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INORGANIC CHEMISTRY[Type the document subtitle][Type the abstract of the document here. The abstract is typically a short summary of the contents of the document. Type the abstract of the document here. The abstract is typically a short summary of the contents of the document.]
[Year]
Meiliza Ekayanti[Type the company name]
[Pick the date]
2
PREFACE
Bismillahirrahmanirrahim.
Assalamu’alaikum Wr. Wb.
Alhamdulillahirabbil’alamin, praise and our thanks gives presence of God
who is praised and most high. that has gave us enjoy islam, enjoy faith and enjoy
healthy that no predicted. Shalawat and greeting not forgets us submit to our big
prophet, Mohammad saw that has brought us from stupidity era to science era.
We render thanks to the whole party that has supported and helped us
until we can finish this paper. Especially to both our old fellow, our lecturer Drs.
Agung Purwanto, M.Si., friends and other parties that can not we mention one by
one.
“Nothing are perfect except God.”. Maybe that words that can depict this
paper. Until we open our liver door to accept criticism and suggestion from all
parties, in order to later, in paper hereinafter can be better next.
And we apologize if existed insuffiency and mistake in this paper making.
Because all that correctness is only God property, and all wrong ones is ours
person.
Akhirul kalam.
Wassalamu’alaikum Wr. Wb.
Jakarta, October 2010
Chemistry 2009
COMPILERS
Atomic Theory and Elements Periodic System
1. Ratna Purnama Sari2. Sri Astriani3. Endah Dianty
Axial component of angular momentum are allowed only two
values, +1 / 2 (h/2π) and -1 / 2 (h/2π). Spin magnetic quantum
number associated with this value (m s = +1 / 2 or -1 / 2). Only the
spin quantum number alone in an amount not rounded.
21
Table 2.3 Numbers of quantum
Name (quantum
number)
symbo
l
Values allowed
Main n 1, 2, 3, ...
Azimuth l 0, 1, 2, 3, ... n -
1
Magnetic m
(ml)
0, ± 1, ± 2 ,...±
l
Magnetic spin ms +1 / 2, -1 / 2
Other symbols as given in Table 2.4 are commonly used instead.
Hydrogen atom energy or hydrogen-like atom is determined only by
the principal quantum number and energy equations that express
identical to that already derived from the theory of Bohr.
Table 2.4 Symbols azimuth quantum number
value
0
1
2
3
4
symb
ol
s
p
d
f
g
22
Electron wave functions called orbital. When the main quantum
number n = 1, there is only one value of l, namely 0. In this case
there is only one orbital, and a collection of the orbital quantum
number for this is (n = 1, l = 0). When n = 2, there are two values
l, 0 and 1, are allowed. In case there are four orbital by defined set
of quantum numbers: (n = 2, l = 0), (n = 2, l = 1, m = -1), (n = 2, l
= 1, m = 0) , (n = 2, l = 1, m = +1).
2.2.5 Configuration Electron Atom
When atoms contained more than two electrons, the interaction
between electrons must be considered, and difficult to solve the
wave equation of this system is very complicated. With the
assumption that each electron in the poly-electron atom will move
in an electric field which is roughly symmetric orbital for each
electron can be defined by three quantum numbers n, l and m and
the number quantum spin m s, as in the case of hydrogen-like atom .
Hydrogen-like atomic energy is determined only by the principal
quantum number n, but for poly-electron atoms are mainly
determined by n and l. When the atoms have the same quantum
number n, the larger l, the higher the energy.
2.2.6 Pauli Exclusion Principle
According to the Pauli exclusion principle, only one electron in
an atom occupy permitted circumstances defined by a particular set
of four quantum numbers, or, at most two electrons can occupy an
23
orbital which is defined by three quantum numbers n, l and m. The
two electrons that must have a value of m s are different, in other
words its spin antiparallel, and the pair of electrons is called with
a pair of electrons.
Groups of electrons with the same value of n is called with the skin
or skin electron. Notation used for the outer electrons are given
in Table 2.5.
Table 2.5 Symbols electron shells.
n
1
2
3
4
5
6
7
symbo
l
K
L
M
N
O
P
Q
Table 2.6 summarizes the maximum number of electrons in each
skin, skin from K to N. When the atoms in the most stable state, the
ground state, electrons will occupy the orbital with lowest energy,
following the Pauli principle.
Table 2.6 Maximum number of electrons that occupy each skin.
n
ski
n
l
symb
ol
Total
max
electron
total
skin
1
K
0
1s 2 (2 =
2x12)
L 2s 2 (8 =
24
2 0 2x22)
1
2p 6
3
M
0
3s 2 (18 =
2x32)
1
3p 6
2
3d 10
4
N
0
4s 2 (32 =
2x42)
1
4p 6
2
4d 10
3
4F 14
By increasing the orbital energy difference between the orbital
energy becomes smaller, and sometimes the sequence into the
outer shell inverted electrons . Configuration clearly change when
the atomic number changes. This is the basic theory of periodic
law.
It should be added here, using the symbols given in Table 2.6, the
electron configuration of atoms can published. For example,
hydrogen atoms in the ground state has one electron skin Diu K and
electron configuration (1s 1). The carbon atom has 2 electrons in the
25
skin of K and 4 electrons in the skin L. electron configuration is (1s 2
2s 2 2p 2).
CHAPTER 3
PERIODIC TABLE
Periodic Table
One of the greatest intellectual achievements in chemistry is
the periodic table of elements. The periodic table can be printed on
one sheet of paper, but what is contained in it and what can be
given to us very much and are invaluable. This table is the result of
tireless efforts, which originated from Greek times, to know the true
nature of matter. It can be said Shem chemical scripture. Value of
the periodic system not only on the information known to the
organization, but also its ability to predict the nature of the
unknown. Actual efficacy of the periodic table is located here.
3.1 The proposals before Mendeleev
26
The concept of elements is a very old concept, since the days
of Greece, the Greek philosopher said, the material is formed of four
elements: earth, water, fire and air. This view was gradually
abandoned, and finally in the 17th century definition of elements
given by the English chemist Robert Boyle (16,271,691) replaces
the old definition earlier. Boyle stated that the element is a
substance that can not be broken down into simpler substances.
Lavoisier proposed a list of elements in his book "Traite de
Chemie Elementire." Although he entered the light and heat in the
list, other members of the list is what we call the elements to date.
Besides, he added to the list of elements that have not been
detected, but he believed existed. For example, chlorine at that
time has not been isolated, but he added that to the table as a
radical from the acid muriatik. Similarly, sodium and potassium is
also in the table.
In the early 19th century, these elements were isolated by
electrolysis, and slowly expanded the list of elements. In the mid-
19th century, spectroscopic analysis, from method introduced
detecting element and accelerate the accretion of this list.
Although welcomed by chemists, emerging problems. One of the
questions was' Is a limited number of elements? " and another
question is' What is the nature of the elements expected to have a
certain order? "
The discovery of new elements catalysis discussions of this
kind. When iodine was found in 1826, German chemist Johann
Wolfgang Döbereiner (1780-1849) noted the similarity between this
27
element with elements that have been known to chlorine and
bromine. He also detects a trio of other similar elements. This is
what is known as the theory triade Döbereiner.
Table 3.1 Triade Döbereiner
lithium
(Li)
calcium
(Ca)
Chlorine
(Cl)
sulfur (S) manganese
(Mn)
Sodium
(Na)
strontium
(Sr)
Bromine
(Br)
selenium
(Se)
chromium
(Cr)
potassium
(K)
barium
(Ba)
iodine (I) tellurium
(Te)
Iron (Fe)
Octave Newland
In 1865, classifying the elements based on atomic mass
increases, but from the properties of these elements he observed a
repetition or periodic nature of the element. The nature of the
elements to 8 similar properties to-1 elements, so on the nature of
Element 9 has similar properties to the elements of the 2nd.
Because of the repetition of such nature, then it is called the Law of
Octaves. The following table law of octaves
Do Re Mi Fa Sol La Si
28
1 2 3 4 5 6 7H LI Be B C N OF NA Mg Al Si P SCl K Ca Cr Ti Mn FeCo,Ni Cu Zn Y In As SeBr Rb Sr Ce,La Zr Di, Mo Ro, RuPd Ag Cd U Sn Sb ITe Cs Ba Ta W Nb AuPt, Ir Os V Tl Pb Bi Th
Meyer
In 1864, Lothar Meyer conducted experiments that looked at
the relationship between the increase in atomic mass with element
properties. This is done by making the curve of atomic volume
mass for the same element. This dilemma is slowly resolved after
the International Chemical Congress (Congress was held in 1860 in
Karlsruhe, Germany. The purpose of this congress to discuss the
problem of unification of atomic mass. Cannizzaro take this
opportunity to introduce the theory of Avogadro.) First, which was
attended by Mendeleev, but difficulties still exist.
By basing on the valence in determining the atomic mass,
Mendeleev bit much to solve the problem.
As has been described Sisem Modern Periodic periodic system is perfected that had been developed by Mendeleev, Similarity properties of the elements with electron configuration of elements, but it also turns out the elements in one group have the same valence electron. Here is a picture of modern periodic system
The alkali metals are a series of chemical elements forming Group 1 (IUPAC style) of the periodic table: lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs), and francium (Fr). (Hydrogen, although nominally also a member of Group 1, very rarely exhibits behavior comparable to the alkali metals). The alkali metals provide one of the best examples of group trends in properties in the periodic table, with well characterized homologous behavior down the group.
Properties
The alkali metals are all highly reactive and are never found in elemental forms in nature. Because of this, they are usually stored immersed in mineral oil or kerosene (paraffin oil). They also tarnish easily and have low melting points and densities.
Physically, the alkali metals are mostly silver-colored, except for metallic caesium, which has a golden tint. These elements are all soft metals of low density. Chemically, all of the alkali metals react aggressively with the halogens to form ionic salts. They all react with water to form strongly alkaline hydroxides. The vigor of reaction increases down the group. All of the atoms of alkali metals have one electron in their outmost electron shells, hence their only way for achieving the equivalent of filled outmost electron shells is to give up one electron to an element with high electronegativity, and hence to become singly charged positive ions, i.e. cations.
When it comes to their nuclear physics, the elements potassium and rubidium are naturally weakly radioactive because they each contain a long half-life radioactive isotope.
The element hydrogen, with its solitary one electron per atom, is usually placed at the top of Group 1 of the periodic table for convenience, but hydrogen is not counted as an alkali metal. Under typical conditions, pure hydrogen exists as a diatomic gas consisting of two atoms per molecule.
The removal of the single electron of hydrogen requires considerably more energy than removal of the outer electron from the atoms of the alkali metals. As in the halogens, only one additional electron is required to fill in
the outermost shell of the hydrogen atom, so hydrogen can in some circumstances behave like a halogen, forming the negative hydride ion. Binary compounds of hydrogen with the alkali metals and some transition metals have been produced in the laboratory, but these are only laboratory curiosities without much practical use. Under extremely high pressures and low temperatures, such as those found at the cores of the planets Jupiter and Saturn, hydrogen does become a metallic element, and it behaves like an alkali metal. (See metallic hydrogen.)
The alkali metals have the lowest ionization potentials in their periods of the periodic table, because the removal of their single electrons from their outmost electron shells gives them the stable electron configuration of inert gases. Another way of stating this is that they all have a high electropositivity. The "second ionization potential" of all of the alkali metals is very high, since removing any electron from an atom having a noble gas configuration is difficult to do.
Series of alkali metals, stored in mineral oil ("natrium" is sodium.)
All of the alkali metals are notable for their vigorous reactions with water, and these reactions become increasingly vigorous when going down their column in the periodic table towards the heaviest alkali metals, such as caesium. Their chemical reactions with water are as follows:
Alkali metal + water → Alkali metal hydroxide + hydrogen gas
For a typical example (M represents an alkali metal):
2 M (s) + 2 H2O (l) → 2 MOH (aq) + H2 (g)
Trends
Like in other columns of the periodic table, the members of the alkali metal family show patterns in their electron configurations, especially their outmost electron shells. This causes similar patterns in their chemical properties:
The alkali metals show a number of trends when moving down the group - for instance: decreasing electronegativity, increasing reactivity, and decreasing melting and boiling point. Their densities generally increase, with the notable exception that potassium is less dense than sodium, and the possible exception of francium being less dense than caesium. (The highly radioactive element francium only exists in microscopic quantities.)
IUPAC has not recommended a specific format for the periodic table, so different conventions are permitted and are often used for group IB. The following d-block transition metals are always considered members of group IB:
When defining the remainder of Group IB, four different conventions may be encountered:
Some tables [1] include lanthanum (La) and actinium (Ac), (the beginnings of the lanthanide and actinide series of elements, respectively) as the remaining members of Group IB. In their most commonly encountered tripositive ion forms, these elements do not possess any partially filled f orbitals, thus resulting in more d-block-like behavior.
