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Aaja mahi aaja mahi mp3 song download 320kbps. Predictperiodic trends including electronegativity, ionization energy, and therelative sizes of ions and atoms (1.1c)
- Electronegativity- How badly it want to 'hog' electrons
Study Guide for Periodic Table and Atomic Theory; Understanding the Periodic Table; Periodic Table Lesson; Chemical Bonds; Chemical Reactions Lesson; Redox Reactions Presentation; Electrolysis Pre-Aice Chemistry; Introduction to the Mole; Molar Calculations Lesson; Molar Volume and Percent Composition; Emprical and Molecular Formula.
- When two elements are joined in a chemical bond,the element that attracts the shared electrons more strongly ismore electronegative
- Elements with low electronegativities (themetallic elements) are said to be electropositive.
- It is important to understand thatelectronegativities are properties of atoms that are chemically bound to eachother; there is no way of measuring the electronegativity of an isolated atom.
- Moreover, the same atom can exhibit different electronegativitiesin different chemical environments, so the 'electronegativity of anelement' is only a general guide to its chemical behavior rather than anexact specification of its behavior in a particular compound.
- Nevertheless, electronegativity is eminentlyuseful in summarizing the chemical behavior of an element. You will make considerable use ofelectronegativity when you study chemical bonding and the chemistry of theindividual elements.
- Because there is no single definition ofelectronegativity, any numerical scale for measuring it must of necessity besomewhat arbitrary.
- Most such scales are themselves based on atomicproperties that are directly measurable and which relate in one way or theother to electron-attracting propensity.
- The most widely used of these scales was devisedby Linus Pauling and is related to ionization energy and electron affinity.
- The Pauling scale runs from 0 to 4; the highestelectron affinity, 4.0, is assigned to fluorine, while cesium has the lowestvalue of 0.7.
- Values less than about 2.2 are usuallyassociated with electropositive, or metallic character. In the representationof the scale shown in figure, the elements are arranged in rows correspondingto their locations in the periodic table.
- The correlation is obvious; electronegativity isassociated with the higher rows and the rightmost columns.
- The location of hydrogen on this scalereflects some of the significant chemical properties of this element.
- Although it acts like a metallic element in manyrespects (forming a positive ion, for example), it can also form hydride-ion (H–)solids with the more electropositive elements, and of course its ability toshare electrons with carbon and other p-block elements gives rise to avery rich chemistry, including of course the millions of organic compound.
- Ionization energy - Energy to remove a electron
- This term always refers to the formation of positive ions. In order to remove an electron from an atom, work must be done to overcome the electrostatic attraction between the electron and the nucleus; this work is called the ionization energy of the atom and corresponds to the exothermic process
M(g) → M+(g) + e–
in which M(g) stands for any isolated (gaseous) atom.
- An atom has as many ionization energies as it has electrons. Electrons are always removed from the highest-energy occupied orbital.
- An examination of the successive ionization energies of the first ten elements (below) provides experimental confirmation that the binding of the two innermost electrons (1s-orbital) is significantly different from that of the n=2 electrons.
- Successive ionization energies of an atom increase rapidly as reduced electron-electron repulsion causes the electron shells to contract, thus binding the electrons even more tightly to the nucleus.
- Ionization energies increase with the nuclear charge Z as we move across the periodic table. They decrease as we move down the table because in each period the electron is being removed from a shell one step farther from the nucleus than in the atom immediately above it. This results in the familiar zig-zag lines when the first ionization energies are plotted as a function of Z.
- Electron Affinity - How much they 'want' electrons
- Defined as the change in energy (in kJ/mole) of a neutral atom (in the gaseous phase) when an electron is added to the atom to form a negative ion. In other words, the neutral atom's likelihood of gaining an electron
- First Electron Affinity
- Ionization energies are always concerned with the formation of positive ions. Electron affinities are the negative ion equivalent, and their use is almost always confined to elements in groups 16 and 17 of the Periodic Table.
- The first electron affinity is the energy released when 1 mole of gaseous atoms each acquire an electron to form 1 mole of gaseous -1 ions. It is the energy released (per mole of X) when this change happens. First electron affinities have negative values
- Electron Affinity of Metals
- When an electron is added to a metal element, energy is needed to gain that electron (endothermic reaction).
- Metals have a less likely chance to gain electrons because it is easier to lose their valance electrons and form cations.
- It is easier to lose their valence electrons because metals' nuclei do not have a strong pull on their valence electrons.
- Thus, metals are known to have lower electron affinities.
- Electron Affinity of Nonmetals
- When nonmetals gain electrons, the energy change is usually negative because they give off energy to form an anion (exothermic process); thus, the electron affinity will be a negative number.
- Nonmetals have a greater electron affinity than metals because of their atomic structures: first, nonmetals have more valence electrons than metals do, thus it is easier for the nonmetals to gain electrons to fulfill a stable octet and
- secondly, the valence electron shell is closer to the nucleus, thus it is harder to remove an electron and it easier to attract electrons from other elements (especially metals).
