9.2 Polarity
PRE-HEALTH PREP
Polar vs. Nonpolar Molecules
General Chemistry 9.2
Lesson Overview
Molecular polarity depends on both bond polarity and three-dimensional molecular geometry — polar bonds can cancel when oriented symmetrically, producing a nonpolar molecule. Key examples including CO2, OCS, CH2Cl2, and water illustrate how electronegativity differences, bond angles, and lone pairs on the central atom determine whether bond dipoles cancel or produce a net dipole moment. A summary table covers all common molecular geometries and identifies the two notable exceptions — square planar and 5-domain linear — where lone pairs are present yet the molecule remains nonpolar.
Table of Contents
- Electronegativity and Polar Bonds
- From Polar Bonds to Polar Molecules
- Example: CO2 — Polar Bonds That Cancel
- Example: OCS — Unequal Bond Dipoles
- Example: CH2Cl2 — Don't Trust the Lewis Structure
- The General Rule: Identical Atoms, No Lone Pairs
- Effect of Lone Pairs on the Central Atom
- 2 Exceptions: Lone Pairs That Don't Cause Polarity
- 13 Molecular Geometries: Polarity Reference Table
- 7 Key Takeaways
Electronegativity and Polar Bonds
Electronegativity Trends
Before identifying polar molecules, you need a firm grip on electronegativity. Recall that fluorine is the most electronegative element on the periodic table. Electronegativity increases as you move up a group or across a period from left to right — in other words, the closer an element sits to fluorine on the periodic table, the more electronegative it is. The two values worth committing to memory are fluorine at #1 and oxygen at #2.
Identifying Polar Bonds
A bond between two identical atoms is nonpolar. The carbon–hydrogen bond is also treated as nonpolar, since carbon and hydrogen have nearly identical electronegativities. As a rule of thumb, a difference in electronegativity below 0.5 on the Pauling scale defines a nonpolar bond — but unless you're given the full Pauling scale, you won't be expected to calculate exact differences. What you will need to do is recognize which bonds are polar and rank them by degree of polarity.
Example: Ranking Bond Polarity — C–I, C–Br, C–Cl, C–F
Which of these four bonds is the most polar? Look for the largest electronegativity difference. Using Pauling values: carbon = 2.5, iodine = 2.5, bromine = 2.8, chlorine = 3.0, fluorine = 4.0.
The C–F bond has by far the largest difference (4.0 − 2.5 = 1.5) and is therefore the most polar. The C–I bond, with a difference of essentially zero, is nearly nonpolar. C–Br and C–Cl fall in between, with C–Cl being more polar than C–Br.
From Polar Bonds to Polar Molecules
Identifying polar bonds is only the first step. The polarity of the molecule depends on two things: the polarity of its individual bonds and how those bonds are oriented in three-dimensional space.
- If a molecule has no polar bonds, it is automatically a nonpolar molecule.
- If a molecule has polar bonds, it may or may not be polar overall — the bond dipoles must be evaluated as vectors.
Each polar bond carries a bond dipole — a vector quantity pointing from the partially positive atom toward the more electronegative atom. When all bond dipoles are added as vectors, the result is the molecule's net dipole moment. If the net dipole moment sums to zero, the molecule is nonpolar. If any residual dipole remains, the molecule is polar. The larger the net dipole moment, the more polar the molecule.
Representation tip: A bond dipole is drawn as an arrow with a small cross (+) at the partially positive end, pointing toward the more electronegative atom. This is equivalent to labeling the atoms with partial charges (δ+ and δ−), but the arrow notation is far more convenient when you need to add dipoles visually.
Crucially, you cannot rely on a Lewis structure to judge bond angles. Lewis structures make all angles look like 90° or 180°, which is almost never accurate. You must apply your knowledge of molecular geometry — VSEPR electron domain geometry, bond angles, and the effect of lone pairs — to determine how bond dipoles are truly oriented.