Some tables [2] include lutetium (Lu) and lawrencium (Lr) as the remaining members of Group IB. These elements terminate the lanthanide and actinide series, respectively. Since the f-shell is nominally full in the ground state electron configuration for both of these metals, they behave most like d-block metals out of all the lanthanides and actinides, and thus exhibit the most similarities in properties with Sc and Y. For Lr, this behavior is expected, but it has not been observed because sufficient quantities are not available. (See also Periodic table (wide) and Periodic table (extended).)
Some tables [3] refer to all lanthanides and actinides by a marker in Group IB. A third and fourth alternative are suggested by this arrangement:
The third alternative is to regard all 30 lanthanide and actinide elements as included in Group IB. Lanthanides, as electropositive trivalent metals, all have a closely related chemistry, and all show many similarities to Sc and Y.
The fourth alternative is to include none of the lanthanides and actinides in Group IB. The lanthanides possess additional properties characteristic of their partially-filled f orbitals which are not common to Sc and Y. Furthermore, the actinides show a much wider variety of chemistry (for instance, in range of oxidation states) within their series than the lanthanides, and comparisons to Sc and Y are even less useful.
Hydrogen is the lightest element. It is by far the most abundant element in the universe and makes up about about 90% of the universe by weight. Hydrogen as water (H2O) is absolutely essential to life and it is present in all organic compounds. Hydrogen is the lightest gas. Hydrogen gas was used in lighter-than-air balloons for transport but is far too dangerous because of the fire risk (Hindenburg). It burns in air to form only water as waste product and if hydrogen could be made on sufficient scale from other than fossil fuels then there might be a possibility of a hydrogen economy.
Note that while normally shown at the top of the Group 1 elements in the periodic table, the term "alkaline metal" refers only to Group 1 elements from lithium onwards.
Table: basic information about and classifications of hydrogen. Name : Hydrogen
The lifting agent for the ill fated Hindenberg ballooon was hydrogen rather than the safer helium. The image below is the scene probably in a way you have not seen it before. This is a "ray-traced" image reproduced with the permission of Johannes Ewers, the artist, who won first place with this image in the March/April 1999 Internet Raytracing Competition. For details of ray-tracing you can't beat the POV-Ray site.
Isolation
Isolation: in the laboratory, small amounts of hydrogen gas may be made by the reaction of calcium hydride with water.
CaH2 + 2H2O → Ca(OH)2 + 2H2
This is quite efficient in the sense that 50% of the hydrogen produced comes from water. Another very convenient laboratory scale experiment follows Boyle's early synthesis, the reaction of iron filings with dilute sulphuric acid.
Fe + H2SO4 → FeSO4 + H2
There are many industrial methods for the production of hydrogen and that used will depend upon local factors such as the quantity required and the raw materials to hand. Two processes in use involve heating coke with steam in the water gas shift reaction or hydrocarbons such as methane with steam.
CH4 + H2O (1100°C) → CO + 3H2
C(coke) + H2O (1000°C) → CO + H2
In both these cases, further hydrogen may be made by passing the CO and steam over hot (400°C) iron oxide or cobalt oxide.
All values of electron binding energies are given in eV. The binding energies are quoted relative to the vacuum level for rare gases and H2, N2, O2, F2, and Cl2 molecules; relative to the Fermi level for metals; and relative to the top of the valence band for semiconductors.
Lithium
Lithium is a Group 1 (IA) element containing just a single valence electron (1s22s1). Group 1 elements are called "alkali metals". Lithium is a solid only about half as dense as water and lithium metal is the least dense metal. A freshly cut chunk of lithium is silvery, but tarnishes in a minute or so in air to give a grey surface. Its chemistry is dominated by its tendency to lose an electron to form Li+. It is the first element within the second period.
Lithium is mixed (alloyed) with aluminium and magnesium for light-weight alloys, and is also used in batteries, some greases, some glasses, and in medicine.
Table: basic information about and classifications of lithium. Name : Lithium
Isolation: lithium would not normally be made in the laboratory as it is so readily available commercially. All syntheses require an electrolytic step as it is so difficult to add an electron to the poorly electronegative lithium ion Li+.
The ore spodumene, LiAl(SiO3)2, is the most important commercial ore containing lithium. The α form is first converted into the softer β form by heating to around 1100°C. This is mixed carefully with hot sulphuric acid and extracted into water to form lithium sulphate, Li2SO4, solution. The sulphate is washed with sodium carbonate, Na2CO3, to form a precipitate of the relatively insoluble lithium carbonate, Li2CO3.
Li2SO4 + Na2CO3 → Na2SO4 + Li2CO3 (solid)
Reaction of lithium carbonate with HCl then provides lithium chloride, LiCl.
Li2CO3 + 2HCl → 2LiCl + CO2 +H2O
Lithium chloride has a high melting point (> 600°C) meaning that it sould be expensive to melt it in order to carry out the electrolysis. However a mixture of LiCl (55%) and KCl (45%) melts at about 430°C and so much less energy and so expense is required for the electrolysis.
Sodium is a Group 1 element (or IA in older labelling styles). Group 1 elements are often referred to as the "alkali metals". The chemistry of sodium is dominated by the +1 ion Na+. Sodium salts impart a characteristic orange/yellow colour to flames and orange street lighting is orange because of the presence of sodium in the lamp.
Soap is generally a sodium salt of fatty acids. The importance of common salt to animal nutrition has been recognized since prehistoric times. The most common compound is sodium chloride, (table salt).
Table: basic information about and classifications of sodium. Name : Sodium
Symbol : Na
Atomic number : 11
Atomic weight : 22.98976928 (2)
Standard state : solid at 298 K
CAS Registry ID : 7440-23-5
Group in periodic table : 1
Group name : Alkali metal
Period in periodic table : 3
Block in periodic table : s-block
Colour : silvery white
Classification : Metallic
The result of adding different metal salts to a burning reaction mixture of potassium chlorate and sucrose. The red colour originates from strontium sulphate. The orange/yellow colour originates from sodium chloride. The green colour originates from barium chlorate and the blue colour originates from copper (I) chloride. The lilac colour that should be evident from the potassium chlorate is washed out by the other colours, all of which are more intense (only to be demonstrated by a professionally qualified chemist following a legally satisfactory hazard asessment). Improperly done, this reaction is dangerous!
The picture above shows the colour arising from adding common salt (NaCl) to a burning mixture of potassium chlorate and sucrose.
Isolation
Isolation: sodium would not normally be made in the laboratory as it is so readily available commercially. All syntheses require an electrolytic step as it is so difficult to add an electron to the poorly electronegative sodium ion Na+.
Sodium is present as salt (sodium chloride, NaCl) in huge quantities in underground deposits (salt mines) and seawater and other natural waters. It is easily recovered as a solid by drying.
Sodium chloride has a high melting point (> 800°C) meaning that it sould be expensive to melt it in order to carry out the electrolysis. However a mixture of NaCl (40%) and calcium chloride, CaCl2 (60%) melts at about 580°C and so much less energy and so expense is required for the electrolysis.
cathode: Na+(l) + e- → Na (l)
anode: Cl-(l) → 1/2Cl2 (g) + e-
The electrolysis is carried out as a melt in a "Downs cell". In practice, the electrolysis process produces calcium metal as well but this is solidified in a collection pipe and returned back to the melt.
Table: valence shell orbital radii for sodium.Orbital Radius [/pm] Radius [/AU]
Potassium is a metal and is the seventh most abundant and makes up about 1.5 % by weight of the earth's crust. Potassium is an essential constituent for plant growth and it is found in most soils. It is also a vital element in the human diet.
Potassium is never found free in nature, but is obtained by electrolysis of the chloride or hydroxide, much in the same manner as prepared by Davy. It is one of the most reactive and electropositive of metals and, apart from lithium, it is the least dense known metal. It is soft and easily cut with a knife. It is silvery in appearance immediately after a fresh surface is exposed.
It oxidises very rapidly in air and must be stored under argon or under a suitable mineral oil. As do all the other metals of the alkali group, it decomposes in water with the evolution of hydrogen. It usually catches fire during the reaction with water. Potassium and its salts impart a lilac colour to flames.
Table: basic information about and classifications of potassium. Name : Potassium
Symbol : K
Atomic number : 19
Atomic weight : 39.0983 (1)
Standard state : solid at 298 K
CAS Registry ID : 7440-09-7
Group in periodic table : 1
Group name : Alkali metal
Period in periodic table : 4
Block in periodic table : s-block
Colour : silvery white
Classification : Metallic
The reaction between potassium metal and water (only to be demonstrated by a professionally qualified chemist).
The picture above shows the colour arising from a burning mixture of potassium chlorate (KClO3) and sucrose (only to be demonstrated by a professionally qualified chemist).Isolation
Isolation: potassium would not normally be made in the laboratory as it is so readily available commercially. All syntheses require an electrolytic step as it is so difficult to add an electron to the poorly electronegative potassium ion K+.
Potassium is not made by the same method as sodium as might have been expected. This is because the potassium metal, once formed by electrolysis of liquid potassium chloride (KCl), is too soluble in the molten salt.
cathode: K+(l) + e- → K (l)
anode: Cl-(l) → 1/2Cl2 (g) + e-
Instead, it is made by the reaction of metallic sodium with molten potassium chloride at 850°C.
Na + KCl ⇌ K + NaCl
This is an equilibrium reaction and under these conditions the potassium is highly volatile and removed from the system in a form relatively free from sodium impurities, allowing the reaction to proceed.
Table: valence shell orbital radii for potassium.Orbital Radius [/pm] Radius [/AU]
Rubidium can be liquid at ambient temperature, but only on a hot day given that its melting point is about 40°C. It is a soft, silvery-white metallic element of the alkali metals group (Group 1). It is one of the most most electropositive and alkaline elements. It ignites spontaneously in air and reacts violently with water, setting fire to the liberated hydrogen. As so with all the other alkali metals, it forms amalgams with mercury. It alloys with gold, caesium, sodium, and potassium. It colours a flame yellowish-violet.
Table: basic information about and classifications of rubidium. Name : Rubidium
Symbol : Rb
Atomic number : 37
Atomic weight : 85.4678 (3) [see note g]
Standard state : solid at 298 K
CAS Registry ID : 7440-17-7
Group in periodic table : 1
Group name : Alkali metal
Period in periodic table : 5
Block in periodic table : s-block
Colour : silvery white
Classification : Metallic
Image adapted with permission from Prof James Marshall's (U. North Texas, USA) Walking Tour of the elements CD.Isolation
Isolation: rubidium would not normally be made in the laboratory as it is available commercially. All syntheses require an electrolytic step as it is so difficult to add an electron to the poorly electronegative rubidium ion Rb+.
Rubidium is not made by the same method as sodium as might have been expected. This is because the rubidium metal, once formed by electrolysis of liquid rubidium chloride (RbCl), is too soluble in the molten salt.
Instead, it is made by the reaction of metallic sodium with hot molten rubidium chloride.
Na + RbCl ⇌ Rb + NaCl
This is an equilibrium reaction and under these conditions the rubidium is highly volatile and removed from the system in a form relatively free from sodium impurities, allowing the reaction to proceed.
The following are calculated values of valence shell orbital radii, Rmax
Table: valence shell orbital radii for rubidium.
Orbital Radius [/pm] Radius [/AU]
s orbital 249.0 4.70616
p orbital no data no data
d orbital no data no data
f orbital no data no data
Effective Nuclear Charges
The following are "Clementi-Raimondi" effective nuclear charges, Zeff. Follow the hyperlinks for more details and for graphs in various formats.
Caesium is known as cesium in the USA. The metal is characterised by a spectrum containing two bright lines in the blue (accounting for its name). It is silvery gold, soft, and ductile. It is the most electropositive and most alkaline element. Caesium, gallium, and mercury are the only three metals that are liquid at or around room temperature. Caesium reacts explosively with cold water, and reacts with ice at temperatures above -116°C. Caesium hydroxide is a strong base and attacks glass.
Table: basic information about and classifications of caesium. Name : Caesium
Symbol : Cs
Atomic number : 55
Atomic weight : 132.9054519 (2)
Standard state : solid at 298 K (but melts only slightly above this temperature)
CAS Registry ID : 7440-46-2
Group in periodic table : 1
Group name : Alkali metal
Period in periodic table: 6
Block in periodic table : s-block
Colour : silvery gold
Classification : Metallic
Image adapted with permission from Prof James Marshall's (U. North Texas, USA) Walking Tour of the elements CD.Isolation
Isolation: caesium (cesium in USA) would not normally be made in the laboratory as it is available commercially. All syntheses require an electrolytic step as it is so difficult to add an electron to the poorly electronegative caesium ion Cs+.
