Source: Chapter 1, Dr Uyeda lecture slides
Tags: electron configuration, electronegativity, ionic bond, covalent bond, polar covalent, octet rule, Pauling scale, valence electrons, shells, orbitals
Difficulty: Introductory Prerequisites: Basic understanding of the periodic table and atomic number.
This is the opening chapter of organic chemistry, but it begins by revisiting general chemistry foundations: how electrons are arranged in atoms and how atoms form bonds. If you are shaky on electron configurations or electronegativity, this is where to shore that up, because every reaction mechanism in later chapters depends on knowing where electrons sit and how willing they are to move. You should already be comfortable reading the periodic table and know what an element's atomic number represents.
Atoms bond by either transferring or sharing electrons to reach a stable, noble-gas-like configuration. The type of bond (ionic, polar covalent, or non-polar covalent) depends on the electronegativity difference between the two atoms. Electrons fill orbitals in a predictable order governed by three rules, and only the outermost (valence) electrons participate in bonding.
Electron shell
An energy level surrounding the nucleus, numbered 1, 2, 3, etc. Each shell holds up to 2n² electrons, where n is the shell number. Think of it as a tier of seats around a stadium: higher tiers are farther out and hold more people.
Orbital
A region of space where an electron is most likely to be found. Orbitals come in types (s, p, d, f) with different shapes and capacities. In simple terms, this is the electron's address within its shell.
s orbital
A spherical orbital. Each shell has one. It holds a maximum of 2 electrons.
p orbital
A dumbbell-shaped orbital with two lobes of opposite phase separated by a node. Each shell from n = 2 onward has three p orbitals (px, py, pz), holding up to 6 electrons total.
Aufbau principle
Electrons fill orbitals starting from the lowest energy level and work upward. Think of it as filling a building from the ground floor up.
Pauli exclusion principle
No more than two electrons can occupy a single orbital, and they must have opposite spins (spin-paired).
Hund's rule
When multiple orbitals of the same energy are available, electrons fill them singly (with parallel spins) before any orbital gets a second electron. In simple terms, electrons spread out before doubling up.
Valence electrons
The electrons in the outermost shell of an atom. These are the only electrons involved in bonding.
Electronegativity
A measure of how strongly an atom attracts shared electrons in a bond. Measured on the Pauling scale. Fluorine is the most electronegative element (4.0).
Ionic bond
A bond formed by the complete transfer of one or more electrons from one atom to another. Occurs when the electronegativity difference (Δ) is greater than 1.9.
Covalent bond
A bond formed by the sharing of electrons between two atoms. In simple terms, both atoms hold on to the electrons together rather than one taking them entirely.
Polar covalent bond
A covalent bond in which electrons are shared unequally because the two atoms have different electronegativities (Δ between 0.5 and 1.9). This creates a bond dipole.
Non-polar covalent bond
A covalent bond in which electrons are shared equally or nearly so (Δ less than 0.5).
Octet rule
Atoms tend to gain, lose, or share electrons until they have eight electrons in their valence shell, mimicking a noble gas configuration. Hydrogen is the exception: it aims for two (a duet).
Atoms have a nucleus (protons + neutrons) surrounded by electrons.
Almost all of chemistry (except nuclear chemistry) is about moving electrons around without altering the nucleus.
Electrons occupy numbered shells (n = 1, 2, 3 …). Higher shells are farther from the nucleus and higher in energy.
Each shell holds a maximum of 2n² electrons: shell 1 holds 2, shell 2 holds 8, shell 3 holds 18.
Within each shell, electrons occupy orbitals of different types:
s orbitals: spherical, one per shell.
p orbitals: dumbbell-shaped, three per shell (from n = 2 onward), oriented along the x, y, and z axes.
d orbitals: five per shell (from n = 3 onward). These appear in later chapters and in transition-metal chemistry.
Orbitals in higher shells are larger and more diffuse (the electron cloud spreads out further from the nucleus).