Example 1: CO2 — Polar Bonds That Cancel
Carbon dioxide contains two C=O double bonds. Oxygen is significantly more electronegative than carbon, so each C=O bond is polar. The bond dipoles point from carbon toward each oxygen.
Here, the Lewis structure actually gets the geometry right: CO2 is linear, with a bond angle of exactly 180°. Because the two bond dipoles are equal in magnitude (both are C=O bonds) and point in exactly opposite directions, they cancel completely — their vector sum is zero.
The "Rope Pull" Analogy
Imagine two equally strong people pulling a rope attached to the central carbon — one tugging directly left, the other directly right, with equal force. The carbon goes nowhere. This is exactly what happens with the bond dipoles in CO2: equal pulls in exactly opposite directions produce zero net movement. The net dipole moment is zero, and the molecule is nonpolar.
Key rule: A nonzero net dipole moment means the molecule is polar; a net dipole moment of zero means the molecule is nonpolar. The larger the net dipole moment, the more polar the molecule.
Example 2: OCS — Unequal Bond Dipoles
Carbonyl sulfide (OCS) looks like carbon dioxide — it is also linear — but one oxygen has been replaced by a sulfur atom (one period lower on the periodic table). The molecule still has two polar bonds and a 180° geometry. What changes is the magnitude of the two bond dipoles.
The C=O bond dipole is significantly larger than the C=S bond dipole, because oxygen is considerably more electronegative than sulfur. Although the two dipoles still point in opposite directions, they are no longer equal in magnitude, so they do not cancel. The larger C=O dipole wins, and the molecule has a net dipole moment pointing toward the oxygen end.
CO2 vs. OCS — Why One Is Nonpolar and the Other Is Polar
Both molecules are linear with two polar bonds. In CO2, the two identical C=O dipoles cancel perfectly (net dipole moment = 0) — nonpolar. In OCS, the C=O and C=S dipoles are unequal in magnitude, so they only partially cancel — the molecule retains a net dipole moment toward oxygen, making it polar.
This distinction matters practically: whether a molecule is polar or nonpolar directly affects its physical properties — boiling point, melting point, viscosity, and solubility — topics you will encounter shortly.
2-Step Decision Strategy
Apply this two-step check to any molecule:
- Does the molecule have any polar bonds? If no, it is nonpolar. If yes, continue to step 2.
- Do the bond dipoles cancel completely? Use actual molecular geometry (not the Lewis structure) to evaluate the vector sum. If the sum is zero, the molecule is nonpolar. If any net dipole moment remains, the molecule is polar.
Example 3: CH2Cl2 — Don't Trust the Lewis Structure
CH2Cl2 (dichloromethane) is a classic trap. Students frequently misidentify it as nonpolar by reasoning: "The two C–Cl bonds are equal in magnitude, and the Lewis structure shows them on opposite sides — so they must cancel, just like CO2." This reasoning has two flaws.
Common mistake: Comparing CH2Cl2 to CO2 and concluding it is nonpolar. The flaw is assuming the two C–Cl bonds are 180° apart. They are not.
Carbon in CH2Cl2 has four bonding electron domains and no lone pairs. Four electron domains means tetrahedral electron domain geometry with bond angles of 109.5° — not 180°. Despite how the Lewis structure appears on paper, no angle in this molecule is 90° or 180°.
Three-Dimensional Structure Reveals the Truth
In the true tetrahedral structure, the two C–Cl bonds are oriented 109.5° apart. When you add their bond dipoles as vectors, they do not cancel — instead, the resultant points down through the center of the molecule (between the two chlorines). The C–H bonds contribute negligibly (carbon–hydrogen bonds are effectively nonpolar), but the two large C–Cl dipoles produce a substantial net dipole moment. CH2Cl2 is definitively polar.
Why CH2Cl2 and CO2 behave differently despite similar Lewis structures
CO2 is linear (180° bond angle) — two equal, opposite dipoles cancel. CH2Cl2 is tetrahedral (109.5° bond angles) — two equal dipoles at 109.5° do not cancel. Molecular geometry makes all the difference.
