Intermolecular Forces — Science Sage Gr10

Intramolecular vs Intermolecular Forces

It is essential to distinguish between two very different types of forces in chemistry:

Intramolecular forces: forces WITHIN a molecule — these are the covalent bonds, ionic bonds, or metallic bonds that hold atoms together within a substance. They determine the chemical properties of the substance (what reactions it undergoes, what products form).
Intermolecular forces (IMF): forces BETWEEN separate molecules — these hold molecules together in a bulk sample of liquid or solid. They do NOT create new bonds; they are attractions between neighbouring molecules. They determine the physical properties of the substance (melting point, boiling point, viscosity, surface tension, solubility).

Key distinction: When water boils, the intermolecular forces (hydrogen bonds between H₂O molecules) are overcome — NOT the O–H covalent bonds within the water molecules. The O–H bonds remain intact in steam.

Type 1 — London Dispersion Forces (Van der Waals Forces)

London dispersion forces (LDF) are present in ALL molecules — both polar and non-polar. They arise because the electron cloud around any atom is constantly moving. At any instant, the electrons may be distributed unevenly, creating a TEMPORARY dipole. This temporary dipole induces a dipole in neighbouring molecules, and a weak attraction results.

Present in ALL substances (the only IMF in non-polar molecules)
Weakest type of intermolecular force
Strength increases with: (1) molar mass — more electrons means larger, more polarisable electron clouds; (2) surface area — more contact between molecules means stronger LDF
Examples: noble gases (liquid He, liquid Ar), non-polar molecules (I₂, Cl₂, Br₂, CH₄, all alkanes)

LDF: strength ∝ molar mass and surface area. More electrons → bigger instantaneous dipoles → stronger LDF

Molecule	Molar Mass (g·mol⁻¹)	BP (°C)	IMF type
CH₄	16	−161	LDF only
C₂H₆	30	−89	LDF only
C₃H₈	44	−42	LDF only
I₂	254	+184	LDF only (very large molecule)

Type 2 — Dipole-Dipole Interactions

Dipole-dipole interactions occur between polar molecules. A polar molecule has a PERMANENT dipole — one end is slightly positive (δ+) and the other is slightly negative (δ−), because of unequal sharing of electrons in a polar bond combined with an asymmetric shape.

When two polar molecules are near each other, they align so that the δ+ end of one molecule is attracted to the δ− end of another. This electrostatic attraction is a dipole-dipole interaction.

Stronger than LDF for molecules of similar molar mass
Only between polar molecules
Examples: HCl, SO₂, CHCl₃ (chloroform), acetone (CH₃COCH₃)

Note: LDF also acts between polar molecules — dipole-dipole interactions are IN ADDITION to LDF, not instead of them.

Type 3 — Hydrogen Bonding (The Strongest IMF)

Hydrogen bonding is a special, unusually strong type of dipole-dipole interaction. It occurs when a hydrogen atom is covalently bonded to a highly electronegative, small atom (N, O, or F). The enormous electronegativity difference creates a very strong δ+ on the H atom. This H atom is then attracted to a lone pair on a N, O, or F atom of an ADJACENT molecule.

X−H ··· Y (where X and Y are N, O, or F; the dashed line ··· is the H-bond)

Only occurs when H is bonded directly to N, O, or F
Much stronger than ordinary dipole-dipole forces
Still weaker than covalent bonds (about 1/10 the strength)
Explains anomalously high boiling points of H₂O, HF, and NH₃

Compound	Molar Mass (g·mol⁻¹)	BP (°C)	IMF
HF	20	+19.5	H-bonding (anomalously high)
HCl	36.5	−85	Dipole-dipole + LDF
HBr	81	−67	Dipole-dipole + LDF
HI	128	−35	Dipole-dipole + LDF (stronger LDF)

Note: HF has a much lower molar mass than HI yet a much higher boiling point. This is because hydrogen bonding in HF is far stronger than the LDF in HI. As a general rule, going from HCl to HBr to HI, BP increases with molar mass (increasing LDF). But HF is the exception because it has strong H-bonding.

