How bonding affects properties

To explain why methane (${\text{CH}}_4$) is a gas at room temperature, focus on what controls melting and boiling.

A substance’s state (solid, liquid, gas) depends mainly on how strongly its particles attract each other between molecules, not on the bonds within each molecule. This fits the core idea in Matter: States and properties of matter about why some substances are gases or liquids at room temperature.

In simple molecular substances, two different kinds of attractions matter:

1. Covalent Bonds (Intramolecular): very strong links holding atoms together inside one ${\text{CH}}_4$ molecule (carbon joined to hydrogen atoms).
2. Intermolecular Forces: much weaker pulls between neighboring ${\text{CH}}_4$ molecules.

When methane melts or boils, the molecule stays intact—the covalent bonds are not broken. Instead, the weak intermolecular forces are the ones that get overcome.

Option A says covalent bonds are weak. That’s not true: covalent bonds are strong chemical bonds. If they were weak, ${\text{CH}}_4$ molecules would split into separate atoms easily, which is not what happens during boiling or when a gas forms.

Option B suggests the atoms aren’t bonded. But methane is defined as a molecule with one carbon atom covalently bonded to four hydrogen atoms. Without those bonds, you wouldn’t have a substance with the formula ${\text{CH}}_4$.

Option D claims there are no electrons, which cannot be correct. Methane has valence electrons that are shared in covalent bonds, giving a stable arrangement (similar in goal to achieving a noble-gas-like configuration). Without electrons, covalent bonding could not occur.

Simple molecular substances like methane have very weak intermolecular forces. Because these attractions are so small, even the thermal energy available at room temperature can be enough to overcome them, letting molecules separate and move freely—so the substance is commonly a gas or a liquid.

The question is asking why simple molecular substances like methane (${\text{CH}}_4$) are usually gases or liquids at room temperature. Melting and boiling happen when particles can separate from each other, which depends on attractions between particles rather than the bonds within particles.

Consider option A: it claims covalent bonds are weak. In Topic 4: Chemical Bonding and Structure, covalent bonding (sharing electron pairs) is taught as a strong type of bonding. If covalent bonds were weak, the ${\text{CH}}_4$ molecule would break apart easily, which is not what occurs during a change of state. So this option can be eliminated.

Option B says the atoms are not bonded, but that conflicts with what a molecule is. Methane ($C{H}_4$) is one carbon atom covalently bonded to four hydrogen atoms. If those bonds were absent, there would be no distinct compound with that formula. Therefore, this option is incorrect.

Option D suggests there are no electrons. From Atomic Structure, you know atoms contain electrons, and electrons are necessary for forming covalent bonds. A stable molecule like methane must involve electrons, so this option cannot be right.

After removing the incorrect choices, option C remains.

In simple molecular substances, the covalent bonds inside each molecule are strong, but the forces between separate molecules (intermolecular forces) are weak. Since only a small amount of thermal energy is needed to overcome these weak forces, the molecules can separate easily—making these substances gases or liquids at room temperature.

To answer this, we must tell apart the two main reasons a substance might conduct electricity:

1. Delocalised electrons: Present in metals (whether solid or molten) and in graphite. These electrons are not tied to one atom, so they can travel through the structure.
2. Mobile ions: Present in ionic compounds only when they are molten or dissolved in water (aqueous). As solids, ionic compounds don’t conduct because their ions are stuck in a rigid lattice.

We need the option where all listed substances conduct using mechanism 1.

Check Option B one by one:

• Aqueous sodium chloride: conducts because ions ($N{a}^{+}$ and $C{l}^{-}$) can move around in solution.
• Molten magnesium oxide: conducts because ions ($M{g}^{2+}$ and $O^{2-}$) are free to move when liquid.
• Iron: a metal, so it conducts using delocalised electrons.

Because the first two rely on ions rather than electrons, Option B does not meet the requirement.

Now evaluate Option C:

• Graphite: carbon atoms bond to three neighbors, leaving one electron per atom available as a delocalised electron for conduction.
• Iron: conducts via delocalised electrons because it is metallic.
• Solid sodium chloride: an ionic solid where ions cannot move in the lattice, so it does not conduct electricity.

