Chemical Bonding — Diagnostic Tests

Unit Tests

UT-1: Ionic Bonding

Question:

(a) Define an ionic bond and state the general rule for which types of elements form ionic bonds together.

(b) Describe, using dot-and-cross diagrams, the formation of an ionic bond between calcium and oxygen to form calcium oxide ($\text{CaO}$).

(c) Explain why ionic compounds have high melting points.

(d) State why solid ionic compounds cannot conduct electricity but molten or dissolved ionic compounds can.

Solution:

(a) An ionic bond is the electrostatic attraction between oppositely charged ions formed when a metal transfers one or more electrons to a non-metal. Ionic bonds form between metals (in Groups 1 and 2) and non-metals (in Groups 6 and 7).

(b) Calcium is in Group 2 with the electron configuration $2, 8, 8, 2$. It loses its two outer electrons to become $\text{Ca}^{2+}$ with configuration $2, 8, 8$. Oxygen is in Group 6 with configuration $2, 6$. It gains two electrons to become $\text{O}^{2-}$ with configuration $2, 8$. The resulting $\text{Ca}^{2+}$ and $\text{O}^{2-}$ ions are held together by strong electrostatic attraction, forming $\text{CaO}$.

(c) Ionic compounds consist of a giant lattice of oppositely charged ions arranged in a regular repeating structure. The electrostatic forces of attraction between the ions are very strong and act in all directions. A large amount of energy is required to overcome these forces, so ionic compounds have high melting points.

(d) In the solid state, the ions are fixed in position in the lattice and cannot move, so electricity cannot be conducted. When molten or dissolved in water, the ions are free to move and carry charge, allowing the compound to conduct electricity.

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UT-2: Covalent Bonding

Question:

(a) Define a covalent bond.

(b) Draw a dot-and-cross diagram for a molecule of water ($\text{H}_2\text{O}$), showing only the outer shell electrons.

(c) Explain the difference between a simple covalent molecule and a giant covalent structure, giving one example of each.

(d) Carbon dioxide ($\text{CO}_2$) has a low boiling point despite containing strong covalent bonds. Explain why.

Solution:

(a) A covalent bond is a shared pair of electrons between two non-metal atoms. Each atom contributes one electron to the shared pair.

(b) Oxygen (Group 6) has 6 outer electrons. Each hydrogen atom has 1 outer electron. Oxygen shares one electron with each hydrogen, forming two O—H covalent bonds. The oxygen atom now has 8 electrons in its outer shell (2 shared + 4 unshared + 2 shared from bonds = 8), and each hydrogen has 2 electrons in its outer shell (satisfying the duet rule).

(c) A simple covalent molecule consists of a small number of atoms held together by strong covalent bonds, with weak intermolecular forces between the molecules. Examples include $\text{H}_2\text{O}$, $\text{CO}_2$, and $\text{CH}_4$. A giant covalent structure has a network of covalent bonds extending throughout the entire structure. Examples include diamond (carbon atoms each bonded to four others in a tetrahedral lattice) and silicon dioxide ($\text{SiO}_2$).

(d) The strong covalent bonds within each $\text{CO}_2$ molecule are not broken when $\text{CO}_2$ boils. Only the weak intermolecular forces (London dispersion forces) between the molecules need to be overcome. Since these intermolecular forces are weak, very little energy is required, so $\text{CO}_2$ has a low boiling point ($-78.5\,^\circ\text{C}$).

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UT-3: Metallic Bonding

Question:

(a) Describe the structure of a metal and explain how metallic bonding works.

(b) Explain why metals are good conductors of electricity.

(c) Explain why metals are malleable and ductile.

(d) Describe and explain the trend in melting point across the metals sodium, magnesium, and aluminium (Period 3).

Solution:

(a) In a metal, atoms are arranged in a regular close-packed lattice. The outer electrons (valence electrons) are delocalised, meaning they are free to move throughout the entire metallic structure. The metal cations (positive ions remaining after electrons are delocalised) are held together by the strong electrostatic attraction between the cations and the sea of delocalised electrons. This is called metallic bonding.

(b) Metals are good conductors of electricity because the delocalised electrons are free to move through the lattice. When a potential difference is applied, these mobile electrons carry charge through the metal.

(c) Metals are malleable (can be hammered into shape) and ductile (can be drawn into wires) because the layers of metal cations can slide over one another without breaking the metallic bonds. The delocalised electrons adjust their positions to maintain the bonding even when the lattice is deformed, unlike ionic compounds which shatter when layers are displaced.