Anomalous Properties of Water

Water (H₂O) is the most important example of hydrogen bonding:

Unusually high boiling point (100°C) for its small molar mass (18 g·mol⁻¹) — due to strong H-bonding
High surface tension — H-bonds hold surface molecules together tightly
High specific heat capacity — requires lots of energy to break H-bonds
Ice is LESS DENSE than liquid water — in ice, H₂O molecules form a rigid hexagonal lattice with H-bonds holding them further apart than in liquid water. This is why ice floats.
Biological importance: DNA base pairs held together by H-bonds (A–T: 2 H-bonds; G–C: 3 H-bonds); protein secondary structure (alpha-helices and beta-sheets); cellulose structure in plant cell walls

Effect of IMF on Physical Properties

Property	Effect of Stronger IMF	Explanation
Boiling point	Higher	More energy needed to separate molecules
Melting point	Higher	Molecules held more firmly in solid structure
Viscosity	Higher	Molecules resist flowing past each other
Surface tension	Higher	Molecules attract each other at the surface more strongly
Vapour pressure	Lower	Fewer molecules have enough energy to escape the liquid surface
Solubility	"Like dissolves like"	Polar dissolves polar/ionic; non-polar dissolves non-polar

"Like dissolves like": Polar solvents (like water) dissolve polar and ionic substances. Non-polar solvents (like hexane) dissolve non-polar substances. Vitamin C (polar) dissolves in water; vitamin A (non-polar) dissolves in fat.

IEB Extension — Ion-Dipole Forces & Applications

Ion-dipole forces occur between an ion (from an ionic compound) and a polar molecule (solvent). When NaCl dissolves in water, the Na⁺ ions attract the δ− oxygen ends of water molecules, and Cl⁻ ions attract the δ+ hydrogen ends. Each ion is surrounded by a "hydration shell" of water molecules. The ion-dipole attraction must overcome the ionic lattice energy — this is possible because many ion-dipole interactions act together, releasing sufficient energy (hydration enthalpy).

Soap and detergent molecules (e.g. sodium stearate) have a dual structure:

Hydrophilic head (ionic carboxylate end, –COO⁻Na⁺): attracted to water molecules via ion-dipole forces
Hydrophobic tail (long non-polar hydrocarbon chain): dissolves non-polar grease/oil via LDF

When washing with soap, the hydrophobic tails dissolve into grease particles while the hydrophilic heads remain in the water. This creates micelles — spherical clusters with grease inside and ionic heads on the outside — that are easily rinsed away in water.

Biological membranes (phospholipid bilayers) use the same principle: phospholipid molecules have a polar head and two non-polar fatty acid tails. In water, they spontaneously form a bilayer with tails facing inward (away from water) and heads facing outward (towards water), forming the basis of all cell membranes.

IMF Visualiser — London Dispersion

Force Type

Force Info

Question 1 (2 marks)

Which type of intermolecular force is present in ALL molecules, including non-polar ones?

London dispersion forces (LDF) arise from temporary dipoles due to electron movement and are present in ALL substances — polar and non-polar. They are the only IMF in non-polar molecules like I₂ or CH₄.

Question 2 (2 marks)

The boiling point of H₂O is 100°C while H₂S (a larger molecule) boils at −60°C. The BEST explanation is:

H₂O forms strong hydrogen bonds (O–H···O) because oxygen is highly electronegative and small. Sulfur is less electronegative and larger, so H₂S only has dipole-dipole and LDF — much weaker. This gives H₂O a much higher BP despite its smaller size.

Question 3 (2 marks)

Which of the following molecules can form hydrogen bonds with water?

Hydrogen bonding requires H bonded to N, O, or F. NH₃ has H atoms bonded to nitrogen (highly electronegative) and lone pairs on N, so NH₃ can both donate and accept H-bonds with water. CH₄, CO₂, and CCl₄ cannot form H-bonds.

Question 4 (1 mark)

Ice is less dense than liquid water because:

In ice, hydrogen bonds lock water molecules into a hexagonal lattice structure where molecules are held further apart than in liquid water. This makes ice less dense, which is why ice floats on water — a property crucial for aquatic life surviving in frozen lakes.

Question 5 (1 mark)

Iodine (I₂) is soluble in hexane (C₆H₁₄) but insoluble in water. This is best explained by:

'Like dissolves like': non-polar I₂ dissolves in non-polar hexane because both have LDF as their only IMF. Water is polar and forms hydrogen bonds — I₂ cannot interact with water molecules strongly enough to dissolve.

Question 6 (2 marks)

Going from HCl to HBr to HI, the boiling point increases. The main reason is:

HCl, HBr, and HI all have dipole-dipole forces and LDF, but NO hydrogen bonding (H is not bonded to N, O or F here). As molar mass increases (HCl 36.5 → HI 128 g·mol⁻¹), the larger electron cloud creates stronger LDF, raising the boiling point.

