How can a molecule with polar bonds be nonpolar?

Bond dipoles are vectors, so direction matters. If the outer atoms are all identical and arranged symmetrically around the central atom, as in linear CO_2, trigonal planar BF_3, or tetrahedral CCl_4, the dipoles point away from each other and cancel exactly. Lone pairs on the central atom, or outer atoms that are not all the same, break that symmetry and leave a net dipole.

Is a hydrogen bond an actual chemical bond?

No. It is an intermolecular attraction between molecules, roughly 5 to 10 percent as strong as a typical covalent bond. Boiling water breaks hydrogen bonds between molecules; it does not break the O-H bonds inside them. The name is historical and misleads a lot of students.

Do polar molecules still have dispersion forces?

Yes. Every substance has London dispersion forces because every substance has moving electrons. Dipole-dipole attraction and hydrogen bonding are added on top of dispersion, never substituted for it. When you list the forces present in a polar molecule, dispersion should always be one of them.

How do I decide which of two substances boils higher?

Work down a checklist. First ask whether either can hydrogen bond (H directly on N, O, or F); that usually decides it for similar-sized molecules. If neither does, compare polarity for dipole-dipole. If polarity is similar or both are nonpolar, compare number of electrons and shape, since bigger and less branched means stronger dispersion. Remember that a much larger nonpolar molecule can outrank a small polar one.

CHEM-4.5

Bond Polarity & Intermolecular Forces

Learn to use electronegativity differences and molecular shape to label bonds and molecules polar or nonpolar, then predict boiling point, solubility, and state.

What you'll do in this lesson

A voice-first session with the Crimsora tutor on Bond Polarity & Intermolecular Forces, then targeted practice and FRQs — with the tutor adapting to where you get stuck.

Start Bond Polarity & Intermolecular Forces free

What this lesson covers

You already know how atoms share or transfer electrons and how VSEPR predicts a molecule's shape. Now comes the payoff: those two ideas together explain why water is a liquid at room temperature while methane is a gas, why oil refuses to mix with vinegar, and why rubbing alcohol evaporates off your skin in seconds.

The logic runs in a chain. Electronegativity difference tells you whether a single bond is polar. Molecular shape tells you whether those bond dipoles cancel or add up to a whole-molecule dipole. Molecular polarity (plus a few structural details) tells you which intermolecular forces act between neighboring molecules. And the strength of those forces sets the boiling point, the solubility, and the state of matter. Master the chain in that order and these questions stop feeling like memorization.

Electronegativity Difference and Bond Character

Electronegativity is an atom's pull on the electrons in a shared bond. Fluorine is the strongest puller at about

4.0

; cesium and francium sit near

0.7

. Electronegativity increases left to right across a period and decreases down a group, so the top right corner of the periodic table (excluding noble gases) holds the greediest atoms.

Subtract the two values to get

\Delta EN

and compare it to the standard cutoffs:

$\Delta EN$	Bond type	What the electrons do
$0$ to $0.4$	Nonpolar covalent	Shared essentially equally
$0.4$ to $1.7$	Polar covalent	Shared unequally; partial charges appear
Greater than $1.7$	Ionic	Transferred; ions form

In a polar covalent bond the more electronegative atom carries a partial negative charge, written

\delta^-

, and its partner carries

\delta^+

. Chemists draw the bond dipole as an arrow pointing toward

\delta^-

with a crossed tail at

\delta^+

.

Two cautions. First, these cutoffs are guidelines, not laws.

HF

has

\Delta EN = 1.9

but is a molecular compound, not an ionic solid, because both atoms are nonmetals. Always check the metal/nonmetal identity too. Second, students often assume every bond between different elements is strongly polar. The

C-H

bond has

\Delta EN \approx 0.4

, which is why hydrocarbons behave as nonpolar substances even though carbon and hydrogen are different elements. That single fact explains most of what oils and greases do.

Polar Bonds Versus Polar Molecules

A molecule can be packed with polar bonds and still be completely nonpolar. Bond dipoles are vectors: they have direction as well as size, so they can cancel.

Carbon dioxide is the classic case. Each

C=O

bond has

\Delta EN = 1.0

and is strongly polar, but the molecule is linear, so the two arrows point in exact opposite directions and sum to zero.

CO_2

is nonpolar. Water has the same two polar bonds to oxygen, but its bent shape (two lone pairs on the central atom) means the dipoles do not oppose each other. They add to a large net dipole, making water one of the most polar small molecules there is.

A reliable two-step test: does the molecule contain at least one polar bond, and are the outer atoms arranged symmetrically around the central atom?

Shape	Identical outer atoms, no lone pairs on center	Result
Linear ( $CO_2$ )	Symmetric	Nonpolar
Trigonal planar ( $BF_3$ )	Symmetric	Nonpolar
Tetrahedral ( $CCl_4$ )	Symmetric	Nonpolar
Bent ( $H_2O$ , $SO_2$ )	Asymmetric	Polar
Trigonal pyramidal ( $NH_3$ )	Asymmetric	Polar

Symmetry breaks the moment the outer atoms differ.