Since one substance in the list is a non-conductor, Option C is not correct.

Consider Option D:

• Diamond: a giant covalent structure where each carbon’s valence electrons are held in strong bonds, so there are no free electrons; it does not conduct.
• Copper: a metal, so it conducts via delocalised electrons.
• Molten sodium chloride: conducts due to mobile ions, not electrons.

Because diamond is an insulator and molten sodium chloride conducts by ions, Option D is incorrect.

Finally, test Option A:

• Copper: a solid metal with a ‘sea’ of delocalised electrons.
• Graphite: has delocalised electrons that can move through the structure.
• Molten aluminium: even when melted, a metal still has delocalised electrons available to carry charge.

All three conduct electricity through delocalised electrons, so Option A is the correct list.

The question is specifically asking for substances that conduct electricity through delocalised electrons.

This type of conduction is typical of:

1. Metals (both solid and molten), due to metallic bonding with a ‘sea’ of delocalised valence electrons.
2. Graphite, where each carbon bonds to three others, leaving one delocalised electron per carbon atom.

So we will cross out any option that includes ionic conduction or a substance that doesn’t conduct at all.

Ionic substances conduct in the molten ($l$) or aqueous ($aq$) state because ions move (e.g., $N{a}^{+}$, $C{l}^{-}$), not because electrons move.

• Option B includes aqueous sodium chloride and molten magnesium oxide, both of which conduct via ions.
• Option D includes molten sodium chloride, which also conducts via ions.

Therefore, Options B and D must be eliminated.

Next, exclude substances that do not conduct electricity:

• Solid ionic compounds: In Option C, solid sodium chloride cannot conduct because its ions are fixed in place.
• Diamond: In Option D, all electrons are locked into covalent bonds, so there are no delocalised electrons to carry charge.

This confirms that Options C and D are not suitable.

Check the remaining list:

1. Copper: metallic bonding provides delocalised electrons.
2. Graphite: has one delocalised electron per carbon atom.
3. Molten aluminium: still behaves as a metal and keeps delocalised electrons when liquid.

So this set conducts electricity by delocalised electrons.

You are being asked which form of energy is needed to melt a giant ionic lattice. In an ionic solid, positive and negative ions are locked into a stiff, repeating arrangement because the electrostatic attractions are very strong. Melting happens only when enough energy is supplied to weaken these attractions so the ions are no longer fixed and can move around in the liquid state.

Now compare the choices.

• Nuclear energy is connected to reactions in the nucleus (such as fission or fusion). A change of state like melting does not involve altering nuclei.
• Light energy is electromagnetic radiation. Although light can deliver energy, it is not the usual term used for providing the large amount of energy needed to melt a bulk solid in normal chemistry contexts.

So the nuclear and light options do not match what melting is.

Electrical energy refers to energy associated with moving charges. Electricity can be used to produce heat, but the process of melting itself is described as adding heat to a substance. In chemistry, electrical energy is more directly linked to processes like electrolysis rather than simply supplying the energy for a phase change.

Melting is an endothermic change, meaning energy must be absorbed. That energy is supplied as heat, which we call thermal energy.

As an ionic solid is heated, the ions gain kinetic energy and vibrate more strongly. When they have enough thermal energy to overcome the strong electrostatic attractions holding the lattice together, the structure breaks down into a liquid.

Because melting requires adding heat to overcome the attractions in the lattice, the correct word to fill the blank is thermal energy.

The question is asking which type of energy is needed to melt a giant ionic lattice. Melting is the change from solid to liquid, which requires enough energy to defeat the strong electrostatic attractions between oppositely charged ions in the lattice.

This fits into Topic: Bonding in MYP Chemistry, where you connect bonding forces to physical properties like melting point.

Nuclear energy comes from changes in the atom’s nucleus and is involved in nuclear reactions. Melting an ionic solid only changes particle spacing and movement; it does not involve any nuclear changes. So nuclear is not correct here.

Electrical energy is energy related to moving charge. Even though molten ionic substances can conduct electricity, “electrical energy” is not the standard way we describe the energy you supply just to melt a solid. The key requirement is increasing the particles’ kinetic energy through heating.

