Atomic radius and electron affinity relationship

Periodic Trends - Chemistry LibreTexts

atomic radius and electron affinity relationship

When looking at the periodic table the atomic radius increases from top to bottom, moving down a column; therefore, the electron affinity. Generally speaking, there is a negative correlation between atomic radius and electron affinity. The atoms with the most electron affinity have valence shells very close to their nuclei, meaning the valence electrons feel a higher effective charge. How does an atomic size and an. Electron affinity is defined as the energy given off when one mole of atoms in the gaseous state each takes in one (or more) electrons to.

This means that an added electron is further away from the atom's nucleus compared with its position in the smaller atom. With a larger distance between the negatively-charged electron and the positively-charged nucleus, the force of attraction is relatively weaker. Therefore, electron affinity decreases. Moving from left to right across a period, atoms become smaller as the forces of attraction become stronger.

atomic radius and electron affinity relationship

This causes the electron to move closer to the nucleus, thus increasing the electron affinity from left to right across a period. Note Electron affinity increases from left to right within a period. This is caused by the decrease in atomic radius. Electron affinity decreases from top to bottom within a group.

This is caused by the increase in atomic radius. Atomic Radius Trends The atomic radius is one-half the distance between the nuclei of two atoms just like a radius is half the diameter of a circle. However, this idea is complicated by the fact that not all atoms are normally bound together in the same way.

atomic radius and electron affinity relationship

Some are bound by covalent bonds in molecules, some are attracted to each other in ionic crystals, and others are held in metallic crystals. Nevertheless, it is possible for a vast majority of elements to form covalent molecules in which two like atoms are held together by a single covalent bond.

This distance is measured in picometers. Atomic radius patterns are observed throughout the periodic table. Atomic size gradually decreases from left to right across a period of elements.

Periodic Trends

This is because, within a period or family of elements, all electrons are added to the same shell. However, at the same time, protons are being added to the nucleus, making it more positively charged. The effect of increasing proton number is greater than that of the increasing electron number; therefore, there is a greater nuclear attraction. This means that the nucleus attracts the electrons more strongly, pulling the atom's shell closer to the nucleus.

The valence electrons are held closer towards the nucleus of the atom. As a result, the atomic radius decreases. The valence electrons occupy higher levels due to the increasing quantum number n. Note Atomic radius decreases from left to right within a period. This is caused by the increase in the number of protons and electrons across a period. Atomic radius increases from top to bottom within a group. This is caused by electron shielding.

  • Electron affinity
  • What is the relationship between atomic radius and ionization energy?
  • 9.9: Periodic Trends: Atomic Size, Ionization Energy, and Metallic Character

Melting Point Trends The melting points is the amount of energy required to break a bond s to change the solid phase of a substance to a liquid. Because temperature is directly proportional to energy, a high bond dissociation energy correlates to a high temperature. Melting points are varied and do not generally form a distinguishable trend across the periodic table. However, certain conclusions can be drawn from the graph below.

Metals generally possess a high melting point. Most non-metals possess low melting points. The non-metal carbon possesses the highest boiling point of all the elements. The semi-metal boron also possesses a high melting point. Chart of Melting Points of Various Elements Metallic Character Trends The metallic character of an element can be defined as how readily an atom can lose an electron.

From right to left across a period, metallic character increases because the attraction between valence electron and the nucleus is weaker, enabling an easier loss of electrons. Metallic character increases as you move down a group because the atomic size is increasing. When the atomic size increases, the outer shells are farther away. The principal quantum number increases and average electron density moves farther from nucleus.

atomic radius and electron affinity relationship

Note Metallic characteristics decrease from left to right across a period. Electron affinities of the elements[ edit ] Electron affinity Eea vs atomic number Z.

Ionization Energy Electron Affinity Atomic Radius Ionic Radii Electronegativity Metallic Character

Note the sign convention explanation in the previous section. Electron affinity data page Although Eea varies greatly across the periodic table, some patterns emerge. Generally, nonmetals have more positive Eea than metals. Atoms whose anions are more stable than neutral atoms have a greater Eea. Chlorine most strongly attracts extra electrons; mercury most weakly attracts an extra electron.

Electron affinity - Wikipedia

The electron affinities of the noble gases have not been conclusively measured, so they may or may not have slightly negative values. Eea generally increases across a period row in the periodic table prior to reaching group This is caused by the filling of the valence shell of the atom; a group 17 atom releases more energy than a group 1 atom on gaining an electron because it obtains a filled valence shell and therefore is more stable.

In group 18, the valence shell is full, meaning that added electrons are unstable, tending to be ejected very quickly. Counterintuitively, Eea does not decrease when progressing down the rows of the periodic table, as can be clearly seen in the group 2 data.