Periodic trends by Pramod Kothari

Periodic Trends Lulu Press, Raleigh, N.C. USA.

Dr. Pramod Kothari Assistant Professor, Department Of Chemistry Government Post Graduate College, Berinag, District â&#x20AC;&#x201C; Pithoragarh Uttarakhand (India)

Copyright ÂŠ Creative Commons Attribution-Share Alike 3.0 //creativecommons.org/licenses/by-sa/3.0/ Disclaimer All the material contained in this book is provided for educational and informational purposes only. No responsibility can be taken for any results or outcomes resulting from the use of this material. While every attempt has been made to provide information that is both accurate and effective, the author does not assume any responsibility for the accuracy or use/misuse of this information.

Dr. Pramod Kothari / Periodic trends

Table of Contents Periodic trends Atomic radius

........................................................................................... 1

………………………………………………………………………………………………………………………..4

Ionization energy…………………………………………………………………………………...10 Electron affinity

.........................................................................................17

Electronegativity………………………..

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.............................................................22

iii

Chapter 1: Periodic trends

The Periodic Trends In chemistry, periodic trends are the tendencies of certain elemental characteristics to increase or decrease as one progresses along a row or column of the periodic table of elements.

All periodic trends of the chemicals are based on Coulomb's law . As distance from the protons in the nucleus to the valence electrons increases values associated with attributes such as electron affinity, ionization energy, and electronegativity decrease. Atomic radius The atomic radius is the distance from the atomic nucleus to the outermost stable electron orbital in an atom that is at equilibrium. The atomic radii tends to decrease across a period from left to right. The atomic radius usually increases while going down a group due to the addition of a new energy level (shell). However, diagonally, the number of electrons has a larger effect than the sizeable radius. For example, lithium (145 picometer) has a smaller [citation needed] atomic radius than magnesium (150 picometer). Atomic radius decreases from left to right across a period, and also increases from top to bottom down a group. Atomic radius can be further specified as: 

Covalent radius: half the distance between two atoms of a diatomic compound, singly bonded.



Van der Waals radius: half the distance between the nuclei of atoms of different molecules in a lattice of covalent molecules.



Metallic radius: half the distance between two adjacent nuclei of atoms in a metallic lattice.



Ionic radius: half the distance between two nuclei

Ionization energy The ionization potential is the minimum amount of energy required to remove one electron from each atom in a mole of atoms in the gaseous state. The first ionization energy is the

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energy required to remove one, the nth ionization energy is the energy required to remove the atom's nth electron, after the (nâ&#x2C6;&#x2019;1) electrons before it have been removed. Trend-wise, ionization energy tends to increase while one progresses across a period because the greater number of protons (higher nuclear charge) attract the orbiting electrons more strongly, thereby increasing the energy required to remove one of the electrons. Ionization [citation needed] energy and ionization potentials are completely different. The potential is an intensive property and it is measured by "volt" ; whereas the energy is an extensive property expressed by "eV" or "kJ/mole". As one progresses down a group on the periodic table, the ionization energy will likely decrease since the valence electrons are farther away from the nucleus and experience a weaker attraction to the nucleus's positive charge. There will be an increase of ionization energy from left to right of a given period and a decrease from top to bottom. As a rule, it requires far less energy to remove an outer-shell electron than an inner-shell electron. As a result the ionization energies for a given element will increase steadily within a given shell, and when starting on the next shell down will show a drastic jump in ionization energy. Simply put, the lower the principal quantum number, the higher the ionization energy for the electrons within that shell. The exceptions are the elements in the boron and oxygen family, which require slightly less energy than the general trend. Helium has the highest ionization energy while Francium has the lowest. Electron affinity The electron affinity of an atom can be described either as the energy gained by an atom when an electron is added to it, or conversely as the energy required to detach an electron from a singly charged anion. The sign of the electron affinity can be quite confusing, as atoms that become more stable with the addition of an electron (and so are considered to have a higher electron affinity) show a decrease in potential energy; i.e. the energy gained by the atom appears to be negative. For atoms that become less stable upon gaining an electron, potential energy increases, which implies that the atom gains energy. In such a case, the atom's electron affinity value is positive. Consequently, atoms with a more negative electron affinity value are considered to have a lower electron affinity (they are more receptive to gaining electrons), and vice versa. However in the reverse scenario where electron affinity is defined as the energy required to detach an electron from an anion, the energy value obtained will be of the same magnitude but have the opposite sign. This is because those atoms with a high electron affinity are less inclined to give up an electron, and so take more energy to remove the electron from the atom. In this case, the atom with the more positive energy value has the higher electron affinity. As one progresses from left to right across a period, the electron affinity will increase. Although it may seem that Fluorine should have the greatest electron affinity, the small size of fluorine generates enough repulsion that Chlorine has the greatest electron affinity. Electronegativity Electronegativity is a measure of the ability of an atom or molecule to attract pairs of electrons in the context of a chemical bond. The type of bond formed is largely determined by the difference in electronegativity between the atoms involved, using the Pauling scale. Trend-wise, as one moves from left to right across a period in the periodic table, the electronegativity increases due to the stronger attraction that the atoms obtain as the nuclear charge increases. Moving down in a group, the electronegativity decreases due to the longer distance between the nucleus and the valence electron shell, thereby decreasing the attraction, making the atom have less of an attraction for electrons or protons.

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In the group 13 elements electronegativity increases from aluminium to thallium. In group 14 electronegativity of lead is higher than that of tin. Valence Electrons Valence electrons are the electrons in the outermost electron shell of an isolated atom of an element. Sometimes, it is also regarded as the basis of ModernPeriodic Table. In a period, the number of valence electrons increases(mostly for light metal/elements)as we move from left to right side. However, in a group this periodic trend is constant, that is the number of valence electrons remains the same. However, this periodic trend is sparsely followed for heavier elements (elements with atomic number greater than 20), especially for lanthanide and actinide series. Summary For atomic number vs. atomic radius, it has a trend of four peaks that quickly decline in value because of an increase in the nuclear charge and also an increase in the number of electrons in the same principal energy level. Increasing the quantity of charge attracts the electrons closer to the nucleus. For atomic number vs. melting point, there is almost a straight line of an upward slope because the number of ionic bonds increase with the atomic number; therefore, requiring a higher temperature. For atomic number vs. ionization energy, it increases as you move from the alkali metal to the noble gas because when the nuclear charge increase and the atomic radius decrease, the increased attraction makes it more difficult to remove an electron. Finally, for atomic number vs. electronegativity, the trend is that it increases as it goes from left to right and decreases as it goes from top to bottom because it depends on the number of atoms. The electronegativity levels peak from 0 to a certain height, then back down to 0 because the electron levels fluctuate up and down between atomic numbers. Moving Left → Right • Atomic Radius Decreases • Ionization Energy Increases • Electronegativity Increases Moving Top → Bottom • Atomic Radius Increases • Ionization Energy Decreases • Electronegativity Decreases Metallic properties Metallic properties increases down the group as the size of an atom increases which leads to lesser attraction between the nuclei and the electrons thus the outer most electrons of a metallic atom is loosely bound which makes metals a good conductor of heat and electricity. As we move across the period we notice that the size of the atom decreases which means the attraction between the nuclei and the electrons increases so the metallic character decreases. Non-metallic properties Non-metallic property increases across a period and decreases down the group due to the same reason. http://www.jstage.jst.go.jp/article/jlve/33/2/33_67/_article

