If you're seeing this message, it means we're having trouble loading external resources on our website.

If you're behind a web filter, please make sure that the domains *.kastatic.org and *.kasandbox.org are unblocked.

# Intramolecular and intermolecular forces

There are two kinds of forces, or attractions, that operate in a molecule—intramolecular and intermolecular. Let's try to understand this difference through the following example.
Figure of towels sewn and Velcroed representing bonds between hydrogen and chlorine atoms
We have six towels—three are purple in color, labeled hydrogen and three are pink in color, labeled chlorine. We are given a sewing needle and black thread to sew one hydrogen towel to one chlorine towel. After sewing, we now have three pairs of towels: hydrogen sewed to chlorine. The next step is to attach these three pairs of towels to each other. For this we use Velcro as shown above.
So, the result of this exercise is that we have six towels attached to each other through thread and Velcro. Now if I ask you to pull this assembly from both ends, what do you think will happen? The Velcro junctions will fall apart while the sewed junctions will stay as is. The attachment created by Velcro is much weaker than the attachment created by the thread that we used to sew the pairs of towels together. A slight force applied to either end of the towels can easily bring apart the Velcro junctions without tearing apart the sewed junctions.
Exactly the same situation exists in molecules. Just imagine the towels to be real atoms, such as hydrogen and chlorine. These two atoms are bound to each other through a polar covalent bond—analogous to the thread. Each hydrogen chloride molecule in turn is bonded to the neighboring hydrogen chloride molecule through a dipole-dipole attraction—analogous to Velcro. We’ll talk about dipole-dipole interactions in detail a bit later. The polar covalent bond is much stronger in strength than the dipole-dipole interaction. The former is termed an intramolecular attraction while the latter is termed an intermolecular attraction.
Figure of towels sewn and Velcroed representing bonds between hydrogen and chlorine atoms, illustrating intermolar and intramolar attractions
So now we can define the two forces:
Intramolecular forces are the forces that hold atoms together within a molecule. Intermolecular forces are forces that exist between molecules.
Figure of intermolecular attraction between two H-Cl molecules and intramolecular attraction within H-Cl molecule

## Types of intramolecular forces of attraction

1. Ionic bond: This bond is formed by the complete transfer of valence electron(s) between atoms. It is a type of chemical bond that generates two oppositely charged ions. In ionic bonds, the metal loses electrons to become a positively charged cation, whereas the nonmetal accepts those electrons to become a negatively charged anion.
Figure of ionic bond forming between Na and Cl
1. Covalent bond: This bond is formed between atoms that have similar electronegativities—the affinity or desire for electrons. Because both atoms have similar affinity for electrons and neither has a tendency to donate them, they share electrons in order to achieve octet configuration and become more stable.
A nonpolar covalent bond is formed between same atoms or atoms with very similar electronegativities—the difference in electronegativity between bonded atoms is less than 0.5.
Figure of covalent bond forming between two Cl molecules
A polar covalent bond is formed when atoms of slightly different electronegativities share electrons. The difference in electronegativity between bonded atoms is between 0.5 and 1.9. Hydrogen chloride, start text, H, C, l, end text; the start text, O, end text, minus, H bonds in water, start text, H, end text, start subscript, 2, end subscript, start text, O, end text; and hydrogen fluoride, start text, H, F, end text, are all examples of polar covalent bonds.
Figure of polar covalent bond forming between H and Cl
1. Metallic bonding: This type of covalent bonding specifically occurs between atoms of metals, in which the valence electrons are free to move through the lattice. This bond is formed via the attraction of the mobile electrons—referred to as sea of electrons—and the fixed positively charged metal ions. Metallic bonds are present in samples of pure elemental metals, such as gold or aluminum, or alloys, like brass or bronze.
Figure of metal with positively charged atoms and mobile valence electrons
The freely moving electrons in metals are responsible for their a reflecting property—freely moving electrons oscillate and give off photons of light—and their ability to effectively conduct heat and electricity.