Rule 1, Aufbau principle: fill from lowest energy to highest. The filling order is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, and so on.
Rule 2, Pauli exclusion principle: each orbital holds at most 2 electrons, which must be spin-paired (opposite spins).
Rule 3, Hund's rule: when orbitals of equal energy are available (e.g. the three 2p orbitals), place one electron in each before pairing any.
Atom | Electrons | Configuration |
|---|---|---|
H | 1 | 1s¹ |
He | 2 | 1s² |
Li | 3 | 1s² 2s¹ |
Be | 4 | 1s² 2s² |
B | 5 | 1s² 2s² 2p¹ |
C | 6 | 1s² 2s² 2p² |
N | 7 | 1s² 2s² 2p³ |
O | 8 | 1s² 2s² 2p⁴ |
F | 9 | 1s² 2s² 2p⁵ |
Ne | 10 | 1s² 2s² 2p⁶ |
Inner (core) electrons are chemically inert. Valence electrons (those in the outermost shell) do all the bonding work.
Carbon, for example, has 6 total electrons. Its configuration is 1s² 2s² 2p². The 1s² electrons are core; the 2s² 2p² electrons (4 total) are valence.
Shorthand notation uses a noble gas core in brackets: carbon = [He] 2s² 2p².
Lewis dot structures show only the valence electrons as dots arranged around the element symbol.
Carbon has four dots (one on each side), representing its four valence electrons. This model, introduced in 1916, does not show orbitals.
Electronegativity measures how strongly an atom pulls shared electrons towards itself in a bond.
It is measured on the Pauling scale (values from roughly 0.7 to 4.0).
Electronegativity increases as you move right across a period (greater nuclear charge pulling on valence electrons) and up a group (valence electrons are closer to the nucleus).
Fluorine (4.0) is the most electronegative element. The underlying physics is Coulomb's law: F = k(q₁q₂)/r². A larger nuclear charge and a shorter distance both increase the pull on electrons.
Δ > 1.9: ionic bond (electrons are transferred entirely from one atom to the other).
1.9 > Δ > 0.5: polar covalent bond (electrons are shared, but unequally; a bond dipole results, written δ+ on the less electronegative atom and δ– on the more electronegative one).
Δ < 0.5: non-polar covalent bond (electrons are shared roughly equally).
Na (electronegativity 0.9) and Cl (3.0): Δ = 2.1, which is greater than 1.9, so the bond is ionic.
Na transfers its one valence electron to Cl. Na becomes Na⁺; Cl becomes Cl⁻. Both achieve noble-gas configurations.
H (2.1) and Cl (3.0): Δ = 0.9, which falls between 0.5 and 1.9, so the bond is polar covalent.
The shared pair sits closer to chlorine. The molecule has a bond dipole: H is δ+, Cl is δ–.
In a Lewis structure this is drawn as H–Cl with three lone pairs on Cl.
Maximum electrons per shell: 2n² (where n = shell number).
Electron filling order: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → …
Coulomb's law: F = k(q₁q₂) / r². Governs the electrostatic attraction between nucleus and electrons, and explains why electronegativity rises with nuclear charge and falls with distance.
Electronegativity difference thresholds: Δ > 1.9 (ionic), 0.5 < Δ < 1.9 (polar covalent), Δ < 0.5 (non-polar covalent).
Ionic compounds like NaCl dissolve in water and conduct electricity, which is why saline solutions are used in medicine and why salt melts road ice. Polar covalent bonds explain why water (H–O bonds, Δ ≈ 1.4) is such an effective solvent for other polar and ionic substances, and why oil (non-polar C–H bonds) and water do not mix.
Students often think that a "polar bond" and a "polar molecule" are the same thing. They are not. A molecule can have polar bonds and still be non-polar overall if the bond dipoles cancel due to symmetry (e.g. CO₂).