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Chapter 2: Atomic radius

Diagram of a helium atom, showing the electron probability density as shades of gray. The atomic radius of a chemical element is a measure of the size of its atoms, usually the mean or typical distance from the nucleus to the boundary of the surrounding cloud of electrons. Since the boundary is not a well-defined physical entity, there are various nonequivalent definitions of atomic radius. Three widely used definitions of atomic radius are Van der Waals radius, ionic radius, and covalent radius. Depending on the definition, the term may apply only to isolated atoms, or also to atoms in condensed matter, covalently bound in molecules, or in ionized and excited states; and its value may be obtained through experimental measurements, or computed from theoretical models. Under some definitions, the value of the radius may depend on the atom's state and context. Electrons do not have definite orbits, or sharply defined ranges. Rather, their positions must be described as probability distributions that taper off gradually as one moves away from the nucleus, without a sharp cutoff. Moreover, in condensed matter and molecules, the electron clouds of the atoms usually overlap to some extent, and some of the electrons may roam over a large region encompassing two or more atoms. Under most definitions the radii of isolated neutral atoms range between 30 and 300 pm (trillionths of a meter), or between 0.3 and 3 angstroms. Therefore, the radius of an atom is more than 10,000 times the radius of its nucleus (1â&#x20AC;&#x201C;10 fm), and less than 1/1000 of the wavelength of visible light (400â&#x20AC;&#x201C;700 nm).

The approximate shape of a molecule of ethanol, CH3CH2OH. Each atom is modeled by a sphere with the element's Van der Waals radius.

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For many purposes, atoms can be modeled as spheres. This is only a crude approximation, but it can provide quantitative explanations and predictions for many phenomena, such as the density of liquids and solids, the diffusion of fluids through molecular sieves, the [citation needed] arrangement of atoms and ions in crystals, and the size and shape of molecules. Atomic radii vary in a predictable and explicable manner across the periodic table. For instance, the radii generally decrease along each period (row) of the table, from the alkali metals to the noble gases; and increase down each group (column). The radius increases sharply between the noble gas at the end of each period and the alkali metal at the beginning of the next period. These trends of the atomic radii (and of various other chemical and physical properties of the elements) can be explained by the electron shell theory of the atom; they provided important evidence for the development and confirmation of quantum theory. The atomic radii decreases across the Periodic Table because as the atomic number increases, the number of protons increases across the period, but the extra electrons are only added to the same quantum shell. Therefore, the effective nuclear charge towards the outermost electrons increases, drawing the outermost electrons closer. As a result, the electron cloud contracts and the atomic radii decreases. History In 1920, shortly after it had become possible to determine the sizes of atoms using X-ray crystallography, it was suggested that all atoms of the same element have the same radii. However, in 1923, when more crystal data had become available, it was found that the approximation of an atom as a sphere does not necessarily hold when comparing the same atom in different crystal structures. Definitions Widely used definitions of atomic radius include: 

Van der Waals radius: in principle, half the minimum distance between the nuclei of two atoms of the element that are not bound to the same molecule.



Ionic radius: the nominal radius of the ions of an element in a specific ionization state, deduced from the spacing of atomic nuclei in crystalline salts that include that ion. In principle, the spacing between two adjacent oppositely charged ions (the length of the ionic bond between them) should equal the sum of their ionic radii.



Covalent radius: the nominal radius of the atoms of an element when covalently bound to other atoms, as deduced from the separation between the atomic nuclei in molecules. In principle, the distance between two atoms that are bound to each other in a molecule (the length of that covalent bond) should equal the sum of their covalent radii.



Metallic radius: the nominal radius of atoms of an element when joined to other [citation needed] atoms by metallic bonds.



Bohr radius: the radius of the lowest-energy electron orbit predicted by Bohr model of the atom (1913). It is only applicable to atoms and ions with a single electron, such as hydrogen, singly ionized helium, and positronium. Although the model itself is now obsolete, the Bohr radius for the hydrogen atom is still regarded as an important physical constant.

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Empirically measured atomic radii The following table shows empirically measured covalent radii for the elements, as −12 published by J. C. Slater in 1964. The values are in picometers (pm or 1×10 m,), with an accuracy of about 5 pm. The shade of the box ranges from red to yellow as the radius increases; gray indicates lack of data. Group

10 11 12 13 14 15 16 17 18

(vertical) Period (horizontal) 1

25 2

145 105

85 70 65 60 50

Na Mg

Al Si P

180 150

125 110 100 100 100

Ca Sc Ti

Cl Ar

Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

220 180 160 140 135 140 140 140 135 135 135 135 130 125 115 115 115 5

Rb Sr Y

Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I

235 200 180 155 145 145 135 130 135 140 160 155 155 145 145 140 140 6

Cs Ba *

Hf Ta W Re Os Ir

Pt Au Hg Tl Pb Bi

Po At Rn

260 215

155 145 135 135 130 135 135 135 150 190 180 160 190

Fr Ra **

Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Fl Uup Lv Uus Uuo

215 Lanthanides *

La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu 195 185 185 185 185 185 185 180 175 175 175 175 175 175 175

Actinides

Ac Th Pa U

Np Pu Am Cm Bk Cf Es Fm Md No Lr

195 180 180 175 175 175 175

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Explanation of the general trends

