what is the periodic properties of elements?

The properties of the elements exhibit trends. These trends can be predicted using the periodic table and can be explained and understood by analyzing the electron configurations of the elements. Elements tend to gain or lose valence electrons to achieve stable octet formation. Stable octets are seen in the inert gases, or noble gases, of Group VIII of the periodic table. In addition to this activity, there are two other important trends. First, electrons are added one at a time moving from left to right across a period. As this happens, the electrons of the outermost shell experience increasingly strong nuclear attraction, so the electrons become closer to the nucleus and more tightly bound to it. Second, moving down a column in the periodic table, the outermost electrons become less tightly bound to the nucleus. This happens because the number of filled principal energy levels (which shield the outermost electrons from attraction to the nucleus) increases downward within each group. These trends explain the periodicity observed in the elemental properties of atomic radius, ionization energy, electron affinity, and electronegativity.

Atomic Radius
The atomic radius of an element is half of the distance between the centers of two atoms of that element that are just touching 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.

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 to form 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 a group (increasing atomic radius). Group I elements have low ionization energies because the loss of an electron forms a stable octet.


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 have filled 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 an electron readily. Elements of other groups have low electron affinities.

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 pull exerted 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.

taken from:http://chemistry.about.com/od/periodictableelements/a/periodictrends.htm





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Properties of elements


The properties of the elements exhibit trends. These trends can be predicted using the periodic table and can be explained and understood by analyzing the electron configurations of the elements. Elements tend to gain or lose valence electrons to achieve stable octet formation. Stable octets are seen in the inert gases, or noble gases, of Group VIII of the periodic table.
In addition to this activity, there are two other important trends. First, electrons are added one at a time moving from left to right across a period. As this happens, the electrons of the outermost shell experience increasingly strong nuclear attraction, so the electrons become closer to the nucleus and more tightly bound to it. Second, moving down a column in the periodic table, the outermost electrons become less tightly bound to the nucleus. This happens because the number of filled principal energy levels (which shield the outermost electrons from attraction to the nucleus) increases downward within each group. These trends explain the periodicity observed in the elemental properties of atomic radius, ionization energy, electron affinity, and electronegative periodic Properties of elementsperiodic Properties of elements

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Atomic Radius
The atomic radius of an element is half of the distance between the centers of two atoms of that element that are just touching 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, but the 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.periodic Properties of elements

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 to form 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 a group (increasing atomic radius). Group I elements have low ionization energies because the loss of an electron forms a stable octet.
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 have filled 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 an electron readily. Elements of other groups have low electron affinities.

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 pull exerted 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.periodic Properties of elementsperiodic Properties of elements

Summary of Periodic Table Trends



Moving Left --> Right
  • Atomic Radius Decreases
  • Ionization Energy Increases
  • Electronegativity Increases
Moving Top --> Bottom
  • Atomic Radius Increases
  • Ionization Energy Decreases
  • Electronegativity Decreases

taken of http://chemistry.about.com/od/periodictableelements/a/periodictrends.htm
organizaton of the periodic table
To understand how the periodic table is organized, imagine that we write down a long horizontal list of the elements in order of their increasing atomic number. It would begin this way:
H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca...
Now if we look at the various physical and chemical properties of these elements, we would find that their values tend to increase or decrease with Z in a manner that reveals a repeating pattern— that is, a periodicity. For the elements listed above, these breaks can be indicated by the vertical bars shown here in color:
H He | Li Be B C N O F Ne | Na Mg Al Si P S Cl Ar | Ca ...
Periods.

To construct the table, we place each sequence in a separate row, which we call a period. The rows are aligned in such a way that the elements in each vertical column possess certain similarities. Thus the first short-period elements H and He are chemically similar to the elements Li and Ne at the beginning and end of the second period. The first period is split in order to place H above Li and He above Ne.
The "block" nomenclature shown above refers to the sub-orbital type (quantum number l, or s-p-d-f classification) of the highest-energy orbitals that are occupied in a given element. For n=1 there is no p block, and the s block is split so that helium is placed in the same group as the other inert gases, which it resembles chemically. For the second period (n=2) there is a p block but no d block; in the usual "long form" of the periodic table it is customary to leave a gap between these two blocks in order to accommodate the d blocks that occur at n=3 and above. At n=4 we introduce an f block, but in order to hold the table to reasonable dimensions the f blocks are placed below the main body of the table.

