Atomic Structure and Periodic Trends | General Chemistry 1

Atomic structure and periodic trends of the elements are studied in this chapter: core and valence electrons, effective nuclear charge, periodicity of atomic radius, ionization energy, and electron affinity, electron configuration of ions, ionic radius, isoelectronic series.

Core and Valence Electrons

Core electrons: 

Core electrons are the electrons located in the inner energy levels of an atom. They are tightly bound to the nucleus and do not participate in chemical bonding. They form the atomic core along with the nucleus and play a role in shielding the valence electrons from the full positive charge of the nucleus.
 

Valence electrons: 

Valence electrons are the electrons located in the outermost occupied shell of an atom. They are farther away from the positive charge of the nucleus and are therefore involved in chemical bonding. 

Valence electrons determine the chemical properties and reactivity of an atom ⇒ Elements of the same period have similar valence electron configurations and therefore behave in the same way during a chemical reaction.
 

Number of valence electrons in an oxygen atom:

Electron configuration of oxygen: 1s2s2p4
Outermost occupied shell: n = 2. There are 2 e- in 2s, 4 e- in 2p ⇒ 6 valence electrons

 

Abbreviated electron configuration:

For the electron configuration, we can use an abbreviated form by replacing the electronic configuration of the core electrons with [previous nearest noble gas].
 

Iron (Z = 26): 1s2s2p3s3p6 4s3d6 = [Ar] 4s3d6

Effective Nuclear Charge

Nuclear charge vs. effective nuclear charge:

  • Nuclear charge (Z):

The nuclear charge is the number of protons in the nucleus of an atom. It determines the overall positive charge of the nucleus.

  • Effective nuclear charge (Zeff):

The effective nuclear charge is the net positive charge experienced by an electron in an atom, taking into account shielding effects from other electrons. Electrons in an atom are attracted to the positively charged nucleus, but are also 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 (Z) increases down a group due to the addition of more protons, the shielding effect also increases. The addition of new electron shells and increased distance from the nucleus results in greater shielding (σ), which more than compensates for the increase in nuclear charge ⇒ Zeff decreases down a group.

Atomic Radius

Atomic radius:

The atomic radius is the distance between the nucleus of an atom and its outermost electrons. It is mainly influenced by the effective nuclear charge (Zeff) and the principal quantum number (n).
 

Trend in the periodic table:

The atomic radius decreases across a period and increases down a group:

  • The effective nuclear charge (Zeff) affects the attraction between the nucleus and the outer electrons. A higher Zeff results in a stronger attraction, causing the outer electrons to be pulled closer to the nucleus ⇒ The atomic radius decreases as Zeff increases across a period.
  • The principal quantum number (n) describes the size of the orbital ⇒ As you move down a group, the value of n increases, leading to larger orbitals and consequently larger atomic radii.

Ionization Energy

Ionization Energy (IE):

Ionization energy is the energy required to remove an electron from a gaseous atom or ion: X (g) → X(g) + e-

The first ionization energy is the minimum energy required to remove an electron from a neutral gaseous atom X (g). The second ionization energy is the minimum energy required to remove an electron from a gaseous cation X+ (g), and so on.
 

The 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:

In general, the trend of ionization energy follows the effective nuclear charge (Zeff): as Zeff increases, electrons are more tightly bound to the nucleus ⇒ Ionization energy increases across a period and decreases down a group.

The general trend of increasing ionization energy across a period may encounter interruptions, which can be explained by the electron configuration of the atom.

 

Other trends in ionization energy values:

  • Core Electrons: Core electrons, which are closer to the nucleus, experience stronger attraction and have higher ionization energies compared to valence electrons ⇒ significant jump in ionization energies occurs after the outermost electrons are removed.

  • Second ionization energy: Each additional electron removed from an ion requires more energy than the previous one due to the increased electrostatic attraction between the remaining electrons and the positively charged nucleus ⇒ the second ionization energy is greater than the first ionization energy.
  • Noble gases: Noble gases have high first ionization energies due to their stable electron configurations and strong effective nuclear charges.

 

Alkali metals: they want to lose one electron to have the same electron configuration as the nearest noble gas ⇒ relatively low ionization energy

Electron Affinity

Electron Affinity (EA):

Electron affinity is the energy released when an electron is added to a neutral atom or molecule in the gaseous state, forming a negatively charged ion: X (g) + e- → X(g). 

The more negative the EA, the greater the tendency of the atom to accept an electron and the more stable the anion that results.

 

Trend in the periodic table:

Similar to ionization energy, electron affinity is influenced by the effective nuclear charge (Zeff). The greater the Zeff, which reflects a higher positive charge of the nucleus, the more readily the atom accepts an additional negatively charged electron ⇒ Electron affinity follows the trend of Zeff: in general, it increases across a period and decreases down a group.

There are periodic interruptions of the upward trend of EA across a period, similar to those observed for IE, although they do not occur for the same elements.

Electron Configuration of Ions

Ions of main group elements:

  • Main group elements tend to lose or gain electrons to match their electron configuration to that of the nearest noble gas.
  • Species with identical electron configurations 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 become isoelectronic with neon.
⇒ F-: 1s2 2s2 2p6

 

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 loses electrons first from the shell with the highest value of n ⇒ transition metal loses electrons from the ns subshell, followed by the (n-1)d subshell. 
  • Order of orbitals losing 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. Configurations with 16 electrons or half-filled orbitals can also occur.

 

  • 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

Ionic Radius

Ionic radius:

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

 

Comparison with atomic radius:

  • Cations have fewer electrons in their outer shell than the corresponding neutral parent atom ⇒ positive charge towards the nucleus leads to a smaller ionic radius for a cation compared to the neutral parent atom.
  • Anions have more electrons in their outer shell than the corresponding neutral parent atom ⇒ the electron-electron repulsion increases, resulting in a larger ionic radius for an anion compared to the neutral parent atom.

 

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

 

Trend in isoelectronic series:

Isoelectronic series consist of species with identical electron configurations but different nuclear charges (Z). A higher nuclear charge leads to a stronger attraction between the nucleus and the electrons, resulting in smaller radii.
 

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-

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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