A graph comparing the calculated atomic radius of elements with atomic numbers 1â&#x20AC;&#x201C;100. Accuracy of Âą5 pm. The way the atomic radius varies with increasing atomic number can be explained by the arrangement of electrons in shells of fixed capacity. The shells are generally filled in order of increasing radius, since the negatively charged electrons are attracted by the positively charged protons in the nucleus. As the atomic number increases along each row of the periodic table, the additional electrons go into the same outermost shell; whose radius gradually contracts, due to the increasing nuclear charge. In a noble gas, the outermost shell is completely filled; therefore, the additional electron of next alkali metal will go into the next outer shell, accounting for the sudden increase in the atomic radius. The increasing nuclear charge is partly counterbalanced by the increasing number of electrons, a phenomenon that is known as shielding; which explains why the size of atoms usually increases down each column. However, there is one notable exception, known as the lanthanide contraction: the 5d block of elements are much smaller than one would expect, due to the shielding caused by the 4f electrons. The following table summarizes the main phenomena that influence the atomic radius of an element: factor

principle

increase with...

electron

quantum mechanics

principal

shells

tend to

and increase

effect on radius increases

down

azimuthal quantum atomic radius each column numbers

nuclear

attractive force acting on atomic number

decrease

charge

electrons

atomic radius each period

protons

decreases along

nucleus shielding

repulsive force acting on number of electron increase outermost shell electrons by shells

reduces the effect

atomic radius of the 2nd factor

inner electrons

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Lanthanide contraction Main article: Lanthanide contraction The electrons in the 4f-subshell, which is progressively filled from cerium (Z = 58) to lutetium (Z = 71), are not particularly effective at shielding the increasing nuclear charge from the sub-shells further out. The elements immediately following the lanthanides have atomic radii which are smaller than would be expected and which are almost identical to the atomic radii of the elements immediately above them. Hence hafnium has virtually the same atomic radius (and chemistry) as zirconium, and tantalum has an atomic radius similar to niobium, and so forth. The effect of the lanthanide contraction is noticeable up to platinum (Z = 78), after which it is masked by a relativistic effect known as the inert pair effect. Due to lanthanide contraction, the 5 following observations can be drawn: 3+

1. The size of Ln ions regularly decreases with atomic number. According to Fajans' 3+ rules, decrease in size of Ln ions increases the covalent character and decreases 3+ â&#x2C6;&#x2019; the basic character between Ln and OH ions in Ln(OH)3. Hence the order of size 3+ of Ln is given: 3+ 3+ 3+ La > Ce > ..., ... > Lu . 2. There is a regular decrease in their ionic radii. 3. There is a regular decrease in their tendency to act as a reducing agent, with increase in atomic number. 4. The second and third rows of d-block transition elements are quite close in properties. 5. Consequently, these elements occur together in natural minerals and are difficult to separate. d-Block contraction Main article: d-block contraction The d-block contraction is less pronounced than the lanthanide contraction but arises from a similar cause. In this case, it is the poor shielding capacity of the 3d-electrons which affects the atomic radii and chemistries of the elements immediately following the first row of the transition metals, from gallium (Z = 31) to bromine (Z = 35). Calculated atomic radii The following table shows atomic radii computed from theoretical models, as published by Enrico Clementi and others in 1967. The values are in picometres (pm). Group

10 11 12 13 14 15 16 17 18

(vertical) Period (horizontal)

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167 112

87 67 56 48 42 38

Na Mg

Al Si P

190 145

118 111 98 88 79 71

Ca Sc Ti

Cl Ar

Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr

243 194 184 176 171 166 161 156 152 149 145 142 136 125 114 103 94 88 5

Rb Sr Y

Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te I

265 219 212 206 198 190 183 178 173 169 165 161 156 145 133 123 115 108 6

Cs Ba *

Hf Ta W Re Os Ir

298 253

208 200 193 188 185 180 177 174 171 156 154 143 135

Fr Ra **

Rf Db Sg Bh Hs Mt Ds Rg Cn Uut Fl Uup Lv Uus Uuo

Lanthanides *

Pt Au Hg Tl Pb Bi

120

La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu 247 206 205 238 231 233 225 228

Actinides

Po At Rn

Ac Th Pa U

226 222 222 217

Np Pu Am Cm Bk Cf Es Fm Md No Lr

See also 

Atomic radii of the elements (data page)



Chemical bond



Covalent radius



Bond length



Steric hindrance

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Chapter 3: Ionization energy

Periodic trend for ionization energy. Each period begins at a minimum for the alkali metals, and ends at a maximum for the noble gases. The ionization energy (IE) of an atom or molecule describes the minimum amount of energy required to remove an electron (to infinity) from the atom or molecule in the gaseous state. +

X + energy → X + e

The term ionization potential has been used in the past but is not recommended. The units for ionization energy vary from discipline to discipline. In physics, the ionization energy is typically specified in electron volts (eV) and refers to the energy required to remove a single electron from a single atom or molecule. In chemistry, the ionization energy is typically specified as a molar quantity (molar ionization energy or enthalpy) and is reported in units of kJ/mol or kcal/mol (the amount of energy it takes for all the atoms in a mole to lose one electron each). th

The n ionization energy refers to the amount of energy required to remove an electron from the species with a charge of (n-1). For example, the first three ionization energies are defined as follows: st

1 ionization energy +

X→X +e

2 ionization energy +

X →X

3 ionization energy X

→X

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Values and trends

Main articles: Molar ionization energies of the elements and Ionization energies of the elements (data page) Generally the (n+1)th ionization energy is larger than the nth ionization energy. Always, the next ionization energy involves removing an electron from an orbital closer to the nucleus. Electrons in the closer orbitals experience greater forces of electrostatic attraction; thus, their removal requires increasingly more energy. Ionization energy becomes greater up and to the right of the periodic table. Some values for elements of the third period are given in the following table: Successive

molar

ionization

energies

kJ/mol

(96.485 kJ/mol = 1 eV/particle) Element First Second Third Fourth Fifth Na

496 4,560

738 1,450

7,730

577 1,816

2,881 11,600

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786 1,577

3,228 4,354 16,100

1,060 1,890

2,905 4,950 6,270 21,200

999.6 2,260

3,375 4,565 6,950 8,490 27,107

1,256 2,295

3,850 5,160 6,560 9,360 11,000

1,520 2,665

3,945 5,770 7,230 8,780 12,000

Large jumps in the successive molar ionization energies occur when passing noble gas configurations. For example, as can be seen in the table above, the first two molar ionization energies of magnesium (stripping the two 3s electrons from a magnesium atom) are much smaller than the third, which requires stripping off a 2p electron from the very stable neon 2+ configuration of Mg . Ionization energy is also a periodic trend within the periodic table organization. Moving left to right within a period or upward within a group, the first ionization energy generally increases with a few discrepancies (aluminum and sulphur). As the nuclear charge of the nucleus increases across the period, the atomic radius decreases and the electron cloud becomes closer towards the nucleus. Ionization energy increases from left to right in a period and decreases from top to bottom in a group. Electrostatic explanation Atomic ionization energy can be predicted by an analysis using electrostatic potential and the Bohr model of the atom, as follows (note that the derivation uses Gaussian units). Consider an electron of charge -e and an atomic nucleus with charge +Ze, where Z is the number of protons in the nucleus. According to the Bohr model, if the electron were to approach and bind with the atom, it would come to rest at a certain radius a. The electrostatic potential V at distance a from the ionic nucleus, referenced to a point infinitely far away, is:

