Ionic bonding

The claim says that simple covalent substances (for example, water or oxygen) melt at low temperatures because the covalent bonds inside the molecules are weak and easy to break. To judge this properly, we must separate what holds atoms together within a molecule from what holds molecules together in a substance.

In simple molecular (simple covalent) substances, there are two different kinds of attractions to keep straight:

1. Intramolecular forces: the covalent bonds joining atoms inside one molecule. These bonds are actually strong.
2. Intermolecular forces: weaker attractions between nearby molecules (for example London dispersion forces, and sometimes hydrogen bonding).

When a simple molecular substance melts or boils, the strong covalent bonds inside each molecule usually stay intact. The energy supplied is mainly used to overcome the intermolecular forces so the molecules can move past each other more freely.

Because these forces between molecules are weak, only a small amount of heat energy is needed, so the melting point is low.

The statement gives the wrong reason: it blames low melting points on weak covalent bonds. Covalent bonds are strong; it’s the weak intermolecular forces that explain the low melting points. So the statement cannot be correct.

Since melting does not involve breaking covalent bonds and those bonds are not weak, the statement is false.

The claim argues that simple covalent substances (such as $C{O}_2$ or $H_{2}O$) have low melting points because the covalent bonds inside the molecules are weak and easy to break. To decide true/false, we should confirm whether covalent bonds are actually weak.

A covalent bond is a strong electrostatic attraction between the shared pair of electrons and the nuclei of the bonded atoms.

So the statement is backwards: covalent bonds are generally very strong, and breaking them takes a lot of energy. If you broke the covalent bonds in water, you would be splitting hydrogen from oxygen—this is far more than simply changing ice into liquid water.

If covalent bonds are strong, why can melting points still be low?

During melting of a simple molecular substance, the covalent bonds within molecules are not broken. Instead, the substance only needs enough energy to overcome weak intermolecular forces between molecules. Because these attractions are weak, the energy required is small, giving low melting and boiling points.

Because the statement incorrectly says low melting points come from “weak covalent bonds,” it is not scientifically accurate. That means we eliminate Option A.

Covalent bonds are strong, and low melting points for simple covalent substances are due to weak intermolecular forces. Therefore, the statement is False.

Substances such as oxygen ($O_2$) and methane ($C{H}_4$) are simple molecular covalent substances. That means they are made of small, separate molecules rather than giant networks.

To explain why they are gases at room temperature, we need to think about what must be overcome when a substance changes state, linking ideas from Topic: Bonding and Topic: Matter.

In these substances, there are two different kinds of attractions:

1. Intramolecular bonds: the covalent bonds that hold atoms together within one molecule, and these are strong.
2. Intermolecular forces (IMFs): weaker attractions between neighboring molecules.

When a liquid evaporates or a substance boils to form a gas, the molecules do not get torn apart into atoms. The energy supplied mainly overcomes the weak intermolecular forces so the molecules can spread out and move freely.

Consider the first and third choices:

• Option A cannot be right because every atom contains electrons. Without electrons, atoms could not form covalent bonds, and molecules like $O_2$ and $C{H}_4$ would not exist.
• Option C is also wrong because simple covalent molecules have definite, predictable shapes. For instance, $C{H}_4$ has a tetrahedral arrangement (explained by VSEPR), so it absolutely does have a defined structure.

Option D reflects a common misunderstanding. Covalent bonds are generally strong, not fragile at room temperature. If covalent bonds broke easily, molecules like oxygen and methane would not remain as stable molecules.

Since $O_2$ and $C{H}_4$ are still intact as molecules at room temperature, the reason they are gases must be about weak attractions between molecules, not the breaking of covalent bonds within molecules.

Because $O_2$ and $C{H}_4$ are small, non-polar molecules, the main intermolecular forces present are very weak (London dispersion forces).

At room temperature, particles have enough thermal energy to overcome these weak forces, so the molecules separate easily and behave as gases. Therefore, the correct reason is that the forces between molecules are weak.

The question asks why simple covalent substances such as oxygen (${\text{O}}_2$) and methane (${\text{CH}}_4$) are gases at room temperature. We can answer by using the Process of Elimination to remove options that conflict with basic chemistry.

Option A claims there are no electrons. From Atomic structure, we know all atoms and molecules contain electrons. For instance, one oxygen atom has 8 electrons, and a methane molecule contains 10 electrons in total. These electrons allow at least temporary attractions (London dispersion forces), so Option A is not scientifically possible.

Option C says the substances lack a defined structure. But in Structure and bonding, we learn simple molecules have fixed shapes: oxygen is linear, and methane has a specific tetrahedral geometry. Since their atoms are arranged in definite ways, Option C must be incorrect.

