risen since 2001, so enrichment tailings containing more than 0.35% uranium-235 are being considered for re-enrichment, driving the price of depleted uranium hexafluoride above $130 per kilogram in July 2007 from $5 in 2001.[96]

text:
The gas centrifuge process, where gaseous uranium hexafluoride (UF
6) is separated by the difference in molecular weight between 235UF6 and 238UF6 using high-speed centrifuges, is the cheapest and leading enrichment process.[30] The gaseous diffusion process had been the leading method for enrichment and was used in the Manhattan Project. In this process, uranium hexafluoride is repeatedly diffused through a silver-zinc membrane, and the different isotopes of uranium are separated by diffusion rate (since uranium 238 is heavier it diffuses slightly slower than uranium-235).[30] The molecular laser isotope separation method employs a laser beam of precise energy to sever the bond between uranium-235 and fluorine. This leaves uranium-238 bonded to fluorine and allows uranium-235 metal to precipitate from the solution.[6] An alternative laser method of enrichment is known as atomic vapor laser isotope separation (AVLIS) and employs visible tunable lasers such as dye lasers.[97] Another method used is liquid thermal diffusion.[8]

text: Human exposure

text: A person can be exposed to uranium (or its radioactive daughters, such as radon) by inhaling dust in air or by ingesting contaminated water and food. The amount of uranium in air is usually very small; however, people who work in factories that process phosphate fertilizers, live near government facilities that made or tested nuclear weapons, live or work near a modern battlefield where depleted uranium weapons have been used, or live or work near a coal-fired power plant, facilities that mine or process uranium ore, or enrich uranium for reactor fuel, may have increased exposure to uranium.[98][99] Houses or structures that are over uranium deposits (either natural or man-made slag deposits) may have an increased incidence of exposure to radon gas. The Occupational Safety and Health Administration (OSHA) has set the permissible exposure limit for uranium exposure in the workplace as 0.25 mg/m3 over an 8-hour workday. The National Institute for Occupational Safety and Health (NIOSH) has set a recommended exposure limit (REL) of 0.2 mg/m3 over an 8-hour workday and a short-term limit of 0.6 mg/m3. At levels of 10 mg/m3, uranium is immediately dangerous to life and health.[100]

text: Most ingested uranium is excreted during digestion. Only 0.5% is absorbed when insoluble forms of uranium, such as its oxide, are ingested, whereas absorption of the more soluble uranyl ion can be up to 5%.[26] However, soluble uranium compounds tend to quickly pass through the body, whereas insoluble uranium compounds, especially when inhaled by way of dust into the lungs, pose a more serious exposure hazard. After entering the bloodstream, the absorbed uranium tends to bioaccumulate and stay for many years in bone tissue because of uranium's affinity for phosphates.[26] Uranium is not absorbed through the skin, and alpha particles released by uranium cannot penetrate the skin.

text: Incorporated uranium becomes uranyl ions, which accumulate in bone, liver, kidney, and reproductive tissues. Uranium can be decontaminated from steel surfaces[101] and aquifers.[102]

text: Effects and precautions

text:
Normal functioning of the kidney, brain, liver, heart, and other systems can be affected by uranium exposure, because, besides being weakly radioactive, uranium is a toxic metal.[26][103][104] Uranium is also a reproductive toxicant.[105][106] Radiological effects are generally local because alpha radiation, the primary form of 238U decay, has a very short range, and will not penetrate skin. Alpha radiation from inhaled uranium has been demonstrated to cause lung cancer in exposed nuclear workers.[107] Uranyl (UO2+
2) ions, such as from uranium trioxide or uranyl nitrate and other hexavalent uranium compounds, have been shown to cause birth defects and immune system damage in laboratory animals.[108] While the CDC has published one study that no human cancer has been seen as a result of exposure to natural or depleted uranium,[109] exposure to uranium and its decay products, especially radon, are widely known and significant health threats.[19] Exposure to strontium-90, iodine-131, and other fission products is unrelated to uranium exposure, but may result from medical procedures or exposure to spent reactor fuel or fallout from nuclear weapons.[110] Although accidental inhalation exposure to a high concentration of uranium hexafluoride has resulted in human fatalities, those deaths were associated with the generation of highly toxic hydrofluoric acid and uranyl fluoride rather than with uranium itself.[111] Finely divided uranium metal presents a fire hazard because uranium is pyrophoric; small grains will ignite spontaneously in air at room temperature.[9]

text: Uranium metal is commonly handled with gloves as a sufficient precaution.[112] Uranium concentrate is handled and contained so as to ensure that people do not inhale or ingest it.[112]

text:
Compilation of 2004 review on uranium toxicity[103]

Body system
Human studies
Animal studies
In vitro

Renal
Elevated levels of protein excretion, urinary catalase and diuresis
Damage to proximal convoluted tubules, necrotic cells cast from tubular epithelium, glomerular changes (renal failure)
No studies

Brain/CNS
Decreased performance on neurocognitive tests
Acute cholinergic toxicity; Dose-dependent accumulation in cortex, midbrain, and vermis; Electrophysiological changes in hippocampus
No studies

DNA
Increased reports of cancers[50][113][114][115][116][117]
Increased mutagenicity (in mice) and induction of tumors
Binucleated cells with micronuclei, Inhibition of cell cycle kinetics and proliferation; Sister chromatid induction, tumorigenic phenotype

Bone/muscle
No studies
Inhibition of periodontal bone formation; and alveolar wound healing
No studies

Reproductive
Uranium miners have more first-born female children
Moderate to severe focal tubular atrophy; vacuolization of Leydig cells
No studies

Lungs/respiratory
No adverse health effects reported
Severe nasal congestion and hemorrhage, lung lesions and fibrosis, edema and swelling, lung cancer
No studies

Gastrointestinal
Vomiting, diarrhea, albuminuria
No studies
No studies

Liver
No effects seen at exposure dose
Fatty livers, focal necrosis
No studies

Skin
No exposure assessment data available
Swollen vacuolated epidermal cells, damage to hair follicles and sebaceous glands
No studies

Tissues surrounding embedded DU fragments
Elevated uranium urine concentrations
Elevated uranium urine concentrations, perturbations in biochemical and neuropsychological testing
No studies

Immune system
Chronic fatigue, rash, ear and eye infections, hair and weight loss, cough. May be due to combined chemical exposure rather than DU alone
No studies
No studies

Eyes
No studies
Conjunctivitis, irritation inflammation, edema, ulceration of conjunctival sacs
No studies

Blood
No studies
Decrease in RBC count and hemoglobin concentration
No studies

Cardiovascular
Myocarditis resulting from the uranium ingestion, which ended six months after ingestion
No effects
No studies