The General Rule: Identical Surrounding Atoms and No Lone Pairs
Most polarity questions you will encounter feature a central atom surrounded by identical atoms with no lone pairs on the central atom. For this common case, a clean rule applies:
If all surrounding atoms are identical AND the central atom has no lone pairs, the bond dipoles always cancel — the molecule is nonpolar.
BCl3 — Trigonal Planar, Nonpolar
Boron in BCl3 has three bonding electron domains and no lone pairs. That gives a trigonal planar geometry with bond angles of exactly 120°. All three B–Cl bonds are identical in polarity, and three equal vectors spaced 120° apart always sum to zero. The molecule is nonpolar.
Lewis structure warning — BCl3: A naive reading of the Lewis structure might suggest two of the dipoles cancel while the third goes uncanceled. This is wrong because the Lewis structure misrepresents the angles as 90° or 180°. The true 120° spacing of the trigonal planar geometry is what produces perfect cancellation.
CF4 — Tetrahedral, Nonpolar
Carbon in CF4 has four bonding domains and no lone pairs — a perfect tetrahedral arrangement with 109.5° bond angles. All four C–F bonds are identical, and four equal vectors in a symmetric tetrahedral arrangement sum to zero. The molecule is nonpolar.
Replace even one fluorine with a different atom (say, hydrogen), and the symmetry breaks. The bond dipoles no longer cancel, and the molecule becomes polar. Perfect symmetry around the central atom is what enables cancellation.
Effect of Lone Pairs on the Central Atom
Lone pairs on the central atom disrupt the symmetric arrangement of bonding domains. When all surrounding atoms are identical but the central atom carries one or more lone pairs, the bond dipoles generally do not cancel, and the molecule is usually polar.
SO2 — Bent Geometry, Polar
Sulfur in SO2 has three electron domains: two bonding and one lone pair. The electron domain geometry is trigonal planar, but the molecular geometry — which ignores the lone pair's "position" — is bent. The two S–O bonds are separated by approximately 120°, not 180°. Even though both bond dipoles are equal in magnitude, a 120° angle is not enough for cancellation. The molecule has a net dipole moment pointing along the molecular axis and is polar.
NF3 — Trigonal Pyramidal, Polar
Nitrogen in NF3 has four electron domains: three bonding and one lone pair. The electron domain geometry is tetrahedral, giving a trigonal pyramidal molecular geometry with bond angles close to 109.5°. If there were a fourth fluorine pointing straight up — as in CF4 — the four bond dipoles would cancel. Without it, the three N–F dipoles all point in roughly the same downward direction and reinforce one another. The molecule has a clear net dipole moment and is polar.
H2O — Bent Geometry, Polar
Water is a quintessential example of a polar molecule. The Lewis structure is frequently drawn in a way that makes the two O–H bonds appear 180° apart, but the true geometry tells a different story. Oxygen has four electron domains (two bonding, two lone pairs), giving a bent molecular geometry with a bond angle slightly less than 109.5°. The two O–H bond dipoles are far from antiparallel, so they add — not cancel — to produce a large net dipole moment. Water is one of the most polar small molecules known.
General pattern with lone pairs: If all surrounding atoms are identical but the central atom has lone pairs, the molecule is most likely polar. There are two important exceptions — keep reading.
2 Exceptions: Lone Pairs That Don't Make a Molecule Polar
Two molecular geometries are exceptions to the "lone pairs on central atom implies polar molecule" generalization. Both arise from high-symmetry arrangements where the bond dipoles still cancel despite the presence of lone pairs.
XeF4 — Square Planar, Nonpolar
Xenon in XeF4 has six electron domains: four bonding and two lone pairs. Six electron domains gives an octahedral electron domain geometry. With four bonding and two lone pairs, the molecular geometry is square planar — the four fluorines occupy the corners of a perfect square, and the two lone pairs sit above and below the plane.