CCl_4

is nonpolar, but

CH_3Cl

is tetrahedral with one atom unlike the others, so the dipoles do not cancel and the molecule is polar. The most common mistake in this topic is stopping after finding polar bonds and declaring the molecule polar. Always draw the shape first.

The Intermolecular Forces

Intermolecular forces (IMFs) are attractions between separate molecules, not the bonds inside them. They are far weaker than covalent bonds, which is why boiling water separates molecules from each other without breaking any

O-H

bonds.

London dispersion forces (LDF) exist between all particles. Electrons move, so at any instant a molecule has a lopsided electron cloud, an instantaneous dipole that induces a matching dipole in its neighbor. LDFs get stronger as the number of electrons (roughly, the molar mass) increases and as the molecule gets longer and more surface-contacting. This is the only force between nonpolar molecules, and it explains why

F_2

is a gas,

Br_2

a liquid, and

I_2

a solid at room temperature.

Dipole-dipole forces act between permanent dipoles of polar molecules: the

\delta^+

end of one aligns with the

\delta^-

end of another. They act in addition to dispersion forces, never instead of them.

Hydrogen bonding is an unusually strong dipole-dipole attraction that occurs only when hydrogen is bonded directly to nitrogen, oxygen, or fluorine, and interacts with a lone pair on N, O, or F in another molecule. Remember the phrase "H on N, O, or F."

CH_3CH_2OH

hydrogen bonds;

CH_3OCH_3

, which has the same molecular formula

C_2H_6O

, does not, because its hydrogens sit on carbon.

General strength ranking for similar-sized molecules: hydrogen bonding is greater than dipole-dipole, which is greater than dispersion. But dispersion wins when the nonpolar molecule is much larger; hexane boils higher than the polar

HCl

Using Forces to Explain Properties

Stronger intermolecular forces mean molecules cling together harder, so more energy is needed to pull them apart.

Boiling and melting points. Boiling separates molecules completely, so boiling point tracks IMF strength almost directly. Water boils at 100 degrees Celsius while

H_2S

, which is heavier but cannot hydrogen bond, boils at about negative 60 degrees Celsius. To rank boiling points, first check for hydrogen bonding, then compare polarity, then compare size and shape for dispersion. Among same-formula isomers, the more branched molecule has less surface contact, weaker dispersion, and a lower boiling point.

State of matter. At room temperature, weak forces leave molecules free to fly apart as a gas; moderate forces give a liquid; strong forces (or ionic bonding) give a solid.

Solubility. "Like dissolves like" is a statement about IMFs, not a magic phrase. A solvent dissolves a solute when solute-solvent attractions can replace the attractions each had on its own. Polar water dissolves polar sugar and ionic salt (through ion-dipole attractions) but not nonpolar oil, because water molecules would have to give up hydrogen bonds and get nothing comparable back.

Vapor pressure and evaporation. These run opposite to boiling point. Weak IMFs mean high vapor pressure and fast evaporation, which is why acetone flashes off a surface much faster than water does.

A frequent error is explaining a low boiling point by saying the covalent bonds are weak. Boiling never breaks covalent bonds. Name the intermolecular force by name in every explanation.

Key terms

Electronegativity.: A measure of how strongly an atom attracts the shared electrons in a chemical bond; it increases across a period and decreases down a group.
Polar covalent bond.: A bond in which electrons are shared unequally, typically with $\Delta EN$ between $0.4$ and $1.7$ , producing partial charges $\delta^+$ and $\delta^-$ .
Dipole moment.: The net separation of charge in a molecule, found by adding all bond dipoles as vectors; a molecule with a nonzero dipole moment is polar.
Intermolecular force.: An attraction between separate molecules, much weaker than the covalent bonds within a molecule; it determines physical properties such as boiling point.
London dispersion force.: The attraction caused by instantaneous, temporary dipoles in electron clouds; present in all substances and stronger for larger, more elongated molecules.
Dipole-dipole force.: The attraction between the permanently positive end of one polar molecule and the negative end of another.
Hydrogen bond.: A strong dipole-dipole attraction occurring when hydrogen bonded to N, O, or F is attracted to a lone pair on N, O, or F in a neighboring molecule.
Like dissolves like.: The rule that substances dissolve best in solvents with similar polarity and intermolecular forces, because comparable attractions can replace the ones being broken.

Worked example

Rank

CH_4

CH_3Cl

CH_3OH

, and

C_5H_{12}

(pentane) from lowest to highest boiling point, and predict which of the four will dissolve in water. Electronegativities: H

2.1

, C

2.5

, Cl

3.0

, O

3.5

Step 1: Classify each bond.

C-H

gives

\Delta EN = 0.4

, essentially nonpolar.

C-Cl

gives

\Delta EN = 0.5

, polar.

C-O

gives

\Delta EN = 1.0

and

O-H

gives

\Delta EN = 1.4

, both polar.

Step 2: Apply shape.

CH_4

is tetrahedral with four identical outer atoms, so it is symmetric and nonpolar.

CH_3Cl

is tetrahedral but one outer atom differs, so the dipoles do not cancel and it is polar.

