Organic Chemistry · Structure and Bonding
Development of Chemical Bonding Theory
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Why do atoms stick together at all? Answering that took chemistry nearly two centuries, and today's answer is layered: a covalent bond A bond formed by sharing an electron pair between two atoms Full entry → is a pair of electrons shared between two atoms (the Lewis picture, 1916), which is also a region of orbital overlap concentrating electron density between the nuclei (the quantum picture, 1927 onward). This topic traces that development — from the electron's discovery and the octet rule Atoms tend to bond until they have eight valence electrons (two for H) Full entry → to the quantum description — because each layer explains what the previous one could not.
Why this matters
Chemical bonding theory is the operating manual for every reaction in this book.
- Lewis structures are the everyday language of organic chemistry: drawing dots and lines for shared pairs predicts how many bonds an atom forms, where lone pairs sit, and which atoms carry formal charge — skills used in every later chapter.
- The octet rule explains bonding patterns. Carbon forms four bonds, nitrogen three, oxygen two, hydrogen one — each atom seeks eight valence electrons (two, for hydrogen).
- Electronegativity explains bond polarity: the electronegativity difference between two atoms decides whether a bond is covalent, polar covalent, or ionic — and polarity controls solubility, reactivity, and acidity.
- The quantum layer fixes Lewis's limitations: Lewis structures don't explain bond shapes and strengths; orbital theory does. Chemists switch between the layers fluidly.
The college version
Core Concepts
Before Lewis: ions and the electron
By the early 1900s chemists knew some compounds are held together by ionic bonds — electrostatic attraction between oppositely charged ions (NaCl being the textbook case). The discoveries of the electron (Thomson, 1897) and the nuclear atom (Rutherford, 1911) made it clear that electrons must be what bonds are made of.
G. N. Lewis and the octet rule (1916)
Gilbert N. Lewis proposed that atoms bond to achieve a filled valence shell — eight electrons for second-row elements (an octet; hydrogen, two). He introduced Lewis structures: the element's symbol surrounded by dots, one per valence electron, with shared pairs shown as lines (each line = one shared pair = one bond).
Key Lewis rules: count valence electrons from the periodic table (column number for main-group elements); form bonds by sharing pairs so each atom (except H) ends with eight electrons; lone pairs are valence electrons not involved in bonding.
Example — methane, CH₄: carbon contributes 4 valence electrons, each hydrogen 1, total 4 + 4(1) = 8. Four C–H bonds use all 8; carbon ends with an octet (four shared pairs) and each hydrogen with two. Methane, with four C–H single bonds, is the archetypal organic molecule.
Ionic vs. covalent: the electronegativity spectrum
Lewis's picture needed refinement for bonds like Na–Cl, where the "shared" pair sits almost entirely on chlorine. Linus Pauling quantified this with electronegativity — an atom's ability to attract electrons in a bond — assigning each element a value (H 2.20, C 2.55, N 3.04, O 3.44, F 3.98, Na 0.93, Cl 3.16).
The difference in electronegativity, ΔEN, classifies the bond:
- ΔEN < 0.5: nonpolar covalent — electrons shared nearly equally (C–H: |2.55 - 2.20| = 0.35).
- 0.5 ≤ ΔEN ≤ 1.9: polar covalent — unequal sharing, partial charges (O–H: |3.44 - 2.20| = 1.24).
- ΔEN > 1.9: ionic — the electron pair is essentially transferred (Na–Cl: |0.93 - 3.16| = 2.23).
The spectrum is continuous — no sharp boundary, only a trend from equal sharing to full transfer.
The quantum layer: from dots to orbitals
In 1927, Walter Heitler and Fritz London showed with quantum mechanics that two hydrogen atoms bond because their 1s orbitals overlap, concentrating electron density between the nuclei. Pauling extended this into valence bond theory Model in which bonds form by overlap of atomic orbitals Full entry →: a covalent bond forms when atomic orbitals overlap, stronger with greater overlap. The next topics develop this into hybridization, which explains the bond angles Lewis structures cannot predict.
How It Works / Step-by-Step Process
To draw a Lewis structure A drawing of a molecule showing atoms, bonds (lines), and lone pairs (dots) Full entry → for a small molecule:
- Sum the valence electrons of all atoms.
- Place the least electronegative atom in the center (H is never central).
- Connect atoms with single bonds (2 electrons per bond).
- Distribute remaining electrons as lone pairs, starting with outer atoms.
- Check octets; if the central atom is short, form double or triple bonds from lone pairs.
- Verify the count matches the starting total.
Common Confusions
| Common Confusion | Correct Understanding |
|---|---|
| "Ionic and covalent bonds are completely different kinds of bonds." | They are the two ends of a continuous spectrum graded by ΔEN; most real bonds are polar covalent. |
| "The octet rule applies to every atom in every molecule." | It is a rule of thumb for main-group elements; hydrogen needs 2, and period-3+ elements (P, S) can exceed an octet. |
| "A lone pair doesn't count as electrons around an atom." | Lone pairs count fully toward the octet and affect shape, basicity, and reactivity. |
| "Electronegativity is the same as electron affinity." | Electronegativity is about attracting electrons in a bond; electron affinity is the energy change when a free atom gains one. Different quantities. |
| "Lewis theory and quantum theory are rivals." | They are complementary: Lewis structures show electron-pair connectivity; orbital overlap explains why, and with what geometry. |
| "Polar bonds make the whole molecule polar." | A molecule's overall polarity also depends on geometry — symmetric molecules like CO₂ have polar bonds but zero net dipole. |
Eli explains
The same idea, in plain words
Explain it like I’m 10
Atoms hold hands to feel complete. Lewis found that atoms are happiest with eight electrons around them, like a table set for eight guests. If two atoms each have an empty chair, they each bring an electron and share a chair — that's a bond, drawn as a line between them. Scientists later found the shared electrons sit in a fuzzy cloud between the two atoms, like two people sharing one umbrella.
