1.5+Periodic+properties+of+elements

PERIODIC PROPERTIES OF ELEMENTS ​The properties of the elements exhibit trends and these trends can be predicted with the help of the periodic table. They can also be explained and understood by analyzing the electron configurations of the elements. This is because, elements tend to gain or lose valence electrons to achieve the stable octet formation. 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. And, as this happens, the electrons of the outermost shell experience increasingly strong nuclear attraction. As a result, the electrons become closer to the nucleus and more tightly bound. The second trend is the moving down a column in the periodic table, where the outermost electrons become less tightly bound to the nucleus. And these trends explain the periodicity observed in the elemental properties of atomic radius, ionization energy, electron affinity, and electronegativity. But, before going into that we need to know a bit more about the above mentioned terms: The atomic radius of an element is half of the distance between the centers of two atoms of an element that are in contact with each other. Generally, the atomic radius decreases across a period, from left to right and increases down a given group. Therefore, the atoms with the largest atomic radii are located in Group I and at the lower half of groups. Ionization energy or ionization potential is the energy required to completely remove an electron from a gaseous atom or ion. And, the closer and more tightly an electron is bound to the nucleus, the more difficult it is to remove and the higher its ionization energy. Ionization energy is also required to remove a second valence electron from the univalent ion to form the divalent ion, and so on. Electron affinity is the energy change that occurs when an electron is added to a gaseous atom. It reflects the ability of an atom to accept an electron. And the atoms with stronger effective nuclear charge have a greater electron affinity. Therefore, some generalizations can be made about the electron affinities of certain groups in the periodic table. The alkaline earths have low electron affinity values. This is because they have filled sub shells. But, the halogens have high electron affinities because of the addition of an electron to an atom results in a completely filled shell. Noble gases have zero electron affinities, since each atom possesses a stable octet and will not accept an electron readily. An atom with higher electro negativity has a great capacity for attracting bonding electrons. Therefore, electro negativity is a measure of the attraction of an atom for the electrons in a [|chemical bond]. It’s related to ionization energy. So, electrons with low ionization energies have low electro negativities because their nuclei do not exert a strong attractive force on electrons. And, elements with high ionization energies have high electro negativities. This is because of the strong pull exerted on electrons by the nucleus. Therefore, electro negativity is dependant on the atomic number. As the atomic number increases, the electro negativity decreases, as a result of increased distance between the valence electron and nucleus. An example of an electropositive element, i.e. one with low electro negativity, is cesium. And an example of a highly electronegative element is fluorin e
 * Atomic Radius **
 * Ionization Energy **
 * Electron Affinity **
 * Electro negativity **

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= Periodic Properties of the 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.

From: http://chemistry.about.com/od/periodictableelements/a/periodictrends.htm By: Julian Jimenez Aguilar

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]. 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. 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 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// sub shells. 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 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.
 * Atomic Radius **
 * Ionization Energy **
 * Electron Affinity **
 * Electronegativity **

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= Atomic Properties = The electrons associated with atoms are found to have measurable properties which exhibit quantization. The electrons are normally found in quantized energy states of the lowest possible energy for the atom, called ground states. The electrons can also exist in higher "excited states", as evidenced by the line spectra (e.g. the hydrogen spectrum) observed when they make transitions back to the ground states. The existence of these excited states can be demonstrated more directly in collision experiments like the Franck-Hertz experiment. Other properties associated with the electron energy levels such as orbital angular momentum and electron spin are also quantized and give rise to the quantum numbers used to characterize the levels. These quantized properties are associated with periodic table of the elements, and the requirements of the Pauli exclusion principle on the quantum numbers can be viewed as the origin of the periodicity. The periodic table provides a convenient framework for cataloging other physical and chemical properties of atoms. While the hydrogen electron energy levels are found to depend only upon the principal quantum number, the energy levels in other atoms are found to have strong dependence upon the orbital quantum number. These levels show a smaller amount of dependence upon the total angular momentum. This dependence may arise from interactions within the atom such as the spin-orbit interaction or may arise only when external fields are applied. When magnetic fields are applied, there is splitting of atomic energy levels from the Zeeman effect, and in response to electric fields there is splitting called the Stark effect. []