- Thus, nonmetals have a higher electron affinity than metals, meaning they are more likely to gain electrons than atoms with a lower electron affinity.
- Nonmetals vs. Metals
- Metals:Metalslike to lose valence electrons to form cations to have a fully stable octet.
- They absorb energy (endothermic) to lose electrons. The electron affinity of metals is lower than that of nonmetals.
- Nonmetals: Nonmetals like to gain electrons to form anions to have a fully stable octet. They release energy (exothermic) to gain electrons to form an anion; thus, electron affinity of nonmetals is higher than that of metals.
- Patterns in Electron Affinity
- Electron affinity increases upward for the groups and from left to right across periods of a periodic table because the electrons added to energy levels become closer to the nucleus, thus a stronger attraction between the nucleus and its electrons.
- Remember that greater the distance, the less of an attraction; thus, less energy is released when an electron is added to the outside orbital.
- In addition, the more valence electrons an element has, the more likely it is to gain electrons to form a stable octet.
- The less valence electrons an atom has, the least likely it will gain electrons.
- Electron affinity decreases down the groups and from right to left across the periods on the periodic table because the electrons are placed in a higher energy level far from the nucleus, thus a decrease from its pull.
- However, one might think that since the number of valence electrons increase going down the group, the element should be more stable and have higher electron affinity. One fails to account for the shielding affect.
- As one goes down the period, the shielding effect increases, thus repulsion occurs between the electrons. This is why the attraction between the electron and the nucleus decreases as one goes down the group in the periodic table.
·Size of Atoms
- Note the Patterns
- Atomic radii decrease from left to rightacross a row and increase from top to bottom down a column.
- Electrons in the same principal shell are notvery effective at shielding one another from the nuclear charge, whereaselectrons in filled inner shells are highly effective at shielding electrons inouter shells from the nuclear charge.
- The increase in atomic sizegoing down a column is also due to electron shielding, but the situation ismore complex because the principal quantum numbernis notconstant.
- Periodic trends in atomic size
- We would expect the size of an atom to depend mainly on the principal quantum number of the highest occupied orbital; in other words, on the 'number of occupied electron shells'.
- Since each row in the periodic table corresponds to an increment in n, atomic radius increases as we move down a column.
- The other important factor is the nuclear charge; the higher the atomic number, the more strongly will the electrons be drawn toward the nucleus, and the smaller the atom.
- This effect is responsible for the contraction we observe as we move across the periodic table from left to right.
- Atomic radii decrease from left to rightacross a row and increase from top to bottom down a column.
- We would expect the size of an atom to depend mainly on the principal quantum number of the highest occupied orbital; in other words, on the 'number of occupied electron shells'.
Predictperiodic trends including electronegativity, ionization energy, and therelative sizes of ions and atoms (1.1c)
- Electronegativity- How badly it want to 'hog' electrons
- When two elements are joined in a chemical bond,the element that attracts the shared electrons more strongly ismore electronegative
- Elements with low electronegativities (themetallic elements) are said to be electropositive.
- It is important to understand thatelectronegativities are properties of atoms that are chemically bound to eachother; there is no way of measuring the electronegativity of an isolated atom.
- Moreover, the same atom can exhibit different electronegativitiesin different chemical environments, so the 'electronegativity of anelement' is only a general guide to its chemical behavior rather than anexact specification of its behavior in a particular compound.
- Nevertheless, electronegativity is eminentlyuseful in summarizing the chemical behavior of an element. You will make considerable use ofelectronegativity when you study chemical bonding and the chemistry of theindividual elements.
- Because there is no single definition ofelectronegativity, any numerical scale for measuring it must of necessity besomewhat arbitrary.
- Most such scales are themselves based on atomicproperties that are directly measurable and which relate in one way or theother to electron-attracting propensity.
- The most widely used of these scales was devisedby Linus Pauling and is related to ionization energy and electron affinity.
- The Pauling scale runs from 0 to 4; the highestelectron affinity, 4.0, is assigned to fluorine, while cesium has the lowestvalue of 0.7.
- Values less than about 2.2 are usuallyassociated with electropositive, or metallic character. In the representationof the scale shown in figure, the elements are arranged in rows correspondingto their locations in the periodic table.
- The correlation is obvious; electronegativity isassociated with the higher rows and the rightmost columns.
- The location of hydrogen on this scalereflects some of the significant chemical properties of this element.
- Although it acts like a metallic element in manyrespects (forming a positive ion, for example), it can also form hydride-ion (H–)solids with the more electropositive elements, and of course its ability toshare electrons with carbon and other p-block elements gives rise to avery rich chemistry, including of course the millions of organic compound.
- Ionization energy - Energy to remove a electron
- This term always refers to the formation of positive ions. In order to remove an electron from an atom, work must be done to overcome the electrostatic attraction between the electron and the nucleus; this work is called the ionization energy of the atom and corresponds to the exothermic process
M(g) → M+(g) + e–
in which M(g) stands for any isolated (gaseous) atom.