Groups.
Each column of the periodic table is known as a group. The elements belonging to a given group bear a strong similarity in their chemical behaviors.
In the past, two different systems of Roman numerals and letters were used to denote the various groups. North Americans added the letter B to denote the d-block groups and A for the others; this is the system shown in the table above. The the rest of the world used A for the d-block elements and B for the others. In 1985, a new international system was adopted in which the columns were simply labeled 1-18. Although this system has met sufficient resistance in North America to slow its incorporation into textbooks, it seems likely that the "one to eighteen" system will gradually take over as older professors (the main hold-outs!) retire.

Families.
Chemists have long found it convenient to refer to the elements of different groups, and in some cases of spans of groups by the names indicated in the table shown below. The two of these that are most important for you to know are the noble gases and the transition metals.

periodic table element families
periodic table element families



taken from:http://www.chem1.com/acad/webtext/atoms/atpt-6.Subcategories

This category has the following 3 subcategories, out of 3 total:
Allotropy(3 C, 11 P)
Isotopes (124 C, 23 P)
Biology and pharmacology of chemical elemnts (1 C, 44 P)



The atomic number does not determine the number of neutrons in an atomic core. As a result, the number of neutrons within an atom can vary. Then atoms that have the same atomic number may differ in atomic mass. Atoms of the same element that differ in atomic mass are called isotopes.
Mainly with the heavier atoms that have a higher atomic number, the number of neutrons within the core may exceed the number of protons.
Isotopes of the same element are often found in nature alternately or in mixtures.
An example: chlorine has an atomic number of 17, which basically means that all chlorine atoms contain 17 protons within their core. There are two isotopes. Three-quarters of the chlorine atoms found in nature contain 18 neutrons and one quarter contains 20 neutrons. The mass numbers of these isotopes are 17 + 18 = 35 and 17 + 20 = 37. The isotopes are written as follows: 35Cl and 37Cl.
When isotopes are noted this way the number of protons and neutrons does not have to be mentioned separately, because the symbol of chlorine .
within the periodic chart (Cl) is set on the seventeenth place. This already indicates the number of protons, so that one can always calculate the number of neutrons easily by means of the mass number.

A great number of isotopes is not stable. They will fall apart during radioactive decay processes. Isotopes that are radioactive are called radioisotopes.


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taken from: http://www.lenntech.com/periodic/properties/chemical-properties.htm#ixzz0iSl4A










The atomic number indicates the number of protons within the core of an atom. The atomic number is an important concept of chemistry and quantum mechanics. An element and its place within the periodic table are derived from this concept.
When an atom is generally electrically neutral, the atomic number will equal the number of electrons in the atom, which can be found around the core. These electrons mainly determine the chemical behaviour of an atom. Atoms that carry electric charges are called ions. Ions either have a number of electrons larger (negatively charged) or smaller (positively charged) than the atomic number.


The name indicates the mass of an atom, expressed in atomic mass units (amu). Most of the mass of an atom is concentrated in the protons and neutrons contained in the nucleus. Each proton or neutron weighs about 1 amu, and thus the atomic mass in always very close to the mass (or nucleon) number, which indicates the number of particles within the core of an atom; this means the protons and neutrons. Each isotope of a chemical element can vary in mass. The atomic mass of an isotope indicates the number of neutrons that are present within the core of the atoms. The total atomic mass of an element is an equivalent of the mass units of its isotopes. The relative occurrence of the isotopes in nature is an important factor in the determination of the overall atomic mass of an element. In reference to a certain chemical element, the atomic mass as shown in the periodic table is the average atomic mass of all the chemical element's stable isotopes. The average is weighted by the relative natural abundances of the element's isotopes.