Students sometimes assume that higher shell number always means higher energy for every orbital. It does not: the 4s orbital fills before the 3d because of shielding effects.
Students frequently confuse the Pauli exclusion principle with Hund's rule. Pauli limits an orbital to two electrons with opposite spins. Hund's rule says that electrons spread across equal-energy orbitals before pairing up. Both apply simultaneously.
Students may think that the electronegativity cutoffs (1.9 and 0.5) are sharp boundaries. They are guidelines, not rigid walls; real bonding sits on a continuum.
⚠️ Writing electron configurations quickly and correctly is a must. Expect to be asked for configurations of second-row elements (C, N, O) repeatedly throughout the course.
⚠️ Know the three filling rules by name (Aufbau, Pauli, Hund's) and be able to identify which rule is violated in a given incorrect configuration.
⚠️ Be able to calculate Δ for any pair of atoms and classify the bond as ionic, polar covalent, or non-polar covalent.
⚠️ Remember the bonding preferences for neutral atoms: C makes 4 bonds, N makes 3, O makes 2, F (and other halogens) makes 1. These shortcuts appear in almost every Lewis-structure question.
True or false: the 3s orbital fills before the 3p orbital.
Fill in the blank: the maximum number of electrons in shell 3 is ___.
True or false: a bond with Δ = 1.0 is ionic.
Fill in the blank: fluorine has an electronegativity of ___ on the Pauling scale.
True or false: Hund's rule states that no orbital can hold more than two electrons.
Answers: 1. True. 2. 18. 3. False (it is polar covalent). 4. 4.0. 5. False (that is the Pauli exclusion principle; Hund's rule is about spreading electrons across degenerate orbitals before pairing).
Q: Write the ground-state electron configuration of oxygen (Z = 8). How many valence electrons does it have?
A: 1s² 2s² 2px² 2py¹ 2pz¹. Oxygen has 6 valence electrons (the 2s and 2p electrons).
Q: Sodium (electronegativity 0.9) bonds with fluorine (electronegativity 4.0). What is the electronegativity difference, and what type of bond forms?
A: Δ = 4.0 – 0.9 = 3.1. Because Δ > 1.9, the bond is ionic. Na transfers its valence electron to F, forming Na⁺ and F⁻.
Q: Carbon (electronegativity 2.5) bonds with oxygen (electronegativity 3.5). Classify the bond and indicate the direction of the bond dipole.
A: Δ = 1.0, so the bond is polar covalent. The bond dipole points from carbon (δ+) toward oxygen (δ–), because oxygen attracts the shared electrons more strongly.
Q: Using Hund's rule, explain why the ground-state configuration of nitrogen places one electron in each of the three 2p orbitals rather than pairing two in one orbital.
A: Hund's rule states that for orbitals of equal energy, electrons occupy them singly with parallel spins before any pairing occurs. The three 2p orbitals are degenerate, so nitrogen's three 2p electrons each go into a separate orbital (2px¹ 2py¹ 2pz¹).
Q: How many valence electrons does a neutral carbon atom have, and how many bonds does carbon typically form in a neutral molecule?
A: Carbon has 4 valence electrons and typically forms 4 bonds to achieve an octet.
This material connects directly to Lewis structures, formal charge, and resonance (covered in the next set of notes for this chapter), because you need to know how many valence electrons each atom brings before you can draw any structure. Electronegativity reappears when predicting molecular polarity and, later in the course, when determining the direction of electron flow in reaction mechanisms.
electron configuration, ground state, orbital, shell, subshell, s orbital, p orbital, Aufbau, Pauli exclusion, Hund's rule, valence electrons, core electrons, electronegativity, Pauling scale, ionic bond, covalent bond, polar covalent, non-polar covalent, bond dipole, octet rule, Lewis dot structure, noble gas configuration, Coulomb's law, electrostatic force, electron transfer, electron sharing, CHEM 25500, organic chemistry Purdue, Chapter 1