Contract Job in Chem

7/28/2019 Contract Job in Chem

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History

In 1829, Johann Wolfgang Dbereinerobserved that many of the

elements could be grouped into t r iads (groups of three) based on their

chemical properties. Lithium, sodium, and potassium, for example,were grouped together as being soft, reactive metals. Dbereiner also

observed that, when arranged by atomic weight, the second member of

each triad was roughly the average of the first and the third. This

became known as the Law of Triads. In 1864, German chemist Julius

Lothar Meyerpublished a table of the 49 known elements arranged by

valency. The table revealed that elements with similar properties often

shared the same valency. English chemist John Newlands produced a

series of papers in 1864 and 1865 that described his own classification

of the elements: he noted that when listed in order of increasing

atomic weight, similar physical and chemical properties recurred at

intervals of eight, which he likened to the octaves of music. Russian

chemistry professorDmitri Ivanovich Mendeleev and German chemist

Julius Lothar Meyer independently published their periodic tables in

1869 and 1870, respectively. They both constructed their tables in a

similar manner: by listing the elements in a row or column in order of

atomic weight and starting a new row or column when the

characteristics of the elements began to repeat.
http://en.wikipedia.org/wiki/Johann_Wolfgang_D%C3%B6bereinerhttp://en.wikipedia.org/wiki/Lithiumhttp://en.wikipedia.org/wiki/Sodiumhttp://en.wikipedia.org/wiki/Potassiumhttp://en.wikipedia.org/wiki/Reactivity_(chemistry)http://en.wikipedia.org/wiki/Law_of_triadshttp://en.wikipedia.org/wiki/Julius_Lothar_Meyerhttp://en.wikipedia.org/wiki/Julius_Lothar_Meyerhttp://en.wikipedia.org/wiki/John_Alexander_Reina_Newlandshttp://en.wikipedia.org/wiki/Octavehttp://en.wikipedia.org/wiki/Dmitri_Mendeleevhttp://en.wikipedia.org/wiki/Dmitri_Mendeleevhttp://en.wikipedia.org/wiki/Dmitri_Mendeleevhttp://en.wikipedia.org/wiki/Dmitri_Mendeleevhttp://en.wikipedia.org/wiki/Dmitri_Mendeleevhttp://en.wikipedia.org/wiki/Octavehttp://en.wikipedia.org/wiki/John_Alexander_Reina_Newlandshttp://en.wikipedia.org/wiki/Julius_Lothar_Meyerhttp://en.wikipedia.org/wiki/Julius_Lothar_Meyerhttp://en.wikipedia.org/wiki/Julius_Lothar_Meyerhttp://en.wikipedia.org/wiki/Julius_Lothar_Meyerhttp://en.wikipedia.org/wiki/Law_of_triadshttp://en.wikipedia.org/wiki/Law_of_triadshttp://en.wikipedia.org/wiki/Reactivity_(chemistry)http://en.wikipedia.org/wiki/Potassiumhttp://en.wikipedia.org/wiki/Sodiumhttp://en.wikipedia.org/wiki/Lithiumhttp://en.wikipedia.org/wiki/Johann_Wolfgang_D%C3%B6bereinerhttp://en.wikipedia.org/wiki/Johann_Wolfgang_D%C3%B6bereiner


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THE OLD PERIODIC TABLE


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Periodic Table and its Features The periodic table is a tabular display of the chemical elements,

organized on a basis of their properties. Elements are presented in

increasing atomic number; while rectangular in general outline,gaps are included in the rows orperiods to keep elements with

similar properties together, such as the halogens and the noble

gases, in columns orgroups, forming distinct rectangular areas

orblocks.

A group orfamily is a vertical column in the periodic table. Groups

are considered the most important method of classifying theelements. In some groups, the elements have very similar

properties and exhibit a clear trend in properties down the group.

Under the international naming system, the groups are numbered

numerically 1 through 18 from the left most column (the alkali

metals) to the right most column (the noble gases).