Since the electron is negatively charged, it is drawn inwards by this positive electrostatic potential. The energy required for the electron to "climb out" and leave the atom is:

This analysis is incomplete, as it leaves the distance a as an unknown variable. It can be made more rigorous by assigning to each electron of every chemical element a characteristic distance, chosen so that this relation agrees with experimental data.

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It is possible to expand this model considerably by taking a semi-classical approach, in which momentum is quantized. This approach works very well for the hydrogen atom, which only has one electron. The magnitude of the angular momentum for a circular orbit is:

The total energy of the atom is the sum of the kinetic and potential energies, that is:

Velocity can be eliminated from the kinetic energy term by setting the Coulomb attraction equal to the centripetal force, giving:

Solving the angular momentum for v and substituting this into the expression for kinetic energy, we have:

This establishes the dependence of the radius on n. That is:

Now the energy can be found in terms of Z, e, and r. Using the new value for the kinetic energy in the total energy equation above, it is found that:

At its smallest value, n is equal to 1 and r is the Bohr radius a0 which equals to . Now, the equation for the energy can be established in terms of the Bohr radius. Doing so gives the result:

Quantum-mechanical explanation According to the more complete theory of quantum mechanics, the location of an electron is best described as a probability distribution. The energy can be calculated by integrating over this cloud. The cloud's underlying mathematical representation is the wavefunction which is built from Slater determinants consisting of molecular spin orbitals. These are related by

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Pauli's exclusion principle to the antisymmetrized products of the atomic or molecular orbitals. In general, calculating the nth ionization energy requires calculating the energies of and electron systems. Calculating these energies exactly is not possible except for the simplest systems (i.e. hydrogen), primarily because of difficulties in integrating the electron correlation terms. Therefore, approximation methods are routinely employed, with different methods varying in complexity (computational time) and in accuracy compared to empirical data. This has become a well-studied problem and is routinely done in computational chemistry. At the lowest level of approximation, the ionization energy is provided by Koopmans' theorem. Vertical and adiabatic ionization energy in molecules

Figure 1. Franckâ&#x20AC;&#x201C;Condon principle energy diagram. For ionization of a diatomic molecule the only nuclear coordinate is the bond length. The lower curve is the potential energy curve of the neutral molecule, and the upper curve is for the positive ion with a longer bond length. The blue arrow is vertical ionization, here from the ground state of the molecule to the v=2 level of the ion. Ionization of molecules often leads to changes in molecular geometry, and two types of (first) ionization energy are defined â&#x20AC;&#x201C; adiabatic and vertical. Adiabatic ionization energy: The adiabatic ionization energy of a molecule is the minimum amount of energy required to remove an electron from a neutral molecule, i.e. the difference between the energy of the vibrational ground state of the neutral species and that of the positive ion. The specific equilibrium geometry of each species does not affect this value.

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Vertical ionization energy: Due to the possible changes in molecular geometry that may result from ionization, additional transitions may exist between the vibrational ground state of the neutral species and vibrational excited states of the positive ion. In other words, ionization is accompanied by vibrational excitation. The intensity of such transitions are explained by the Franck–Condon principle, which predicts that the most probable and intense transition corresponds to the vibrational excited state of the positive ion that has the same geometry as the neutral molecule. This transition is referred to as the "vertical" ionization energy since it is represented by a completely vertical line on a potential energy diagram (see Figure). For a diatomic molecule, the geometry is defined by the length of a single bond. The removal of an electron from a bonding molecular orbital weakens the bond and increases the bond length. In Figure 1, the lower potential energy curve is for the neutral molecule and the upper surface is for the positive ion. Both curves plot the potential energy as a function of bond length. The horizontal lines correspond to vibrational levels with their associated vibrational wave functions. Since the ion has a weaker bond, it will have a longer bond length. This effect is represented by shifting the minimum of the potential energy curve to the right of the neutral species. The adiabatic ionization is the diagonal transition to the vibrational ground state of the ion. Vertical ionization involves vibrational excitation of the ionic state and therefore requires greater energy. In many circumstances, the adiabatic ionization energy is often a more desirable physical quantity since it describes the difference in energy between the two potential energy surfaces. However, due to experimental limitations, the adiabatic ionization energy is often difficult to determine, whereas the vertical detachment energy is easily identifiable and measurable. Analogs of Ionization Energy to Other Systems While the term ionization energy is largely used only for gas-phase atomic or molecular species, there are a number of analogous quantities that consider the amount of energy required to remove an electron from other physical systems. Electron binding energy: A generic term for the ionization energy that can be used for species with any charge state. For example, the electron binding energy for the chloride ion is the minimum amount of energy required to remove an electron from the chlorine atom when it has a charge of -1. In this particular example, the electron binding energy has the same magnitude as the electron affinity for the neutral chlorine atom. In another example, the electron binding energy refers the minimum amount of energy required to remove an electron from the dicarboxylate dianion O2C(CH2)8CO2 . Work function: The minimum amount of energy required to remove an electron from a solid surface. See also 

Electron affinity — a closely related concept describing the energy released by adding an electron to a neutral atom or molecule.



Work function is the energy required to strip an electron from a solid to just outside its surface.



Electronegativity is a number that shares some similarities with ionization energy.

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

Koopmans' theorem, regarding the predicted ionization energies in Hartree–Fock theory.



Di-tungsten tetra(hpp) has the lowest recorded ionization energy for a stable chemical compound.

