Intermolecular Forces and Properties of Substances
Why does a puddle of water take hours to dry, but rubbing alcohol on your skin vanishes in seconds? Why can you melt a candle on a stove, yet table salt stays solid in a 500 °C oven? The answer is in the forces between particles. In this page you will learn to read those forces and use them to predict how a substance behaves.
Forces inside vs. forces between
There are two different "levels" of attraction in matter.
An intramolecular force (or bond) holds atoms together inside a molecule or formula unit. Examples are covalent and ionic bonds.
An intermolecular force (IMF) is an attraction between separate molecules. It is much weaker than a bond.
Take water. The O–H covalent bonds inside each $\text{H}_2\text{O}$ molecule are intramolecular. The hydrogen bonds that pull one water molecule toward its neighbors are intermolecular.
Why does this matter? When water boils, the molecules separate from each other, but each molecule stays $\text{H}_2\text{O}$. Boiling overcomes intermolecular forces. It does not break the covalent bonds.
Explain the difference between an intramolecular force and an intermolecular force. Give one example of each using water.
Explain the difference between an intramolecular force and an intermolecular force. Give one example of each using water.
Three types of intermolecular forces
All IMFs come from electrical attraction between positive and negative regions. There are three types, listed here from weakest to strongest (for molecules of similar size).
| Force | Where it acts | Example |
|---|---|---|
| London dispersion | All molecules, especially nonpolar ones | $\text{CH}_4$, $\text{I}_2$ |
| Dipole-dipole | Polar molecules | $\text{HCl}$, $\text{CH}_3\text{Cl}$ |
| Hydrogen bonding | H bonded to N, O, or F, attracted to a lone pair on N, O, or F of another molecule | $\text{H}_2\text{O}$, $\text{NH}_3$, $\text{HF}$ |
London dispersion forces. Electrons move constantly. For a moment, one side of a molecule may have slightly more electrons, creating a temporary dipole. That dipole shifts the electrons of the neighbor, and the two briefly attract. Bigger molecules have bigger electron clouds that are easier to distort (more polarizable), so their dispersion forces are stronger.
Dipole-dipole forces. A polar molecule has a permanent partial positive end ($\delta^+$) and partial negative end ($\delta^-$). The $\delta^+$ end of one molecule is attracted to the $\delta^-$ end of the next.
Hydrogen bonding. This is an especially strong dipole-dipole attraction. It needs H attached to N, O, or F. These atoms are small and very electronegative, so the H is left almost bare and is strongly attracted to a lone pair on a neighbor.
Name the strongest type of intermolecular force present in each substance: (a) $\text{HCl}$, (b) $\text{CH}4$, (c) $\text{H}2\text{O}$.
Name the strongest type of intermolecular force present in each substance: (a) $\text{HCl}$, (b) $\text{CH}_4$, (c) $\text{H}_2\text{O}$.
Polarity from shape
A molecule is polar only if its bond dipoles do not cancel. Two steps:
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Are the bonds polar? (Compare electronegativities.)
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Is the shape symmetric? If yes, the dipoles cancel and the molecule is nonpolar.
$\text{CO}_2$ has two polar C=O bonds, but the molecule is linear, so the dipoles point in opposite directions and cancel. $\text{H}_2\text{O}$ is bent, so its dipoles add up and the molecule is polar. Use the shapes from the VSEPR page to decide.
Types of substances and their properties
The bonding type tells you what particles are present and what forces hold them. That predicts the properties.
| Type | Particles | Held by | Melting point | Conducts electricity? |
|---|---|---|---|---|
| Ionic (e.g., NaCl, MgO) | Cations and anions in a lattice | Ionic bonds (strong) | High | Not as a solid; yes when molten or dissolved |
| Covalent (molecular) (e.g., $\text{CO}_2$, $\text{CH}_3\text{OH}$) | Molecules | IMFs between molecules (weak) | Low | No |
| Metallic (e.g., Cu, Fe) | Cations in a "sea" of mobile electrons | Metallic bonds | Usually high | Yes, solid and liquid |
Why ionic solids conduct only when molten. Conducting needs charged particles that can move. In solid NaCl, ions are locked in place. When melted or dissolved, ions are free to flow.
Why metals conduct and bend. The valence electrons are free to move through the whole metal, carrying current. When you hammer a metal, layers of cations slide past each other and the electron sea keeps them bonded, so the metal is malleable and ductile (can be drawn into wire). In an ionic crystal, a shift lines up like charges, which repel, and the crystal shatters.