## Relative strength of the intramolecular forces

Intramolecular forceBasis of formationRelative strength
Metallic bondMetal cations to delocalized electrons1, strongest
Ionic bondCations to anions2
Polar covalent bondPartially charged cation to partially charged anion3
Nonpolar covalent bondNuclei to shared electrons4, weakest

## Intermolecular forces of attraction

Now let’s talk about the intermolecular forces that exist between molecules. Intermolecular forces are much weaker than the intramolecular forces of attraction but are important because they determine the physical properties of molecules like their boiling point, melting point, density, and enthalpies of fusion and vaporization.

## Types of intermolecular forces that exist between molecules

1. Dipole-dipole interactions: These forces occur when the partially positively charged part of a molecule interacts with the partially negatively charged part of the neighboring molecule. The prerequisite for this type of attraction to exist is partially charged ions—for example, the case of polar covalent bonds such as hydrogen chloride, start text, H, C, l, end text. Dipole-dipole interactions are the strongest intermolecular force of attraction.
Figure of H-Cl to H-Cl dipole-dipole attraction
1. Hydrogen bonding: This is a special kind of dipole-dipole interaction that occurs specifically between a hydrogen atom bonded to either an oxygen, nitrogen, or fluorine atom. The partially positive end of hydrogen is attracted to the partially negative end of the oxygen, nitrogen, or fluorine of another molecule. Hydrogen bonding is a relatively strong force of attraction between molecules, and considerable energy is required to break hydrogen bonds. This explains the exceptionally high boiling points and melting points of compounds like water, start text, H, end text, start subscript, 2, end subscript, start text, O, end text, and hydrogen fluoride, start text, H, F, end text. Hydrogen bonding plays an important role in biology; for example, hydrogen bonds are responsible for holding nucleotide bases together in start text, D, N, A, end text and start text, R, N, A, end text.
Figure of intramolecular polar covalent bonding within H20 molecules and hydrogen bonding between O and H atoms.
1. London dispersion forces, under the category of van der Waal forces: These are the weakest of the intermolecular forces and exist between all types of molecules, whether ionic or covalent—polar or nonpolar. The more electrons a molecule has, the stronger the London dispersion forces are. For example, bromine, start text, B, r, end text, start subscript, 2, end subscript, has more electrons than chlorine, start text, C, l, end text, start subscript, 2, end subscript, so bromine will have stronger London dispersion forces than chlorine, resulting in a higher boiling point for bromine, 59 start superscript, start text, o, end text, end superscriptC, compared to chlorine, –35 start superscript, start text, o, end text, end superscriptC. Also, the breaking of London dispersion forces doesn’t require that much energy, which explains why nonpolar covalent compounds like methane—start text, C, H, end text, start subscript, 4, end subscript—oxygen, and nitrogen—which only have London dispersion forces of attraction between the molecules—freeze at very low temperatures.
Figure of intramolecular nonpolar covalent bonding between Cl atoms and Long dispersion forces between Cl-Cl molecules

## Relative strength of intermolecular forces of attraction

Intermolecular forceOccurs between …Relative strength
Dipole-dipole attractionPartially oppositely charged ionsStrong
Hydrogen bondingstart text, H, end text atom and start text, O, end text, start text, N, end text/ or start text, F, end text atomStrongest of the dipole-dipole attractions
London dispersion attractionTemporary or induced dipolesWeakest