Caesium is not made by the same method as sodium as might have been expected. This is because the caesium metal, once formed by electrolysis of liquid caesium chloride (CsCl), is too soluble in the molten salt.
Instead, it is made by the reaction of metallic sodium with hot molten caesium chloride.
Na + CsCl ⇌ Cs + NaCl
This is an equilbrium reaction and under these conditions the caesium is highly volatile and removed from the system in a form relatively free from sodium impurities, allowing the reaction to proceed. It can be purified by distillation.
The following are calculated values of valence shell orbital radii, Rmax
Table: valence shell orbital radii for caesium.
Orbital Radius [/pm] Radius [/AU]
s orbital 282.4 5.33704
p orbital no data no data
d orbital no data no data
f orbital no data no data
Effective Nuclear Charges
The following are "Clementi-Raimondi" effective nuclear charges, Zeff. Follow the hyperlinks for more details and for graphs in various formats.
Francium occurs as a result of α disintegration of actinium. Francium is found in uranium minerals, and can be made artificially by bombarding thorium with protons. It is the most unstable of the first 101 elements. The longest lived isotope, 223Fr, a daughter of 227Ac, has a half-life of 22 minutes. This is the only isotope of francium occurring in nature, but at most there is only 20-30 g of the element present in the earth's crust at any one time. No weighable quantity of the element has been prepared or isolated. There are about 20 known isotopes.
Table: basic information about and classifications of francium. Name : Francium
This sample of uraninite contains some francium because of a steady-state decay chain. An estimate suggests there is about 10-20
grammes of francium (about 1 atom!) at any one time. Image adapted with permission from Prof James Marshall's (U. North Texas, USA) Walking Tour of the elements CD.Isolation
Isolation: francium is vanishingly rare and is found only as very small traces in some uranium minerals. It has never been isolated as the pure element. As it is so radioactive, any amount formed would decompose to other elements.
Actinium decays by β decay most of the time but about 1% of the decay is by α decay. The "daughter" element of this reaction, which used to be called actinium-K, is now recognized as 223
87Fr - the longest-lived isotope of actinium with a half life of about 22 minutes.
The following are calculated values of valence shell orbital radii, Rmax
Scandium is a silvery-white metal which develops a slightly yellowish or pinkish cast upon exposure to air. It is relatively soft, and resembles yttrium and the rare-earth metals more than it resembles aluminium or titanium. Scandium reacts rapidly with many acids.
Scandium is apparently a much more abundant element in the sun and certain stars than on earth.
Table: basic information about and classifications of scandium.
Image adapted with permission from Prof James Marshall's (U. North Texas, USA) Walking Tour of the elements CD.Isolation
Isolation: preparation of metallic samples of scandium is not normally necessary given that it is commercially avaialable. In practice littel scandium is produced. The mineral thortveitite contains 35-40% Sc2O3 is used to produce scandium metal but another important source is as a byproduct from uranium ore processing, even though these only contain 0.02% Sc2O3.
Brief description: yttrium has a silvery-metallic lustre. Yttrium turnings ignite in air. Yttrium is found in most rare-earth minerals. Moon rocks contain yttrium and yttrium is used as a "phosphor" to produce the red colour in television screens.
Scandium: orbital properties
Valence shell orbital radii
The following are calculated values of valence shell orbital radii, Rmax
This sample is from The Elements Collection, an attractive and safely packaged collection of the 92 naturally occurring elements that is available for sale.Isolation
Isolation: yttrium metal is available commercially so it is not normally necesary to make it in the laboratory. Yttrium is found in lathanoid minerals and the extraction of the yttrium and the lanthanoid metals from the ores is highly complex. Initially, the metals are extractedas salts from the ores by extraction with sulphuric acid (H2SO4), hydrochloric acid (HCl), and sodium hydroxide (NaOH). Modern purification techniques for these lanthanoid salt mixtures involve selective complexation techniques, solvent extractions, and ion exchange chromatography.
Pure yttrium is available through the reduction of YF3 with calcium metal.
2YF3 + 3Ca → 2Y + 3CaF2
Yttrium: orbital properties
Valence shell orbital radii
The following are calculated values of valence shell orbital radii, Rmax
Table: valence shell orbital radii for yttrium.Orbital Radius [/pm] Radius [/AU]
Brief description: pure metal lutetium has been isolated only in recent years and is one of the more difficult to prepare. It can be prepared by the reduction of anhydrous LuCl3 or LuF3 by an alkali or alkaline earth metal.
The metal is silvery white and relatively stable in air. It is a rare earth metal and perhaps the most expensive of all rare elements. It is found in small amounts with all rare earth metals, and is very difficult to separate from other rare elements.
Table: basic information about and classifications of lutetium. Name : Lutetium
This sample is from The Elements Collection, an attractive and safely packaged collection of the 92 naturally occurring elements that is available for sale.Isolation
Isolation: lutetium metal is available commercially so it is not normally necessary to make it in the laboratory, which is just as well as it is difficult to isolate as the pure metal. This is largely because of the way it is found in nature. The lanthanoids are found in nature in a number of minerals. The most important are xenotime, monazite, and bastnaesite. The first two are orthophosphate minerals LnPO4 (Ln deonotes a mixture of all the lanthanoids except promethium which is vanishingly rare) and the third is a fluoride carbonate LnCO3F. Lanthanoids with even atomic numbers are more common. The most comon lanthanoids in these minerals are, in order, cerium, lanthanum, neodymium, and praseodymium. Monazite also contains thorium and ytrrium which makes handling difficult since thorium and its decomposition products are radioactive.
For many purposes it is not particularly necessary to separate the metals, but if separation into individual metals is required, the process is complex. Initially, the metals are extracted as salts from the ores by extraction with sulphuric acid (H2SO4), hydrochloric acid (HCl), and sodium hydroxide (NaOH). Modern purification techniques for these lanthanoid salt mixtures are ingenious and involve selective complexation techniques, solvent extractions, and ion exchange chromatography.
Pure lutetium is available through the reduction of LuF3 with calcium metal.
2LuF3 + 3Ca → 2Lu + 3CaF2
This would work for the other calcium halides as well but the product CaF2 is easier to handle under the reaction conditions (heat to 50°C above the melting point of the element in an argon atmosphere). Excess calcium is removed from the reaction mixture under vacuum.
1. D.R. Lide, (Ed.) in Chemical Rubber Company handbook of chemistry and physics, CRC Press, Boca Raton, Florida, USA, 81st edition, 2000.
2. E. Clementi and D.L.Raimondi, J. Chem. Phys. 1963, 38, 2686.
3. E. Clementi, D.L.Raimondi, and W.P. Reinhardt, J. Chem. Phys. 1967, 47, 1300.
4. J. A. Bearden and A. F. Burr, "Reevaluation of X-Ray Atomic Energy Levels," Rev. Mod. Phys., 1967, 39, 125.
5. J.B. Mann, Atomic Structure Calculations II. Hartree-Fock wave functions and radial expectation values: hydrogen to lawrencium, LA-3691, Los Alamos Scientific Laboratory, USA, 1968.
6. J. C. Fuggle and N. Mårtensson, "Core-Level Binding Energies in Metals," J. Electron Spectrosc. Relat. Phenom., 1980, 21, 275.
7. Gwyn Williams WWW table of values
8. M. Cardona and L. Ley, Eds., Photoemission in Solids I: General Principles (Springer-Verlag, Berlin) with additional corrections, 1978.
Some of the physical properties usually associated with metals—hardness,
high m.p. and b.p.—are noticeably lacking in these metals, but they all have
a metallic appearance and are good electrical conductors. Group II metals
are more dense, are harder and have higher m.p. and b.p. than the
corresponding Group I metals. bonding generally depends on the ratio
(number of electrons available for bonding)/(atomic size). The greater this
ratio is, the stronger are the bonds between the metal atoms. In the pre-
transition metals, this ratio is small and at a minimum in Group I with only
one bonding electron. Metallic bond strength is greater in Group II but there
are still only two bonding electrons available, hence the metals are still
relatively soft and have low melting and boiling points. Hardness, m.p. and
b.p. all decrease steadily down Group I, the metallic bond strength
decreasing with increasing atomic radius. These changes are not so well
marked in Group II but note that beryllium and, to a lesser extent,
magnesium are hard metals, as a result of their small atomic size; this
property, when coupled with their low density, makes them of some
technological importance.
Group II (alkaline earths) are known as s-block elements because their
valence (bonding) electrons are in s orbitals.
PHYSICAL PROPERTIES OF IIA ELEMENTS
Element Atomic
number
Outer
electrons
Density
(g/ )
m.p.
(K)
b.p.
(K)
Hardness
(Brinell)
Be
Mg
Ca
Sr
Ba
4
12
20
38
56
2
3
4
5
1.86
1.75
1.55
2.6
3.59
1553
924
1124
1073
998
3243
1380
1760
1639
1910
-
30-40
23
20.
-
153
6
Element lonisatio
n
energy*
(kJ/mol)
Metallic
radius
(nm)
Ionic
radius
(nm)
Heat of
vaporisatio
n
at 298 K
(kJ/mol)
Hydrat ion
energy of
gaseous ion
(kJ/mol)
Be
Mg
Ca
Sr
Ba
2657
2187
1735
1613
1467
0.112
0.160
0.197
0.215
0.221
0.031
0.065
0.099
0.113
0.135
326
149
177
164
178
2494
1921
1577
1443
1305
Atomic Radius Increases down each group electrons are in shells further
from the nucleus
Ionic Size Increases down the group The size of positive ions is less than
the original atom because the nuclear charge exceeds the electronic charge.
Melting Points Decrease down each group metallic bonding gets weaker
due to increased size
Each atom contributes two electrons to the delocalised cloud. Melting points
tend
not to give a decent trend as different crystalline structures affect the melting
point.
CHEMICAL PROPERTIES OF THE ELEMENTSOverall Reactivity increases down the Group due to the ease of cation formation
Oxygen • react with increasing vigour down the groupWater • react with increasing vigour down the group
154
OXIDES OF GROUP II ELEMENTS
Properties • ionic solids; EXC. beryllium oxide which has covalent character
Hydroxides• basic strength also increases down group• this is because the solubility increases• the metal ions get larger so charge density decreases• there is a lower attraction between the OH¯ ions and larger dipositive ions• the ions will split away from each other more easily• there will be a greater concentration of OH¯ ions in water
Uses of hydroxides Ca(OH) • used in agriculture to neutralise acid soilsCARBONATES
Properties • insoluble in water• undergo thermal decomposition to oxide and carbon dioxide
We note first that the elements are all electropositive, having relatively
low ionisation energies, and are, in consequence, very reactive. The enthalpy
change required for the process M(metal) -»M + (g) for Group I, or M(metal) -
> M2+(g) for Group II is at a maximum at the top of each group, and it is,
therefore, not surprising
to find that lithium, beryllium and, to some extent, magnesium do form some
covalent compounds. Most solid compounds of Group 1 and II elements,
however, have ionic structures and the properties associated with such
structures—high m.p. and b.p., solubility in water rather than in organic
solvents and electrical conductance when molten.
IONS IN SOLUTION
The hydration energies (strictly, hydration enthalpies) fall, as expected, as
we descend either Group, and are larger for Group II than for Group I ions.
The solubilities of the salts of Groups I and II are determined by a balance
155
between lattice energy, hydration energy and the entropy change in going
from solid to solution, and
only a few generalisations are possible. Thus high charge and low ionic radii
tend to produce insolubility (for example salts of lithium, beryllium and
magnesium, especially those with doubly charged anions such as carbonate
COa~). At the other end of the scale, low charge and large radii also produce
low solubility (for example salts of potassium, rubidium and caesium
containing large anions such as the tetraphenylborate anion (p. 136). In
between, solubility is the rule for all Group I salts, and for most Group II salts
containing singly-charged negative ions; for many Group II salts with
doublyor triply-charged anions (for example COj", SOj", PO^ ) insolubility is
often observed. The decreasing tendency to form salts with water of
crystallization (as a group is descended) is again in line with the falling
hydration
energy. For example, both sodium sulphate and carbonate form hydrates but
neither of the corresponding potassium salts do; the sulphates of Group II
elements show a similar trend MgSO4 , 7H2O, CaSO4 . 2H2O, BaSO4. For the
most part, however, the chemistry of the Group I and II elements is that of
the metal and the ions M +
for Group I and M2* for Group II. As already noted the two head elements,
lithium and beryllium, tend to form covalent compounds; the beryllium ion
Be2 + , because of its very small radius and double charge, has also some
peculiar properties in solution, which are examined later
OCCURRENCE AND EXTRACTION
Of the Group II metals (beryllium to barium) beryllium, the rarest, occurs as
the aluminatesilicate, beryl \magnesium is found as the carbonate and (with
calcium) as the double carbonate dolomite', calcium, strontium and barium
all occur as carbonates, calcium carbonate being very plentiful as limestone.
