Periodic Trends of the Elements | General Chemistry 1

Periodic trends of the elements are studied in this chapter: effective nuclear charge, periodicity of atomic radius, ionization energy and electron affinity, electron configuration of ions, sizes of ions, isoelectronic series

Effective Nuclear Charge

Nuclear charge vs. effective nuclear charge

Nuclear charge (Z): number of protons in the nucleus of an atom

Effective nuclear charge (Zeff): actual magnitude of positive charge that is "experienced" by an electron in the atom. Electrons in an atom are simultaneously attracted to the positively charged nucleus and repelled by other negatively charged electrons ⇒ an electron in a many-electron atom is partially shielded from the positive charge of the nucleus by the other electrons, mainly the core electrons
 

Zeff = Z - σ

Zeff = effective nuclear charge
Z = nuclear charge
σ = shielding constant

 

Trend in the periodic table:

The effective nuclear charge increases across a period and decreases down a group

  • Atoms in the same period of the periodic table have the same number of core electrons (σ remains the same) but an increased number of protons (Z increases) ⇒ Zeff increases from left to right across a period
  • Although the nuclear charge increases down a group, the shielding effect more than counters its effect ⇒ Zeff decreases down a group

Periodicity in Properties of Elements

The properties of elements depend on their valence electrons. Elements of the same period have similar valence electron configurations and therefore behave in the same way during a chemical reaction


Atomic radius:

The distance between the nucleus and the outermost electrons

  • The higher the effective nuclear charge Zeff, the more attracted the outer electrons are to the nucleus and the closer they are to the it. Zeff increases across a period ⇒ the atomic radius decreases across a period
  • The principal quantum number n describes the size of the orbital: the larger n, the larger the orbital is ⇒ the atomic radius increases down a group

 

Ionization Energy (IE):

The energy required to remove an electron from a gaseous atom or ion

  • First ionization energy: the minimum energy required to remove an electron from a neutral gaseous atom X (g)
  • Second ionization energy:  the minimum energy required to remove an electron from a gaseous ion X+ (g)
     

First and second ionization energies of Mg:

Mg (g) → Mg+ (g) + e-     (first ionization)
Mg+ (g) → Mg2+ (g) + e-     (second ionization)

 

Trend in the periodic table: 

  • The closer an electron is to the nucleus, the more difficult it is to remove
  • The greater the Zeff, the more tightly bound an electron is to an atom ⇒ ionization energy follows the trend of Zeff: IE increases across a period and decreases down a group
  • Each additional electron is more difficult to remove (less e-- e- repulsion) ⇒ the second ionization energy is greater than the first ionization energy
  • Core electrons are very attracted by the nucleus ⇒ high IE ⇒ significant jump in ionization energies occurs after the outermost electrons are removed
     

Noble gases: great Zeff values and no valence electrons (very stable electronic structures) ⇒ high first ionization energy
Alkali metals: they want to lose one electron to have the same electronic configuration as the nearest noble gas ⇒ relatively low ionization energy

 

Electron Affinity (EA):

The amount of energy released when an electron is attached to a neutral atom or molecule in the gaseous state:
X (g) + e- → X(g)

  • The greater the Zeff, the higher the positive charge of the nucleus, the easier it is to add a negatively charged electron ⇒ electron affinity follows the trend of Zeff: EA increases across a period and decreases down a group

Electron Configuration of Ions

Main group elements ions:

The main group elements tend to lose or gain the number of electrons needed to get the same number of electrons as the nearest noble gas. Species with identical electron configuration are called isoelectronic species
 

F: 1s2 2s2 2p5   The nearest noble gas is neon Ne: 1s2 2s2 2p6
⇒ Fluorine tends to gain one electron to have the number of electrons as neon
⇒ F-: 1s2 2s2 2p6   (10 electrons total, isoelectronic with Ne)


How to write the electron configuration of a main group element ion:

  1. Write the electron configuration of the corresponding neutral atom
  2. Identify the nearest noble gas
  3. Add or remove the appropriate number of electrons

 

Transition metal ions:

An atom always first loses electrons from the shell with the highest value of n ⇒ when a transition metal becomes an ion, it first loses electrons from the ns subshell, then from the (n-1)d subshell 
Order of orbitals that lose electrons first: 1s < 2s < 2p < 3s < 3p < 3d < 4s < 4p < 4d < 5s

A relatively stable electron configuration for transition metals is one with 18 electrons in the outer shell. Electron configurations with 16 electrons or half-filled orbitals can also be observed
 

Ag: Z = 47 ⇒ 1s2s2p3s3p3d10 4s4p4d10 5s1
Ag tends to lose the e- in the 5s orbital ⇒ Ag+ ⇒ 18 electrons in the outer shell

Ni: Z = 28 ⇒ 1s2s2p3s3p3d8 4s2
Ni tends to lose the 2 e- in the 4s orbital ⇒ Ni2+ ⇒ 16 electrons in the outer shell

Sizes of Ions

Ionic radius:

