Organic Chemistry | Purdue University | Lecture 2
Difficulty: Introductory Prerequisites: Basic understanding of atomic structure (protons, neutrons, electrons) and the periodic table.
This material is foundational for the entire organic chemistry course. Before you can understand how molecules form, react, or behave, you need to know how electrons are arranged around atoms and why atoms bond in the first place. Electronic configuration tells you where the electrons sit; the octet rule tells you what arrangement atoms are trying to reach. If you missed general chemistry or need a refresher on atomic structure, start there. Everything in organic chemistry, from drawing Lewis structures to predicting reaction mechanisms, builds on what is covered here.
Electrons fill orbitals in a specific order (aufbau principle) and each orbital holds at most two electrons with opposite spins (Pauli exclusion principle). Atoms are most stable when their outer shell has eight electrons (the octet rule), which is why they form ions or share electrons through covalent bonds.
Electronic configuration (electron configuration)
The notation describing how electrons are distributed among an atom's orbitals and energy levels, e.g. carbon is 1s² 2s² 2p². In simple terms, it is the address system for every electron in an atom.
Aufbau principle
Electrons fill orbitals starting from the lowest energy level and work upward. Think of it as filling seats in a theatre from the front row back.
Pauli exclusion principle
No two electrons in the same atom can have the same set of four quantum numbers. In practice, each atomic orbital can hold a maximum of two electrons, and they must have opposite spins (spin up and spin down). In simple terms, each orbital is a seat for two, and the two occupants must face opposite directions.
Core electrons
The inner-shell electrons that are not involved in bonding. For chlorine (1s² 2s² 2p⁶ 3s² 3p⁵), the core electrons are 1s² 2s² 2p⁶.
Valence shell (valence electrons)
The outermost occupied electron shell of an atom. These are the electrons that participate in chemical bonding. Think of it as the outer ring of electrons that other atoms can actually interact with.
Octet rule
Atoms tend to gain, lose, or share electrons until they have eight electrons in their valence shell, matching the electron configuration of the nearest noble gas. In simple terms, atoms want a full outer shell of eight, and they will bond or ionise to get there.
Noble gas configuration
An electron arrangement matching that of a noble gas (He, Ne, Ar, etc.), with a completely filled valence shell. This is the most stable arrangement.
Ionisation
The process of an atom gaining or losing electrons to form an ion. For example, lithium (1s² 2s¹) loses one electron to become Li⁺ with a 1s² configuration (same as helium). Think of it as an atom shedding or picking up electrons to reach a noble gas configuration.
Electrons fill orbitals in order of increasing energy: 1s → 2s → 2p → 3s → 3p → 4s → 3d (the aufbau principle).
Each orbital holds a maximum of two electrons with opposite spins (Pauli exclusion principle).
Example: carbon (⁶C) has six electrons, written as 1s² 2s² 2p².
Example: chlorine (¹⁷Cl) is 1s² 2s² 2p⁶ 3s² 3p⁵.
Core electrons (fully filled inner shells): 1s² 2s² 2p⁶
Valence shell: 3s² 3p⁵ (seven valence electrons)
The most stable electron configuration has eight electrons in the valence shell.
This matches the configuration of a noble gas (He, Ne, Ar, Kr, etc.).
Noble gases are inert (unreactive) precisely because their valence shells are already full.
Example: neon (¹⁰Ne) has the configuration 1s² 2s² 2p⁶, giving it a full valence shell of eight electrons.
There are two main routes:
1. Ionisation
An atom loses or gains electrons to match the nearest noble gas.
Example: lithium (¹³Li) has configuration 1s² 2s¹. It loses one electron to become Li⁺, leaving 1s² (the same configuration as helium).
Neutral lithium (Li⁰) loses an electron to become the Li⁺ ion.
2. Covalent bonding
Atoms share electrons so that each atom can count the shared pair toward its own octet.
