Refresher: Atoms consist of a nucleus with positively charged protons and neutral neutrons surrounded by shells of electrons. Elements in the periodic table are organized into periods and groups. Periods run across the table horizontally, while groups are the vertical groupings. Elements in the same period have the same number of electron shells, while elements in the same group have the same number of valence electrons.
How do you find the atomic radius?
The atomic radius is calculated by measuring the distance between the nuclei of two identical atoms bonded together. Half this distance is the atomic radius.
Why does the atomic radius increase as you progress down a group?
As we progress down a group in the periodic table, the number of electrons increases, and so does the number of shells that those electrons are organized into. This increases the atomic radius as the electrons in the outermost shell are further away from the nucleus.
Why does the atomic radius decrease across a period?
Across a period, the number of electron shells remains the same, while the number of electrons increases. Thus, as the nuclear charge increases across a period, the protons in the nucleus can attract the higher number of electrons more closely due to the attraction of the positive protons to the negative electrons. This means the electrons are pulled more closely to the nucleus, reducing the size of the atomic radius.
Which element has the smallest atomic radius?
Helium has the smallest atomic radius at 31 picometers. Helium is in the top period and the farthest right group, which follows the patterns of atomic radius on the periodic table.
The atomic radius of a particular element is an important characteristic as it helps us to understand many properties of atoms and how they react. Atomic radius is the distance from the atom’s nucleus to the outermost electron orbital, and a lot of trends in the periodic table rely on this property due to its relationship to other atomic properties such as nuclear charge and shielding. Because the boundaries where the electron shells end on an atom can be a bit unclear, the actual definition of atomic radius is one-half the distance between the nuclei of two identical bonded atoms. The atomic radius is measured in picometers, which is one trillionth of a meter or 1x10-12.
As we progress down a group in the periodic table, the number of electrons increases, and so does the number of shells that those electrons are organized into. This increases the atomic radius as the electrons in the outermost shell are further away from the nucleus. The size of the nucleus also increases as you move down the group. Thus, as you move down the periodic table, both the size of the nucleus and the number of shells grows, increasing the total radius.
For example the atomic radius of Lithium is 152 picometers, but if we progress down to caesium, its atomic radius is 262 picometers. This is because caesium not only has a greater number of protons, but also 6 electron shells Due to the large nucleus and the large number of shells, the outer valence electron is much further away, meaning it’s atomic radius is larger.
Things are a little different when you look across the periods, though. While the number of protons and the nuclear size still does increase across a period, the atomic radius actually decreases. This has to do with the number of electron shells. Across a period, the number of electron shells remains the same, while the number of electrons increases. Thus, as the nuclear charge increases across a period, the protons in the nucleus can attract the higher number of electrons more closely due to the attraction of the positive protons to the negative electrons. This means the electrons are pulled more closely to the nucleus, reducing the size of the atomic radius.
For example, Sodium in period 3 has an atomic radius of 186 picometers and chlorine in the same period has an atomic radius of 99 picometers. This is because Chlorine has a larger number of protons and a higher nuclear charge, with no additional shells to put the electrons further away. This increased nuclear charge attracts the electrons more strongly to the nucleus, making the radius smaller.
