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:

Atoms in the same period of the periodic table have the same number of core electrons ⇒ Zeff increases from left to right across a period because Z increases and σ remains the same

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 similarly 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 nucleus. Zeff increases across a period ⇒ atomic radius decreases across a period
  • The principal quantum number n designates the size of the orbital: the larger n, the larger the orbital ⇒ atomic radius increases across a group

 

Ionization Energy (IE):

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

  • The closer an electron is to the nucleus, the more difficult it is to remove
  • The greater the Zeff, the more tightly an electron is bonded to an atom ⇒ IE increases across the period, decreases across the group
  • Each additional electron is more difficult to remove (less e-- e- repulsion)
  • Core electrons are very attracted by the nucleus ⇒ high IE ⇒ significant jump in ionization energies occurs after the outermost electrons are removed


First ionization energy: the minimum energy required to remove an electron from a neutral gaseous atom A
A(g) → A+(g) + e-     (first ionization)

Second ionization energy:  the minimum energy required to remove an electron from a gaseous ion A+
A+(g) → A2+(g) + e-     (second ionization)

 

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:
A(g) + e- → A(g)

  • The greater the Zeff, the higher the positive charge of the nucleus, the easier to add a negatively charged electron ⇒ EA increases across the period, decreases across the group

Electron Configuration of Ions

Main group elements ions:

Main group elements tend to either lose or gain the number of electrons needed to achieve the same number of electrons as the nearest noble gas. Species with identical electron configurations are called isoelectronic
 

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 electron configuration of main group elements ions:

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

 

Transition metal ions:

An atom always loses electrons first from the shell with the highest value of n ⇒ when a transition metal becomes an ion, it loses electrons first from the ns subshell and 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: no (or less) electrons in the outer shell of 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: more electrons in the outer shell of the corresponding neutral parent atom ⇒ electron-electron repulsion increases ⇒ the ionic radius of anions is greater 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

 

Isoelectronic series:

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, 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-