The general characteristics of all these elements generally preclude their
extraction by any method involving aqueous solution. For the lighter, less
volatile metals (Li, Na, Be, Mg, Ca) electrolysis of a fused salt (usually the
156
chloride), or of a mixture of salts, is used. The heavier, more volatile metals
in each group can all be similarly obtained by electrolysis, but it is usually
more convenient to take advantage of their volatility and obtain them from
their oxides or chlorides by displacement, i.e. by general reactions such as
3M2O + 2Mm -* M2
mO3 4- 6M|
MCI + M1 ~» M!C1 + M|
Thus potassium is obtained by heating potassium chloride with sodium, and
barium by reduction of barium oxide with aluminium. Sodium is important in
many technical processes and is therefore prepared in considerable quantity.
Almost all of it is now made by electrolysis of the fused sodium chloride,
using the Downs cell. The graphite anode is cylindrical and is surrounded by
the steel gauze diaphragm and the concentric cylindrical cathode (also of
steel). The electrolyte is usually a mixture of sodium chloride and calcium
chloride; the latter is added to reduce the m.p. of the
sodium chloride to approximately 800 K. (Some calcium is therefore
liberated with the sodium.) The gap between anode and cathode is kept as
small as possible to reduce resistance: the heat developed by the current
maintains the temperature of the cell. Chlorine is set free at the anode
surface, rises into the nickel cone and can be collected. Sodium, liberated at
the cathode, is prevented by the diaphragm from passing into the anode
region; the molten sodium collects under the circular hood and rises up the
pipe, being assisted if necessary by the puddle-rod. The calcium, being
almost immiscible with sodium and much more dense, can readily be
separated from
the molten sodium. The graphite anode wears away and must be renewed
from time to time.
Element Be Mg Ca Sr Ba
Reaction Does not react with
Very slowl with water, readily
React with cold water. vigour of reaction increasing.
157
Conditions
Basic
properties
of product
with water,
Be(OH)2
amphoteric
with
steam
MgO
insoluble
Slightly soluble M soluble
Base strength increasing
USES
Beryllium is added to copper to produce an alloy with greatly increased wear
resistance; it is used for current-carrying springs and non-sparking safety
tools. It is also used as a neutron moderator and reflector in nuclear reactors.
Much magnesium is used to prepare light metal alloys; other uses include
the extraction of titanium and in the removal of oxygen and sulphur from
steels; calcium finds a similar use.
BIOLOGICAL IMPORTANCE
Sodium and potassium ions are found in all animal cells and, usually, the
concentration of potassium ions inside the cell is greater than that of sodium.
In many cells, this concentration difference is maintained by a 'sodium
pump', a process for which the energy is supplied by the hydrolysis of
adenosine triphosphate (ATP). Diffusion of excess potassium ions outwards
through the cell wall gives the inside of the cell a net negative charge (due
to the anions present) and a potential difference is established across the
cell wall. In a nerve cell, a momentary change in the permeability of the cell
wall to sodium ions can reverse the sign of this potential difference, and this
produces the electrical impulse associated with the action of the nerve.
158
The ability of living organisms to differentiate between the chemically similar
sodium and potassium ions must depend upon some difference between
these two ions in aqueous solution. Essentially, this difference is one of size
of the hydrated ions, which in turn means a difference in the force of
electrostatic (coulombic) attraction between the hydrated cation and a
negatively-charged site in the cell membrane; thus a site may be able to
accept the smaller ion Na+(aq) and reject the larger K+(aq). This same
mechanism of selectivity operates in other 'ion-selection' processes, notably
in ionexchange resins. All organisms seem to have an absolute need for
magnesium. In plants, the magnesium complex chlorophyll is the prime
agent in photosynthesis. In animals, magnesium functions as an enzyme
activator; the enzyme which catalyses the ATP hydrolysis mentioned above
is an important example. Calcium plays an important part in structure-
building in living organisms, perhaps mainly because of its ability to link
together phosphate-containing materials. Calcium ions in the cell play a vital
part in muscle contraction.
THE HYDRIDES
All Group I and II elements, except beryllium, form hydrides by direct
combination with hydrogen. The hydrides of the metals except those of
beryllium and magnesium, are white mainly ionic solids, all Group I hydrides
having the sodium chloride lattice
structure. All the hydrides are stable in dry air but react with water, the
vigour of the reaction increasing with the molecular weight of the hydride for
any particular group.
Group II metals also form halides by direct combination. The trends in heat of
formation and m.p., however, whilst following the general pattern of the
corresponding Group I compounds, are not so regular.
* Lithium bromide and iodide probably have some degree of covalency but
this
does not affect the general conclusion.
159
As a consequence of the high ionisation energy of beryllium its halides are
essentially covalent, with comparatively low m.p., the melts being non-
conducting and (except beryllium fluoride) dissolving in many organic
solvents. The lower members in Group II form essentially ionic halides, with
magnesium having intermediate properties, and both magnesium bromide
and iodide dissolve in organic solvents. The lattice energies of the Group II
fluorides are generally greater than those for the corresponding Group I
fluorides; consequently all but beryllium fluoride are insoluble. (The solubility
of
beryllium fluoride is explained by the high hydration energy of the beryllium
ion, cf. LiF.) The high hydration energy of the Be2+ ion* results in hydrolysis
in neutral or alkaline aqueous solution; in this reaction the beryllium halides
closely resemble the aluminium halides.
The magnesium ion having a high hydration energy also shows hydrolysis
but to a lesser extent, The chloride forms several hydrates which decompose
on heating to give a basic salt. Other Group II halides are essentially ionic
and therefore have relatively high m.p., the melts acting as conductors, and
they are soluble in water but not in organic solvents.
a. Beryllium
Beryllium was discovered by Louis-Nicholas Vauquelin in 1798. Vauquelin
found beryllia (BeO) in emeralds and in the mineral beryl (beryllium
aluminum cyclosilicate). Beryllium was first isolated by Friederich Wöhler in
1828. Wöhler reacted potassium with beryllium chloride in a platinum
crucible yielding potassium chloride and beryllium.
Unlike most metals, beryllium is virtually transparent to x-rays and hence it
is used in radiation windows for x-ray tubes. Beryllium alloys are used in the
aerospace industry as light-weight materials for high performance aircraft,
satellites and spacecraft. Beryllium is also used in nuclear reactors as a
reflector and absorber of neutrons, a shield and a moderator.
In periodicity, there are many elements that classification based on its properties. For example are metal and non metal. The elements are classification into groups and period. Periodicity has 144 elements with 16 groups (8 groups A and 8 groups B), Aktanida and Lantanida.
This paper will discuss about elements in groups VA and VB. Groups VA consist of Nitrogen, Phospor, Arsenic, Antimony , and Bismut. Groups VB consist of Vanadium, Niobium, Tantalum, and Dubnium.
The elements in group VA and VB has application in daily activity. Nitrogen has use for freeze. Phospor uses for toxic of rat and makes alloy. Bismut uses for makes tender iron, and catalyst for make akrilat fiber. Antanium(Sb) uses as semiconductor and batere. Arsen as additive of Ge and Si.
Vanadium used in nuclear reactors. Niobium used in surgical implants because they do not react with human tissue. Tantalum resists corrosion and is almost impervious to chemical attack.
217
CHAPTER II
MATTER
2.1. Elements of group VA
VA group consists of five elements, namely Nitrogen, Phospor, Bismut, Antimony, and Arsen. Some properties of the elements of this group we can see in the table bellow :
Nitrogen (pronounced) is a chemical element that has the symbol N, atomic number of 7 and atomic mass 14.00674 u. Elemental nitrogen is a colorless, odorless, tasteless and mostly inert diatomic gas at standard conditions, constituting 78.08% by volume of Earth's atmosphere.
Many industrially important compounds, such as ammonia, nitric acid, organic nitrates (propellants and explosives), and cyanides, contain nitrogen. The extremely strong bond in elemental nitrogen dominates nitrogen chemistry, causing difficulty for both organisms and industry in breaking the bond to convert the N2 into useful compounds, but releasing large amounts of often useful energy, when these compounds burn, explode, or decay back into nitrogen gas.
The element nitrogen was discovered by Scottish physician Daniel Rutherford in 1772. Nitrogen occurs in all living organisms. It is a constituent element of amino acids and thus of proteins, and of nucleic acids (DNA and RNA). It resides in the chemical structure of almost all neurotransmitters, and is a defining component of alkaloids, biological molecules produced by many organisms.
1. History
Nitrogen is formally considered to have been discovered by Daniel Rutherford in 1772, who called it noxious air or fixed air. That there was a fraction of air that did not support combustion was well-known to the late 18th century chemist. Nitrogen was also studied at about the same time by Carl Wilhelm Scheele, Henry Cavendish, and
Joseph Priestley, who referred to it as burnt air or phlogisticated air. Nitrogen gas was inert enough that Antoine Lavoisier referred to it as "mephitic air" or azote, from the Greek word (azotos) meaning "lifeless".Animals died in it, and it was the principal component of air in which animals had suffocated and flames had burned to extinction. Lavoisier's name for nitrogen is used in many languages (French, Russian, etc.) and still remains in English in the common names of many compounds, such as hydrazine and compounds of the azide ion. Compounds of nitrogen were known in the Middle Ages.
The alchemists knew nitric acid as aqua fortis (strong water). The mixture of nitric and hydrochloric acids was known as aqua regia (royal water), celebrated for its ability to dissolve gold (the king of metals). The earliest military, industrial and agricultural applications of nitrogen compounds involved uses of saltpeter (sodium nitrate or potassium nitrate), notably in gunpowder, and later as fertilizer. In 1910, Lord Rayleigh discovered that an electrical discharge in nitrogen gas produced "active nitrogen", an allotrope considered to be monatomic. The "whirling cloud of brilliant yellow light" produced by his apparatus reacted with quicksilver to produce explosive mercury nitride.
2. Properties
Nitrogen is a nonmetal, with an electronegativity of 3.04. It has five electrons in its outer shell and is therefore trivalent in most compounds. The triple bond in molecular nitrogen (N2) is the strongest. The resulting difficulty of converting N2 into other compounds, and the ease (and associated high energy release) of converting nitrogen compounds into elemental N2, have dominated the role of nitrogen in both nature and human economic activities.
At atmospheric pressure molecular nitrogen condenses (liquefies) at 77 K (−195.8 °C) and freezes at 63 K (−210.0 °C) into the beta hexagonal close-packed crystal allotropic form. Below 35.4 K (−237.6 °C) nitrogen assumes the alpha cubic crystal allotropic form. Liquid nitrogen, a fluid resembling water in appearance, but with 80.8% of the density (the density of liquid nitrogen at its boiling point is 0.808 g/mL), is a common cryogen.
Unstable allotropes of nitrogen consisting of more than two nitrogen atoms have been produced in the laboratory, like N3 and N4. Under extremely high pressures (1.1 million atm) and high temperatures (2000 K), as produced using a diamond anvil cell,
nitrogen polymerizes into the single-bonded cubic gauche crystal structure. This structure is similar to that of diamond, and both have extremely strong covalent bonds. N4 is nicknamed "nitrogen diamond.
a. Isotopes
There are two stable isotopes of nitrogen: 14N and 15N. By far the most common is 14N (99.634%), which is produced in the CNO cycle in stars. Of the ten isotopes produced synthetically, 13N has a half-life of ten minutes and the remaining isotopes have half-lives on the order of seconds or less. Biologically mediated reactions (e.g., assimilation, nitrification, and denitrification) strongly control nitrogen dynamics in the soil. These reactions typically result in 15N enrichment of the substrate and depletion of the product.
A small part (0.73%) of the molecular nitrogen in Earth's atmosphere is the isotopologue 14N15N, and almost all the rest is 14N2.
Radioisotope 16N is the dominant radionuclide in the coolant of pressurized water reactors during normal operation. It is produced from 16O (in water) via (n,p) reaction. It has a short half-life of about 7.1 s, but during its decay back to 16O produces high-energy gamma radiation (5 to 7 MeV). Because of this, the access to the primary coolant piping must be restricted during reactor power operation. 16N is one of the main means used to immediately detect even small leaks from the primary coolant to the secondary steam cycle.
Electromagnetic spectrum:
A 1×5 cm vial of glowing ultrapure nitrogen Nitrogen discharge (spectrum) tube.