the periodic table of elements  The **periodic table of the chemical elements** (also **Mendeleev's table**, **periodic table of the elements** or just **periodic table**) is a tabular display of the chemical elements. Although precursors to this table exist, its invention is generally credited to Russian chemist Dmitri Mendeleev in 1869, who intended the table to illustrate recurring ("periodic") trends in the properties of the elements. The layout of the table has been refined and extended over time, as new elements have been discovered, and new theoretical models have been developed to explain chemical behavior. 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]. 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. 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 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// sub shells. 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 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.
 * Atomic Radius **
 * Ionization Energy **
 * Electron Affinity **
 * Electronegativity **

[]



= Atomic Properties = The electrons associated with atoms are found to have measurable properties which exhibit quantization. The electrons are normally found in quantized energy states of the lowest possible energy for the atom, called ground states. The electrons can also exist in higher "excited states", as evidenced by the line spectra (e.g. the hydrogen spectrum) observed when they make transitions back to the ground states. The existence of these excited states can be demonstrated more directly in collision experiments like the Franck-Hertz experiment. Other properties associated with the electron energy levels such as orbital angular momentum and electron spin are also quantized and give rise to the quantum numbers used to characterize the levels. These quantized properties are associated with periodic table of the elements, and the requirements of the Pauli exclusion principle on the quantum numbers can be viewed as the origin of the periodicity. The periodic table provides a convenient framework for cataloging other physical and chemical properties of atoms. While the hydrogen electron energy levels are found to depend only upon the principal quantum number, the energy levels in other atoms are found to have strong dependence upon the orbital quantum number. These levels show a smaller amount of dependence upon the total angular momentum. This dependence may arise from interactions within the atom such as the spin-orbit interaction or may arise only when external fields are applied. When magnetic fields are applied, there is splitting of atomic energy levels from the Zeeman effect, and in response to electric fields there is splitting called the Stark effect. []

the periodic table of elements  The **periodic table of the chemical elements** (also **Mendeleev's table**, **periodic table of the elements** or just **periodic table**) is a tabular display of the chemical elements. Although precursors to this table exist, its invention is generally credited to Russian chemist Dmitri Mendeleev in 1869, who intended the table to illustrate recurring ("periodic") trends in the properties of the elements. The layout of the table has been refined and extended over time, as new elements have been discovered, and new theoretical models have been developed to explain chemical behavior.

by julian peñarredonda #19 7b what are the periodic properties ? The main value of the periodic table is the ability to predict the chemical properties of an element based on its location on the table. It should be noted that the properties vary differently when moving vertically along the columns of the table than when moving horizontally along the rows.

Trends of groups
Modern quantum mechanical theories of atomic structure explain group trends by proposing that elements within the same group have the same electron configurations in their valence shell, which is the most important factor in accounting for their similar properties. Elements in the same group also show patterns in their atomic radius, ionization energy, and electronegativity. From top to bottom in a group, the atomic radii of the elements increase. Since there are more filled energy levels, valence electrons are found farther from the nucleus. From the top, each successive element has a lower ionization energy because it is easier to remove an electron since the atoms are less tightly bound. Similarly, a group will also see a top to bottom decrease in electronegativity due to an increasing distance between valence electrons and the nucleus

Trends of periods
Elements in the same period show trends in atomic radius, ionization energy, electron affinity, and electronegativity. Moving left to right across a period, atomic radius usually decreases. This occurs because each successive element has an added proton and electron which causes the electron to be drawn closer to the nucleus. This decrease in atomic radius also causes the ionization energy to increase when moving from left to right across a period. The more tightly bound an element is, the more energy is required to remove an electron. Similarly, electronegativity will increase in the same manner as ionization energy because of the amount of pull that is exerted on the electrons by the nucleus. Electron affinity also shows a slight trend across a period. Metals (left side of a period) generally have a lower electron affinity than nonmetals (right side of a period) with the exception of the noble gases.

taken from : [] pedro juan mejia aguilar



<span style="color: #ff0000; font-family: 'Comic Sans MS',cursive; font-size: 120%;">periodic properties of elements

by julian peñarredonda #19