A per iod is a horizontal row in the periodic table. Although groups

are the most common way of classifying elements, there are

regions where horizontal trends are more significant than vertical

group trends, such as the f-block, where

the lanthanides and actinides form two substantial horizontal

series of elements.
http://en.wikipedia.org/wiki/Chemical_elementhttp://en.wikipedia.org/wiki/Atomic_numberhttp://en.wikipedia.org/wiki/Period_(periodic_table)http://en.wikipedia.org/wiki/Halogenhttp://en.wikipedia.org/wiki/Noble_gashttp://en.wikipedia.org/wiki/Noble_gashttp://en.wikipedia.org/wiki/Group_(periodic_table)http://en.wikipedia.org/wiki/Block_(periodic_table)http://en.wikipedia.org/wiki/F-blockhttp://en.wikipedia.org/wiki/Lanthanidehttp://en.wikipedia.org/wiki/Actinidehttp://en.wikipedia.org/wiki/Actinidehttp://en.wikipedia.org/wiki/Lanthanidehttp://en.wikipedia.org/wiki/F-blockhttp://en.wikipedia.org/wiki/F-blockhttp://en.wikipedia.org/wiki/F-blockhttp://en.wikipedia.org/wiki/Block_(periodic_table)http://en.wikipedia.org/wiki/Group_(periodic_table)http://en.wikipedia.org/wiki/Noble_gashttp://en.wikipedia.org/wiki/Noble_gashttp://en.wikipedia.org/wiki/Halogenhttp://en.wikipedia.org/wiki/Period_(periodic_table)http://en.wikipedia.org/wiki/Atomic_numberhttp://en.wikipedia.org/wiki/Chemical_element


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A block of the periodic table of elements is a set

of adjacent groups. The respective highest-

energy electrons in each element in a block

belong to the same atomic orbital type. Each

block is named after its characteristic orbital: s,

p, d, f, g (no elements belonging to the g-block

have been observed).


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Trends in the Periodic TableAtomic Radius

The atomic radius of an element is half of the distance

between the centers of two atoms of that element that are justtouching each other. Generally, the atomic radius decreases

across a period from left to right and increases down a given

group. The atoms with the largest atomic radii are located in

Group I and at the bottom of groups. Moving from left to right

across a period, electrons are added one at a time to the

outer energy shell. Electrons within a shell cannot shield

each other from the attraction to protons. Since the number

of protons is also increasing, the effective nuclear charge

increases across a period. This causes the atomic radius to

decrease. Moving down a group in the periodic table, the

number of electrons and filled electron shells increases, butthe number of valence electrons remains the same. The

outermost electrons in a group are exposed to the same

effective nuclear charge, but electrons are found farther from

the nucleus as the number of filled energy shells increases.

Therefore, the atomic radii increase.


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Ionization Energy

The ionization energy, or ionization potential, is the energy

required to completely remove an electron from a gaseous

atom or ion. The closer and more tightly bound an electron is

to the nucleus, the more difficult it will be to remove, and the

higher its ionization energy will be. The first ionization energy

is the energy required to remove one electron from the parent

atom. The second ionization energy is the energy required to

remove a second valence electron from the univalent ion toform the divalent ion, and so on. Successive ionization

energies increase. The second ionization energy is always

greater than the first ionization energy. Ionization energies

increase moving from left to right across a period (decreasing

atomic radius). Ionization energy decreases moving down agroup (increasing atomic radius). Group I elements have low

ionization energies because the loss of an electron forms a

stable octet.


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Electron Affinity

Electron affinity reflects the ability of an atom to accept an

electron. It is the energy change that occurs when an electron

is added to a gaseous atom. Atoms with stronger effective

nuclear charge have greater electron affinity. Some

generalizations can be made about the electron affinities of

certain groups in the periodic table. The Group IIA elements,

the alkaline earths, have low electron affinity values. These

elements are relatively stable because they havefilled s subshells. Group VIIA elements, the halogens, have

high electron affinities because the addition of an electron to

an atom results in a completely filled shell. Group VIII

elements, noble gases, have electron affinities near zero, since

each atom possesses a stable octet and will not accept anelectron readily. Elements of other groups have low electron

affinities.


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Electronegativity

Electronegativity is a measure of the attraction of an atom for

the electrons in a chemical bond. The higher the

electronegativity of an atom, the greater its attraction for

bonding electrons. Electronegativity is related to ionization

energy. Electrons with low ionization energies have low

electronegativities because their nuclei do not exert a strong

attractive force on electrons. Elements with high ionization

energies have high electronegativities due to the strong pullexerted on electrons by the nucleus. In a group, the

electronegativity decreases as atomic number increases, as a

result of increased distance between the valence electron and

nucleus (greater atomic radius). An example of an

electropositive (i.e., low electronegativity) element is cesium;an example of a highly electronegative element is fluorine.