Molecular nitrogen (14N2) is largely transparent to infrared and visible radiation because it is a homonuclear molecule and thus has no dipole moment to couple to electromagnetic radiation at these wavelengths. Significant absorption occurs at extreme ultraviolet wavelengths, beginning around 100 nanometers. This is associated with electronic transitions in the molecule to states in which charge is not distributed evenly between nitrogen atoms. Nitrogen absorption leads to significant absorption of ultraviolet radiation in the Earth's upper atmosphere and the atmospheres of other planetary bodies. For
similar reasons, pure molecular nitrogen lasers typically emit light in the ultraviolet range.
Nitrogen also makes a contribution to visible air glow from the Earth's upper atmosphere, through electron impact excitation followed by emission. This visible blue air glow (seen in the polar aurora and in the re-entry glow of returning spacecraft) typically results not from molecular nitrogen, but rather from free nitrogen atoms combining with oxygen to form nitric oxide (NO).
b. Reactions
Structure of [Ru(NH3)5(N2)]2+.
Nitrogen is generally unreactive at standard temperature and pressure. N2 reacts spontaneously with few reagents, being resilient to acids and bases as well as oxidants and most reductants. When nitrogen reacts spontaneously with a reagent, the net transformation is often called nitrogen fixation.
Nitrogen reacts with elemental lithium. Lithium burns in an atmosphere of N2 to give lithium nitride:
6 Li + N2 → 2 Li3N
Magnesium also burns in nitrogen, forming magnesium nitride.
3 Mg + N2 → Mg3N2
N2 forms a variety of adducts with transition metals. The first example of a dinitrogen complex is [Ru(NH3)5(N2)]2+ (see figure at right). Such compounds are now numerous, other examples include IrCl(N2)(PPh3)2, W(N2)2(Ph2CH2CH2PPh2)2, and [(η5-C5Me4H)2Zr]2(μ2, η²,η²-N2). These complexes illustrate how N2 might bind to the metal(s) in nitrogenase and the catalyst for the Haber process. A catalytic process to reduce N2 to ammonia with the use of a molybdenum complex in the presence of a proton source was published in 2005. (See nitrogen fixation).
The starting point for industrial production of nitrogen compounds is the Haber process, in which nitrogen is fixed by reacting
N2 and H2 over an iron(III) oxide (Fe3O4) catalyst at about 500 °C and 200 atmospheres pressure. Biological nitrogen fixation in free-living cyanobacteria and in the root nodules of plants also produces ammonia from molecular nitrogen. The reaction, which is the source of the bulk of nitrogen in the biosphere, is catalyzed by the nitrogenase enzyme complex which contains Fe and Mo atoms, using energy derived from hydrolysis of adenosine triphosphate (ATP) into adenosine diphosphate and inorganic phosphate (−20.5 kJ/mol).
3. Occurrence
Nitrogen is the largest single constituent of the Earth's atmosphere (78.082% by volume of dry air, 75.3% by weight in dry air). It is created by fusion processes in stars, and is estimated to be the 7th most abundant chemical element by mass in the universe.
Molecular nitrogen and nitrogen compounds have been detected in interstellar space by astronomers using the Far Ultraviolet Spectroscopic Explorer. Molecular nitrogen is a major constituent of the Saturnian moon Titan's thick atmosphere, and occurs in trace amounts in other planetary atmospheres.
Nitrogen is present in all living organisms, in proteins, nucleic acids and other molecules. It typically makes up around 4% of the dry weight of plant matter, and around 3% of the weight of the human body. It is a large component of animal waste (for example, guano), usually in the form of urea, uric acid, ammonium compounds and derivatives of these nitrogenous products, which are essential nutrients for all plants that cannot fix atmospheric nitrogen.
4. Applications
Liquid : Nitrogen has use for freeze.
Nitrogen occurs naturally in many minerals, such as saltpetre
(potassium nitrate), Chile saltpetre (sodium nitrate) and sal
ammoniac (ammonium chloride). Most of these are
uncommon, partly because of the minerals' ready solubility in
water. See also Nitrate minerals and Ammonium minerals.
The discovery of phosphorus is credited to the German alchemist Hennig Brand in 1669, although other chemists might have discovered phosphorus around the same time. Brand experimented with urine, which contains considerable quantities of dissolved phosphates from normal metabolism. Working in Hamburg, Brand attempted to create the fabled philosopher's stone through the distillation of some salts by evaporating urine, and in the process produced a white material that glowed in the dark and burned brilliantly. It was named phosphorus mirabilis ("miraculous bearer of light"). His process originally involved letting urine stand for days until it gave off a terrible smell. Then he boiled it down to a paste, heated this paste to a high temperature, and led the vapours through water, where he hoped they would condense to gold. Instead, he obtained a white, waxy substance that glowed in the dark. Brand had discovered phosphorus, the first element discovered since antiquity. We now know that Brand produced ammonium sodium hydrogen phosphate, (NH4)NaHPO4. While the quantities were essentially correct (it took about 1,100 L of urine to make about 60 g of phosphorus), it was unnecessary to allow the urine to rot. Later scientists would discover that fresh urine yielded the same amount of phosphorus.
Since that time, phosphors and phosphorescence were used loosely to describe substances that shine in the dark without burning. However, as mentioned above, even though the term phosphorescence was originally coined as a term by analogy with the glow from oxidation of elemental phosphorus, is now reserved for another fundamentally different process—re-emission of light after illumination.
Brand at first tried to keep the method secret,[24] but later sold the recipe for 200 thaler to D Krafft from Dresden,[4] who could now make it as well, and toured much of Europe with it, including England, where he met with Robert Boyle. The secret that it was made from urine leaked out and first Johann Kunckel (1630–1703) in Sweden (1678) and later Boyle in London (1680) also managed to make phosphorus. Boyle states that Krafft gave him no information as to the preparation of phosphorus other than that it was derived from "somewhat that belonged to the body of man". This gave Boyle a valuable clue, however, so that he, too, managed to make phosphorus, and published the method of its manufacture. Later he improved Brand's process by using sand in the reaction (still using urine as base material),
4 NaPO3 + 2 SiO2 + 10 C → 2 Na2SiO3 + 10 CO + P4
Robert Boyle was the first to use phosphorus to ignite sulfur-tipped wooden splints, forerunners of our modern matches, in 1680.
In 1769 Johan Gottlieb Gahn and Carl Wilhelm Scheele showed that calcium phosphate (Ca3(PO4)2) is found in bones, and they obtained phosphorus from bone ash. Antoine Lavoisier recognized phosphorus as an element in 1777. Bone ash was the major source of phosphorus until the 1840s. Phosphate rock, a mineral containing
calcium phosphate, was first used in 1850 and following the introduction of the electric arc furnace in 1890, this became the only source of phosphorus. Phosphorus, phosphates and phosphoric acid are still obtained from phosphate rock. Phosphate rock is a major feedstock in the fertilizer industry.
2. Occurrence
Due to its reactivity with air and many other oxygen-containing substances, phosphorus is not found free in nature but it is widely distributed in many different minerals. Phosphate rock, which is partially made of apatite (an impure tri-calcium phosphate mineral), is an important commercial source of this element. About 50 percent of the global phosphorus reserves are in the Arab nations. Large deposits of apatite are located in China, Russia, Morocco, Florida, Idaho, Tennessee, Utah, and elsewhere. Albright and Wilson in the United Kingdom and their Niagara Falls plant, for instance, were using phosphate rock in the 1890s and 1900s from Connetable, Tennessee and Florida; by 1950 they were using phosphate rock mainly from Tennessee and North Africa. In the early 1990s Albright and Wilson's purified wet phosphoric acid business was being adversely affected by phosphate rock sales by China and the entry of their long-standing Moroccan phosphate suppliers into the purified wet phosphoric acid business.
In 2007, at the current rate of consumption, the supply of phosphorus was estimated to run out in 345 years. However, scientists are now claiming that a "Peak Phosphorus" will occur in 30 years and that "At current rates, reserves will be depleted in the next 50 to 100 years. The stability of the +5 oxidation state is illlustrated by the wide range of phosphate materials available in the earth.
3. Applications
Match striking surface made of a mixture of red phosphorus, glue and ground glass. The glass powder is used to increase the friction.
Phosphorus, being an essential plant nutrient, finds its major use as a constituent of fertilizers for agriculture and farm production in the form of concentrated phosphoric acids, which can consist of 70% to 75% P2O5. Global demand for fertilizers led to large increase in phosphate (PO4
3–) production in the second half of the 20th century. Due to the essential nature of phosphorus to living organisms, the low solubility of natural phosphorus-containing compounds, and the slow natural cycle of phosphorus, the agricultural industry is heavily reliant on fertilizers that contain phosphate, mostly in the form of superphosphate of lime. Superphosphate of lime is a mixture of two phosphate salts, calcium dihydrogen phosphate Ca(H2PO4)2 and calcium sulfate dihydrate CaSO4·2H2O produced by the reaction of sulfuric acid and water with calcium phosphate.
Phosphorus is widely used to make organophosphorus
compounds, through the intermediates phosphorus chlorides
and two phosphorus sulfides: phosphorus pentasulfide, and
Phosphoric acid made from elemental phosphorus is used in
food applications such as some soda beverages. The acid is also
a starting point to make food grade phosphates. These include
mono-calcium phosphate that is employed in baking powder and
sodium tripolyphosphate and other sodium phosphates. Among
other uses these are used to improve the characteristics of
processed meat and cheese. Others are used in toothpaste.
Trisodium phosphate is used in cleaning agents to soften water
and for preventing pipe/boiler tube corrosion.
Phosphorus sesquisulfide is used in heads of strike-anywhere
matches.
In trace amounts, phosphorus is used as a dopant for n-type
semiconductors.
32P and 33P are used as radioactive tracers in biochemical
laboratories (see Isotopes).
Phosphate is a strong complexing agent for the hexavalent
uranyl (UO22+) species and this is the reason why apatite and
other natural phosphates can often be very rich in uranium.
Tributylphosphate is an organophosphate soluble in kerosene
and used to extract uranium in the Purex process applied in the
reprocessing of spent nuclear fuel.
2.1.3. Arsen
Arsenic is the chemical element that has the symbol As, atomic number 33 and atomic mass 74.92. Arsenic was first documented by Albertus Magnus in 1250. Arsenic is a notoriously poisonous metalloid with many allotropic forms, including a yellow (molecular non-metallic) and several black and grey forms (metalloids). Three metalloidal forms of arsenic, each with a different crystal structure, are found free in nature (the minerals arsenic sensu stricto and the much rarer arsenolamprite and pararsenolamprite). However, it is more commonly found as arsenide and in arsenate compounds, several hundred of which are known. Arsenic and its compounds are used as pesticides, herbicides, insecticides and in various alloys.
Arsenic sulfides (orpiment, realgar) and oxides have been known and used since ancient times. Zosimos (circa 300 AD) describes roasting sandarach (realgar) to obtain cloud of arsenic (arsenious oxide) which he then reduces to metallic arsenic. As the symptoms of arsenic poisoning were somewhat ill-defined, it was frequently used for murder until the advent of the Marsh test, a sensitive chemical test for its presence. (Another less sensitive but more general test is the Reinsch test.) Owing to its use by the ruling class to murder one another and its potency and discreetness, arsenic has been called the Poison of Kings and the King of Poisons.
During the Bronze Age, arsenic was often included in bronze, which made the alloy harder (so-called "arsenical bronze"). Albertus Magnus (Albert the Great, 1193–1280) is believed to have been the first to isolate the element in 1250 by heating soap together with arsenic trisulfide. In 1649, Johann Schröder published two ways of preparing arsenic.
Alchemical symbol for arsenic
Cadet's fuming liquid (impure cacodyl), the first organometallic compound, was synthesized in 1760 by Louis Claude Cadet de Gassicourt by the reaction of potassium acetate with arsenic trioxide.
2. Properties
a. Isotopes :
Naturally occurring arsenic is composed of one stable isotope, 75As. As of 2003, at least 33 radioisotopes have also been synthesized, ranging in atomic mass from 60 to 92. The most stable of these is 73As with a half-life of 80.3 days. Isotopes that are lighter than the stable 75As tend to decay by β + decay , and those that are heavier tend to decay by β - decay , with some exceptions.
At least 10 nuclear isomers have been described, ranging in atomic mass from 66 to 84. The most stable of arsenic's isomers is 68mAs with a half-life of 111 seconds.
