Source: CHEM1061 CLUE Worksheet 3.3, University of Minnesota Twin Cities
Difficulty: Introductory | Prerequisites: Basic understanding of atomic structure, electron configuration, and the periodic table (Worksheets 3.1 and 3.2 recommended).
This worksheet sits at the intersection of two big ideas in general chemistry: what holds atoms together within a molecule (intramolecular bonding, specifically covalent bonds) and what holds separate molecules or atoms near each other in a bulk substance (intermolecular forces, specifically London dispersion forces). You need both concepts to explain something as basic as why carbon is a solid at room temperature while nitrogen is a gas, even though they are neighbours on the periodic table. If you are comfortable with Lewis structures and electronegativity from earlier worksheets, you have the background you need here.
Elements that form extended covalent networks (like carbon, boron, beryllium, lithium) have dramatically higher melting and boiling points than elements that exist as small discrete molecules or lone atoms (like neon, helium, hydrogen). London dispersion forces are present in every substance, but they are the only attractive force between noble gas atoms and nonpolar small molecules, which is why those substances boil at very low temperatures.
Covalent bond
A chemical bond formed when two atoms share one or more pairs of electrons. In simple terms, atoms hold onto the same electrons rather than one stealing them from the other.
London dispersion forces (LDFs)
Weak, temporary attractive forces between all atoms and molecules, caused by momentary fluctuations in electron distribution that create instantaneous dipoles. Think of it as a brief flicker of uneven charge that nudges nearby atoms or molecules closer for an instant. Also called van der Waals forces or dispersion forces.
Intermolecular forces (IMFs)
Attractive forces that act between separate molecules or atoms in a substance. These are not the bonds within a molecule; they are the much weaker attractions between molecules that determine physical properties like boiling point.
Extended network (network covalent solid)
A structure in which atoms are connected by covalent bonds in a continuous three-dimensional lattice, with no discrete small molecules. Diamond (carbon) is the classic example. In simple terms, every atom is bonded to its neighbours in an enormous, unbroken web.
Small molecules (discrete molecules)
Substances that exist as individual, separate units with a fixed number of atoms (e.g. N₂, O₂, F₂, H₂). The atoms within each molecule are covalently bonded, but the molecules themselves are held together only by weak intermolecular forces.
Boiling point
The temperature at which a substance changes from liquid to gas. A higher boiling point means stronger forces must be overcome to pull particles apart. In simple terms, it tells you how "sticky" the particles in a liquid are.
Melting point
The temperature at which a solid becomes a liquid. Like boiling point, a higher melting point indicates stronger attractive forces between particles in the solid phase.
Metallic bonding
Bonding found in metals (e.g. lithium, beryllium) where valence electrons are delocalised across a lattice of metal cations. Think of it as a "sea of electrons" shared by all the metal atoms, which is why metals conduct electricity and have high melting points.
Ionic bonding
Bonding formed by the complete transfer of electrons from one atom to another, creating oppositely charged ions held together by electrostatic attraction. Lithium in some contexts is described this way, though elemental lithium is metallic.
The worksheet provides melting and boiling point data for ten elemental forms across Period 2 (plus hydrogen). The data splits cleanly into two groups.
Low boiling point group (below ~100 K): He (4.22 K), H₂ (20.28 K), Ne (27.07 K), N₂ (77.36 K), O₂ (90.20 K), F₂ (85.03 K)
These are all either noble gases (He, Ne) or diatomic molecules (H₂, N₂, O₂, F₂).
They exist as small, discrete particles. Between those particles, the only attractive forces are London dispersion forces.
LDFs are weak, so very little energy is needed to pull these particles apart, hence the low boiling points.
High boiling point group (above 1000 K): Li (1615 K), Be (2744 K), B (4273 K), C (4098 K)
These are all solids at room temperature.
Lithium and beryllium are metals held together by metallic bonding (delocalised electrons across a cation lattice).
Boron and carbon form extended covalent network solids, where atoms are linked by strong covalent bonds throughout the entire structure.
Breaking these structures apart requires overcoming very strong bonds, so boiling points are extremely high.
This distinction is the single most important concept in the worksheet. Whether a substance exists as small molecules or an extended network determines its physical properties.