In the square planar arrangement, each Xe–F bond dipole is directly opposed by the bond dipole across the square (180° apart). The two pairs cancel perfectly, and all four bond dipoles sum to zero. Despite the two lone pairs on xenon, XeF4 is nonpolar.
Do not overapply the lone-pair rule: The rule is that no lone pairs guarantees nonpolar (when surrounding atoms are identical). The presence of lone pairs makes a molecule likely polar — but square planar and the linear case below are the two exceptions professors love to test.
XeF2 — Linear Molecular Geometry, Nonpolar
Xenon in XeF2 has five electron domains: two bonding and three lone pairs. Five electron domains gives a trigonal bipyramidal electron domain geometry. With only two bonding domains and three lone pairs, the molecular geometry is linear. The two Xe–F bonds sit 180° apart — exactly opposite each other — and their bond dipoles cancel perfectly. XeF2 is nonpolar.
13 Molecular Geometries: Polarity Reference Table
The table below summarizes all 13 common molecular geometries and their polarity outcome when all surrounding atoms are identical. Geometries with no lone pairs on the central atom are always nonpolar; geometries with lone pairs are usually polar, with the two exceptions noted above.
| Molecular Geometry | Electron Domains | Lone Pairs on Central Atom | Polarity (identical surrounding atoms) |
|---|---|---|---|
| Linear (2 ED) | 2 | 0 | Nonpolar |
| Trigonal Planar | 3 | 0 | Nonpolar |
| Bent (3 ED) | 3 | 1 | Polar |
| Tetrahedral | 4 | 0 | Nonpolar |
| Trigonal Pyramidal | 4 | 1 | Polar |
| Bent (4 ED) | 4 | 2 | Polar |
| Trigonal Bipyramidal | 5 | 0 | Nonpolar |
| Seesaw | 5 | 1 | Polar |
| T-Shape | 5 | 2 | Polar |
| Linear (5 ED) | 5 | 3 | Nonpolar (exception) |
| Octahedral | 6 | 0 | Nonpolar |
| Square Pyramidal | 6 | 1 | Polar |
| Square Planar | 6 | 2 | Nonpolar (exception) |
The 2 exceptions professors love to test: Square planar (6 electron domains, 2 lone pairs) and the linear geometry derived from 5 electron domains (3 lone pairs) are the only geometries where lone pairs on the central atom are present yet the molecule is still nonpolar — provided all surrounding atoms are identical. Expect to see these on exams.
7 Key Takeaways
- No polar bonds = nonpolar molecule. If every bond in a molecule is nonpolar (same atoms bonded together, or C–H bonds), the molecule is automatically nonpolar.
- Polar bonds alone are not enough to make a molecule polar. The orientation of bond dipoles in three-dimensional space determines whether they cancel. A net dipole moment of zero means nonpolar; any nonzero net dipole moment means polar.
- Never use bond angles from Lewis structures. Lewis structures distort geometry, showing all angles as 90° or 180°. Always apply actual molecular geometry (tetrahedral = 109.5°, trigonal planar = 120°, etc.) when evaluating dipole cancellation.
- Identical surrounding atoms + no lone pairs on central atom = nonpolar. This applies to all symmetric geometries (linear, trigonal planar, tetrahedral, trigonal bipyramidal, octahedral).
- Lone pairs on the central atom usually cause polarity by breaking geometric symmetry — but square planar and the 5-domain linear geometry are the two exceptions where bond dipoles still cancel.
- Different surrounding atoms usually produce a polar molecule, because bond dipoles of unequal magnitude rarely cancel — as demonstrated by OCS and CH2Cl2.
- Molecular polarity drives physical properties. Whether a substance is polar or nonpolar directly influences its boiling point, melting point, viscosity, and solubility — making this concept foundational to all of the intermolecular forces topics that follow.