Organic Chemistry · Polar Covalent Bonds; Acids and Bases
Polar Covalent Bonds and Electronegativity
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In a covalent bond, two atoms share a pair of electrons — but they rarely share it equally. electronegativity How strongly an atom in a molecule attracts shared bonding electrons Full entry → is the measure of an atom's ability to attract the shared electrons of a bond to itself. When two atoms of different electronegativity bond, the electron density shifts toward the more electronegative atom, creating a polar covalent bond A bond with unevenly shared electrons, giving partial charges δ+/δ− Full entry → with partial charges: the more electronegative atom carries a partial negative charge (δ−), and the less electronegative atom carries a partial positive charge (δ+).
Bond polarity is not an either/or property but a spectrum. At one extreme, identical atoms share electrons evenly (H–H, C–C): nonpolar covalent bonds. At the other extreme, the electron transfer is complete, producing ions (Na⁺Cl⁻): ionic bonds. Polar covalent bonds sit in the middle. Organic chemists use the difference in electronegativity between two bonded atoms, ΔEN, to place a bond on that spectrum and to predict where electron density — and therefore reactivity — will concentrate.
Why this matters
Polarity is the root cause of most organic reactivity. A bond's partial charges tell you which atom is electron-poor (δ+, electrophilic) and which is electron-rich (δ−, nucleophilic). The carbonyl carbon of a ketone, for example, is δ+ because oxygen is far more electronegative than carbon — that single fact explains why nucleophiles attack carbonyl carbons throughout the book. Polarity also controls solubility ("like dissolves like"), boiling points, hydrogen bonding, drug–receptor interactions, and the appearance of characteristic absorptions in infrared spectroscopy. Mastering electronegativity now makes every reaction-mechanism chapter more intuitive.
The college version
Core Concepts
What electronegativity is (and is not)
Electronegativity is a relative, dimensionless number — it has no units. The most common scale, the Pauling scale, runs from about 0.7 (cesium) to 4.0 (fluorine, the most electronegative element). In organic chemistry the values you must know by heart are:
| Element | Electronegativity |
|---|---|
| H | 2.1 |
| C | 2.5 |
| N | 3.0 |
| O | 3.5 |
| F | 4.0 |
| Cl | 3.0 |
| Br | 2.8 |
| I | 2.5 |
| S | 2.5 |
| P | 2.1 |
Electronegativity is not the same as electron affinity Energy change when a free atom gains an electron Full entry →: electron affinity is the energy change when a free atom gains an electron, while electronegativity describes how strongly an atom in a molecule attracts shared bonding electrons.
Periodic trends in electronegativity
Electronegativity increases left to right across a period (nuclear charge rises while the shell stays the same, so the pull on shared electrons grows) and decreases down a group (the atom gets larger, and the bonding electrons sit farther from the nucleus, shielded by more core electrons). The most electronegative elements — F, O, N, Cl — cluster in the upper right of the periodic table. Two practical organic consequences: fluorine is the most electronegative element of all, and hydrogen (2.1) is close to carbon (2.5), which is why C–H bonds are essentially nonpolar.
The polarity spectrum and the ΔEN guide
The polarity of a bond depends on the difference between the two atoms' electronegativities:
ΔEN = | EN(A) - EN(B) |
As a rule of thumb:
- ΔEN ≈ 0 to 0.4 → nonpolar covalent (H–H, C–C, C–H).
- ΔEN ≈ 0.4 to 1.7 → polar covalent (C–O, C–N, O–H, C–Cl).
- ΔEN > 1.7 → largely ionic (Na–Cl, K–O).
This scale is a guideline, not a law — many compounds with ΔEN above 1.7 still behave covalently (e.g., Be–Cl), and no bond is purely one type. In organic compounds, bonds between carbon and the common heteroatoms are polar covalent, never truly ionic.
Partial charges and the dipole arrow
In a polar bond the electron cloud is displaced toward the more electronegative atom, giving it a partial negative charge (δ−) and the other atom a partial positive charge (δ+). These are not full charges — no electron transfer has occurred. Chemists indicate the direction of electron pull with a dipole arrow Crossed-tail arrow pointing from δ+ toward δ− Full entry →: an arrow with a crossed tail (like a plus sign) pointing from the δ+ end toward the δ− end, drawn along the bond.
Polar bonds in organic molecules
The bonds that dominate organic chemistry, ranked by polarity:
| Bond | ΔEN | Character |
|---|---|---|
| C–C | 0.0 | Nonpolar |
| C–H | 0.4 | Essentially nonpolar |
| C–N | 0.5 | Weakly polar |
| C–Cl | 0.5 | Weakly polar |
| C–O | 1.0 | Polar |
| C–F | 1.5 | Very polar |
| O–H | 1.4 | Very polar |
| N–H | 0.9 | Polar |
The polar bonds create the reactive "hot spots" of a molecule: the δ+ carbon of a C=O, the δ+ carbon of a C–Cl (electrophilic site for substitution), and the δ− oxygen or nitrogen that can act as a nucleophile or hydrogen-bond acceptor.