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Chapter 4: Electron affinity In chemistry and atomic physics, the electron affinity of an atom or molecule is defined as the energy change when an electron is added to a neutral atom or molecule to form a negative ion. −

−

X + e → X + energy In solid state physics, the electron affinity for a surface is defined somewhat differently (see below). Measurement and use of electron affinity This property is measured for atoms and molecules in the gaseous state only, since in the solid or liquid states their energy levels would be changed by contact with other atoms or molecules. A list of the electron affinities was used by Robert S. Mulliken to develop an electronegativity scale for atoms, equal to the average of the electron affinity and ionization potential. Other theoretical concepts that use electron affinity include electronic chemical potential and chemical hardness. Another example, a molecule or atom that has a more positive value of electron affinity than another is often called an electron acceptor and the less positive an electron donor. Together they may undergo charge-transfer reactions. Sign convention To use electron affinities properly, it is essential to keep track of sign. For any reaction that releases energy, the change ΔE in total energy has a negative value and the reaction is called an exothermic process. Electron capture for almost all non-noble gas atoms involves the release of energy and thus are exothermic. The positive values that are listed in tables of Eea are amounts or magnitudes. It is the word, released within the definition energy released that supplies the negative sign. Confusion arises in mistaking Eea for a change in energy, ΔE, in which case the positive values listed in tables would be for an endo- not exo-thermic process. The relation between the two is Eea = −ΔE(attach). However, if the value assigned to Eea is negative, the negative sign implies a reversal of direction, and energy is required to attach an electron. In this case, the electron capture is an endothermic process and the relationship, Eea = −ΔE(attach) is still valid. Negative values typically arise for the capture of a second electron, but also for the nitrogen atom. The usual expression for calculating Eea when an electron is attached is Eea = (Einitial − Efinal)attach = −ΔE(attach) This expression does follow the convention ΔX = X(final) − X(initial) since −ΔE = −(E(final) − E(initial)) = E(initial) − E(final). Equivalently, electron affinity can also be defined as the amount of energy required to detach an electron from a singly charged negative ion, i.e. the energy change for the process −

X →X+e

−

If the same table is employed for the forward and reverse reactions, without switching signs, care must be taken to apply the correct definition to the corresponding direction, attachment

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(release) or detachment (require). Since almost all detachments (require +) an amount of energy listed on the table, those detachment reactions are endothermic, or ΔE(detach) > 0. Electron affinities of the elements

EA vs Atomic Number Main article: Electron affinity (data page) Although Eea varies greatly across the periodic table, some patterns emerge. Generally, nonmetals have more positive Eea than metals. Atoms whose anions are more stable than neutral atoms have a greater Eea. Chlorine most strongly attracts extra electrons; mercury most weakly attracts an extra electron. The electron affinities of the noble gases have not been conclusively measured, so they may or may not have slightly negative values. Eea generally increases across a period (row) in the periodic table. This is caused by the filling of the valence shell of the atom; a Group 17 atom releases more energy than a Group 1 atom on gaining an electron because it obtains a filled valence shell and therefore is more stable. A trend of decreasing Eea going down the groups in the periodic table would be expected. The additional electron will be entering an orbital farther away from the nucleus. Since this electron is farther from the nucleus it is less attracted to the nucleus and would release less energy when added. However, a clear counterexample to this trend can be found in Group 2, and this trend only applies to Group 1 atoms. Electron affinity follows the trend of electronegativity. Fluorine (F) has a higher electron affinity than oxygen and so on. The following data are quoted in kJ/mol. Elements marked with an asterisk are expected to have electron affinities close to zero on quantum mechanical grounds. Electron affinities in the periodic table Group → 1

12 13 14

↓ Period 1

Li Be

60 *

27 122 *

Dr. Pramod Kothari / Periodic trends

141 328 *

Page 18

Na Mg

53 *

42 134 72

Ca Sc Ti V

48 2 5

18 8

Cr Mn Fe

51 65 *

Co Ni 64

30 41 86 72 *

112 119 *

Cs Ba *

Hf Ta W Re Os Ir

46 14 7

101 110 54

31 79 *

41 119 79 Sn Sb

Au Hg Tl

Pb Bi

36 35

Ds Rg Cn Uut Fl

La Ce Pr Nd Pm Sm Eu Gd Tb

Dy Ho Er

45 92 ** Actinides

Np Pu Am Cm Bk

200 349 * Se Br

195 324 * Te

Po At

91 Uup Lv

Rn *

Uus Uuo

Tm Yb Lu 99

Ac Th Pa U

39 107 101 190 295 *

126 *

104 150 205 223 *

Fr Ra ** Rf Db Sg Bh Hs Mt

* Lanthanides

Cu Zn Ga Ge As

Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In 47 5

Cf Es Fm Md No Lr

Legend The number mentioned is Electron affinity in kJ/mol (rounded) For the equivalent value in eV, see: Electron affinity (data page) * Denotes elements that are expected to have electron affinities close to zero on quantum mechanical grounds

black=Solid

green=Liquid

red=Gas

grey=Unknown Color of the atomic number shows state of matter (at 0 째C and 1 atm)

Primordial From

Synthetic Border shows natural occurrence of the element

decay

Dr. Pramod Kothari / Periodic trends

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Background color shows subcategory in the metalâ&#x20AC;&#x201C;nonmetal range: Metal

Metalloid Nonmetal

Alkali Alkaline Lanmetal earth

thanide

Actinide Transition Poor metal

metal

Unknown

chemical Polyatomic Diatomic Noble properties nonmetal nonmetal gas

metal

Molecular electron affinities The electron affinity of molecules is a complicated function of their electronic structure. For instance the electron affinity for benzene is negative, as is that of naphthalene, while those of anthracene, phenanthrene and pyrene are positive. In silico experiments show that the electron affinity of hexacyanobenzene surpasses that of fullerene. "Electron affinity" as defined in solid state physics

Band diagram of semiconductor-vacuum interface showing electron affinity EEA, defined as the difference between near-surface vacuum energy Evac, and near-surface conduction band edge EC. Also shown: Fermi level EF, valence band edge EV, work function W. In the field of solid state physics, the electron affinity is defined differently from in chemistry. For a semiconductor-vacuum interface (that is, the surface of a semiconductor), electron affinity, typically denoted by EEA or Ď&#x2021;, is defined instead as the energy obtained by moving an electron from the vacuum just outside the semiconductor to the bottom of the conduction band just inside the semiconductor:

In an intrinsic semiconductor at absolute zero, this concept is analogous to the chemistry definition of electron affinity, since an added electron will go to the bottom of the conduction band. At nonzero temperature, and for other materials (metals, semimetals, heavily doped semiconductors), the analogy does not hold since an added electron will instead go to the Fermi level on average. In any case, the electron affinity of a solid substance is very different from the chemistry electron affinity of that substance in gas phase. For example, a silicon