Worked example
Example 1: Lewis structure of water, H₂O
Step 1 — count valence electrons: oxygen 6, each hydrogen 1:
total valence electrons = 6 + 2(1) = 8
Step 2 — skeleton: oxygen central, H on each side: H–O–H.
Step 3 — bonded electrons: two O–H bonds = 4 electrons; remaining 8 - 4 = 4 electrons go on oxygen as two lone pairs.
Result: H–O–H with two lone pairs on oxygen. Oxygen has an octet (two bonds × 2 + two lone pairs × 2 = 8) and each hydrogen its duet. The lone pairs explain water's bent shape (104.5°), hydrogen bonding, and solvent power.
Example 2: Classifying bonds by electronegativity difference
Classify C–H, O–H, and Na–Cl using Pauling values (H 2.20, C 2.55, O 3.44, Cl 3.16, Na 0.93).
The formula before substitution:
ΔEN = |ENA - ENB|
Substituting:
ΔEN(C–H) = |2.55 - 2.20| = 0.35 (nonpolar covalent)
ΔEN(O–H) = |3.44 - 2.20| = 1.24 (polar covalent)
ΔEN(Na–Cl) = |0.93 - 3.16| = 2.23 (ionic)
Interpretation: C–H in methane is nearly nonpolar (hydrocarbons are greasy and water-repelling), O–H is polar (hydrogen carries a partial positive charge — the origin of hydrogen bonding and carboxylic acid acidity), and Na–Cl is ionic, dissolving in water through ion–dipole attractions.
Example 3: How many electrons surround nitrogen in ammonia?
Ammonia is NH₃: nitrogen contributes 5 valence electrons, three hydrogens 3:
total = 5 + 3(1) = 8
Three N–H bonds use 6 electrons; the remaining 2 sit on nitrogen as a lone pair, completing nitrogen's octet (3 × 2 + 2 = 8). The lone pair makes ammonia a base (it can donate to a proton) — a fact that returns in the acid–base and amines chapters.
Key takeaways
- A covalent bond = a shared electron pair (Lewis) = a region of orbital overlap (quantum).
- Octet rule: main-group atoms bond to surround themselves with 8 valence electrons (H: 2).
- Lewis recipe: count valence electrons → share pairs as bonds → satisfy octets → place lone pairs.
- Electronegativity (Pauling scale) measures electron-attracting power; the bond type follows ΔEN.
- Rule of thumb: ΔEN < 0.5 nonpolar; 0.5–1.9 polar; > 1.9 ionic.
- Bond polarity → partial charges → controls solubility, boiling points, and reactivity.
- Heitler–London and Pauling added the quantum layer: bonds form from orbital overlap; stronger overlap → stronger bond.
- Lewis theory explains what bonds exist; orbital theory explains why and what shapes.
Check yourself
6 review questions from the chapter. Try each one, then open the answer.
State the octet rule and explain why hydrogen is the exception.
Show answer
Main-group atoms bond to surround themselves with eight valence electrons; hydrogen completes at two (a "duet") because its only shell holds two electrons.
Draw (in words) the Lewis structure of methane and count the electrons around carbon.
Show answer
Methane: carbon central, four C–H single bonds, no lone pairs; carbon is surrounded by 4 × 2 = 8 electrons (an octet); each H has 2.
Classify the C–O bond using Pauling values (C 2.55, O 3.44).
Show answer
ΔEN = |3.44 - 2.55| = 0.89 → polar covalent bond (between 0.5 and 1.9).
What did Heitler and London contribute to bonding theory?
Show answer
They showed with quantum mechanics that two H atoms bond because their 1s orbitals overlap, concentrating electron density between the nuclei — founding valence bond theory.
Why does ammonia act as a base?
Show answer
Nitrogen in NH₃ has a lone pair it can donate to a proton, so ammonia is a base.
A bond has ΔEN = 2.0. What type of bond is it, and what does that imply about electron location?
Show answer
Ionic (above 1.9): the electron pair sits almost entirely on the more electronegative atom — effectively transferred.
Study tools & related lessonsKey vocabulary · Related
Key vocabulary
- covalent bond
- A bond formed by sharing an electron pair between two atoms
- ionic bond
- A bond from electrostatic attraction between oppositely charged ions
- octet rule
- Atoms tend to bond until they have eight valence electrons (two for H)
- Lewis structure
- A drawing of a molecule showing atoms, bonds (lines), and lone pairs (dots)
- lone pair
- A pair of valence electrons not involved in bonding
- electronegativity (EN)
- An atom's ability to attract electrons in a chemical bond
- polar covalent bond
- A bond with unequal electron sharing, creating partial charges
- valence bond theory
- Model in which bonds form by overlap of atomic orbitals
Sources & references
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