## How forces of attraction affect properties of compounds

Polar covalent compounds—like hydrogen chloride, start text, H, C, l, end text, and hydrogen iodide, start text, H, I, end text—have dipole-dipole interactions between partially charged ions and London dispersion forces between molecules. Nonpolar covalent compounds—like methane start text, C, H, end text, start subscript, 4, end subscript and nitrogen gas, start text, N, end text, start subscript, 2, end subscript)—only have London dispersion forces between molecules. The rule of thumb is that the stronger the intermolecular forces of attraction, the more energy is required to break those forces. This translates into ionic and polar covalent compounds having higher boiling and melting points, higher enthalpy of fusion, and higher enthalpy of vaporization than covalent compounds.
Boiling and melting points of compounds depend on the type and strength of the intermolecular forces present, as tabulated below:
Type of compoundIntermolecular forces presentRelative order of boiling and melting points
Ionic compoundsIon to ion attraction between ions, London dispersion forces1, highest)
Covalent compounds containing hydrogen bondsHydrogen bonds, London dispersion forces2
Polar covalent compoundsDipole-dipole attraction between dipoles created by partially charged ions, London dispersion forces3
Nonpolar covalent compoundsLondon dispersion forces4, lowest
Let’s try to identify the different kinds of intermolecular forces present in some molecules.
1. start text, H, end text, start subscript, 2, end subscript, start text, S, end text—London dispersion force—by default every compound will have this force of attraction between molecules—and dipole-dipole attraction
Figure of H2S London dispersion force and dipole-dipole attraction
1. start text, C, H, end text, start subscript, 3, end subscript, start text, O, H, end text—London dispersion force, dipole-dipole attraction, and hydrogen bonding
Figure of CH3OH London dispersion force, dipole-dipole attraction and hydrogen bonding
1. start text, C, end text, start subscript, 2, end subscript, start text, H, end text, start subscript, 6, end subscript—London dispersion forces—it’s a nonpolar covalent compound— and no other intermolecular attractions
Figure of C2H6 London dispersion forces

## Want to join the conversation?

• Then what are dipole-induced dipole forces, ion-dipole forces, and ion-induced dipole forces?
• A dipole-induced dipole attraction is a weak attraction that results when a polar molecule induces a dipole in an atom or in a nonpolar molecule by disturbing the arrangement of electrons in the nonpolar species.
An ion-dipole force is an attractive force that results from the electrostatic attraction between an ion and a neutral molecule that has a dipole.
An ion-induced dipole attraction is a weak attraction that results when the approach of an ion induces a dipole in an atom or in a nonpolar molecule by disturbing the arrangement of electrons in the nonpolar species.
• I thought ionic bonds were much weaker than covalent bonds, for example the lattice structure of a carbon diamond is much stronger than a crystal lattice structure of NaCl. Is this table of bond strength wrong?
• I initially thought the same thing, but I think there is a difference between bond strengths, and intramolecular forces. The intramolecular force strength is relative to the electronegativity of the 2 atoms in the molecule. The bond strength relates to the stability of the bond in it's energy state.
• isnt hydrogen bonding stronger than dipole-dipole ??
• Hydrogen bonding is the strongest form of dipole-dipole interaction.
• Why can't we say that H2S also has Hydrogen bond along with London dispersion bond and dipole-dipole attraction ?
• Always remember to *(H)Hold the phone (*FON)
• #3 (C2H6) says that Van Der Waal Forces are found in non polar compounds. OK that i understand. What i'm not so clear on is the reasoning why #2 has Van Der Waal Forces.
If Van Der Waal forces are found in any compound irrespective of bond or polarity, can I assume that the compound has van der waal forces if it has dipole or hydrogen bonds?
• LDFs exist in everything, regardless of polarity.
However nonpolar compounds can only have LDFs.
Polar compounds can have LDFs, in addition to dipole-dipoles. Maybe H bonding if O-H F-H N-H bonds exist.

Edit: One more thing, LDFs can only happen everywhere, so when drawing molecules, LDFs can be drawn randomly as long as they are connecting the two molecules.

And to answer your question, you can say a compound has Van der waal forces if it has dip-dip bonds, H bonding, or LDFs.
• difference between inter and intramolecular bonds?
• Intermolecular bonds are the forces between the molecules. Intramolecular are the forces within two atoms in a molecule.
• The article said dipole-dipole interactions and hydrogen bonding are equally strong and hydrogen bonding is a type of dipole-dipole interaction, so how come covalent compounds containing hydrogen bonds have higher boiling and melting points than polar covalent compounds?
• *Hydrogen bonding is the strongest form of dipole-dipole interaction.*
So naturally, compounds that can hydrogen bond have higher bp since the H bonds are harder to overcome than other types of dipole-dipole interactions.
• In CH3OH (Methanol) Is there really a hydrogen bond between the carbon atom and the top left oxygen atom?
Oh by the way wonderful drawings!
• You are correct that would be impossible, but that isn't what the figure shows.

The red dashed line on the left is labeled as a dipole-dipole bond.
• Can an ionic bond be classified as an intermolecular and an intramolecular bond? If so, how?
• What is diff between diploid-diploid and London dispersion ?