Which effect on atomic size is more significant

Previous Video 8. However, orbitals do not describe a confined space, but rather the statistical probability of where an electron can be found. So how is the atomic size defined, and what influences it? An atomic radius can be described in two ways. Nonbonding atomic radius, or van der Waals radius of an atom, is one-half of the distance between adjacent nuclei in the atomic solid.

Conversely, a bonding atomic radius, or covalent radius, distinguishes between metals and nonmetals. In metals, the radius is described for atoms in their crystal structure as one-half of the distance between the centers of two neighboring atoms.

In nonmetal, diatomic molecules, the radius is described as one-half of the distance between the centers of bonded atoms. The periodic table depicts variations in covalent radii that are often called atomic radii, which are influenced by two factors; the number of principal energy levels of valence electrons, and the effective nuclear charge.

Moving down a group, the principal quantum number, n , increases by one for each element. Thus, as outer electrons get farther from the nucleus, the atomic radius increases down the group. For example, moving down group 1, the atomic radius increases from lithium to cesium. This trend is demonstrated by the entire periodic table.

Further, the plot reveals that the atomic radius is maximum for each alkali metal and falls to a minimum with each noble gas across the period. The decreasing atomic radii across a period can be explained by the effective nuclear charge. Recall the concept of an effective nuclear charge. In any multi-electron atom, the inner shell electrons partially shield the outer shell electrons from the pull of the nucleus.

Thus, the effective nuclear charge, the charge felt by an outer electron is lesser than the actual nuclear charge. Electrons in the same valence shell do not shield one another very effectively. Across the period, the nuclear charge increases while the number of inner shell electrons remains constant.

Thus, as the effective nuclear charge increases steadily, the shielding of outer electrons becomes less, and this leads to a decrease in atomic radii. The radii of most transition elements, however, stay roughly constant across each row. This is because the number of electrons in the outermost principal energy level is nearly constant. The elements in groups of the periodic table exhibit similar chemical behavior.

This similarity occurs because the members of a group have the same number and distribution of electrons in their valence shells. Going across a period from left to right, a proton is added to the nucleus and an electron to the valence shell with each successive element. Going down the elements in a group, the number of electrons in the valence shell remains constant, but the principal quantum number increases by one each time.

An understanding of the electronic structure of the elements allows us to examine some of the properties that govern their chemical behavior. These properties vary periodically as the electronic structure of the elements changes. The quantum mechanical picture makes it difficult to establish a definite size of an atom. However, there are several practical ways to define the radius of atoms and, thus, to determine their relative sizes that give roughly similar values.

The atomic radius in metals is one-half of the distance between the centers of two neighboring atoms. It is one-half of the distance between the centers of bonded atoms for elements that exist as diatomic molecules. Moving across a period from left to right, generally, each element has a smaller atomic radius than the element preceding it.

This might seem counterintuitive because it implies that atoms with more electrons have a smaller atomic radius. This can be explained based on the concept of an effective nuclear charge. Thus, the effective nuclear charge, the charge felt by an electron, is lesser than the actual nuclear charge Z and can be estimated by the following:.

Atomic radii vary in a predictable manner across the periodic table. Radii generally decrease along each period row of the table from left to right and increase down each group column.

These trends in atomic radii as well as trends in various other chemical and physical properties of the elements can be explained by considering the structure of the atom. As the atomic number increases along each row of the periodic table, the additional electrons go into the same outermost principal energy level also known as valence level. This can be predicted to lead to.

Experiments have shown that the first case is what happens: the increase in nuclear charge overcomes the repulsion between the additional electrons in the valence level.

Therefore, the size of atoms decreases as one moves across a period from left to right in the periodic table. The principal energy levels hold electrons at increasing radii from the nucleus. Metals also form basic oxides; the more basic the oxide, the higher the metallic character. As you move across the table from left to right, the metallic character decreases, because the elements easily accept electrons to fill their valance shells.

Therefore, these elements take on the nonmetallic character of forming anions. As you move up the table, the metallic character decreases, due to the greater pull that the nucleus has on the outer electrons. This greater pull makes it harder for the atoms to lose electrons and form cations.

Uses in knowing the Periodic Properties of Elements. Certain properties—nota bly atomic radius, ionization energies, and electron affinities - can be qualitatively understood by the positions of the elements on the periodic table. Atomic size is the distance from the nucleus to the valence shell where the valence electrons are located. Atomic radius is a more definite and measurable way of defining atomic size.

It is the distance from the center of one atom to the center of another atom in a homonuclear diatomic molecule. There are three factors that help in the prediction of the trends in the Periodic Table: number of protons in the nucleus, number of shells, and shielding effect.

The atomic size increases from the top to the bottom in any group as a result of increases in all of the three factors. As the number of energy levels increases, the size must increase. Going across a period from left to right , the number of protons increases and therefore the nuclear charge increases.

Going across a period, the number of electron energy levels remains the same but the number of electrons increases within these energy levels. Therefore the electrons are pulled in closer to the nucleus. Ionization energy is the energy required to remove the most loosely held electron from a gaseous atom or ion. Ionization energy generally increases across a period and decreases down a group.

The effective nuclear charge is the charge of the nucleus felt by the valence electron. Electron affinity is the energy required or released when an electron is added to a gaseous atom or ion.

Electron affinity generally increases going up a group and increases left to right across a period. Non-metals tend to have the highest electron affinities. Learning Objectives Be able to state how certain properties of atoms vary based on their relative position on the periodic table. Si or S S or Te Solution Si is to the left of S on the periodic table, so it is larger because as you go across the row, the atoms get smaller.

S is above Te on the periodic table, so Te is larger because as you go down the column, the atoms get larger.