Small molecules (He, Ne, H₂, N₂, O₂, F₂):
Each unit is a self-contained particle (one or two atoms).
Covalent bonds hold atoms together within each molecule, but molecules interact with each other only through weak LDFs.
Result: gases or liquids at room temperature, very low boiling points.
Extended networks (B, C) and metallic lattices (Li, Be):
There are no discrete small molecules. The bonding extends continuously through the material.
To boil the substance, you must break strong covalent or metallic bonds, not just weak intermolecular forces.
Result: solids at room temperature with very high boiling points.
Every substance in the table has LDFs. This is a universal force. Noble gases, diatomic molecules, metals, network solids: all of them experience London dispersion forces. The question is whether LDFs are the dominant interaction (as in noble gases and small molecules) or a minor addition on top of much stronger bonding (as in network solids and metals).
This is a classic exam comparison. Both nitrogen and carbon are Period 2 nonmetals, but they behave very differently at room temperature.
Nitrogen (N₂): exists as diatomic molecules. Each molecule has a strong triple bond (N≡N) holding the two atoms together, but between separate N₂ molecules the only attraction is LDFs. Room temperature (~298 K) is well above nitrogen's boiling point (77.36 K), so nitrogen is a gas.
Carbon (C): forms an extended covalent network (diamond or graphite). Every carbon atom is bonded to its neighbours by strong covalent bonds throughout the entire solid. There are no separate small molecules to pull apart. You have to break covalent bonds to melt or boil carbon, which takes enormous energy (boiling point 4098 K). Carbon is therefore a solid at room temperature.
The key insight: the strength of the bond within an N₂ molecule is irrelevant to its boiling point. What matters is the force between molecules. A molecule can have very strong internal bonds and still be a gas if the forces between molecules are weak.
Element | Formula | Melting Point (K) | Boiling Point (K) | Bonding Type | Structure |
|---|---|---|---|---|---|
Hydrogen | H₂ | 13.81 | 20.28 | Covalent | Small molecules |
Helium | He | 0.95 | 4.22 | None (noble gas) | Small molecules |
Lithium | Li(s) | 453.65 | 1615 | Metallic | Extended network |
Beryllium | Be(s) | 1560 | 2744 | Metallic | Extended network |
Boron | B(s) | 2348 | 4273 | Covalent (network) | Extended network |
Carbon | C(s) | 3823 | 4098 | Covalent (network) | Extended network |
Nitrogen | N₂ | 63.15 | 77.36 | Covalent | Small molecules |
Oxygen | O₂ | 54.36 | 90.20 | Covalent | Small molecules |
Fluorine | F₂ | 53.53 | 85.03 | Covalent | Small molecules |
Neon | Ne | 24.56 | 27.07 | None (noble gas) | Small molecules |
The reason diamond is one of the hardest known materials is precisely because it is an extended covalent network: breaking it means breaking C-C bonds throughout the lattice, which requires enormous energy. Graphite (another form of carbon) is used in pencils because its layers, while covalently bonded within each sheet, slide over each other due to weak LDFs between the sheets.
Nitrogen's low boiling point is why liquid nitrogen is used as a cryogenic coolant. It boils at 77 K (-196 °C), well below room temperature, so it evaporates readily and absorbs a large amount of heat as it does so. The weakness of LDFs between N₂ molecules is what makes this possible.
"Stronger bonds within the molecule means a higher boiling point." This is the most common mistake. N₂ has one of the strongest bonds in chemistry (the triple bond), yet its boiling point is only 77 K. Boiling point depends on forces between molecules, not the strength of bonds within them.
"Noble gases have no forces acting on them." Incorrect. Noble gases experience London dispersion forces. That is why helium and neon have nonzero boiling points (4.22 K and 27.07 K, respectively). Those forces are just very weak.
"LDFs only exist in nonpolar molecules." LDFs are present in every substance, polar or nonpolar, molecular or network. They are universal. In polar molecules, LDFs exist alongside dipole-dipole forces.
"Carbon is a solid because it has covalent bonds, and nitrogen is a gas because it has covalent bonds that are weaker." Both have covalent bonds, but the critical difference is structural. Carbon forms an extended network; nitrogen forms discrete N₂ molecules. The bonding pattern, not bond strength alone, determines the phase.
