Mastering Orbital Diagrams and Electron Configuration in Atoms

The video tutorial explains orbital diagrams, focusing on the Aufbau principle, Pauli Exclusion Principle, and Hund's Rule. It demonstrates how electrons are distributed in an atom's orbitals, using iron as an example. The tutorial also provides a method to remember the order of filling orbitals and discusses the rules governing electron placement.

What is the primary concept behind the Aufbau principle?

• Electrons fill orbitals based on atomic number.

• Electrons fill the highest energy orbitals first.

• Electrons fill the lowest energy orbitals first.

• Electrons fill orbitals randomly.

Which orbital is filled immediately after the 2s orbital according to the Aufbau diagram?

• 3p orbital

• 2p orbital

• 4s orbital

• 3s orbital

What is the maximum number of electrons that can occupy a single orbital according to the Pauli Exclusion Principle?

• Three

• Two

• Four

• One

According to Hund's Rule, how do electrons fill orbitals of equal energy?

• They pair up immediately.

• They fill one orbital completely before moving to the next.

• They fill orbitals based on atomic mass.

• They occupy all orbitals singly before pairing up.

Which principle explains why electrons in the same orbital must have opposite spins?

• Hund's Rule

• Bohr's Model

• Aufbau Principle

• Pauli Exclusion Principle

What is the main reason electrons prefer to occupy separate orbitals of the same energy before pairing up?

• They are positively charged.

• They are attracted to each other.

• They are negatively charged and repel each other.

• They are neutral.

How many electrons does iron have in its ground state?

• 24

• 27

• 25

• 26

In the orbital diagram for iron, which orbital is filled after the 3p orbital?

• 4p orbital

• 4s orbital

• 5s orbital

• 3d orbital

How many orbitals are present in the 3d sublevel?

• Four

• Three

• Six

• Five

What is the total number of electrons that can be accommodated in the 2p sublevel?

• Two

• Four

• Six

• Eight