Dr. Pramod Kothari / Periodic trends

Page 20

crystal surface has electron affinity 4.05 eV, whereas an isolated silicon atom has electron affinity 1.39 eV. The electron affinity is a similar concept to the work function, but distinct. The work function is the thermodynamic work that can be obtained by reversibly, isothermally moving an electron from the vacuum to the material; this thermodynamic electron goes to the Fermi level on average, not the conduction band edge: . While the work function of a semiconductor can be changed by doping, the electron affinity ideally does not change with doping and so it is closer to being a material constant. However, the electron affinity does depend on the surface termination (crystal face, surface chemistry, etc.). In certain circumstances, the electron affinity may become negative. Often negative electron affinity is desired to obtain efficient cathodes that can supply electrons to the vacuum with little energy loss. The observed electron yield as a function of various parameters such as bias voltage or illumination conditions can be used to describe these structures with band diagrams in which the electron affinity is one parameter. For one illustration of the apparent effect of surface termination on electron emission, see Figure 3 in Marchywka Effect. In semiconductor physics, the primary use of the electron affinity is not actually in the analysis of semiconductor-vacuum junctions, but rather in heuristic electron affinity rules for estimating the band bending that occurs at the interface of two materials. See also 

Ionization energy — a closely related concept describing the energy required to remove an electron from a neutral atom or molecule



One-electron reduction



Electron-capture mass spectrometry



Electronegativity



Valence electron



Vacuum level



Electron donor



Tro, Nivaldo J. (2008). Chemistry: A Molecular Approach (2nd Edn.). New Jersey: Pearson Prentice Hall. ISBN 0-13-100065-9. pp. 348–349.

External links 

Electron affinity, definition from the IUPAC Gold Book

Dr. Pramod Kothari / Periodic trends

Page 21

Chapter 5: Electronegativity

This electrostatic potential map shows how the oxygen atom has a more negative charge (red) than the positive (blue) hydrogen atoms of a water molecule . Electronegativity, symbol Ď&#x2021;, is a chemical property that describes the tendency of an atom or a functional group to attract electrons (or electron density) towards itself. An atom's electronegativity is affected by both its atomic number and the distance that its valence electrons reside from the charged nucleus. The higher the associated electronegativity number, the more an element or compound attracts electrons towards it. First proposed by Linus Pauling in 1932 as a development of valence bond theory, it has been shown to correlate with a number of other chemical properties. Electronegativity cannot be directly measured and must be calculated from other atomic or molecular properties. Several methods of calculation have been proposed, and although there may be small differences in the numerical values of the electronegativity, all methods show the same periodic trends between elements. The most commonly used method of calculation is that originally proposed by Linus Pauling. This gives a dimensionless quantity, commonly referred to as the Pauling scale, on a relative scale running from around 0.7 to 3.98 (hydrogen = 2.20). When other methods of calculation are used, it is conventional (although not obligatory) to quote the results on a scale that covers the same range of numerical values: this is known as an electronegativity in Pauling units. Electronegativity, as it is usually calculated, is not strictly a property of an atom, but rather a property of an atom in a molecule. Properties of a free atom include ionization energy and electron affinity. It is to be expected that the electronegativity of an element will vary with its chemical environment, but it is usually considered to be a transferable property, that is to say that similar values will be valid in a variety of situations. On the most basic level, electronegativity is determined by factors like the nuclear charge (the more protons an atom has, the more "pull" it will have on negative electrons) and the number/location of other electrons present in the atomic shells (the more electrons an atom has, the farther from the nucleus the valence electrons will be, and as a result the less positive charge they will experienceâ&#x20AC;&#x201D;both because of their increased distance from the nucleus, and because the other electrons in the lower energy core orbitals will act to shield the valence electrons from the positively charged nucleus). The opposite of electronegativity is electropositivity: a measure of an element's ability to donate electrons. Electronegativities of the elements Periodic table of electronegativity using the Pauling scale

Dr. Pramod Kothari / Periodic trends

Page 22

→ Atomic radius decreases → Ionization energy increases → Electronegativity increases → Group → 1

↓ Period 1

2.20 2

0.98 1.57

2.04 2.55 3.04 3.44 3.98

Na Mg

0.93 1.31

1.61 1.90 2.19 2.58 3.16

Ca Sc

Mn Fe

Co Ni

Cu Zn

Ga Ge As

0.82 1.00 1.36 1.54 1.63 1.66 1.55 1.83 1.88 1.91 1.90 1.65 1.81 2.01 2.18 2.55 2.96 3.00 5

Rb Sr

Nb Mo Tc

Ru Rh Pd

Cd In

0.82 0.95 1.22 1.33 1.6 2.16 1.9 2.2 2.28 2.20 1.93 1.69 1.78 1.96 2.05 2.1 2.66 2.60 6

Re Os Ir

Hg Tl

0.79 0.89

1.3 1.5 2.36 1.9 2.2 2.20 2.28 2.54 2.00 1.62 1.87 2.02 2.0 2.2 2.2

Ra **

Db Sg

Rg Cn Uut Fl

Uup Lv

Uus Uuo

Tm Yb

0.7 0.9

* Lanthanoids

Ce Pr

Nd Pm Sm Eu

Gd Tb

Ho Er

1.1 1.12 1.13 1.14 1.13 1.17 1.2 1.2 1.1 1.22 1.23 1.24 1.25 1.1 1.27 ** Actinoids

Np Pu

Am Cm Bk

Fm Md No Lr

1.1 1.3 1.5 1.38 1.36 1.28 1.13 1.28 1.3 1.3 1.3 1.3 1.3 1.3 1.3 Values are given for the elements in their most common and stable oxidation states. See also: Electronegativities of the elements (data page)