Electro negativity measures the inclination of an atom to pull the electronic cloud in its direction during chemical bonding with another atom.
Pauling's scale is a widely used method to order chemical elements according to their electro negativity. Nobel prize winner Linus Pauling developed this scale in 1932.
The values of electro negativity are not calculated, based on mathematical formula or a measurement. It is more like a pragmatic range.
Pauling gave the element with the highest possible electro negativity, fluorine, a value of 4,0. Francium, the element with the lowest possible electro negativity, was given a value of 0,7. All of the remaining elements are given a value of somewhere between these two extremes.


The density of an element indicates the number of units of mass of the element that are present in a certain volume of a medium. Traditionally, density is expressed through the Greek letter ro (written as r).Within the SI system of units density is expressed in kilograms per cubic meter (kg/m3). The density of an element is usually expressed graphically with temperatures and air pressures, because these two properties influence density.


The melting point of an element or compound means the temperatures at which the solid form of the element or compound is at equilibrium with the liquid form. We usually presume the air pressure to be 1 atmosphere.
For example: the melting point of water is 0 oC, or 273 K.


The boiling point of an element or compound means the temperature at which the liquid form of an element or compound is at equilibrium with the gaseous form. We usually presume the air pressure to be 1 atmosphere.
For example: the boiling point of water is 100 oC, or 373 K.
At the boiling point the vapour pressure of an element or compound is 1 atmosphere.


Even when two atoms that are near one another will not bind, they will still attract one another. This phenomenon is known as the Vanderwaals interaction.
The Vanderwaals forces cause a force between the two atoms. This force becomes stronger, as the atoms come closer together. However, when the two atoms draw too near each other a rejecting force will take action, as a consequence of the exceeding rejection between the negatively charged electrons of both atoms. As a result, a certain distance will develop between the two atoms, which is commonly known as the Vanderwaals radius.
Through comparison of Vanderwaals radiuses of several different pairs of atoms, we have developed a system of Vanderwaals radiuses, through which we can predict the Vanderwaals radius between two atoms, through addition.


Ionic radius is the radius that an ion has in an ionic crystal, where the ions are packed together to a point where their outermost electronic orbitals are in contact with each other. An orbital is the area around an atom where, according to orbital theory, the probability of finding an electron is the greatest.


The atomic number does not determine the number of neutrons in an atomic core. As a result, the number of neutrons within an atom can vary. Then atoms that have the same atomic number may differ in atomic mass. Atoms of the same element that differ in atomic mass are called isotopes.
Mainly with the heavier atoms that have a higher atomic number, the number of neutrons within the core may exceed the number of protons.
Isotopes of the same element are often found in nature alternately or in mixtures.
An example: chlorine has an atomic number of 17, which basically means that all chlorine atoms contain 17 protons within their core. There are two isotopes. Three-quarters of the chlorine atoms found in nature contain 18 neutrons and one quarter contains 20 neutrons. The mass numbers of these isotopes are 17 + 18 = 35 and 17 + 20 = 37. The isotopes are written as follows: 35Cl and 37Cl.
When isotopes are noted this way the number of protons and neutrons does not have to be mentioned separately, because the symbol of chlorine within the periodic chart (Cl) is set on the seventeenth place. This already indicates the number of protons, so that one can always calculate the number of neutrons easily by means of the mass number.

A great number of isotopes is not stable. They will fall apart during radioactive decay processes. Isotopes that are radioactive are called radioisotopes.


The electronic configuration of an atom is a description of the arrangement of electrons in circles around the core. These circles are not exactly round; they contain a wave-like pattern. For each circle the probability of an electron to be present on a certain location is described by a mathematic formula. Each one of the circles has a certain level of energy, compared to the core. Commonly the energy levels of electrons are higher when they are further away from the core, but because of their charges, electrons can also influence each another's energy levels. Usually the middle circles are filled up first, but there may be exceptions due to rejections.
The circles are divided up in shells and sub shells, which can be numbered by means of quantities.


The ionisation energy means the energy that is required to make a free atom or molecule lose an electron in a vacuum. In other words; the energy of ionisation is a measure for the strength of electron bonds to molecules. This concerns only the electrons in the outer circle.


