Atomic Radius is defined as half the distance between the nuclei of two atoms. But, this concept is complicated because not all atoms are usually bonded in the same manner.
Covalent bonds hold some in molecules—others by ionic crystals and others by metallic crystals. But, many elements can form covalent molecules in which two atoms of the same kind are bonded together by a single covalent bond.
These covalent radii of molecules are often referred to as atomic radii. The length of this distance is measured in picometers.
The size of atoms falls from left to right across a period of elements. All electrons are added to the same shell within a period or family of elements.
Also, protons are added to the nucleus, increasing its positive charge. The effect of raising the number of protons is larger than increasing the number of electrons. Hence, there is a greater nuclear attraction.
It indicates that the nucleus exerts a stronger attraction on the electrons, drawing the atom’s shell closer to the nucleus. The valence electrons are maintained closer to the atomic nucleus. So, the atomic radius decreases.
Atomic radii rise as you go down a group. Valence electrons now fill higher levels because of the rising quantum number (n). As ‘n’ grows, the valence electrons move away from the nucleus. An atomic radius is larger because shielding prevents these outer electrons from attaching to the nucleus.
X-rays or other types of spectroscopy methods measure the atomic radius of an atom. On the periodic table, the sizes of the atoms follow a set pattern. We can figure out why this is happening by looking at the charge and energy level of the nuclei.
As we move from left to right in a period, the atomic radius gets smaller. As we move down a group, it gets larger. The valence electrons are all in the same outer shell in periods.
In a move from left to right, the atomic number goes up in the same amount of time, making the effective nuclear charge go up. When attractive forces get stronger, the atomic radius of an element gets smaller.
It was amazing to see how the attraction between electrons and protons affects whether the atomic radius gets bigger or smaller.
First Atomic Radius Trend: (Decrease From Left to Right Across a Period)
The first periodic trend in atomic radius is a reduction in atomic size as one moves from left to right throughout a period. Each extra electron within a period of elements is added to the same shell.
When an electron is added to a nucleus, a new proton is added. It gives the nucleus a higher positive charge and a stronger nuclear force.
The nucleus acquires a higher positive charge, which attracts the electrons more and draws them closer to the atom’s nucleus. As electrons are closer to the nucleus, the atom’s radius shrinks.
It compared carbon (C) with an atomic number of 6 with the atomic number of 9 of fluorine (F). We can see that carbon (C) will be larger in diameter than fluorine (F) because the three extra protons in fluorine will pull its electrons closer to the nucleus and reduce the fluorine’s radius.
This is because of atomic number trends. As it turns out, the average atomic radius of fluorine is about half that of carbon, at 50 pm against 70 pm.
Second Atomic Radius Trend: (Increase as You Move Down a Group)
Another atomic radius periodic trend is that atomic radii rise as you move downward in a group in the periodic table. Each time you go down a group, the atom gets another electron shell. The atomic radius grows as each new shell moves out from the atom’s nucleus.
While you might expect valence electrons (those in the outermost shell) to draw the nucleus, electron shielding prevents this. When an atom contains more than one electron shell, electron shielding refers to a reduced attraction between outer electrons and the nucleus. As a result of electron shielding, valence electrons can’t reach as close to the atom’s center as they would like. consequence, the atom has a bigger radius.
As an example, potassium (K) has a greater average atomic radius than sodium (Na) does (220 pm) (180 pm). The potassium atom contains a more electron shell compared to the Na atom. It means that its valence electrons are away from the nucleus. It results in a greater atomic radius.
Exception of atomic radius trend
Exceptions include noble gases. In group 18 of the periodic table, six noble gases: were helium, neon, argon, krypton, xenon, and radon. These noble gases are an exception because they bind differently than other atoms.
Noble gas atoms do not get as close when they bond. The atomic radius is equal to half the distance between the nuclei of two atoms. The proximity of the atoms impacts the atomic radius.
Each noble gas fill the outermost electron shell. It indicates that Van der Waals forces rather than bonds hold many noble gas atoms together. Two atoms joined by Van der Waals forces don’t get as close together as linked atoms.
Because Van der Waals forces are weaker than covalent bonds, we can never get an empirical radius for any of the noble gases. As a result, their radii deviate from the atomic radius trends since their radii would be overestimated.
- From left to right in a period, the atomic radius gets smaller. The number of protons and electrons keeps going up across the period. One proton is more powerful than one electron, so electrons are attracted toward the nucleus. This makes the radius of the atom smaller.
- The atomic radius grows from the top to the bottom within a group. This happens because electrons get in the way.