An Introduction to Electrons

The Exclusion Principle
Simply states that no two electrons in a single atom can have the same quantum numbers.

The Effect of Nuclear Charge (Z) on Orbital Energy
Greater nuclear charge lowers orbital energy (more negative number), which makes the electron harder to remove.

Shielding: The Effect of Electron Repulsion on Orbital Energy
Within an atom, the electrons "feel" an attraction to the nucleus and they also feel the repulsion of other electrons. These repulsive "feeling" effect make it easier to remove the electrons and thus change the nuclear charge to an effective nuclear charge (Zeff).

Penetration: The Effect of Orbital Shape on Orbital Energy
When an electron instantaneously moves very close to the nucleus for a small period of time thus increasing its radial probability and thus lowering the energy level. This explains why the 2s orbital has a lower energy than the 2d orbital.

Orbital Diagrams
Orbital diagrams are pictorial representations of the electrons in their different energy levels. Alternatively, one can instead write out the entire electron configuration with the shorthand notation.

Note: Copper (Cu) and Chromium (Cr) are exception to the standard trend.
Note: Due to shielding and penetration, the 4s orbital has a lower energy level than the 3d and hence comes before it; the 3d orbital is then followed by the 4p orbital.

Categories of Electrons
The elements have three categories of electrons:
1. Inner (core) electrons: those seen in the previous noble gas and any completed transition series. They fill all the lower energy levels of an atom.
2. Outer electrons: are those in the highest energy level (highest n value). They spend most of the time farthest from the nucleus.
3. Valence electrons: are those involved in forming compounds. Among the main group elements, the valence electrons are the outer electrons. Among the transition elements, the (n-1)d electrons are counted among the valence electrons because some or all of them are often involved in bonding.

Group and Period Numbers
1. Among the main-group elements (A group), the group number equals the number of outer electrons.
2. The period number is the "n" value of the highest energy level.
3. The "n" value squared (n2) gives the total number of orbitals in that energy level.

Trends in Atomic Size
Influences:
1. As the principle quantum number "n" increases, the probability that the outer electron will spend more time father from the nucleus increases as well; thus, the atoms are larger.
2. As the effective nuclear charge (Zeff) increases, the positive charge "felt" by an electron increases and the outer electrons are pulled closer to the nucleus; thus the atoms are smaller (from left to right).

Down a Column:
1. Down a group "n" dominates. As we move down a main group, each member has one more level of inner electrons that shield the outer electrons very effectively. Atomic radius generally increases in a group from top to bottom.
2. Across a period, Zeff dominates. As we move across a period main-group elements, electrons are added to the same outer level, so the shielding by inner electrons does not change. Atomic radius generally decreases in a period from left to right.

Trends in Ionization Energy
As the electron orbit decreases in size it becomes harder and harder to remove an electron and consequently requires more and more energy. The typical trends are as follows:
1. Down a group: as we move down a main group, the orbital's "n" value increases and so does atomic size. As the distance from the nucleus to outermost electron increases, the attraction between them lessens, which makes the electron easier to remove. Ionization energy generally decreases down a group.
2. Across a period. As we move left to right across a period, the orbital's "n" value stays the same, so Zeff increases and atomic size decreases. As a result the ionization energy increases.

Note: B, Al, O and S elements are misfits from the general trends.
Note: When there is a great jump in the ionization energy required to remove an electron it simply means that you are trying to remove an inner or core electron (all valence electrons have been removed).

Electron Affinity
Electron affinity is the energy change (kJ) accompanying the addition of 1mol of electrons to 1mol of gaseous atoms or ions.
In most cases, energy is released when the first electron is added because it is attracted to the atom's nuclear charge. The second electron affinity, on the other hand, is always positive because energy must be absorbed in order to overcome electrostatic repulsions and add another electron to a negative ion.

Magnetic Properties
Paramagnetism: it is attracted by an external magnetic field and is caused by having many unpaired electrons.
Diamagnetism: it is not attracted (and, in fact, is slightly repelled) by a magnetic field.
Cations: are smaller than their parent atoms. When a cation forms, electrons are removed from the out level. The resulting decrease in electron repulsion allows the nucleus charge to pull the remaining electrons closer.
Anions: are larger than their parent atoms. When an anion forms, electrons are added to the outer level. The increase in repulsion causes the electrons to occupy more space.