Structure of yellow arsenic As4 and white phosphorus P4
Like phosphorus, arsenic is an excellent example of an element that exhibits allotropy, as its various allotropes have strikingly different properties. The three most common allotropes are metallic grey, yellow and black arsenic. The most common allotrope of arsenic is grey arsenic. It has a similar structure to black phosphorus (β-metallic phosphorus) and has a layered crystal structure somewhat resembling that of graphite. It consists of many six-membered rings which are interlinked. Each atom is bound to three other atoms in the layer and is coordinated by each 3 arsenic atoms in the upper and lower layer. This relatively close packing leads to a high density of 5.73 g/cm3.
Yellow arsenic (As4) is soft and waxy, somewhat similar to P4. Both have four atoms arranged in a tetrahedral structure in which each atom is bound to each of the other three atoms by a single bond, resulting in very high ring strain and instability. This form of arsenic is the least stable, most reactive, most volatile, least dense, and most toxic of all the allotropes. Yellow arsenic is produced by rapid cooling of arsenic vapour with liquid nitrogen. It is rapidly transformed into the grey arsenic by light. The yellow form has a density of 1.97 g/cm3.
b. Chemical properties
The most common oxidation states for arsenic are −3 (arsenides: usually alloy-like intermetallic compounds), +3 (arsenates(III) or arsenites, and most organoarsenic compounds), and +5 (arsenates: the most stable inorganic arsenic oxycompounds). Arsenic also bonds readily to itself, forming square As3−4 ions in the arsenide skutterudite. In the +3 oxidation state, the stereochemistry of arsenic is affected by the presence of a lone pair of electrons.
Arsenic is very similar chemically to its predecessor in the Periodic Table, phosphorus. Like phosphorus, it forms colourless, odourless, crystalline oxides As2O3 and As2O5 which are hygroscopic and readily soluble in water to form acidic solutions. Arsenic(V) acid is a weak acid. Like phosphorus, arsenic forms an unstable, gaseous hydride: arsine (AsH3). The similarity is so great that arsenic will partly substitute for phosphorus in biochemical reactions and is thus poisonous. However, in subtoxic doses, soluble arsenic compounds act as stimulants, and were once popular in small doses as medicine by people in the mid 18th century.
When heated in air, arsenic oxidizes to arsenic trioxide; the fumes from this reaction have an odour resembling garlic. This odour can be detected on striking arsenide minerals such as arsenopyrite with a hammer. Arsenic (and some arsenic compounds) sublimes upon heating at atmospheric pressure, converting directly to a gaseous form without an intervening liquid state. The liquid state appears at 20 atmospheres and above, which explains why the melting point is higher than the boiling point.
Arsenopyrite, also unofficially called mispickel, (FeAsS) is the most common arsenic-bearing mineral. In the lithosphere, the minerals of the formula M(II)AsS, with M(II) being mostly Fe, Ni and Co, are the dominant arsenic minerals.
Orpiment and realgar were formerly used as painting pigments, though they have fallen out of use owing to their toxicity and reactivity. Although arsenic is sometimes found native in nature, its main economic source is the mineral arsenopyrite mentioned above; it is also found in arsenides of metals such as silver, cobalt (cobaltite: CoAsS and skutterudite: CoAs3) and nickel, as sulfides, and when oxidised as arsenate minerals such as mimetite, Pb5(AsO4)3Cl and erythrite, Co3(AsO4)2·8H2O, and more rarely arsenites ('arsenite' = arsenate(III), AsO3
3− as opposed to arsenate (V), AsO43−).
In addition to the inorganic forms mentioned above, arsenic also occurs in various organic forms in the environment.Other naturally occurring pathways of exposure include volcanic ash, weathering of the arsenic-containing mineral and ores as well as groundwater. It is also found in food, water, soil and air.
4. Applications
a. Wood preservation :
The toxicity of arsenic to insects, bacteria, and fungi led to its use as a wood preservative. In the 1950s a process of treating wood with chromated copper arsenate (also known as CCA or Tanalith) was invented, and for decades this treatment was the most extensive industrial use of arsenic. Due to an increased understanding of arsenic's high level of toxicity, most countries banned the use of CCA in consumer products. The European Union and United States led this ban, beginning in 2004.
As of 2002, US-based industries consumed 19,600 metric tons of arsenic. 90% of this was used for treatment of wood with CCA. In 2007, 50% of the 5,280 metric tons of consumption was still used for this purpose. In the United States, the use of arsenic in consumer products was discontinued for residential and general consumer construction on December 31, 2003 and alternative chemicals are now used, such as Alkaline Copper Quaternary, borates, copper azole, cyproconazole, and propiconazole.
b. Other uses of arsen :
Various agricultural insecticides, termination and poisons. For
example Lead hydrogen arsenate was used well into the 20th
through a process of methylation. For example, the mold Scopulariopsis brevicaulis produce significant amounts of trimethylarsine if inorganic arsenic is present. The organic compound arsenobetaine is found in some marine foods such as fish and algae, and also in mushrooms in larger concentrations. The average person's intake is about 10–50 µg/day. Values about 1000 µg are not unusual following consumption of fish or mushrooms. But there is little danger in eating fish because this arsenic compound is nearly non-toxic. Some species of bacteria obtain their energy by oxidizing various fuels while reducing arsenate to arsenite. The enzymes involved are known as arsenate reductases (Arr).
5. Occupational exposures
Industries that use inorganic arsenic and its compounds include wood preservation, glass production, nonferrous metal alloys, and electronic semiconductor manufacturing. Inorganic arsenic is also found in coke oven emissions associated with the smelter industry. Occupational exposure and poisoning may occur in persons working in these industries
2.1.4.Antimony
Antimony (pronounced Latin: stibium) is a chemical element with the symbol Sb and an atomic number of 51. It has two stable isotopes, one with seventy neutrons, and the other with seventy-two. A silvery lustrous grey metalloid, it is found mainly as antimony sulfide, commonly known as stibnite. Elemental antimony has applications in electronics and as an alloy with other metals it is used for small arms ammunition.
1. History
One of the alchemical symbols for antimony
Antimony's sulfide compound, antimony(III) sulfide, Sb2S3 was recognized in antiquity, at least as early as 3000 BC. An artifact made of antimony dating to about 3000 BC was found at Tello, Chaldea (part of present-day Iraq), and a copper object plated with antimony dating between 2500 BC and 2200 BC has been found in Egypt. There is some uncertainty as to the description of the artifact from Tello. Although it is sometimes reported to be a vase, a recent detailed discussion reports it to be rather a fragment of indeterminate purpose. The first
European description of a procedure for isolating antimony is in the book De la pirotechnia of 1540 by Vannoccio Biringuccio, written in Italian. This book precedes the more famous 1556 book in Latin by Agricola, De re metallica, even though Agricola has been often incorrectly credited with the discovery of metallic antimony. A text describing the preparation of metallic antimony that was published in Germany in 1604 purported to date from the early fifteenth century, and if authentic it would predate Biringuccio. The book, written in Latin, was called "Currus Triumphalis Antimonyi" (The Triumphal Chariot of Antimony), and its putative author was a certain Benedictine monk, writing under the name Basilius Valentinus. Already in 1710 Wilhelm Gottlob Freiherr von Leibniz, after careful inquiry, concluded that the work was spurious, that there was no monk named Basilius Valentinus, and the book's author was its ostensible editor, Johann Thölde (ca. 1565-ca. 1624). There is now agreement among professional historians that the Currus Triumphalis.. was written after the middle of the sixteenth century and that Thölde was likely its author. An English translation of the "Currus Triumphalis" appeared in English in 1660, under the title The Triumphant Chariot of Antimony. The work remains of great interest, chiefly because it documents how followers of the renegade German physician, Philippus Theophrastus Paracelsus von Hohenheim (of whom Thölde was one), came to associate the practice of alchemy with the preparation of chemical medicines.
According to the traditional history of Middle Eastern alchemy, pure antimony was well known to Jābir ibn Hayyān, sometimes called "the Father of Chemistry", in the 8th century. Here there is still an open controversy: Marcellin Berthelot, who translated a number of Jābir's books, stated that antimony is never mentioned in them, but other authors claim that Berthelot translated only some of the less important books, while the more interesting ones (some of which might describe antimony) are not yet translated, and their content is completely unknown.
The first natural occurrence of pure antimony ('native antimony') in the Earth's crust was described by the Swedish scientist and local mine district engineer Anton von Swab in 1783. The type-sample was collected from the Sala Silvermine in the Bergslagen mining district of Sala, Västmanland, Sweden.
2. Properties
a. Physical properties :
A vial containing a black allotrope of antimony Native massive antimony with oxidation products
There are four known allotropes of antimony: a stable metallic form, and three meta-stable forms which are explosive, black and yellow. Each has its own distinct physical properties, the most common of which is metallic antimony, a brittle, silver-white shiny metal. When molten antimony is slowly cooled to metallic antimony, it forms with an hexagonal crystal structure, isomorphic with that of the grey form of arsenic.
The explosive form of antimony is formed from the electrolysis of antimony(III) trichloride, under specific temperatures and concentration. In a bath of hydrochloric with an antimony anode and platinum foil cathode, explosive antimony is deposited on the latter. When scratched with a sharp implement, an exothermic reaction occurs and white fumes given off as metallic antimony is formed; alternatively, when rubbed with a pestle in a mortar, an strong detonation occurs. Black antimony is formed when gaseous metallic antimony is rapidly cooled. It oxidies in air and is sometimes spontaneously combustible. At 100 °C, it gradually transforms into the stable form. Finally, the yellow allotrope of antimony is the most unstable. While it cannot be produced as the black allotrope by rapid cooling, it can only be formed by introducing oxygen into antimony hydride at -90 °C. Above this temperature and in ordinary light, it transforms into the stabler black allotrope.
b. Chemistry
Antimony trioxide (Sb4O6) is formed when antimony is burnt in an excess of air. In the gas phase, this compound exists as Sb4O6, a species that is retained when cooled to its solid, cubic form. However, in the rhombic form, the molecules polymerise to form chains of [Sb2O3]x. Antimony pentoxide, (Sb4O10) can only be formed by oxidation by concentrated nitric acid. Antimony also forms a mixed-valence oxide, antimony tetroxide (Sb2O4), where it is found in both the +3 and +5 oxidation states. Unlike phosphorus and arsenic, these various oxides are amphoteric and do not form well-defined oxoacids and react with acids to form antimony salts. Antimony trioxide dissolves in concentrated acid to form antimony oxo- (antimonyl) compounds such as SbOCl and (SbO)2SO4. The hypothetical antimonous acid Sb(OH)3 only exists as its salts,[19]:763 such as sodium antimonyte ([Na3SbO3]4), formed by fusing sodium oxide and Sb4O6. Transition metal antimonytes are best described as mixed metal oxides. Antimonyc acid exists only as the hydrate HSb(OH)6, forming salts containing the antimonate anion Sb(OH)−6. Dehydrating metal salts containing this anion yields mixed oxides.
Many antimony ores are sulfides, including stibnite (Sb2S3), pyrargyrite (Ag3SbS3), zinkenite, jamesonite, and boulangerite. Antimony pentasulfide is known, but is non-stoichiometric and contains only antimony in the +3 oxidation state. Several complex anions of antimony and sulfur are known, such as [Sb6S10]2− and [Sb8S13]2−.
Antimony forms two series of halides: SbX3 and SbX5, where X is one of the halogens. The trihalides SbF3, SbCl3, SbBr3, and SbI3 are all
molecular compounds having trigonal pyramidal molecular geometry. The trifluoride SbF3 is prepared by the reaction of Sb2O3 with HF.
Sb2O3 + HF → 2 SbF3 + 3 H2O
It is a strong Lewis acid that readily accepts fluoride ions to form the complex anions SbF−4 and SbF2−5. Molten SbF3 is a weak electrical conductor.
The trichloride SbCl3 is prepared by dissolving Sb2S3 in hydrochloric acid:
Sb2S3 + HCl → 2 SbCl3 + 3 H2S
The pentahalides SbF5 and SbCl5 have trigonal bipyramidal molecular geometry in the gas phase, but in the liquid phase, SbF5 is polymeric, whereas SbCl5 is monomeric. SbF5 is a powerful Lewis acid used to make the superacid fluoroantimonyc acid (HSbF6), and is an important solvent used in the study of noble gas compounds.
Antimony forms antimonydes with metals, such as indium antimonyde (InSb), and silver antimonyde (Ag3Sb). Treating antimonydes with acid produces the unstable toxic gas stibine, SbH3.
Sb3− + 3 H+ → SbH3
Stibine may also be produced by reacting Sb3+ salts with sources of the hydride ion H−. Antimony does not react with hydrogen directly to form stibine.
