Covalent bonding

The sentence argues that simple covalent substances (such as water or oxygen) have low melting points because the covalent bonds inside their molecules are weak and easy to snap. 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 substances there are two different kinds of forces to think about:

1. Intramolecular forces: the covalent bonds joining atoms within one molecule. These are generally strong.
2. Intermolecular forces: weaker attractions between nearby molecules (for instance, London dispersion forces or hydrogen bonding).

When a simple molecular substance melts or boils, the molecules themselves stay intact—so we are not breaking the strong covalent bonds. The energy supplied is mainly used to overcome the intermolecular forces that keep molecules close together in the solid or liquid.

Because those intermolecular forces are weak, only a small amount of heat is needed to separate the molecules, giving a low melting point.

The claim is wrong because it blames low melting points on weak covalent bonds. Covalent bonds are strong; it is the weak forces between molecules that explain the low melting point. So the option that treats the statement as true should be rejected.

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

The statement claims that simple covalent substances (like $C{O}_2$ or $H_{2}O$) melt easily because the covalent bonds inside the molecules are weak and easy to break. To decide true/false, we should verify whether covalent bonds are actually weak.

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

This means covalent bonds are typically very strong, and breaking them needs 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 turning ice into liquid water.

So if covalent bonds are strong, why do many simple molecular substances have low melting points?

During melting, the covalent bonds within each molecule remain unbroken. The change of state happens because we overcome the weak intermolecular forces between separate molecules. Since these forces are weak, only a small amount of energy is required, leading to low melting and boiling points.

Because the statement incorrectly says low melting points are due to “weak covalent bonds,” it is not scientifically accurate. That means Option A should be eliminated.

Covalent bonds are strong, and low melting points in simple molecular substances are explained by weak intermolecular forces. Therefore, the statement is False.

The question is about conductivity in a solid copper wire. The key idea is to recognise copper as a metal. That means we should use what we know about Bonding, especially structure and bonding in metals, to explain how charge moves in it.

In metallic bonding, a metal is pictured as a repeating lattice of positive metal ions (cations) with a surrounding ‘sea’ of delocalised valence electrons.

These electrons are shared by the whole structure rather than belonging to one specific atom, so they can move through the solid.

Electrical conduction needs charged particles that are free to move.

In solid copper:

• The copper ions stay in fixed positions in the lattice, so they do not travel through the wire.
• The delocalised electrons can drift through the metal when a potential difference (voltage) is applied.

So, the delocalised electrons are the mobile charge carriers.

Here is why the other answers don’t work:

• Dissolved salts: They only help conduct in solutions (or melts) where ions can move, not inside a solid metal wire.
• Copper ions (${\text{Cu}}^{2+}$): These ions exist in the metallic lattice but are not free to flow through the solid.
• Protons: Protons are in the nucleus and are not the particles that move to produce an electric current.

Using the metallic bonding idea, solid copper conducts because its delocalised valence electrons can move through the lattice and carry charge.

The question is asking which particles carry charge in a solid copper wire. With elimination, we think about which particles are genuinely mobile inside a solid metal lattice, from Bonding: Structure and bonding.

Dissolved salts conduct when their ions are free to move in an aqueous solution or a molten state. A solid copper wire isn’t a solution or a liquid, so there are no dissolved ions moving around. This option must be wrong.

Protons are trapped in the atom’s nucleus and held by very strong forces. They do not move during normal electrical conduction. If protons were moving through the wire, it would involve nuclear changes and the copper would no longer remain the same element.

Copper metal can be described as positive ions (${\text{Cu}}^{2+}$) in a fixed crystal lattice. In the solid, these ions can vibrate but cannot migrate through the wire to form a current. Ions only become mobile charge carriers when a substance is molten or dissolved.

After eliminating the others, delocalised electrons are left. In metals, these outer electrons are not tied to individual atoms; they form a mobile ‘sea’ of negative charge that can drift through the solid lattice when a voltage is applied.

Substances such as oxygen (${\text{O}}_2$) and methane (${\text{CH}}_4$) are classified as simple molecular covalent substances. That means they are made of small, separate molecules rather than a huge connected network.

To explain why they are gases, we focus on what must happen when particles spread out during a change of state. This links ideas from Topic: Bonding and Topic: Matter.