Besides the energy of the first ionisation, which indicates how difficult it is to remove the first electron
from an atom, there is also an energy measure for second ionisation. This energy of second ionisation indicates the degree of difficulty to remove the second atom.

As such, there is also the energy of a third ionisation, and sometimes even the energy of a fourth or fifth ionisation.


The standard potential means the potential of a redox reaction, when it is at equilibrium, in relation to zero. When the standard potential exceeds zero, we are dealing with an oxidation reaction. When the standard potential is below zero, we are dealing with a reduction reaction. The standard potential of electrons is expressed in volt (V), by the symbol V0.

Read more: http://www.lenntech.com/periodic/properties/chemical-properties.htm#ixzz0j7zcVJTY







What are the properties of metal elements in the periodic table?
Metals
Metals makeup more than 75% of the elements in the periodic table. Metals are characterized by the following physical properties.
1. They have metallic shine or luster.
2. They are usually solids at room temperature.
3. They are malleable. Malleable means that metals can be hammered, pounded, or pressed into different shapes without breaking.
4. They are ductile meaning that they can be drawn into thin sheets or wires without breaking. 5. They are good conductors of heat and electricity.
Nonmetals
There are 17 nonmetals in the periodic table, and they are characterized by four major physical properties.
1. They rarely have metallic luster.
2. They are usually gases at room temperature.
3. Nonmetallic solids are neither malleable nor ductile.
4. They are poor conductors of heat and electricity.
Metalloids
The six metalloids are B, Si, Ge, As, Sb, and Te. The properties of the metalloids have characteristics in between that of the metals and the nonmetals. They are good conductors of heat and electricity, but they are not good conductors or insulators.
Periodic Properties
The periodic table also has certain properties characteristic of certain regions in the periodic table.
Alkali Metals
These are the metals in the first column of the periodic table. They are soft shiny metals that usually combine with group VIIA nonmetals in chemical compounds in a 1:1 ratio.
Alkaline Earth Metals
These are the elements in the second column of the periodic table, and they are very similar to the alkali metals. They combine with the group VIIA nonmetals in a 1:2 ratio.
Halogens
The halogens are found in group VIIA. They are fluorine, chlorine, bromine, and iodine. The halogens exist as diatomic molecules in nature.
Noble Gases
The noble gases are also called rare gas elements, and they all occur in nature as gases. The noble gases make up the group VIIIA which is the last column in the periodic table. The noble gases fulfill the octet rule by having a full outer level with 8 valence electrons. Therefore, they do not undergo chemical reactions because they do not accept any electrons.
Transition Metals
The transition metals are the metals located between columns IIA and IIIA in the periodic table. The elements also have valence electrons in two shells instead of one.
taken from: wiki/properties of elements .com

Periodic Properties of elements

Atomic weight:The atomic weight is the number assigned to each chemical element to specify the average mass of atoms.
Melting point : The melting point is the temperature at which matter changes from solid to liquid is melting
Boiling point: is the temperature at which matter changes from liquid to gas
Density: is a quantity refers to the amount of mass contained in a given volume.
Electronegativity: is the ability of an atom to attract electrons to it, or electron density, it forms a covalent bond in a molecule
Ionization potential: is the energy required to remove the outermost electron of an atom in a gaseous
Atomic number: is the positive integer that equals the total number of protons in the nucleus of the atom
The oxidation state: is part of a compound, is regarded as the apparent burden under which the element is working in the compound. The oxidation states may be positive, negative, zero, integers and fractions.
Electronic configuration: is the manner in which electrons in an atom structure

this properties are show in this link (open de link)


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taken from:
http://es.wikipedia.org/wiki/Configuración_electrónica http://es.wikipedia.org/wiki/Estado_de_oxidación http://es.wikipedia.org/wiki/Número_atómico http://es.wikipedia.org/wiki/Energía_de_ionización http://es.wikipedia.org/wiki/Electronegatividad http://es.wikipedia.org/wiki/Densidad http://es.wikipedia.org/wiki/Punto_de_ebullición
http://es.wikipedia.org/wiki/Punto_de_fusión http://es.wikipedia.org/wiki/Peso_atómico