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Groups

1. Alkali Metals(Group 1)

-do not occur elementally in nature- have one valence electron

- have low ionization energies

- form colorless ions, each with a +1 charge

- are reactive metals obtained by reducing the +1 ions in

their natural compounds- are stored under kerosene or other hydrocarbon solvent

because they react with water vapor or oxygen in air

- form water-soluble bases

- are strong reducing agents

- are good conductors of electricity and heat- are ductile, malleable, and soft enough to be cut with a

knife

- have a silvery luster, low density, and a low melting point


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2. Alkaline Earth Metals(Group 2)

-do not occur elementally in nature

-occur most commonly as carbonates, phosphates,

silicates and sulfates- occur naturally as compounds that are either insoluble or

only slightly soluble in water

- contain two valence electrons

- tend to lose two electrons per atom, forming ions with a

+2 charge- are less reactive than alkali metals

- primarily form ionic compounds

- react with water to form bases and hydrogen gas

- are good conductors of heat and electricity

- are ductile and malleable

- have a silvery luster

- include the naturally radioactive element radium


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3. Transition Metals(Group 3-12)

- consist of metals in groups 3 through 12

- contain one or two valence electrons

- are usually harder and more brittle than metals in groups1 and 2

- have higher melting and boiling points that metals in

groups 1 and 2

- are good conductors of heat and electricity

- are malleable and ductile- have a silvery luster, except copper and gold

- include radioactive elements 89 through 109

- include mercury, the only metal that is liquid at room

temperature

- have chemical properties that tend to differ from each

other

- tend to have two or more common oxidation states

- often form colored compounds

- may form complex ions


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4. Boron Family(Group 13)

- do not occur elementally in nature

- are scarce in nature (except aluminum, which is the most

abundant metallic element)- have three valence electrons

- are metallic (except boron, which is a solid metalloid)

- are soft and have low melting points (except boron, which

is hard and has a high melting point)

- are chemically reactive at moderate temperatures (exceptboron)


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5. Carbon Family(Group 14)

- includes a nonmetal (carbon), two metalloids (silicon and

germanium) and two metals (tin and lead)

- vary greatly in both physical and chemical properties- occur in nature in both combined and elemental forms

- have four valence electrons

- are relatively unreactive

- tend to form covalent compounds (tin and lead also form

ionic compounds)


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6. Nitrogen Family(Group 15)

- consists of two nonmetals (nitrogen and phosphorus), two

metalloids (arsenic and antimony), and one metal

(bismuth)- nitrogen is most commonly found as atmospheric N2,

phosphorus as phosphate rock, and arsenic, antimony, and

bismuth as sulfides or oxides; antimony and bismuth are

also found elementally

- range from very abundant elements (nitrogen andphosphorus) to relatively rare elements (arsenic, antimony,

and bismuth)

- have five valence electrons

- tend to form covalent compounds, most commonly with

oxidation numbers of +3 or +5

- are solids at room temperature, except nitrogen


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7. Oxygen Family(Group 16)

- occur elementally in nature and in combined states

- consists of three nonmetals (oxygen, sulfur, and

selenium), one metalloid (tellurium), and one metal(polonium)

- have six valence electrons

- tend to form covalent compounds with other elements

- tend to exist as diatomic and polyatomic molecules, such

as O2, O3, S6, S8, and Se8- commonly exist in compounds with the -2 oxidation state,

but often exhibit other oxidation states


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8. Halogen Family(Group 17)

- are nonmetals and occur in combined form in nature,

mainly as metal halides

- are found in the rocks of Earth's crust and dissolved insea water

- range from fluorine, the 13th most abundant element, to

astatine, which is one of the rarest

- exist at room temperature as a gas (F2 and Cl2), a liquid

(Br2), and a solid (I2 and At)- have seven valence electrons

- tend to gain one electron to form a halide, X- ion, but also

share electrons and have positive oxidation states

- are reactive, with fluorine being the most reactive of all

nonmetals


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9. Noble Gases(Group 18)

- includes He, Ne, Ar, Kr, Xe, Rn

- not reactive

- have a full outer energy level- are all gases

- are all nonmetals


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