Light energy is made of photons. While intense light could heat a material, it is not the typical energy category used when describing the general melting of a chemical lattice in the lab or in industry.

After eliminating the other choices, thermal energy (heat) is left. Adding heat raises the ions’ kinetic energy, increasing their vibrations until they can escape their fixed positions in the lattice and the solid becomes a liquid.

So the correct answer is thermal energy.

This question is asking you to recall the structural type of silicon dioxide, ${\text{SiO}}_2$. In IB Chemistry, ${\text{SiO}}_2$ is one of the classic examples you’re expected to match to its structure type. It is commonly known as silica and is found as quartz in rocks and sand.

Silicon (Si) and Oxygen (O) are treated as non-metals in bonding questions. With no metal present, we can rule out two possibilities straight away:

• Metallic structures need metal atoms and delocalized electrons.
• Ionic structures usually form from a metal + non-metal combination.

So the ionic-lattice description and the metallic-lattice description do not fit ${\text{SiO}}_2$.

It can be tempting to think ${\text{SiO}}_2$ behaves like ${\text{CO}}_2$, which exists as separate molecules. But ${\text{SiO}}_2$ does not form individual molecules. Silicon is in Period 3, so its atoms are larger than carbon and it tends to form four single covalent bonds rather than making stable double bonds. That bonding pattern produces a repeating structure instead of separate particles.

From the bonding-and-structure knowledge in the syllabus (Topic 4: Chemical Bonding and Structure), silicon dioxide is classified as a giant covalent network (also called a macromolecular structure).

Key features are:

• Each silicon atom forms covalent bonds to four oxygen atoms.
• Each oxygen atom connects to two silicon atoms.
• Strong covalent bonds extend throughout the whole solid, explaining why it is hard and has a very high melting point.

To work out the structure of silicon dioxide (${\text{SiO}}_2$), begin with the elements involved. Silicon (Si) is a metalloid/non-metal in Group 14, and Oxygen (O) is a non-metal in Group 16.

That combination points toward covalent bonding (electrons shared) rather than metallic bonding or ionic bonding. This fits the syllabus idea in Topic: Bonding - Structure and bonding.

Metallic bonding happens in metals where electrons are delocalized across a lattice of metal atoms. Because Si and O are not metals, ${\text{SiO}}_2$ cannot be described as a metallic lattice, so that choice can be removed.

Ionic compounds typically form when a metal transfers electrons to a non-metal, producing ions. Although Si and O have different electronegativities, the bonding in ${\text{SiO}}_2$ is mainly electron sharing. Also, ${\text{SiO}}_2$ does not show typical ionic behavior such as dissolving in water or conducting electricity when molten, so the ionic-lattice description is not suitable.

Substances made of discrete molecules (for example ${\text{CO}}_2$) are held together by weak intermolecular forces, so they usually have low melting points and are often gases or liquids at room temperature.

But silicon dioxide (quartz) is a very hard solid and melts at a very high temperature (around $1710^{\circ }\text{C}$). That strongly suggests the bonding must be strong throughout the entire solid, not just weak forces between molecules.

After eliminating metallic, ionic, and simple molecular structures, the only option left is a giant covalent network. In this arrangement, each silicon is covalently bonded to four oxygen atoms in a tetrahedral pattern, and each oxygen links two silicon atoms. A huge network of strong covalent bonds explains both the hardness and the high melting point.

A helpful way to understand silicon dioxide (${\text{SiO}}_2$) is to compare it with carbon dioxide (${\text{CO}}_2$), since carbon and silicon are both in Group 14. They each have 4 valence electrons, but their compounds can end up with very different structures because carbon and silicon have different atomic sizes.

This links to your syllabus ideas about the Periodic table and Periodic trends.

Carbon is small enough to form effective $p\pi -p\pi$ double bonds with oxygen. That leads to separate $\text{O}=\text{C}=\text{O}$ molecules. Since those molecules only attract each other through weak intermolecular forces, ${\text{CO}}_2$ is a gas at room temperature.