Option D suggests covalent bonds are weak and break at room temperature. This is a common misconception. Covalent bonds (inside the molecule) are strong; if they broke easily, molecules would split into separate atoms. During boiling or evaporation, the covalent bonds remain intact, so Option D is wrong.

After eliminating A, C, and D, Option B remains. Simple molecular substances have low boiling points because only the intermolecular forces (weak attractions between molecules) need to be overcome. Room temperature provides enough energy for molecules to move apart and remain as gases, so the correct answer is that the intermolecular forces are weak.

The clue is the wording about the “ability of an atom to attract a shared pair of electrons.” That wording is used when talking about atoms joined in a covalent bond, and it links to periodic trends studied in Topic: Periodic table.

Polarity describes how charge is unevenly spread in a bond or molecule. It can happen because atoms attract electrons differently, but “polarity” names the outcome in the bond, not the atom’s ability to pull on the shared electrons.

Electron affinity is about the energy change when a gaseous atom gains an electron to form a negative ion:

$X(g)+e^{-}\to X^{-}(g)$

That situation involves adding an electron to an isolated atom, not describing how an atom attracts a shared pair inside a covalent bond.

Ionization energy refers to the energy needed to remove an electron from a gaseous atom:

$X(g)\to X^{+}(g)+e^{-}$

This is about taking an electron away (electron loss), so it does not match the idea of pulling on a shared electron pair.

From standard IUPAC and IB wording, electronegativity is the term for the relative ability of an atom to attract the shared pair of electrons in a covalent bond. It is commonly compared using values such as those on the Pauling scale.

We need the name of a property that measures how well an atom attracts a shared pair of electrons specifically when the atoms are connected by a covalent bond. The best approach is to remove choices that describe different ideas.

Ionization energy is the energy required to remove one mole of electrons from one mole of gaseous atoms.

Because it deals with an electron being taken from an isolated atom rather than attraction within a shared bond, it does not fit and can be eliminated.

Electron affinity is the energy change when a neutral gaseous atom gains an electron to become a negative ion.

Even though it involves attraction for an electron, it applies to a single atom in the gas phase, not to two atoms sharing a pair of electrons in a covalent bond.

Polarity describes how electrical charge is distributed across atoms in a bond (for example, whether a bond is polar covalent).

Polarity is what we observe as a consequence of different electron-pulling strengths, but it is not the name of the pulling-strength property itself.

After removing the other options, the remaining choice is Electronegativity. This term is defined as the relative measure of an atom’s ability to attract the shared electron pair in a covalent bond.

Periodic table comparisons often use scales (such as the Pauling scale) to show how electronegativity differs between elements.

To tackle this, recall how metallic bonding is described in Topic: Bonding. A metal is a repeating lattice of positive metal ions, and around them is a shared pool (a “sea”) of delocalised electrons. These electrons are not attached to one specific atom, so they can move throughout the whole metal.

Choice C talks about ionic bonds. But ionic bonding is found between a metal and a non-metal in compounds such as $NaCl$. A pure metal is not held together by ionic bonds; it is held by metallic bonding. So this option does not fit a metal’s structure and can be removed.

Choice D suggests there are big spaces between ions. Metals are actually closely packed in a giant lattice, which helps make the structure strong and dense. If there were large gaps, heat would not pass easily (gases with lots of space are poor conductors), so this is not a good explanation.

Choice A mentions vibrations of particles in fixed positions. That does happen in solids: when one part is heated, particles vibrate more and can pass energy along. However, this occurs in non-metals too and is not fast enough to explain why metals are especially good conductors. Metals are different because they also have mobile electrons that can move energy rapidly.

The best explanation is the fast movement of delocalised electrons through the lattice. These electrons pick up kinetic energy in the hot region and move through the structure quickly, transferring energy by collisions far more efficiently than lattice vibrations alone.

Metals conduct heat extremely well because of something special about their structure. Using Bonding: Structure and bonding, we can choose the best answer by crossing out options that are wrong for metals or that don’t explain why metals outperform most other solids.

All solids have particles that vibrate in place, and heating increases those vibrations so energy can be passed along. But materials like glass and wood also rely on this and still conduct heat poorly. So vibrations contribute, but they are not the main reason metals are so effective.

Ionic bonding happens when metals react with non-metals to form salts. A metal on its own is not an ionic lattice; it is a metallic lattice with metallic bonding. Therefore, the statement about ionic bonds throughout the lattice is not correct for a pure metal.

Metals are usually dense because the ions are arranged in a tightly packed, regular pattern. Large gaps would reduce particle collisions and act more like insulation (similar to trapped air in warm clothing), which would make heat transfer worse, not better.

The remaining choice is the movement of delocalised electrons. In metals, these electrons are free to travel through the lattice. When heated, they gain kinetic energy and move rapidly, colliding with ions and other electrons and spreading thermal energy much faster than vibrations alone can.