Common Confusions
| Do not confuse | With | Difference |
|---|---|---|
| Electronegativity | Electron affinity | EN is about attracting shared electrons in a bond; EA is the energy change for adding an electron to a free atom |
| A polar bond | A polar molecule | A molecule's overall polarity also depends on geometry — a molecule can have polar bonds yet be nonpolar overall (CO₂; see the next topic) |
| ΔEN > 1.7 | "Always ionic" | It is a rule of thumb, not a law; organic C–X bonds stay covalent, and some compounds above 1.7 behave covalently |
| C–H bonds | "Nonpolar, so hydrogens never matter" | True that C–H is essentially nonpolar, but acidic C–H bonds (near carbonyls, alkynes) still matter in acid–base chemistry |
| "The dipole arrow points to the positive end" | The arrow's direction | The crossed tail sits at the δ+ end and the arrowhead points toward the δ− end — the direction of electron pull |
Eli explains
The same idea, in plain words
Explain it like I’m 10
Imagine two kids holding a rope. If both pull equally, the rope stays in the middle — that is a nonpolar bond. If one kid is much stronger, the rope slides toward that kid — that is a polar bond, and the strong kid gets a tiny minus sign and the weak kid a tiny plus sign. Fluorine is the strongest kid in the whole playground, so whenever it holds a rope with anyone else, the rope almost always ends up on fluorine's side.
Worked example
Example 1: Classifying bonds by ΔEN
Classify the C–O, C–H, and Na–Cl bonds using Pauling electronegativities (C 2.5, O 3.5, H 2.1, Na 0.9, Cl 3.0).
Step 1 — write the formula: ΔEN = |EN(A) - EN(B)|.
Step 2 — substitute for each bond:
ΔEN(C–O) = |2.5 - 3.5| = 1.0
ΔEN(C–H) = |2.5 - 2.1| = 0.4
ΔEN(Na–Cl) = |0.9 - 3.0| = 2.1
Step 3 — classify: C–O (1.0) is polar covalent; C–H (0.4) sits at the nonpolar edge and is treated as essentially nonpolar in organic chemistry; Na–Cl (2.1) is ionic. This is why NaCl is a salt while methanol (CH₃OH) is a covalently bonded liquid — the same pair of atoms, different ΔEN, different bonding.
Example 2: Predicting the electron-rich end of a bond
Identify the δ− and δ+ ends of the C=O bond in acetone (CH₃COCH₃, SMILES CC(=O)C).
Step 1 — compare electronegativities: O (3.5) is more electronegative than C (2.5).
Step 2 — the π and σ electrons of the double bond are pulled toward oxygen, so oxygen is δ− and the carbonyl carbon is δ+.
Step 3 — consequence: the carbon is electron-poor and electrophilic. When a nucleophile (electron-rich species) approaches acetone, it attacks the δ+ carbon — the exact step that opens Chapter 19 on nucleophilic addition to carbonyls.
Example 3: Ranking bonds by polarity
Rank C–F, C–Cl, C–Br, and C–I from least to most polar (F 4.0, Cl 3.0, Br 2.8, I 2.5).
Step 1 — write the formula: ΔEN = |EN(C) - EN(X)|.
Step 2 — substitute: C–F: |2.5 − 4.0| = 1.5; C–Cl: |2.5 − 3.0| = 0.5; C–Br: |2.5 − 2.8| = 0.3; C–I: |2.5 − 2.5| = 0.0.
Step 3 — rank: C–I (0.0) < C–Br (0.3) < C–Cl (0.5) < C–F (1.5). Fluorine pulls hardest, which is why the C–F bond is the most polar of the carbon–halogen bonds — an important trend in organohalide chemistry (Chapter 10).
Key takeaways
- Electronegativity increases across a period and decreases down a group; F (4.0) is the most electronegative element.
- Memorable values: H 2.1, C 2.5, N 3.0, O 3.5, F 4.0, Cl 3.0, Br 2.8, I 2.5, S 2.5, P 2.1.
- ΔEN roughly 0–0.4 nonpolar, 0.4–1.7 polar covalent, >1.7 ionic — a guideline, not a law.
- In a polar bond the more electronegative atom is δ−; the other is δ+; the dipole arrow points from δ+ toward δ−.
- C–H is essentially nonpolar; C–O, C=O, O–H, N–H, C–N, and C–Cl are polar.
- The δ+ carbon of a carbonyl is the electrophilic site attacked by nucleophiles.
- Electronegativity ≠ electron affinity.
Check yourself
5 review questions from the chapter. Try each one, then open the answer.
Define electronegativity and give the trend across a period and down a group.
Show answer
Electronegativity is an atom's ability to attract shared bonding electrons. It increases left to right across a period and decreases down a group.
What is the ΔEN rule of thumb for nonpolar, polar covalent, and ionic bonds?
Show answer
ΔEN ≈ 0–0.4 nonpolar; 0.4–1.7 polar covalent; >1.7 largely ionic (guideline only).
Rank these bonds from least to most polar: C–N, C–O, C–F, C–C.
Show answer
C–C (0.0) < C–N (0.5) < C–O (1.0) < C–F (1.5).
Which atom is δ− in an O–H bond, and why?
Show answer
Oxygen is δ−: O (3.5) is more electronegative than H (2.1), so the shared electrons sit closer to oxygen.
Why is the carbonyl carbon of a ketone electrophilic?
Show answer
Oxygen (3.5) pulls the double-bond electrons toward itself, leaving the carbon electron-poor (δ+), so nucleophiles attack it.
Study tools & related lessonsKey vocabulary · Related
Key vocabulary
- electronegativity
- How strongly an atom in a molecule attracts shared bonding electrons
- polar covalent bond
- A bond with unevenly shared electrons, giving partial charges δ+/δ−
- nonpolar covalent bond
- A bond where electrons are shared evenly (identical or similar EN atoms)
- ionic bond
- Complete transfer of electrons; full + and − ions
- partial charge (δ)
- A fractional charge from uneven electron sharing, less than a full charge
- dipole arrow
- Crossed-tail arrow pointing from δ+ toward δ−
- electron affinity
- Energy change when a free atom gains an electron
Sources & references
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