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

1. Intramolecular bonds: the covalent bonds that strongly hold atoms together within each molecule.
2. Intermolecular forces (IMFs): much weaker attractions that act between one molecule and another.

When a liquid turns into a gas (boiling/evaporating), the molecules do not need to be torn apart into atoms. The energy supplied mainly overcomes the intermolecular forces, allowing whole molecules to separate and move freely.

Consider these options:

• Option A cannot be right because atoms always contain electrons. Without electrons, covalent bonding would not be possible, and matter as we know it would not form.
• Option C is also wrong because simple covalent molecules have definite shapes. For instance, ${\text{CH}}_4$ has a specific tetrahedral geometry (explained using VSEPR), so it absolutely does have a defined structure.

Option D describes a common misunderstanding. Covalent bonds are generally strong. If covalent bonds were weak enough to break at room temperature, ${\text{O}}_2$ and ${\text{CH}}_4$ would not remain as stable molecules—they would split into separate atoms.

Because these substances exist as intact molecules at room temperature, we know the covalent bonds are not being broken. Their gaseous state must instead be explained by weak attractions between molecules.

${\text{O}}_2$ and ${\text{CH}}_4$ are small, non-polar molecules, so the main attractions between their molecules are very weak (London dispersion forces).

At room temperature, particles have enough thermal energy to overcome these weak intermolecular attractions, so the molecules spread out and move around as a gas. Therefore, the correct reason is that the forces between molecules are weak.

The question asks why simple covalent substances like oxygen (${\text{O}}_2$) and methane (${\text{CH}}_4$) are gases at room temperature. We can find the correct choice by systematically removing options that conflict with basic chemistry facts.

Option A claims there are no electrons. From Atomic structure, we know every atom contains electrons. An oxygen atom has 8 electrons, and a methane molecule contains 10 electrons in total. Those electrons allow at least temporary attractions (London dispersion forces), so Option A is scientifically incorrect.

Option C says there is no defined structure. But in Structure and bonding, we learn that molecules have specific shapes: oxygen is linear, and methane has a precise tetrahedral arrangement. Since their atoms are arranged in definite geometries, Option C must be wrong.

Option D suggests covalent bonds break easily at room temperature. That is incorrect: covalent bonds (inside a molecule) are strong. If they broke under everyday conditions, the molecules would not remain as ${\text{O}}_2$ and ${\text{CH}}_4$. During boiling/evaporation, molecules separate from each other while the covalent bonds stay intact.

After eliminating A, C, and D, Option B remains. Simple molecular substances have low boiling points because the intermolecular forces between molecules are weak and need little energy to overcome. Room temperature provides enough energy for the molecules to stay far apart and move freely as a gas.

You are being asked for the term that means an atom’s pull on a pair of electrons that is shared in a covalent bond. That wording points to a bonding property describing how atoms compete for the bonding electrons, which links to periodic trends in the Topic: Periodic table.

Polarity describes whether charge is unevenly spread in a bond or molecule. It can happen because atoms attract bonding electrons differently, but polarity is the outcome, not the term that names the ability to attract the shared electrons.

Electron affinity refers to the energy change when a gaseous atom gains an extra electron to become a negative ion:

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

This is about adding an electron to a single atom, not about pulling on a shared pair of electrons in a covalent bond.

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

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

That describes electron removal from an atom, which is different from attracting electrons that are being shared in a bond.

From standard IUPAC/IB wording, electronegativity is the term for the relative strength with which an atom attracts the shared electron pair in a covalent bond. It is commonly compared using values such as the Pauling scale.

The wording points to a property describing an atom’s pull on a shared pair of electrons specifically when the atoms are joined by a covalent bond. So we can eliminate choices that describe energy changes for isolated atoms instead.

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

Because it deals with taking electrons away from separate atoms (not attracting a bonding pair), it does not fit and can be eliminated.

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

Even though it involves attraction to an electron, it is about a lone atom gaining an electron, not two atoms sharing a pair in a covalent bond, so it can be ruled out.

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

Polarity happens because atoms have different pulls on electrons, but it names the bond’s uneven charge pattern, not the atom’s ability itself.

After eliminating the other terms, Electronegativity remains. This is the accepted definition for the relative measure of how strongly an atom attracts the shared electron pair in a covalent bond.

On the periodic table, scales such as the Pauling scale are used to compare electronegativities between elements.