Orbitals and Periodic Trends

Definition: An orbital is a region of space within an electron subshell where an electron with a specific energy is most likely to be found.
Shells and subshells indicate the net energy of an electron, which can be known precisely.
Orbitals indicate the location of an electron, but not with precise accuracy; rather, they represent volumes where electrons are likely to be found.
This concept arises from the quantum mechanical model.
Chemists prioritize knowing the energy of an electron precisely over its exact location.

Four of the five d orbitals resemble a clover shape but are oriented in different planes.
The fifth d orbital looks like a p orbital with an inner tube around it.
F orbitals have complex shapes, described as resembling "sea monsters in horror movies."
There are three p orbitals, each on a different axis (px, py, pz), positioned 90 degrees apart.

Understanding the geometries of orbitals is crucial in chemistry for determining how atoms bond and the shapes of molecules.

In orbital diagrams, electrons are symbolically placed in orbitals without drawing the actual shapes of the orbitals.
The Aufbau principle states that electrons are added to the lowest energy subshells first.
Hund's rule states that every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin.

The lowest energy orbitals are on the left, and the highest energy orbitals are on the right.
Each line represents an orbital.
Labeling the orbitals is essential on tests.
Core electrons fill the lower energy orbitals.
Chromium prefers a half-filled d orbital.

Orbital diagrams can be drawn vertically with the lowest energy at the bottom and the highest energy at the top.
Copper prefers a completely filled d subshell.

Paramagnetic: Contains one or more unpaired electrons and is attracted to a magnetic field.
Diamagnetic: All electrons are paired, and it is not attracted to a magnetic field.

Chromium is paramagnetic because it has unpaired electrons.
Zinc is diamagnetic because all its electrons are paired.
Noble gases are diamagnetic.
Elements in group 2A and 2B are likely diamagnetic because they have filled s subshells.

The periodic table is based on electronic structure and properties.
The group number indicates the number of valence electrons but not for transition metals.

Metals are typically shiny, conductive, and malleable.
Metalloids (e.g., silicon) are semiconductors, conducting electricity under certain conditions and are next to the darker line in the periodic table.
Hydrogen is a nonmetal, while lithium is a metal.
Beryllium is a metal.
Examples of metalloids include silicon, germanium, and antimony.

Noble gases (Group 8A) such as helium, neon, and argon are stable and exist as single atoms.
Helium's valence shell is shell 1, neon's is shell 2, and argon's is shell 3.
Noble gases have stable electron configurations.
The octet rule refers to having eight valence electrons, except for helium, which has two.

Atoms form ions to obtain a noble gas configuration.
Group 1A elements form +1 cations, Group 2A form +2 cations, Group 6A form -2 anions, and Group 7A form -1 anions.

Metallic character increases as you go down a group and decreases as you go from left to right across a period.

Ionization energy is the energy required to remove an electron from an atom in the gas phase.
A(g) + energy \rightarrow A^+(g) + e^-
Metallic character usually corresponds with how easy it is to lose electrons

As you go down the group, generally it is easier to lose electrons. It decreases as you go down the group and increases as you go across the period.