Everything about Atomic Radius totally explained
Atomic radius, and more generally the
size of an atom, isn't a precisely defined
physical quantity, nor is it constant in all circumstances. The value assigned to the
radius of a particular
atom will always depend on the definition chosen for "atomic radius", and different definitions are more appropriate for different situations.
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.
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. However the electrons don't have definite positions—although they're more likely to be in certain regions than others—and the electron cloud doesn't have a sharp edge.
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.
Periodic trends in atomic radius
Atomic radius tends to increase as one proceeds down any group of the periodic table. This satisfies simple intuition: atoms with more electrons have larger radii. However as one proceeds across any row of the periodic table, a deeper intuition is required: atoms with more numerous electrons exhibit decreasing radius. 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.
| factor |
principle |
increase with... |
tend to |
effect |
| electron shells |
quantum mechanics |
Principal Quantum Number, Azimuthal Quantum Number |
atomic radius↑ |
increase on descending a group |
| nuclear charge |
attractive force acting on electrons by protons in nucleus |
atomic number |
atomic radius↓ |
decrease on passing along a period |
| shielding |
repulsive force acting on outermost shell electrons by inner electrons |
number of electron shells |
atomic radius↑ |
reduce the effect of the 2nd factor |
The increasing nuclear charge is partly counterbalanced by the increasing number of electrons in a phenomenon that's 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.
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's masked by a
relativistic effect known as the
inert pair effect.
d-Block contraction
The d-block contraction is less pronounced than the lanthanide contraction but arises from a similar cause. In this case, it's 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).
Empirically measured atomic radius
Empirically measured atomic radius
in
picometres (pm) to an accuracy of about 5 pm
Periodic table of empirically measured atomic radius
in picometres (pm) to an accuracy of about 5 pm
See also Periodic table
Reference: Bob Jones, J. Chem. Phys. 1964, 41, 3199.
Calculated atomic radius
Calculated atomic radius in picometres (pm)
Periodic table of calculated atomic radius
in picometres (pm)
See also Periodic table
Reference: E. Clementi, D.L.Raimondi, and W.P. Reinhardt, J. Chem. Phys. 1967, 47, 1300.
Further Information
Get more info on 'Atomic Radius'.
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