3. Occurrence and Production
The abundance of antimony in the Earth's crust is estimated at 0.2 to 0.5 parts per million, comparable to thallium at 0.5 parts per million and silver at 0.07 ppm. Even though this element is not abundant, it is found in over 100 mineral species. Antimony is sometimes found native, but more frequently it is found in the sulfide stibnite (Sb2S3) which is the predominant ore mineral. Commercial forms of antimony are generally ingots, broken pieces, granules, and cast cake. Other forms are powder, shot, and single crystals.
In 2005, China was the top producer of antimony with about 84% world share followed at a distance by South Africa, Bolivia and Tajikistan, reports the British Geological Survey. The mine with the largest deposits in China is Xikuangshan mine in Hunan Province with a estimated deposit of 2.1 million metric tons. Antimony is isolated from its ore by a reduction with scrap iron:
Sb2S3 + 3Fe → 2Sb + 3FeSIsolating antimony from its oxide, is performed by a charcoal reduction:
2Sb2O3 + 3C → 4Sb + 3CO2
4. Applications
a. Elemental antimony and alloys
Elemental antimony is increasingly being used in the
semiconductor industry as a dopant for ultra-high conductivity n-
type silicon wafers in the production of diodes, infrared detectors,
and Hall-effect devices. It is also used as an alloy, to increase lead's
hardness and mechanical strength, as in lead-acid batteries, which is
the most common use of antimony. It is used in antifriction alloys, such
as Babbit metal. It is used as an alloy in small arms ammunition,
buckshot, tracer ammunition, cable sheathing, type metal (e.g.
for linotype printing machines), solder – some "lead-free" solders
contain 5% Sb in pewter, and in hardening alloys with low tin content
in the manufacturing of organ pipes.
In the 1950s, tiny beads of a lead-antimony alloy were used to
dope the emitters and collectors of NPN alloy junction transistors
with antimony. A coin made of antimony was issued in the Keichow
Province of China in 1931. The coins were not popular, being too soft
and they wore quickly when in circulation. After the first issue no
others were produced. Elemental antimony as an antimony pill was
once used as a medicine. It could be reused by others after ingestion.
Treatments principally containing are known as antimonyals
and are used as emetics.Antimony compounds are used as
antiprotozoan drugs. Antimony potassium tartrate, or tartar
emetic, has been used in the past as an anti-schistosomal drug, later
replaced by praziquantel.
Antimony and its compounds are used in several veterinary preparations like anthiomaline or lithium antimony thiomalate, which is used as a skin conditioner in ruminants.
Antimony has a nourishing or conditioning effect on keratinized tissues, at least in animals. Antimony-based drugs, such as meglumine antimonyate, are also considered the drugs of choice for treatment of leishmaniasis in domestic animals. Unfortunately, as well as having low therapeutic indices, the drugs are poor at penetrating the bone marrow, where some of the Leishmania amastigotes reside, and so cure of the disease – especially the visceral form – is very difficult.
b. Other uses of antimony :
In the heads of some safety matches in nuclear reactors together with beryllium in startup neutron sources; in the form of antimony oxides, antimony sulfides, and antimony trichloride are used in the making of flame-proofing compounds, ceramic enamels, glass, paints, and pottery. Antimony trioxide is the most important of the antimony compounds and is primarily used in flame-retardant formulations. These flame-retardant applications include such markets
as children's clothing, toys, aircraft and automobile seat covers. It is also used in the fiberglass composites industry as an additive to polyester resins for such items as light aircraft engine covers. The resin will burn while a flame is held to it but will extinguish itself as soon as the flame is removed.
2.1.5.Bismut
Bismuth is a chemical element that has the symbol Bi and atomic number 83. This trivalent poor metal chemically resembles arsenic and antimony. Bismuth is heavy and brittle; it has a silvery white color with a pink tinge owing to the surface oxide. Bismuth is the most naturally diamagnetic of all metals, and only mercury has a lower thermal conductivity. It is generally considered to be the last naturally occurring stable, non-radioactive element on the periodic table, although it is actually slightly radioactive. Its only non-synthetic isotope bismuth-209 decays via alpha decay into thallium-205, with an extremely long half-life of 1.9 × 1019 years.
Bismuth compounds are used in cosmetics, medicines, and in medical procedures. As the toxicity of lead has become more apparent in recent years, alloy uses for bismuth metal as a replacement for lead have become an increasing part of bismuth's commercial importance.
Bismuth crystal with an iridescent oxide surface Bismuth crystals
Bismuth is a brittle metal with a white, silver-pink hue, often occurring in its native form with an iridescent oxide tarnish showing many colors from yellow to blue. The spiral stair stepped structure of a bismuth crystal is the result of a higher growth rate around the outside edges than on the inside edges. The variations in the thickness of the oxide layer that forms on the surface of the crystal causes different wavelengths of light to interfere upon reflection, thus displaying a rainbow of colors. When combusted with oxygen, bismuth burns with a blue flame and its oxide forms yellow fumes. Its toxicity is much lower than that of its neighbors in the periodic table such as lead, tin, tellurium, antimony, and polonium.
Although ununpentium is theoretically more diamagnetic, no other metal is verified to be more naturally diamagnetic than bismuth. (Superdiamagnetism is a different physical phenomenon). Of any metal, it has the second lowest thermal conductivity (after mercury) and the highest Hall coefficient. It has a high electrical resistance. When deposited in sufficiently thin layers on a substrate, bismuth is a semiconductor, rather than a poor metal.
Elemental bismuth is one of very few substances of which the liquid phase is denser than its solid phase (water being the best-known example). Bismuth expands 3.32% on solidification; therefore, it was long an important component of low-melting typesetting alloys, where it compensated for the contraction of the other alloying components.
Though virtually unseen in nature, high-purity bismuth can form distinctive hopper crystals. These colorful laboratory creations are typically sold to collectors. Bismuth is relatively nontoxic and has a low melting point just above 271°C, so crystals may be grown using a household stove, although the resulting crystals will tend to be lower quality than lab-grown crystals.
1. History
Bismuth (New Latin bisemutum from German Wismuth, perhaps from weiße Masse, "white mass") was confused in early times with tin and lead because of its resemblance to those elements. Bismuth has been known since ancient times, and so no one person is credited with its discovery. Agricola, in De Natura Fossilium states that bismuth is a distinct metal in a family of metals including tin and lead in 1546 based on observation of the metals and their physical properties. Claude François Geoffroy demonstrated in 1753 that this metal is distinct from lead and tin.
"Artificial bismuth" was commonly used in place of the actual metal. It was made by hammering tin into thin plates, and cementing them by a mixture of white tartar, saltpeter, and arsenic, stratified in a crucible over an open fire.
Bismuth was also known to the Incas and used (along with the usual copper and tin) in a special bronze alloy for knives.
In the Earth's crust, bismuth is about twice as abundant as gold. It is not usually economical to mine it as a primary product. Rather, it is usually produced as a byproduct of the processing of other metal ores, especially lead, tungsten (China), tin, copper, and also silver (indirectly) or other metallic elements.
The most important ores of bismuth are bismuthinite and bismite. In 2005, China was the top producer of bismuth with at least 40% of the world share followed by Mexico and Peru, reports the British Geological Survey. Native bismuth is known from Australia, Bolivia, and China.
3. Recycling
While bismuth is most available today as a byproduct, its sustainability is more dependent on recycling. Bismuth is mostly a byproduct of lead smelting, along with silver, zinc, antimony, and other metals, and also of tungsten production, along with molybdenum and tin, and also of copper production. Recycling bismuth is difficult in many of its end uses, primarily because of scattering. Probably the easiest to recycle would be bismuth-containing fusible alloys in the form of larger objects, then larger soldered objects. Half of the world solder consumption is in electronics (i.e., circuit boards). As the soldered objects get smaller or contain little solder or little bismuth, the recovery gets progressively more difficult and less economic, although solder with a sizable silver content will be more worth recovering. Next in recycling feasibility would be sizeable catalysts with a fair bismuth content, perhaps as bismuth phosphomolybdate, and then bismuth used in galvanizing and as a free-machining metallurgical additive. Finally, the bismuth in the uses where it gets scattered the most, in stomach medicines (bismuth subsalicylate), paints bismuth vanadate on a dry surface, pearlescent cosmetics (bismuth oxychloride), and bismuth-containing bullets.
The most important sustainability fact about bismuth is its byproduct status, which can either improve sustainability (i.e., vanadium or manganese nodules) or, for bismuth from lead ore, constrain it; bismuth is constrained. The extent that the constraint on bismuth can be ameliorated or not is going to be tested by the future of the lead storage battery, since 90% of the world market for lead is in storage batteries for gasoline or diesel-powered motor vehicles.
The life-cycle assessment of bismuth will focus on solders, one of the major uses of bismuth, and the one with the most complete information. The average primary energy use for solders is around 200 MJ per kg, with the high-bismuth solder (58% Bi) only 20% of that value, and three low-bismuth solders (2% to 5% Bi) running very close
to the average. The global warming potential averaged 10 to 14 kg carbon dioxide, with the high-bismuth solder about two-thirds of that and the low-bismuth solders about average. The acidification potential for the solders is around 0.9 to 1.1 kg sulfur dioxide equivalent, with the high-bismuth solder and one low-bismuth solder only one-tenth of the average and the other low-bismuth solders about average. There is very little life-cycle information on other bismuth alloys or compounds.
4. Chemistry properties
Bismuth forms trivalent and pentavalent compounds. The trivalent compounds are more common. Many of its chemical properties are similar to other elements in its group; namely, arsenic and antimony, although it is less toxic than those elements.
Bismuth is stable to both dry and moist air at ordinary temperatures. At elevated temperatures, the vapours of the metal combine rapidly with oxygen, forming the yellow trioxide, Bi2O3 On reaction with base, this oxide forms two series of oxyanions: BiO−2, which is polymeric and forms linear chains, and BiO3−3. The anion in Li3BiO3 is actually a cubic octameric anion, Bi8O24−24, whereas the anion in Na3BiO3 is tetrameric.
Bismuth sulfide, Bi2S3, occurs naturally in bismuth ores. It is also produced by the combination of molten bismuth and sulfur. Unlike earlier members of group 15 elements such as nitrogen, phosphorus, and arsenic, and similar to the previous group 15 element antimony, bismuth does not form a stable hydride analogous to ammonia and phosphine. Bismuth hydride, bismuthine (BiH3), is an endothermic compound that spontaneously decomposes at room temperature. It is stable only below −60°C.
The halides of bismuth in low oxidation states have been shown to have unusual structures. What was originally thought to be bismuth(I) chloride, BiCl, turns out to be a complex compound consisting of Bi5+9 cations and BiCl2−5 and Bi2Cl2−8 anions. The Bi5+9 cation has a distorted tricapped trigonal prismic molecular geometry, and is also found in Bi10Hf3Cl18, which is prepared by reducing a mixture of hafnium(IV) chloride and bismuth chloride with elemental bismuth, having the structure [Bi+][Bi5+9][HfCl2−6]3. Other polyatomic bismuth cations are also known, such as Bi2+8, found in Bi8(AlCl4)2. Bismuth also forms a low-valence bromide with the same structure as "BiCl". There is a true monoiodide, BiI, which contains chains of Bi4I4 units. BiI decomposes upon heating to the triiodide, BiI3, and elemental bismuth. A monobromide of the same structure also exists.
In oxidation state +3, bismuth forms trihalides with all of the halogens: BiF3, BiCl3, BiBr3, and BiI3. All of these, except BiF3, are hydrolysed by water to form the bismuthyl cation, BiO+, a commonly encountered bismuth oxycation.[14] Bismuth(III) chloride reacts with hydrogen chloride in ether solution to produce the acid HBiCl4.
Bismuth dissolves in nitric acid to form bismuth(III) nitrate, Bi(NO3)3. In the presence of excess water or the addition of a base, the
Bi3+ ion reacts with the water to form BiO+, which precipitates as (BiO)NO3. The oxidation state +5 is less frequently encountered. One such compound is BiF5, a powerful oxidising and fluorinating agent. It is also a strong fluoride acceptor, reacting with xenon tetrafluoride to form the XeF+3 cation:
BiF5 + XeF4 → XeF + 3BiF−6
The dark red bismuth(V) oxide, Bi2O5, is unstable, liberating O2 gas upon heating. In aqueous solution, the Bi3+ ion exists in various states of hydration, depending on the pH:
pH range
Species
<3 Bi(H2O)3+6
0-4 Bi(H2O)5OH2+
1-5Bi(H2O)4(OH)2+
5-14Bi(H2O)3(OH)3
>11Bi(H2O)2(OH)4−
These mononuclear species are in equilibrium. Polynuclear species also exist, the most important of which is BiO+, which exists in hexameric form as the octahedral complex [Bi6O4(OH)4]6+ (or 6 [BiO+]·2 H2O).