Dr. Pramod Kothari / Periodic trends

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Methods of calculation

Pauling electronegativity Pauling first proposed the concept of electronegativity in 1932 as an explanation of the fact that the covalent bond between two different atoms (A–B) is stronger than would be expected by taking the average of the strengths of the A–A and B–B bonds. According to valence bond theory, of which Pauling was a notable proponent, this "additional stabilization" of the heteronuclear bond is due to the contribution of ionic canonical forms to the bonding. The difference in electronegativity between atoms A and B is given by:

where the dissociation energies, Ed, of the A–B, A–A and B–B bonds are expressed in –½ electronvolts, the factor (eV) being included to ensure a dimensionless result. Hence, the difference in Pauling electronegativity between hydrogen and bromine is 0.73 (dissociation energies: H–Br, 3.79 eV; H–H, 4.52 eV; Br–Br 2.00 eV) As only differences in electronegativity are defined, it is necessary to choose an arbitrary reference point in order to construct a scale. Hydrogen was chosen as the reference, as it forms covalent bonds with a large variety of elements: its electronegativity was fixed first at 2.1, later revised to 2.20. It is also necessary to decide which of the two elements is the more electronegative (equivalent to choosing one of the two possible signs for the square root). This is usually done using "chemical intuition": in the above example, hydrogen bromide + – dissolves in water to form H and Br ions, so it may be assumed that bromine is more electronegative than hydrogen. However, in principle, since the same electronegativities should be obtained for any two bonding compounds, the data is in fact overdetermined, and the signs are unique once a reference point is fixed (usually, for H or F). To calculate Pauling electronegativity for an element, it is necessary to have data on the dissociation energies of at least two types of covalent bond formed by that element. A. L. Allred updated Pauling's original values in 1961 to take account of the greater availability of thermodynamic data, and it is these "revised Pauling" values of the electronegativity that are most often used. The essential point of Pauling electronegativity is that there is an underlying, quite accurate, semi-empirical formula for dissociation energies, namely:

or sometimes, a more accurate fit

This is an approximate equation, but holds with good accuracy. Pauling obtained it by noting that a bond can be approximately represented as a quantum mechanical superposition of a covalent bond and two ionic bond-states. The covalent energy of a bond is approximately, by quantum mechanical calculations, the geometric mean of the two energies of covalent bonds of the same molecules, and there is an additional energy that comes from ionic factors, i.e. polar character of the bond.

Dr. Pramod Kothari / Periodic trends

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The geometric mean is approximately equal to the arithmetic mean - which is applied in the first formula above - when the energies are of the similar value, e.g., except for the highly electropositive elements, where there is a larger difference of two dissociation energies; the geometric mean is more accurate and almost always gives a positive excess energy, due to ionic bonding. The square root of this excess energy, Pauling notes, is approximately additive, and hence one can introduce the electronegativity. Thus, it is this semi-empirical formula for bond energy that underlies Pauling electronegativity concept. The formulas are approximate, but this rough approximation is in fact relatively good and gives the right intuition, with the notion of polarity of the bond and some theoretical grounding in quantum mechanics. The electronegativities are then determined to best fit the data. In more complex compounds, there is additional error since electronegativity depends on the molecular environment of an atom. Also, the energy estimate can be only used for single, not for multiple bonds. The energy of formation of a molecule containing only single bonds then can be approximated from an electronegativity table, and depends on the constituents and sum of squares of differences of electronegativities of all pairs of bonded atoms. Such a formula for estimating energy typically has relative error of order of 10%, but can be used to get a rough qualitative idea and understanding of a molecule. Mulliken electronegativity

The correlation between Mulliken electronegativities (x-axis, in kJ/mol) and Pauling electronegativities (y-axis). Robert S. Mulliken proposed that the arithmetic mean of the first ionization energy (Ei) and the electron affinity (Eea) should be a measure of the tendency of an atom to attract electrons. As this definition is not dependent on an arbitrary relative scale, it has also been termed absolute electronegativity, with the units of kilojoules per mole or electronvolts.

However, it is more usual to use a linear transformation to transform these absolute values into values that resemble the more familiar Pauling values. For ionization energies and electron affinities in electronvolts,

and for energies in kilojoules per mole,

Dr. Pramod Kothari / Periodic trends

Page 25

The Mulliken electronegativity can only be calculated for an element for which the electron affinity is known, fifty-seven elements as of 2006. The Mulliken electronegativity of an atom is sometimes said to be the negative of the chemical potential. By inserting the energetic definitions of the ionization potential and electron affinity into the Mulliken electronegativity, it is possible to show that the Mulliken chemical potential is a finite difference approximation of the electronic energy with respect to the number of electrons., i.e.,

Allred–Rochow electronegativity

–2

The correlation between Allred–Rochow electronegativities (x-axis, in Å ) and Pauling electronegativities (y-axis). A. Louis Allred and Eugene G. Rochow considered that electronegativity should be related to the charge experienced by an electron on the "surface" of an atom: The higher the charge per unit area of atomic surface the greater the tendency of that atom to attract electrons. The effective nuclear charge, Zeff experienced by valence electrons can be estimated using Slater's rules, while the surface area of an atom in a molecule can be taken to be proportional to the square of the covalent radius, rcov. When rcov is expressed in picometres,

Dr. Pramod Kothari / Periodic trends

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Sanderson electronegativity equalization

The correlation between Sanderson electronegativities (x-axis, arbitrary units) and Pauling electronegativities (y-axis). Sanderson has also noted the relationship between Mulliken electronegativity and atomic size, and has proposed a method of calculation based on the reciprocal of the atomic volume. With a knowledge of bond lengths, Sanderson's model allows the estimation of bond energies in a wide range of compounds. Sanderson's model has also been used to calculate molecular geometry, s-electrons energy, NMR spin-spin constants and other parameters for organic compounds. This work underlies the concept of electronegativity equalization, which suggests that electrons distribute themselves around a molecule to minimize or to equalize the Mulliken electronegativity. This behavior is analogous to the equalization of chemical potential in macroscopic thermodynamics. Allen electronegativity

The correlation between Allen electronegativities (x-axis, in kJ/mol) and Pauling electronegativities (y-axis). Perhaps the simplest definition of electronegativity is that of Leland C. Allen, who has proposed that it is related to the average energy of the valence electrons in a free atom,

Dr. Pramod Kothari / Periodic trends

Page 27

where εs,p are the one-electron energies of s- and p-electrons in the free atom and ns,p are the number of s- and p-electrons in the valence shell. It is usual to apply a scaling factor, −3 1.75×10 for energies expressed in kilojoules per mole or 0.169 for energies measured in electronvolts, to give values that are numerically similar to Pauling electronegativities. The one-electron energies can be determined directly from spectroscopic data, and so electronegativities calculated by this method are sometimes referred to as spectroscopic electronegativities. The necessary data are available for almost all elements, and this method allows the estimation of electronegativities for elements that cannot be treated by the other methods, e.g. francium, which has an Allen electronegativity of 0.67. However, it is not clear what should be considered to be valence electrons for the d- and f-block elements, which leads to an ambiguity for their electronegativities calculated by the Allen method. In this scale neon has the highest electronegativity of all elements, followed by fluorine, helium, and oxygen. 



