Periodic Trends (College Board AP® Chemistry): Study Guide

Periodic Trends

• Periodicity refers to the recurring and predictable trends in the properties of elements as you move across the periodic table

• These trends can be understood through principles such as Coulomb's law, the shell model, and the concept of shielding/effective nuclear charge

Ionization Energy

• Ionization energy is the energy required to remove one mole of electrons in gaseous state from one mole of neutral gaseous atoms, forming one mole of positively charged gaseous ions

  • e.g. The equation below shows the first ionization energy of sodium (Na)

Na_(g)    →   Na⁺_(g)   +   e^-

• As you move across a period (from left to right), the ionization energy generally increases

  • There is an increase in the effective nuclear charge because the number of protons increase

  • Therefore, the coulombic attraction between the nucleus and the valence electrons is stronger, so it is harder to remove an electron

  • The shielding does not play an important role along the periods, since atoms in the same group have the same shielding effect because the number of inner shells is the same

  • The distance between the nucleus and the valence electrons is reasonably the same

A graph showing the ionization energies of the elements hydrogen to sodium

There is a general increase in ionization energy from left to right

• As you move down a group (from top to bottom), the ionization energy decreases

  • This occurs because the shielding effect increases because the number of inner shells increases

  • The distance between the nucleus and the valence electrons increases, making the Coulombic attraction considerably weaker

  • Therefore, the effective nuclear charge decreases and its easier to remove an electron

A Beryllium and Magnesium Atom

Comparison, between beryllium and magnesium, of the factors that affect the ionization energy 

Atomic Radius

• Atomic radius is the size of an atom

  • It can also be defined as the distance from the nucleus to the outermost electron

Atomic Radius

The atomic radius of an atom is the typical distance between the nucleus and the outermost electron shell

• Atomic radius and ionization energy have opposite trends

• As you move across a period (from left to right), the atomic radius generally decreases

  • There is an increase in the effective nuclear charge because the number of protons increase

  • Therefore, the Coulombic attraction between the nucleus and the valence electrons is stronger, which results in smaller atoms

  • The shielding does not play an important role along the periods, since atoms in the same group have the same shielding effect because the number of inner shells is the same

• As you move down a group, the atomic radius increases

  • This occurs because the shielding effect increases because the number of inner shells increases

  • The distance between the nucleus and the valence electrons increases, making the Coulombic attraction considerably weaker

  • Therefore, the effective nuclear charge decreases and the valence electrons are not pulled strongly, meaning bigger atoms

Atomic Radii in the Periodic Table

Trends in the atomic radii across a period and down a group

Ionic Radius

• Ionic radius is the size of an ion, which can be larger or smaller than the corresponding neutral atom

• Down a group it follows the same pattern as atomic radius

  • Ionic radii increase as energy shells are added, decreasing the effective nuclear charge and the Coulombic attraction from the nucleus

• However, there is a different periodic trend for positive ions (cations) and negative ions (anions)

• The ionic radius increase with an increasing negative charge

  • Negative ions are formed when electrons are accepted, while the nuclear charge is the same

  • Adding electrons creates extra repulsion with the other valence electrons, which means a bigger ionic radius

  • The greatest the negative charge, the largest the ionic radius

• The ionic radius decreases with an increasing positive charge

  • Positive ions are formed when electrons are lost, while the nuclear charge is the same

  • Removing electrons decreases the repulsion within the valence electrons, which means a smaller ionic radius

  • The greatest the positive charge, the smallest the ionic radius

 Sizes of Atoms and their Ions in pm

Trends in the ionic radii across a period and down a group

Electron Affinity

• Electron affinity is the energy change when one mole of electrons are added to one mole of a neutral atom in gaseous state, forming one mole of negatively charged ions

  • E.g. The equation below shows the first ionization energy of sodium (Na)
    

• Electron affinity is an exothermic process, this means that energy is released

  • When a process release energy, the sign of the energy is negative

• Elements on the Group 7 (F, Cl, Br, I) of the periodic table have the most negative electron affinities

  • This occurs because these elements tend to be stable as ions with a -1 charge, by completing the p subshell

• Electron affinity generally decreases down the group

  • This occurs because the shielding effect increases because the number of inner shells increases

  • The distance between the nucleus and the valence electrons increases, making the Coulombic attraction considerably weaker

  • Therefore, the effective nuclear charge decreases and the coulombic attraction to additional electrons is less

  • An exception to this rule is fluorine since an additional electron in the 2p subshell create a considerable repulsion with the other valence electrons

Electron Affinities

Electron affinities down Group 17 from F to I

Electronegativity

• Electronegativity measures the ability of an atom to attract a pair of electrons when it forms a covalent bond

• Electronegativity generally increases across a period (from left to right)

  • There is an increase in the effective nuclear charge because the number of protons increase

  • For that reason, the Coulombic attraction between the nucleus and the electrons from the covalent bond is stronger

  • The shielding does not play an important role along the periods, since atoms in the same group have the same shielding effect because the number of inner shells is the same

• Down a group, electronegativity decreases

  • This occurs because the shielding effect increases because the number of inner shells increases

  • The distance between the nucleus and the valence electrons increases, making the Coulombic attraction considerably weaker

  • Therefore, the effective nuclear charge decreases and the coulombic attraction to the electrons from a covalent bond decrease

• Electronegativity does not apply to noble gasses since they do not form covalent bonds

Trends in Electronegativity

Periodicity of electronegativity - there is an increase from left to right and a decrease from top to bottom