Atomic Structure and Periodic Trends | General Chemistry 1

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: 1s² 2s² 2p⁴
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): 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d⁶ = [Ar] 4s² 3d⁶

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 (Z_eff):

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).
 

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) ⇒ Z_eff 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 ⇒ Z_eff decreases down a group.

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 (Z_eff) 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 (Z_eff) affects the attraction between the nucleus and the outer electrons. A higher Z_eff results in a stronger attraction, causing the outer electrons to be pulled closer to the nucleus ⇒ The atomic radius decreases as Z_eff 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 (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 (Z_eff): as Z_eff 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 (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 (Z_eff). 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.

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 ⇒ 1s² 2s² 2p⁶ 3s² 3p⁶ 3d¹⁰ 4s² 4p⁶ 4d¹⁰ 5s¹
  Ag tends to lose the e^- in the 5s orbital ⇒ Ag⁺ ⇒ 18 electrons in the outer shell.
• Ni: Z = 28 ⇒ 1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁸ 4s²
  Ni tends to lose the 2 e^- in the 4s orbital ⇒ Ni²⁺ ⇒ 16 electrons in the outer shell

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 ⇒ 1s² 2s² 2p⁶ 3s²; Mg²⁺ ⇒ 1s² 2s² 2p⁶ ⇒ 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: 1s² 2s² 2p⁶ 3s² 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^-