Electronegativity using the Allen scale Group → 1

↓ Period 1

2.300

4.160

0.912 1.576

2.051 2.544 3.066 3.610 4.193 4.789

0.869 1.293 4

1.613 1.916 2.253 2.589 2.869 3.242 Sc

Mn Fe

Co Ni

Cu Zn

0.734 1.034 1.19 1.38 1.53 1.65 1.75 1.80 1.84 1.88 1.85 1.59 1.756 1.994 2.211 2.434 2.685 2.966 5

Nb Mo Tc

Ru Rh Pd

Cd In

0.706 0.963 1.12 1.32 1.41 1.47 1.51 1.54 1.56 1.59 1.87 1.52 1.656 1.824 1.984 2.158 2.359 2.582 6

Re Os Ir

Hg Tl

0.659 0.881 1.09 1.16 1.34 1.47 1.60 1.65 1.68 1.72 1.92 1.76 1.789 1.854 2.01 2.19 2.39 2.60

Dr. Pramod Kothari / Periodic trends

Page 28

0.67 0.89 See also: Electronegativities of the elements (data page)

Correlation of electronegativity with other properties

2â&#x20AC;&#x201C;

The variation of the isomer shift (y-axis, in mm/s) of [SnX6]

anions, as measured by

119

MĂśssbauer spectroscopy, against the sum of the Pauling electronegativities of the halide substituents (x-axis). The wide variety of methods of calculation of electronegativities, which all give results that correlate well with one another, is one indication of the number of chemical properties which might be affected by electronegativity. The most obvious application of electronegativities is in the discussion of bond polarity, for which the concept was introduced by Pauling. In general, the greater the difference in electronegativity between two atoms the more polar the bond that will be formed between them, with the atom having the higher electronegativity being at the negative end of the dipole. Pauling proposed an equation to relate "ionic character" of a bond to the difference in electronegativity of the two atoms, although this has fallen somewhat into disuse. Several correlations have been shown between infrared stretching frequencies of certain bonds and the electronegativities of the atoms involved: however, this is not surprising as such stretching frequencies depend in part on bond strength, which enters into the calculation of Pauling electronegativities. More convincing are the correlations between electronegativity and chemical shifts in NMR spectroscopy or isomer shifts in MĂśssbauer spectroscopy (see figure). Both these measurements depend on the s-electron density at the nucleus, and so are a good indication that the different measures of electronegativity really are describing "the ability of an atom in a molecule to attract electrons to itself".

Dr. Pramod Kothari / Periodic trends

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Trends in electronegativity

Periodic trends

The variation of Pauling electronegativity (y-axis) as one descends the main groups of the periodic table from the second period to the sixth period In general, electronegativity increases on passing from left to right along a period, and decreases on descending a group. Hence, fluorine is the most electronegative of the elements (not counting noble gases), whereas caesium is the least electronegative, at least of those elements for which substantial data is available. There are some exceptions to this general rule. Gallium and germanium have higher electronegativities than aluminium and silicon, respectively, because of the d-block contraction. Elements of the fourth period immediately after the first row of the transition metals have unusually small atomic radii because the 3d-electrons are not effective at shielding the increased nuclear charge, and smaller atomic size correlates with higher electronegativity (see Allred-Rochow electronegativity, Sanderson electronegativity above). The anomalously high electronegativity of lead, in particular when compared to thallium and bismuth, appears to be an artifact of data selection (and data availability)â&#x20AC;&#x201D;methods of calculation other than the Pauling method show the normal periodic trends for these elements. Variation of electronegativity with oxidation number In inorganic chemistry it is common to consider a single value of the electronegativity to be valid for most "normal" situations. While this approach has the advantage of simplicity, it is clear that the electronegativity of an element is not an invariable atomic property and, in particular, increases with the oxidation state of the element. Allred used the Pauling method to calculate separate electronegativities for different oxidation states of the handful of elements (including tin and lead) for which sufficient data was available. However, for most elements, there are not enough different covalent compounds for which bond dissociation energies are known to make this approach feasible. This is particularly true of the transition elements, where quoted electronegativity values are usually, of necessity, averages over several different oxidation states and where trends in electronegativity are harder to see as a result.

Dr. Pramod Kothari / Periodic trends

Page 30

Acid

Formula Chlorine pKa oxidation state

Hypochlorous acid HClO

+7.5

Chlorous acid

HClO2

+2.0

Chloric acid

HClO3

–1.0

Perchloric acid

HClO4

–10

The chemical effects of this increase in electronegativity can be seen both in the structures of oxides and halides and in the acidity of oxides and oxoacids. Hence CrO 3 and Mn2O7 are acidic oxides with low melting points, while Cr 2O3 is amphoteric and Mn2O3 is a completely basic oxide. The effect can also be clearly seen in the dissociation constants of the oxoacids of chlorine. The effect is much larger than could be explained by the negative charge being shared among a larger number of oxygen atoms, which would lead to a difference in pKa of log10(¼) = –0.6 between hypochlorous acid and perchloric acid. As the oxidation state of the central chlorine atom increases, more electron density is drawn from the oxygen atoms onto the chlorine, reducing the partial negative charge on the oxygen atoms and increasing the acidity. Group electronegativity In organic chemistry, electronegativity is associated more with different functional groups than with individual atoms. The terms group electronegativity and substituent electronegativity are used synonymously. However, it is common to distinguish between the inductive effect and the resonance effect, which might be described as σ- and πelectronegativities, respectively. There are a number of linear free-energy relationships that have been used to quantify these effects, of which the Hammett equation is the best known. Kabachnik parameters are group electronegativities for use in organophosphorus chemistry. Electropositivity Electropositivity is a measure of an element's ability to donate electrons, and therefore form positive ions; thus, it is opposed to electronegativity. Mainly, this is an attribute of metals, meaning that, in general, the greater the metallic character of an element the greater the electropositivity. Therefore the alkali metals are most electropositive of all. This is because they have a single electron in their outer shell and, as this is relatively far from the nucleus of the atom, it is easily lost; in other words, these metals have low ionization energies. While electronegativity increases along periods in the periodic table, and decreases down groups, electropositivity decreases along periods (from left to right) and increases down groups.

Dr. Pramod Kothari / Periodic trends

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Electropositive shark repellent utilizes electropositive metals as shark repellents, since they generate measurable voltages in a seawater electrolyte relative to a shark. See also 

Electronegativities of the elements (data page)



Chemical polarity

Bibliography 

Jolly, William L. (1991). Modern Inorganic Chemistry (2nd ed.). New York: McGrawHill. pp. 71–76. ISBN 0-07-112651-1.

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