CH_3OH

has a bent

C-O-H

arrangement and a strong net dipole. Pentane is a nonpolar chain of

C-H

and

C-C

bonds.

Step 3: Identify the forces.

CH_4

: dispersion only, and it is a tiny molecule with only 10 electrons, so those forces are very weak.

CH_3Cl

: dispersion plus dipole-dipole.

CH_3OH

: dispersion, dipole-dipole, and hydrogen bonding, because H sits on O. Pentane: dispersion only, but it has 42 electrons and a long chain, so contact area is large.

Step 4: Rank. Lowest is

CH_4

(weakest dispersion), then

CH_3Cl

(small, dipole-dipole), then pentane (dispersion so large it beats the small polar molecule; pentane boils at about 36 degrees Celsius versus about negative 24 for

CH_3Cl

), then

CH_3OH

highest at about 65 degrees Celsius because of hydrogen bonding.

Step 5: Solubility. Water is polar and hydrogen bonds.

CH_3OH

can hydrogen bond directly with water and is fully miscible.

CH_3Cl

is slightly soluble.

CH_4

and pentane are nonpolar and essentially insoluble, since water gains nothing to replace the hydrogen bonds it must break.

Practice questions

Which molecule contains polar bonds but is itself nonpolar?

$NH_3$
$CCl_4$
$H_2O$
$HCl$

Answer: $CCl_4$

Each

C-Cl

bond is polar (

\Delta EN = 0.5

), but

CCl_4

is tetrahedral with four identical chlorine atoms arranged symmetrically, so the four bond dipoles cancel and the net dipole is zero.

NH_3

is trigonal pyramidal and

H_2O

is bent, so in both cases the lone pairs break the symmetry and the dipoles add.

HCl

has only one bond, which cannot cancel with anything.

Ammonia (

NH_3

) boils at about negative 33 degrees Celsius, while phosphine (

PH_3

) boils at about negative 88 degrees Celsius even though phosphine has the greater molar mass. Explain this difference in terms of intermolecular forces.

Answer: $NH_3$ molecules hydrogen bond to one another because hydrogen is bonded directly to nitrogen; $PH_3$ cannot hydrogen bond, so it relies on weaker dipole-dipole and dispersion forces, and hydrogen bonding outweighs phosphine's slightly larger dispersion forces.

Normally boiling point rises with molar mass because dispersion forces grow. Hydrogen bonding is the exception that overrides the trend. Nitrogen is small and highly electronegative, so the

N-H

bond is strongly polarized and nitrogen's lone pair is compact enough to attract a neighboring

\delta^+

hydrogen tightly. Phosphorus has nearly the same electronegativity as hydrogen, so

P-H

bonds are barely polar and no hydrogen bonding network forms. A complete answer names the specific force, states which substance has it, and connects force strength to the energy required to separate molecules.

A student says iodine (

I_2

) must be polar because it is a solid at room temperature while chlorine (

Cl_2

) is a gas. Identify the error and give the correct explanation.

Answer: $I_2$ is nonpolar; the difference comes from London dispersion forces, which are much stronger in iodine because it has far more electrons and a larger, more polarizable electron cloud.

Both molecules are diatomic elements, so

\Delta EN = 0

and neither has any bond dipole. State of matter reflects total intermolecular attraction, not polarity alone. Iodine's 106 electrons form a large, loosely held cloud that distorts easily, creating strong instantaneous dipoles; chlorine's 34 electrons do not. This is the key evidence that dispersion forces are real attractions, not just a footnote for nonpolar substances.

FAQ

How can a molecule with polar bonds be nonpolar?: Bond dipoles are vectors, so direction matters. If the outer atoms are all identical and arranged symmetrically around the central atom, as in linear $CO_2$ , trigonal planar $BF_3$ , or tetrahedral $CCl_4$ , the dipoles point away from each other and cancel exactly. Lone pairs on the central atom, or outer atoms that are not all the same, break that symmetry and leave a net dipole.
Is a hydrogen bond an actual chemical bond?: No. It is an intermolecular attraction between molecules, roughly 5 to 10 percent as strong as a typical covalent bond. Boiling water breaks hydrogen bonds between molecules; it does not break the $O-H$ bonds inside them. The name is historical and misleads a lot of students.
Do polar molecules still have dispersion forces?: Yes. Every substance has London dispersion forces because every substance has moving electrons. Dipole-dipole attraction and hydrogen bonding are added on top of dispersion, never substituted for it. When you list the forces present in a polar molecule, dispersion should always be one of them.
How do I decide which of two substances boils higher?: Work down a checklist. First ask whether either can hydrogen bond (H directly on N, O, or F); that usually decides it for similar-sized molecules. If neither does, compare polarity for dipole-dipole. If polarity is similar or both are nonpolar, compare number of electrons and shape, since bigger and less branched means stronger dispersion. Remember that a much larger nonpolar molecule can outrank a small polar one.

Learn this with a teacher, not a page

The Crimsora tutor teaches Bond Polarity & Intermolecular Forces live — explaining on a whiteboard, asking you questions, and adapting to where you get stuck.

Try a live session free