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"Atomic radius"

"Atomic radius"
Atomic radius, and more generally the size of an atom, is not a precisely defined physical quantity, nor is it constant in all circumstances.[1] The value assigned to the radius of a particular atom always depends on the definition chosen for "atomic radius," and the appropriate definition depends on the context.
The term "atomic radius" itself is problematic: it may be restricted to the size of free atoms, or it may be used as a general term for the different measures of the size of atoms, both bound in molecules and free. In the latter case, which is the approach adopted here, it should also include ionic radius, as the distinction between covalent and ionic bonding is itself somewhat arbitrary.[2]
The atomic radius is determined entirely by the electrons: The size of the atomic nucleus is measured in femtometres, 100,000 times smaller than the cloud of electrons. Electrons, however, do not have definite positions—although they are more likely to be in certain regions than others—and the electron cloud does not have sharp edges.
Despite (or maybe because of) these difficulties, many different attempts have been made to quantify the size of atoms (and ions), based both on experimental measurements and calculation methods. It is undeniable that atoms do behave as if they were spheres with a radius of 30–300 pm, that atomic size varies in a predictable and explicable manner across the periodic table and that this variation has important consequences for the chemistry of the elements. Atomic radii decreases from Alkali Earth metals to the Noble Gases on the far right side of the periodic table. This is determined by the effective nuclear charge that increases with added electrons in a period. The atomic radii increases as you go down each column with added electrons in a period.
Periodic trends in atomic radius
Atomic radius tends to increase as one proceeds down any group of the periodic table. This trend is intuitively satisfying: atoms with more electrons have larger radii. As one proceeds across any row of the periodic table, a more complex rationale is required: atoms contract in size within a period from left to right. This contraction results from the increasing number of protons in the nucleus. Protons make little contribution to the size of the atom, but they increase the positive charge of the nucleus, which draws the electrons into tighter orbitals, this is also known as the nuclear charge of the atom.
factor principle increase with... tend to effect
electron shells quantum mechanics Principal Quantum Number, Azimuthal Quantum Number atomic radius increase increase when going down on group trend
nuclear charge attractive force acting on electrons by protons in nucleus atomic number atomic radius decrease decrease when move from left to right on Periodic Table
shielding repulsive force acting on outermost shell electrons by inner electrons number of electron shells atomic radius increase reduce the effect of the 2nd factor
The increasing nuclear charge is partly counterbalanced by the increasing number of electrons in a phenomenon that is known as shielding, which is why the size of atoms usually increases as a group is descended. However, there are two occasions where shielding is less effective: in these cases, the atoms are smaller than would otherwise be expected. MA7
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.[3] 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.
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).[3]
electron affinity
The electron affinity, Eea, of an atom or molecule is the amount of energy required to detach an electron from a singly charged negative ion,[1] i.e., the energy change for the processX- → X + e−
An equivalent definition is the energy released (Einitial − Efinal) when an electron is attached to a neutral atom or molecule. It should be noted that the sign convention for Eea is the opposite to most thermodynamic quantities: a positive electron affinity indicates that energy is released on going from atom to anion.
All elements whose EA have been measured using modern methods have a positive electron affinity, but older texts mistakenly report that some elements such as alkaline earth metals have negative Eea, meaning they would repel electrons.[citation needed] This is not recognized by modern chemists. The electron affinities of the noble gases have not been conclusively measured, so they may or may not have slightly negative EAs. Atoms whose anions are relatively more stable than neutral atoms have a greater Eea. Chlorine most strongly attracts extra electrons; mercury most weakly attracts an extra electron. Eea of noble gases are close to 0.
Although Eea vary in a chaotic manner across the table, some patterns emerge. Generally, nonmetals have more positive Eea than metals.
Molecular electron affinities
Eea is not limited to the elements but also applies to molecules. 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 [2].
Electronegativity, symbol χ, is a chemical property that describes the ability of an atom (or, more rarely, a functional group) to attract electrons (or electron density) towards itself in a covalent bond.[1] First proposed by Linus Pauling in 1932 as a development of valence bond theory,[2] 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 Pauling. This gives a dimensionless quantity, commonly referred to as the Pauling scale, on a relative scale running from 0.7 to 4.0 (hydrogen = 2.2). 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 an atomic property, but rather a property of an atom in a molecule:[3] the equivalent property of a free atom is its electron affinity. It is to be expected that the electronegativity of an element will vary with its chemical environment,[4] 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.
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 π-electronegativities respectively. There are a number of linear free-energy relationships which 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.
Ionization Energy
The ionization energy is the minimum energy required to remove an electron from the ground state of the isolated gaseous atom
As stated in earlier reading, ionization energies (IE) have to do with things called ions. Ions are atoms which have gained or lost electrons. The ionization energy is the amount of energy it takes to detach one electron from a neutral atom. Some elements actually have several ionization energies. When this is the case, we refer to them as the "first ionization energy" or 'I', "second ionization energy" or 'I2', and so on. Notice that the energy variable follows Ii where i is the orbital from which the electron is lost. Ionization is endothermic meaning that the atom or molecule increases its internal energy (takes energy from an outside source). The equation for the first ionization energy is shown below:
Na --> Na+ + e-
The equation for the second ionization energy is:
Na+ --> Na2+ + e-
Ionization energy values are typically very high and follow trends throughout the periodic table. The IE increase from bottom to top and left to right in the periodic table.

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