- An atom has as many ionization energies as it has electrons. Electrons are always removed from the highest-energy occupied orbital.
- An examination of the successive ionization energies of the first ten elements (below) provides experimental confirmation that the binding of the two innermost electrons (1s-orbital) is significantly different from that of the n=2 electrons.
- Successive ionization energies of an atom increase rapidly as reduced electron-electron repulsion causes the electron shells to contract, thus binding the electrons even more tightly to the nucleus.
- Ionization energies increase with the nuclear charge Z as we move across the periodic table. They decrease as we move down the table because in each period the electron is being removed from a shell one step farther from the nucleus than in the atom immediately above it. This results in the familiar zig-zag lines when the first ionization energies are plotted as a function of Z.
- Electron Affinity - How much they 'want' electrons
- Defined as the change in energy (in kJ/mole) of a neutral atom (in the gaseous phase) when an electron is added to the atom to form a negative ion. In other words, the neutral atom's likelihood of gaining an electron
- First Electron Affinity
- Ionization energies are always concerned with the formation of positive ions. Electron affinities are the negative ion equivalent, and their use is almost always confined to elements in groups 16 and 17 of the Periodic Table.
- The first electron affinity is the energy released when 1 mole of gaseous atoms each acquire an electron to form 1 mole of gaseous -1 ions. It is the energy released (per mole of X) when this change happens. First electron affinities have negative values
- Electron Affinity of Metals
- When an electron is added to a metal element, energy is needed to gain that electron (endothermic reaction).
- Metals have a less likely chance to gain electrons because it is easier to lose their valance electrons and form cations.
- It is easier to lose their valence electrons because metals' nuclei do not have a strong pull on their valence electrons.
- Thus, metals are known to have lower electron affinities.
- Electron Affinity of Nonmetals
- When nonmetals gain electrons, the energy change is usually negative because they give off energy to form an anion (exothermic process); thus, the electron affinity will be a negative number.
- Nonmetals have a greater electron affinity than metals because of their atomic structures: first, nonmetals have more valence electrons than metals do, thus it is easier for the nonmetals to gain electrons to fulfill a stable octet and
- secondly, the valence electron shell is closer to the nucleus, thus it is harder to remove an electron and it easier to attract electrons from other elements (especially metals).
- Thus, nonmetals have a higher electron affinity than metals, meaning they are more likely to gain electrons than atoms with a lower electron affinity.
- Nonmetals vs. Metals
- Metals:Metalslike to lose valence electrons to form cations to have a fully stable octet.
- They absorb energy (endothermic) to lose electrons. The electron affinity of metals is lower than that of nonmetals.
- Nonmetals: Nonmetals like to gain electrons to form anions to have a fully stable octet. They release energy (exothermic) to gain electrons to form an anion; thus, electron affinity of nonmetals is higher than that of metals.
- Patterns in Electron Affinity
- Electron affinity increases upward for the groups and from left to right across periods of a periodic table because the electrons added to energy levels become closer to the nucleus, thus a stronger attraction between the nucleus and its electrons.
- Remember that greater the distance, the less of an attraction; thus, less energy is released when an electron is added to the outside orbital.
- In addition, the more valence electrons an element has, the more likely it is to gain electrons to form a stable octet.
- The less valence electrons an atom has, the least likely it will gain electrons.
- Electron affinity decreases down the groups and from right to left across the periods on the periodic table because the electrons are placed in a higher energy level far from the nucleus, thus a decrease from its pull.
- However, one might think that since the number of valence electrons increase going down the group, the element should be more stable and have higher electron affinity. One fails to account for the shielding affect.
- As one goes down the period, the shielding effect increases, thus repulsion occurs between the electrons. This is why the attraction between the electron and the nucleus decreases as one goes down the group in the periodic table.
·Size of Atoms
- Note the Patterns
- Atomic radii decrease from left to rightacross a row and increase from top to bottom down a column.
- Electrons in the same principal shell are notvery effective at shielding one another from the nuclear charge, whereaselectrons in filled inner shells are highly effective at shielding electrons inouter shells from the nuclear charge.
- The increase in atomic sizegoing down a column is also due to electron shielding, but the situation ismore complex because the principal quantum numbernis notconstant.
- Periodic trends in atomic size
- We would expect the size of an atom to depend mainly on the principal quantum number of the highest occupied orbital; in other words, on the 'number of occupied electron shells'.
- Since each row in the periodic table corresponds to an increment in n, atomic radius increases as we move down a column.
- The other important factor is the nuclear charge; the higher the atomic number, the more strongly will the electrons be drawn toward the nucleus, and the smaller the atom.
- This effect is responsible for the contraction we observe as we move across the periodic table from left to right.
- Atomic radii decrease from left to rightacross a row and increase from top to bottom down a column.
- We would expect the size of an atom to depend mainly on the principal quantum number of the highest occupied orbital; in other words, on the 'number of occupied electron shells'.
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