The distance between the nucleus and the valence shell of a cation or anion

 

Ionic radius vs atomic radius

  • Cations have fewer electrons in their outer shell than the corresponding neutral parent atom ⇒ positive charge towards the nucleus ⇒ the ionic radius of cations is smaller than the atomic radius of the neutral parent atom
  • Anions have more electrons in their outer shell than the corresponding neutral parent atom ⇒ electron-electron repulsion increases ⇒ the ionic radius of anions is larger than the atomic radius of the neutral parent atom
     

Mg: Z = 12 ⇒ 1s2s2p3s2; Mg2+ ⇒ 1s2s2p6 ⇒ no more electrons in the shell n = 3
Mg2+ is smaller than Mg

 

Trend of the ionic radii for an isoelectronic series:

An isoelectronic series is a series of 2 or more species that have identical electron configurations, but different nuclear charges Z. The greater the nuclear charge Z , the greater the attraction between the nucleus and the electrons, and thus the smaller the radius
 

K+, Ar, Cl- have the same electron configuration: 1s2s2p3s2 3p6 ⇒ they form an isoelectronic series
Their nuclear charges are + 19 (K+), + 18 (Ar), and + 17 (Cl-) ⇒ In order of increasing radius: K+ < Ar < Cl-

Check your knowledge about this Chapter

The nuclear charge (Z) is the number of protons in the nucleus of an atom, while the effective nuclear charge (Zeff) is the actual magnitude of positive charge "experienced" by an electron in the atom. The effective nuclear charge experienced by an electron is also called the core charge

Zeff = Z - σ

Zeff = effective nuclear charge
Z = nuclear charge
σ = shielding constant

The effective nuclear charge increases across a period and decreases down a group

  • Atoms in the same period of the periodic table have the same number of core electrons (σ remains the same) but an increased number of protons (Z increases) ⇒ Zeff increases from left to right across a period
  • Although the nuclear charge increases down a group, the shielding effect more than counters its effect ⇒ Zeff decreases down a group

The atomic radius of an element is the distance between the nucleus and the outermost electrons

The higher the effective nuclear charge Zeff, the more attracted the outer electrons are to the nucleus and the closer they are to it. Zeff increases across a period ⇒ the atomic radius decreases across a period

The principal quantum number n describes the size of the orbital: the larger n, the larger the orbital is ⇒ the atomic radius increases down a group

Ionization energy (IE) is the energy required to remove an electron from a gaseous atom or ion

The greater the Zeff, the more tightly bound an electron is to an atom ⇒ Ionization energy follows the trend of Zeff: IE increases from left to right across the periodic table and decreases as we go down a group

The first ionization energy is the minimum energy required to remove an electron from a neutral gaseous atom X (g), while the second ionization energy is the minimum energy required to remove an electron from a gaseous ion X(g)

A cation has fewer electrons than the corresponding neutral atom and therefore less electron-electron repulsion: each additional electron is more difficult to remove à the second ionization energy is greater than the first ionization energy

Electron affinity (EA) is the amount of energy released when an electron is attached to a neutral atom or molecule in the gaseous state: X (g) + e- → X- (g)

The greater the Zeff, the higher the positive charge of the nucleus, the easier it is to add a negatively charged electron ⇒ electron affinity follows the trend of Zeff: EA increases across a period and decreases down a group

The main group elements tend to lose or gain the number of electrons needed to get the same number of electrons as the nearest noble gas

  1. Write the electron configuration of the corresponding neutral atom
  2. Identify the nearest noble gas
  3. Add or remove the appropriate number of electrons

Isoelectronic species are species whose electron configuration is identical

F- and Ne have identical electron configuration [1s2 2s2 2p6] ⇒ they are isoelectronic species

An atom always first loses electrons from the shell with the highest value of n ⇒ when a transition metal becomes an ion, it first loses electrons from the ns subshell, then from the (n-1)d subshell 

Order of orbitals that lose electrons first: 1s < 2s < 2p < 3s < 3p < 3d < 4s < 4p < 4d < 5s

Transition metals generally form positive ions due to the nature of the d-orbital electrons that are weakly bound to the transition metal nucleus

A relatively stable electron configuration for transition metals is one with 18 electrons in the outer shell. Electron configurations with 16 electrons or half-filled orbitals can also be observed

The ionic radius is the distance between the nucleus and the valence shell of a cation or anion

Cations have fewer electrons in their outer shell than the corresponding neutral parent atom ⇒ positive charge towards the nucleus ⇒ the ionic radius of cations is smaller than the atomic radius of the neutral parent atom

Anions have more electrons in their outer shell than the corresponding neutral parent atom ⇒ electron-electron repulsion increases ⇒ the ionic radius of anions is larger than the atomic radius of the neutral parent atom

An isoelectronic series is a series of 2 or more species that have identical electron configurations, but different nuclear charges Z. The greater the nuclear charge Z , the greater the attraction between the nucleus and the electrons, and thus the smaller the radius