Covered in detail in the companion notes on covalent bonding.
Students often think the octet rule applies universally. It does not. Hydrogen only needs two electrons (a duet), and elements in period 3 and below (like phosphorus and sulphur) can exceed eight electrons by using d-orbitals.
Confusing core electrons with valence electrons is a common error. Only the outermost shell electrons are valence electrons. For chlorine, the seven electrons in 3s² 3p⁵ are valence electrons, not the ten underneath.
Some students assume that losing electrons always makes an atom "smaller." In fact, cations (positive ions) are always smaller than the neutral atom because the remaining electrons are held more tightly by the same nuclear charge.
The octet rule describes a tendency, not a law. Many stable molecules violate it (e.g. BF₃ with six electrons on boron, or SF₆ with twelve on sulphur).
⚠️ Writing electron configurations quickly and correctly is a baseline skill. Expect it on early exams and as a building block for Lewis structures and molecular orbital theory later.
⚠️ Know the difference between core and valence electrons for any given element. Questions often ask you to identify the valence shell.
⚠️ Be able to explain why an atom ionises to reach a noble gas configuration and write the resulting ion's electron configuration.
⚠️ The octet rule is the foundation for drawing Lewis structures, which dominate the next several lectures.
True or False: The aufbau principle states that electrons fill higher-energy orbitals before lower-energy ones.
False. Electrons fill the lowest-energy orbitals first.
Fill in the blank: Each atomic orbital can hold a maximum of ____ electrons.
Two.
True or False: Neon is unreactive because it has eight electrons in its valence shell.
True.
Fill in the blank: Lithium achieves a noble gas configuration by losing ____ electron(s) to become Li⁺.
One.
True or False: Core electrons are the electrons that participate in bonding.
False. Valence electrons participate in bonding.
Q: Write the full electronic configuration for chlorine (Z = 17). Identify the core and valence electrons.
A: 1s² 2s² 2p⁶ 3s² 3p⁵. Core electrons: 1s² 2s² 2p⁶ (10 electrons). Valence electrons: 3s² 3p⁵ (7 electrons).
Q: Why is the noble gas configuration considered the most stable arrangement?
A: Noble gases have a completely filled valence shell (eight electrons, or two for helium). A full valence shell means there is no energetic drive to gain, lose, or share electrons, so noble gases are inert.
Q: Lithium has the configuration 1s² 2s¹. Explain how it achieves a noble gas configuration and write the resulting species.
A: Lithium loses its single 2s electron to form Li⁺, which has the configuration 1s². This matches the electron configuration of helium, a noble gas.
Q: State the Pauli exclusion principle and explain its consequence for orbital filling.
A: No two electrons in the same atom can share the same set of four quantum numbers. The practical consequence is that each orbital holds at most two electrons, and those two must have opposite spins (one spin-up, one spin-down).
Q: A student writes the configuration of carbon as 1s² 2s² 2p⁴. What is wrong?
A: Carbon has six electrons. 1s² 2s² 2p⁴ accounts for eight electrons (that configuration belongs to oxygen, Z = 8). The correct configuration for carbon is 1s² 2s² 2p².
This connects directly to covalent bonding (Lecture 2, Part 2), where atoms share electrons to satisfy the octet rule rather than losing or gaining them outright. Understanding valence electrons is essential for drawing Lewis structures, which you will use throughout organic chemistry to visualise molecular geometry and predict reactivity. Later in the course, molecular orbital theory builds on the same orbital-filling principles covered here, extending them to bonding and antibonding orbitals between atoms.
electron configuration, aufbau principle, Pauli exclusion principle, orbital filling order, core electrons, valence electrons, valence shell, octet rule, noble gas configuration, ionisation, cation, anion, Lewis dot structure, energy levels, atomic orbitals, 1s 2s 2p, spin up spin down, inert gases, stable electron arrangement, organic chemistry review, Purdue organic chemistry