Why molecular solids melt easily. Melting a molecular substance only overcomes weak IMFs, not covalent bonds. That is why a candle melts around 60 °C.
Classify each substance as ionic, covalent (molecular), or metallic: (a) $\text{NaCl}$, (b) $\text{Cu}$, (c) $\text{CO}2$, (d) $\text{MgO}$.
Classify each substance as ionic, covalent (molecular), or metallic: (a) $\text{NaCl}$, (b) $\text{Cu}$, (c) $\text{CO}_2$, (d) $\text{MgO}$.
Aaliyah heats two white solids in the lab. Solid A melts at $801\,^\circ\text{C}$ and conducts electricity when molten but not when solid. Solid B melts at…
Aaliyah heats two white solids in the lab. Solid A melts at $801\,^\circ\text{C}$ and conducts electricity when molten but not when solid. Solid B melts at $60\,^\circ\text{C}$ and does not conduct electricity in either state. Classify each solid as ionic or molecular covalent, and explain your reasoning in terms of particles and forces.
Explain why copper can be drawn into wire and conducts electricity, while a crystal of sodium chloride shatters when struck and does not conduct as a solid. Refer to the…
Explain why copper can be drawn into wire and conducts electricity, while a crystal of sodium chloride shatters when struck and does not conduct as a solid. Refer to the type of bonding and the motion of charged particles in each case.
Predicting boiling and melting points
Stronger attractions between particles mean more energy is needed to separate them, so the melting and boiling points are higher. Use this order of strength:
dispersion (small molecules) < dipole-dipole < hydrogen bonding << ionic bonds
Compare substances in this order:
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Decide the type of substance (ionic, molecular, metallic, network).
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For molecular substances, find the strongest IMF.
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If the IMF is the same, compare size: more electrons means stronger dispersion.
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For ionic compounds, higher ion charge and smaller ions mean a stronger attraction.
The table lists boiling points at 1 atm. | Substance | Boiling point (°C) | |---|---| | $\text{CH}4$ | $-162$ | | $\text{NH}3$ | $-33$ | | $\text{H}2\text{O}$ | $100$ |…
The table lists boiling points at 1 atm.
| Substance | Boiling point (°C) |
|---|---|
| $\text{CH}_4$ | $-162$ |
| $\text{NH}_3$ | $-33$ |
| $\text{H}_2\text{O}$ | $100$ |
Identify the main intermolecular force in each substance and use it to explain the trend in boiling points.
The noble gases and halogens show increasing boiling points down a group. | Substance | Molar mass (g/mol) | Boiling point (K) | |---|---|---| | $\text{F}2$ | 38.0 | 85…
The noble gases and halogens show increasing boiling points down a group.
| Substance | Molar mass (g/mol) | Boiling point (K) |
|---|---|---|
| $\text{F}_2$ | 38.0 | 85 |
| $\text{Cl}_2$ | 70.9 | 239 |
| $\text{Br}_2$ | 159.8 | 332 |
| $\text{I}_2$ | 253.8 | 457 |
(a) Identify the intermolecular force acting in these nonpolar molecules. (b) Explain the trend using electron cloud size and polarizability. (c) Predict whether $\text{HF}$ (molar mass 20.0 g/mol, boiling point 293 K) follows this trend, and explain what additional force accounts for your conclusion.
Marcus must rank four substances by melting point: $\text{MgO}$, $\text{NaCl}$, $\text{CH}3\text{OH}$ (methanol), and $\text{CH}4$. (a) Classify each by bonding type and…
Marcus must rank four substances by melting point: $\text{MgO}$, $\text{NaCl}$, $\text{CH}_3\text{OH}$ (methanol), and $\text{CH}_4$. (a) Classify each by bonding type and name the force that must be overcome to melt it. (b) Rank them from lowest to highest melting point and justify the order. For $\text{MgO}$ versus $\text{NaCl}$, use ion charges (Mg$^{2+}$, O$^{2-}$ versus Na$^+$, Cl$^-$) in your argument. (c) State which of the four you expect to conduct electricity when molten.
Back to the start: water dries slowly because its strong hydrogen bonds hold molecules in the liquid, while rubbing alcohol has weaker IMFs overall and evaporates fast. Salt survives a hot oven because melting it means breaking a lattice of strong ionic bonds.