Question 7 (3 marks)

Butane (C₄H₁₀, non-polar, molar mass 58 g·mol⁻¹) has a boiling point of −1°C. Propan-1-ol (C₃H₇OH, molar mass 60 g·mol⁻¹) has a boiling point of 97°C. Despite the two substances having almost identical molar mass, propan-1-ol's boiling point is far higher. What is the BEST explanation?

Because the molar masses are almost identical, LDF strength should be similar for both — molar mass cannot explain the huge BP gap. The real difference is the O–H group in propan-1-ol: oxygen is small and highly electronegative, so propan-1-ol molecules hydrogen-bond to each other. Butane is completely non-polar (only C–H and C–C bonds) and has no N, O or F for H bonding, so it relies only on the much weaker LDF — hence its far lower boiling point.

Question 8 (2 marks)

Two liquids, X and Y, have almost the same molar mass. At the same temperature, liquid X has a much higher vapour pressure than liquid Y. What can you conclude about their intermolecular forces, and why?

Stronger IMF hold molecules in the liquid more tightly, so fewer molecules can escape into the vapour phase, giving a LOWER vapour pressure. Since X has the HIGHER vapour pressure (and similar molar mass, so LDF should be similar), X's molecules must be held together more weakly than Y's — X has the weaker overall intermolecular forces of the two.

IEB Question 1 IEB (2 marks)

When sodium chloride (NaCl) dissolves in water, the process is called hydration. What type of intermolecular force is responsible for holding the water molecules around the Na⁺ and Cl⁻ ions?

Ion-dipole forces act between the charged ions (Na⁺ or Cl⁻) and the polar water molecules (which have δ+ on H and δ− on O). Na⁺ attracts the δ− oxygen of water; Cl⁻ attracts the δ+ hydrogen. Many of these ion-dipole interactions collectively provide enough energy to overcome the lattice energy of NaCl, allowing it to dissolve.

IEB Question 2 IEB (1 mark)

A soap molecule has a long non-polar hydrocarbon tail and an ionic carboxylate head (−COO⁻Na⁺). How does this dual structure allow soap to remove grease from surfaces when washed with water?

Soap molecules form micelles: the hydrophobic (non-polar) hydrocarbon tails dissolve into the grease particle through London dispersion forces, while the hydrophilic (ionic) carboxylate heads remain in the surrounding water through ion-dipole attractions. The grease is trapped inside the micelle and the whole structure is water-soluble — easily rinsed away.

Answer in full sentences. Use diagrams where helpful (described in words).

Question 1

Rank the following substances from LOWEST to HIGHEST boiling point and justify your ranking based on intermolecular forces: CH₄, H₂O, HCl, Ne.

Question 2

Water has many anomalous properties compared to other small covalent molecules. (a) State THREE physical properties of water that are anomalously high or unusual. (b) Explain ONE of these properties in terms of hydrogen bonding.

Question 3

Identify the strongest type of intermolecular force present in each of the following substances: (a) HF (b) CO₂ (c) CH₃OH (methanol) (d) Ar (e) NH₃

Question 4

Explain, using the concept of 'like dissolves like', why: (a) sodium chloride (NaCl) dissolves in water but not in hexane; (b) iodine (I₂) dissolves in hexane but not in water.

Question 5

Draw a diagram (described in words) showing three water molecules held together by hydrogen bonding. Label: (a) the O–H covalent bonds within each molecule; (b) the O···H–O hydrogen bonds between molecules; (c) the δ+ and δ− partial charges. Explain why these hydrogen bonds make water's boiling point much higher than expected.

Question 6

The table shows the molar mass and boiling point of the four halogen elements (all non-polar diatomic X₂ molecules).

Molecule	Molar mass (g·mol⁻¹)	Melting point (°C)	Boiling point (°C)
F₂	38	−220	−188
Cl₂	71	−101	−34
Br₂	160	−7	59
I₂	254	114	184

(a) Describe the trend between molar mass and boiling point shown in the table.
(b) All four molecules are non-polar. Which single type of intermolecular force must be responsible for holding them together as liquids/solids, and why does it get stronger down the group?
(c) Astatine (At₂) is the next halogen below iodine, with a molar mass of about 420 g·mol⁻¹. Using the trend in the table, predict whether At₂'s boiling point would be higher or lower than I₂'s, and justify your answer using the data.
(d) At room temperature (25°C), F₂ and Cl₂ are gases, Br₂ is a liquid, and I₂ is a solid. Use the boiling point (and melting point) data to explain why the four halogens are in different physical states at the same temperature.