5. Applications
Bismuth oxychloride is sometimes used in cosmetics. Bismuth subnitrate and bismuth subcarbonate are used in medicine.a) Health
Bismuth subsalicylate (the active ingredient in Pepto-Bismol and
(modern) Kaopectate) is used as an antidiarrheal and to treat
some other gastro-intestinal diseases (oligodynamic effect). The
means by which this appears to work is still not well-
documented. It is thought to be some combination of:
Killing some bacteria that cause diarrhea
Reducing inflammation/irritation of stomach and intestinal
VB group consists of four elements, namely Vanadium (V), niobium (Nb), Tantalum (Ta). Some properties of the elements of this group we can see in the following :
Outermost elektron V ( 3d3 4S2 ), Nb ( 4d4 4S1 ), Ta ( rd3 4S2 ). Nomer oxsidation varied, stability of +5 oxsidation state increases from V-Nb-Ta. Thus V+5 easily reduced to V+2 is Nb+5 and Ta+5 remained stable, V+5 is a good oxidant. The unique nature of each elements diminished view of the reduced size of cations. Thus, the unique nature of V+4 > V+3 > V+2. Consequently VC14 covalen character. Oxide properties , V205 amphoter but more acidic, while Nb205 and Ta205 fever bases.
Unreactive at room temperature but on heating reacts to from VC15, VC14, VC13, dan VI3. Nb and Ta are formed only halide type MX5. All halides are covalen and volatile. With H2 from non-stoichiometric compuonds, VH0,7 ; NbH0,86 and TaH0,76. Tendency from a complex : V > Nb > Ta. Radius (ionic and atomic) Nb with Ta is almost the same as a result lanthanida contraction, so that bolt are almost the same properties.
Compounds these metals with an oxidation number lower winnowing colored because of the d orbitals that contain partially.
2.2.1. Vanadium
1. History and discovery :
In 1831, Swedish chemist, Niel Grabiol Sefstrom discover new elements in the iron ore in Sweden. The elements was named Vanadium as a means goddess Vanadis gorgeous. Year 1865 Roscor and Thorpe found this elements to be with layers of copper and lawer
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standstone of Cheshire. Vanadium compounds are abundant in scattered earth’s crust. Some prominent vanadium minerals is :
vanadinite : 3 Pb3(VO4)2 . PbCl2
carnotite : K2O . 2UO3 . V2O53H2O
patronite : V2S5 . 3CuS2
Vanadium also occurs in the clay, rocks, coal and crude oil with a small degree.
How to get the vanadium them is by extraction from several compounds, namey :
a) From vanadinite, extraction from ore involves several stages:
Separation of PbCl2
Ore is reacted with concentrated HCl, PbCl2 will settle, dioxovandium chlotida (VO2Cl) remained in solution.
Making V2O5
PbCl2 after inseparable, solution plus NH4C1 and saturated with NH3, forming NH4VO3 which when formed V205.
Reduction of V205.
V205 reduced by Ca in the 900 - 950 ° C to obtain pure vanadium.
b) From carnotite.
The making of sodium orthovanadate.
Carnotite melted with Na2C03, the liquid obtain extracted with water to precipitate Fe(OH)3, evaporated and cooled the importance of the Na3V04.
Making V205.
Solution containing Na3V04 given NH4C1 and saturated NH3, forming NH4V03 (amonium metavanadate), heated to obtain V205.
Reduction of V205.
Rich Mardenand way of vanadium metal was obtained pure.
2. Manufacture of metals :
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This metal is very hard to obtain in a pure state because the melting and high reactivity toward O2, N2 and C at high temperature.
Vanadium ± 99 % can be obtained by V205 with Al (thermit
process).
Pure Vanadium was obtained by reducing VC13 with Na or
with H2 at temperature 900 ° C. VC13 obtained from V205 with
S2C12 at 300 °C.
VC14 reduction with Mg can be obtained 99,3 % vanadium.
3. Aliase vanadium :
Commercial products are mainly as aliase vanadium,
Ferro vanadium and Cupro vanadium
Both are made by reducing vanadium mixed oxide with Fe or Cu with carbon preformance in the electric furnace.
Nikelo vanadium, made by healting a mixture of V205 + NiO.
Obalto vanadium, created by mixing sediment (from the
reactive Na-vanadate solutions with cobalto sulphate) with
Na2C03 electric furnace.
4. Usage :
Addition of 0,1 to 0,3 % V steel will increase the power range.
Vanadium is important for the tools of high speed steel.
V205 used as a catalyst in oxidation naphtalen and also in
making H2S04 contact process.
5. Properties :
Healted H2 (no other gases) at 1100°C to vanadium hydride
is stable.
These metal reactive in cold conditions, when heated formed
V20 (brown), hested hold formed V203 (black), V204 (blue),
end the V205 (orange). These metals burn with a bright flame
with oxygen.
When heated with Cl2 fromed dry VC14.
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This metal does not reactive with bromine water, HCl/winter,
releasing H2 with HF and forming a green solution.
6. Substances :
Vanadium form compounds with oxidation numbr +5, +4, +3 and +2. Compounds with lower oxidation number is reducing agent,is unique and colorful.
1. Compounds V+5 (colorless) is reducted with a reducing agent
according to the following changes :
a. Vanadium pentoksida, V2O5
Made from:
Oxidation/ metal or oxide , toasting with low oxidation states.
V2O5 as the final result.
Hyidrolysis VOCl3.
Heating amonium vanadate.
Usage :
As a catalyst in the oxidation of SO2 → SO3, in sulfuric acid.
V2O5
2SO2 + O2 ↔ 2SO3
Catalyst in a alcohol oxidation and hydrogenation of olefin.
b. Vanadium pentaflourida, VF5.
This compounds is expressed as sublimat pure white. Be made with heating VF4 in a nitrogen environment, at a temperature 350°C – 650°C. This compounds is very solube in water or organic solvents.
c. Vanadium oxitrikhlorida, VOCl3.
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This compounds is made by passing Cl2 dried at VO3to be heated. This compounds is yellow with a translucent boiling point 127 ° C.
d. Vanadium pentasulfida, V2S5.
This compounds is made by heating mixture of vanadium trisulfida, with sulfur without air at 400 ° C. This compounds form black powder.
2. Compound V+4.
Compounds with a +4 oxidation state is stable, easy to manufacture.
a. Vanadium titroksida, V2O4 atau VO2.
Created by heating a mixture of vanadium trioksida and vanadium pentoksida without air with a molar amount same. These compounds form a dark blue cystal, easily soluble acid or alkaline.
b. Vanadium titraflourida, VF4.
Created from the reaction of HF anhydride with VCl4. Start temperature –28°C and rose slowly to 0°C. Flouride is a brownish yellow powder, soluble in water forming a blue solution.
3. Compounds Vanadil
This compound contains a cation vanadil (VO+2) where the number oxidation +4, is unique, colored blue. Vanadil chloride made from the hydrolysis VCl4
VCl4 + H2O → VOCl2 + 2HCl
Or from the hydrolysis V2O5 with HCl
V2O5 + HCl → 2VOCl2 +3H2O + Cl2
VOCl2 compounds are strong reducing agents that are used commercial in coloration. Only the E° of VO+2/ VO3 is –1 volt.
4. Compound V+3.
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a. Vanadium trioxside, V2O3.
Created by reducing V2O5 with hydrogen. V2O3 is alkaline, soluble in acid gives hezaquo ions, V(H2O)6
3+.
b. oxihalida halide and Vanadium.
Vanadium triflourida, VF3.3H2O created when the V2O3 dissolved HF. Trihalida the other is the VCl3 and VBr3, moderate VI3 not known. Vanadium oxihalida is known VOCl and VOBr. Both are insoluble in water but soluble in acid.
5. Compound V+2.
Compound V+2 colored and paramagnetic ion V+2 is a strong reducing agent. V+2 dilute solution (violet) to reduce water liberate H2.
V+2 + H+ → V+3 + ½ H2
(violet) (green)
6. Compound V+1, V-1 and V°.
Oxidation is not common, stabilized by the ligand acid Π. Oxidation numer +1 was found in compound V(CO)6
-1.
2.2.2.Niobium and Tantalum
Although these two elements are metals in some respects, chemical properties such as non logam Although these two elements has no cation real, but formed several anions. Halides and very oxihalida hydrolysed volatile and easily.
Niobium Tantalum
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1. History of Niobium and Tantalum
Niobium was discovered by British chemist from the Charles Hatchett in the year 1801. Hatchett found in samples of niobium minerals columbite and called the new element columbium. Columbium discovered by Hatchett in a mixture with Tantalum. Tantalum itself was discovered by a chemist Anders Ekeberg Sweden.
Then confusion arises between the two element are both are the same element or not. Then in 1809, English chemist William Hyde Wollaston comparing shapes oxidation that occurred columbium – columbite, with a density 5918 g/ml, and tantalum – tantalite with a density 7935 g/ml, and concluded that both elements are different because the course has different densities, so the name became niobium and columbium be Tantalum fixed Tantalum.
2. Properties :
Both metals are very difficult to separate. Niobium metal is thin, soft, grayish, shiny, can be bent, high melting point (Nb= 2468 ° C). Tantalum metal is dark, dense, can be bent, the haerder than Niobium, electrical conductivity and high heat, high melting point (Ta = 2996° C), very resistant to acid. Both can be dissolved with HNO3, HF and dissolve very slowly in the alkaline liquid.
3. Compounds :
a) Compound Nb+5 and Ta+5
Nb2O5 and Ta2O5
Created by dihydroksioksida hydrate (often called acid niobat or tantalat), or by roasting a particular compound with excess oxygen. Both these compound are in powder from dense, relatively inert chemically, almost did not react with acid except for concentrated HF. These compound can be dissolved with melted with alkaline hydrogen sulphate, alkali carbonate or alkali hydroxide.
NbX5 and TaX5 (X = halide).
Compounds NbF5 and TaF5 prepared by reaction direct flourinasi metal or pentachlorida. Both are solid white. It’s easy yawn. Melting point of Nb = 80 ° C, Ta = 95 ° C. boiling point Nb = 235 ° C, Ta = 229 ° C, forming a colorless liquid and vapor. Compound other
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halide yellow to brown, made with the reaction direct metal with excess halogen. Halide is dotted halides liquid and boiling point between 200 - 300 ° C, soluble in organic solvents such us ether, CC14, and so forth.
b) Compound Nb and Ta with low oxidation state.
Nb02Oxide and TaOx (x = 2 s.d 2,5)
Tetrahalida.
All halides known except TaF4. Compound NbF4 not very volatile, paramagnetic. Tetrakhlorida and colored tetrakronida dark brown ao black. Nbl4 can be obtained easily by heating NbI5 up to 300° C. these compounds diamagnetism.
4. Applications
Niobium
• As a material for nuclear powder plant construction.
• As a mixture of rust resistant metal (example Niobium foil),
caused by the presence of Niobium carbide and compound
Niobium Nitrit, with the concentration of Niobium in the
compound approximately 0.1%.
• As a superconducting magnet (3 tesla clinical Magnetic
resonance imaging scanner), and superconducting radio
frequency.
• The manufacture of coin currency (example Austria 2003, Latvia
2004).
• In medical equipment, Pace maker.
• In the manufacture of jewelry.
Tantalum
• Used in the manufacture of children in a loboratory scale.
• Used in making electronic devices.
• In making the camera lens.
• To produce variations of the alloy that has a point high boiling
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and the forces of good.
• Preparation of equipment made of metal carbide.
• Used in the manufacture of jet engine components.
2.2.3.Dubnium
Dubnium is an element of group Vb transition metals are made though nuclesr fusion reactions. This eleemnt was discovered by Albert Ghiorso in 1970. Because of very large nuclei dubnium then dubnium constituents stable and not readily disintegrate.
Dubnium elements can br created by firing element amerisium with atoms of neon, adn produces isotopes dubnium isotopes, and quickly decays by emitting in the from of electromagnetic radition. Reaction as follows :
Compound s that can be formed for example Db205 (Dubnium pentoxide), DbX5 (Dubnium Halide), halide complexes Db04
3" , DbF6", DbF8
3. Other information about the elements Dubnium not clesrly known.
CHAPTER III
SUMMARY
1. Groups VA consist of Nitrogen, Phospor, Arsenic, Antimony , and
Bismut. Groups VB consist of Vanadium, Niobium, Tantalum, and
Dubnium.
2. Nitrogen has use for freeze.
3. Match striking surface made of a mixture of red phosphorus, glue
and ground glass. The glass powder is used to increase the friction.
4. The toxicity of arsenic to insects, bacteria, and fungi led to its use
as a wood preservative.
5. Elemental antimony is increasingly being used in the