⚠️ Expect a question asking you to explain why two elements with similar atomic numbers have very different boiling points. The answer always comes back to structure: extended network vs small molecules.
⚠️ "Do all substances have LDFs?" is a common exam question. The answer is yes, always, for every substance.
⚠️ Be prepared to classify each Period 2 element by bonding type (metallic, network covalent, molecular covalent, or none/noble gas) and by structure (extended network or small molecules).
⚠️ The nitrogen vs carbon comparison is a favourite exam question. Know why N₂ is a gas and C is a solid, and be able to explain it in terms of structure, not just "bond strength."
True or False: London dispersion forces are present only in nonpolar substances. (False: LDFs are present in all substances.)
Fill in the blank: Carbon is a solid at room temperature because it forms a(n) __________ rather than discrete molecules. (extended covalent network)
True or False: A substance with very strong covalent bonds within its molecules will always have a high boiling point. (False: boiling point depends on intermolecular forces, not intramolecular bond strength.)
Fill in the blank: He, Ne, H₂, N₂, O₂, and F₂ all have low boiling points because the dominant attractive force between their particles is __________. (London dispersion forces)
True or False: Lithium and beryllium have high boiling points because of metallic bonding. (True)
Q: Why does helium have the lowest boiling point of all the elements in the table?
A: Helium is a noble gas with no covalent bonds and exists as single atoms. The only attractive force between helium atoms is London dispersion forces, and helium has the fewest electrons of any element (just 2), making its LDFs the weakest. Very little energy is needed to separate helium atoms into the gas phase.
Q: Explain why boron (boiling point 4273 K) and fluorine (boiling point 85.03 K) have such different boiling points, even though both contain covalent bonds.
A: Boron forms an extended covalent network solid. To boil boron, you must break strong covalent bonds throughout the entire lattice, which requires enormous energy. Fluorine exists as discrete F₂ molecules. The covalent bond within each F₂ molecule is strong, but the forces between separate F₂ molecules are only weak LDFs. Boiling fluorine only requires overcoming those weak intermolecular forces.
Q: A student claims that nitrogen has a low boiling point because the N≡N triple bond is easy to break. Is this correct? Explain.
A: This is incorrect. The N≡N triple bond is one of the strongest bonds in chemistry (bond energy ~945 kJ/mol). The boiling point is low not because the triple bond is weak, but because N₂ molecules interact with each other only through weak London dispersion forces. Boiling does not break the N≡N bond; it only separates N₂ molecules from one another.
Q: Do all substances in the table experience London dispersion forces? Justify your answer.
A: Yes. LDFs arise from temporary fluctuations in electron distribution and are present in every substance that has electrons, which is all matter. Noble gases, diatomic molecules, metals, and network solids all experience LDFs. In substances like He and Ne, LDFs are the only attractive force. In metals and network solids, LDFs are present but negligible compared to metallic or covalent bonding.
Q: Classify each of the following as existing in an extended network or as small molecules: Li, Ne, C, O₂, Be, F₂.
A: Extended network: Li, C, Be. Small molecules: Ne, O₂, F₂.
This material connects directly to intermolecular forces (dipole-dipole, hydrogen bonding) covered later in the course. LDFs are the baseline; once you understand them, the other IMF types layer on top for polar molecules. It also ties into the concept of electronegativity and polarity from earlier worksheets, because the type of bond (ionic, covalent, metallic) depends on the electronegativity difference between atoms.
The network solid concept reappears when studying materials science and solid-state chemistry, particularly allotropes of carbon (diamond, graphite, graphene, fullerenes), silicon dioxide, and other ceramics.
Covalent bonding, London dispersion forces, LDFs, van der Waals forces, dispersion forces, intermolecular forces, IMFs, network covalent solid, extended network, discrete molecules, small molecules, diatomic molecules, metallic bonding, boiling point, melting point, phase of matter, noble gases, Period 2 elements, CHEM 1061, general chemistry, intramolecular vs intermolecular, bond strength vs boiling point, nitrogen vs carbon, diamond structure, graphite, allotropes