Silicon atoms are significantly larger than carbon atoms. The larger size makes orbital overlap less effective, so silicon-oxygen double bonds ($\text{Si}=\text{O}$) are not stable.

Instead of forming two double bonds, silicon becomes more stable by making four single covalent bonds to four different oxygen atoms, and each oxygen atom then bonds to two silicon atoms.

Because silicon forms four single bonds that continue linking from atom to atom, ${\text{SiO}}_2$ cannot exist as small, separate molecules like ${\text{CO}}_2$. It builds a huge three-dimensional tetrahedral framework, known as a giant covalent network (macromolecular structure), comparable in idea to diamond’s extended bonding.

Using the ${\text{CO}}_2$ vs ${\text{SiO}}_2$ comparison, we can judge the options:

• Option A doesn’t fit because it describes separate molecules with weak forces (more like ${\text{CO}}_2$), not the extended structure of ${\text{SiO}}_2$.
• Option C doesn’t fit because Si and O share electrons; ${\text{SiO}}_2$ is not an ionic solid.
• Option D doesn’t fit because ${\text{SiO}}_2$ is not a metal and does not have delocalized electrons like a metallic lattice.

The option describing a giant network of atoms connected by strong covalent bonds is exactly the structure we determined for ${\text{SiO}}_2$.

Start with diamond. In diamond, each carbon atom is connected to four other carbon atoms in a strong, repeating 3D arrangement.

That makes diamond a giant covalent network (macromolecular). Since the whole crystal is held together by an unbroken framework of covalent bonds, changing it from solid to liquid would require breaking those chemical bonds themselves.

Now consider ice (H${}_2$O). Ice is not one huge bonded network; it is a simple molecular substance.

It is made of separate H${}_2$O molecules. The atoms within each molecule are joined by covalent bonds, but the attraction between different molecules in the solid is much weaker intermolecular forces (in this case, hydrogen bonding).

Melting happens when particles can move past each other because the forces fixing them in place are overcome:

1. For Diamond: Liquefying diamond would mean breaking the strong covalent bonds holding the carbon lattice together, explaining its extremely high melting point.
2. For Ice: Turning ice into liquid water only needs enough energy to overcome the intermolecular forces between H${}_2$O molecules; the covalent bonds inside each H${}_2$O molecule do not break.

Using the structure types, we can remove wrong choices:

• Option A is wrong because melting ice does not require breaking covalent bonds.
• Option B is wrong because diamond is not made of separate molecules held together by intermolecular forces.
• Option D is wrong because it reverses which substance breaks which type of force.

The only option that correctly pairs each solid with the type of bonding/forces overcome during melting is Option C.

Before checking the options, decide what each substance is like:

• Diamond is a giant covalent structure consisting only of carbon atoms, each strongly covalently bonded to four neighbors.
• Ice ($H_{2}O$) is a simple molecular substance made of individual water molecules, with molecules attracted to one another by relatively weak intermolecular forces (hydrogen bonds).

When melting occurs, energy is used to overcome whatever holds the solid together:

1. In giant covalent structures (such as diamond), melting requires breaking covalent bonds throughout the lattice.
2. In simple molecular substances (such as ice), melting only overcomes intermolecular forces between molecules; the covalent bonds within the molecules remain unchanged.

Option A claims both substances melt by breaking covalent bonds. That fits diamond, but not ice. If the $H-O$ covalent bonds in water were broken during melting, the substance would no longer be water; it would form hydrogen and oxygen instead of becoming liquid water.

Option B says both melt by overcoming weak intermolecular forces. That cannot be true for diamond because diamond is not made of separate molecules. Since it is one continuous covalent network, there are no intermolecular forces to overcome—only covalent bonds.

Option D swaps the correct ideas. Ice does not melt by breaking covalent bonds, and diamond does not melt by overcoming intermolecular forces. Diamond is a covalent lattice, while ice is molecular with intermolecular attractions.

After eliminating the incorrect statements, Option C is left and it matches the science: diamond’s melting requires breaking covalent bonds, while ice melts by overcoming intermolecular forces, which is why ice has a much lower melting point.