(d) From sodium to magnesium to aluminium, the melting point increases significantly. This is because the charge on the metal ion increases from $+1$ (Na) to $+2$ (Mg) to $+3$ (Al), while the ionic radius decreases (greater nuclear charge pulling electrons in more tightly). The electrostatic attraction between the cations and delocalised electrons therefore increases, requiring more energy to break the metallic bonds. Sodium melts at $98\,^\circ\text{C}$, magnesium at $650\,^\circ\text{C}$, and aluminium at $660\,^\circ\text{C}$ (though magnesium’s structure gives a particularly large jump due to changes in crystal packing as well as charge).

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Integration Tests

IT-1: Comparing Bonding Types

Question:

(a) Diamond and sodium chloride both have high melting points. Explain the different reasons for this, referring to the types of bonding and structure in each.

(b) Explain why diamond does not conduct electricity, whereas graphite does, even though both are forms of carbon.

(c) A student has three unnamed substances: Substance A melts at $801\,^\circ\text{C}$, conducts electricity when molten but not when solid, and dissolves in water. Substance B melts at $3550\,^\circ\text{C}$, does not conduct electricity, and is insoluble in water. Substance C melts at $-184\,^\circ\text{C}$, does not conduct electricity, and is a gas at room temperature. Identify the type of bonding and structure in each substance.

(d) Explain why sodium chloride dissolves in water but diamond does not.

Solution:

(a) Diamond has a high melting point because it is a giant covalent structure in which every carbon atom is covalently bonded to four others. Many strong covalent bonds must be broken to melt diamond. Sodium chloride has a high melting point because it is a giant ionic lattice with strong electrostatic forces of attraction between the oppositely charged ions. The key difference is that diamond’s high melting point comes from covalent bonds between atoms, whereas NaCl’s comes from ionic bonds between ions.

(b) In diamond, each carbon atom uses all four of its outer electrons to form four covalent bonds. There are no delocalised electrons available to carry charge, so diamond cannot conduct electricity. In graphite, each carbon atom forms only three covalent bonds with other carbon atoms (in layered hexagonal sheets). The fourth outer electron of each carbon is delocalised and free to move between the layers, allowing graphite to conduct electricity.

(c)

• Substance A: Ionic bonding, giant ionic lattice. The high melting point, electrical conductivity when molten (free-moving ions), and solubility in water (ions attracted to polar water molecules) are all characteristic of ionic compounds. This is likely sodium chloride.
• Substance B: Giant covalent structure. The extremely high melting point, lack of conductivity, and insolubility in water point to a network covalent substance. This is likely diamond.
• Substance C: Simple covalent molecule. The very low melting point, lack of conductivity, and gaseous state at room temperature indicate weak intermolecular forces between small covalent molecules. This could be chlorine ($\text{Cl}_2$) or a similar simple molecular gas.

(d) Sodium chloride dissolves in water because water molecules are polar. The slightly positive hydrogen atoms of water are attracted to the $\text{Cl}^-$ ions, and the slightly negative oxygen atoms are attracted to the $\text{Na}^+$ ions. The water molecules surround and separate the ions, overcoming the ionic bonds in the lattice. Diamond does not dissolve because it is a giant covalent network; there are no ions to be solvated and breaking the strong covalent bonds to separate atoms would require far more energy than water can provide.

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IT-2: Bonding and Properties of Key Compounds

Question:

(a) Explain the difference in electrical conductivity between copper (metallic), silicon dioxide (giant covalent), and carbon dioxide (simple covalent).

(b) Sodium is a metal that reacts vigorously with chlorine gas. Write a balanced chemical equation for the reaction and describe the type of bonding in the product. Explain what happens to the electrons during the reaction.

(c) A student claims that covalent compounds always have low melting points. Evaluate this claim with supporting examples.

(d) Graphite is used as a lubricant. Explain this use in terms of its bonding and structure.

Solution:

(a)

• Copper (metallic): Conducts electricity due to delocalised electrons that are free to move throughout the lattice.
• Silicon dioxide ($\text{SiO}_2$, giant covalent): Does not conduct electricity because all electrons are involved in covalent bonds and are not free to move.
• Carbon dioxide ($\text{CO}_2$, simple covalent): Does not conduct electricity because it consists of neutral molecules with no charged particles or delocalised electrons.

(b) The balanced equation is:

$2\text{Na} + \text{Cl}_2 \rightarrow 2\text{NaCl}$

Each sodium atom loses one electron to become $\text{Na}^+$, and each chlorine atom gains one electron to become $\text{Cl}^-$. The product is sodium chloride, which has ionic bonding — the strong electrostatic attraction between $\text{Na}^+$ and $\text{Cl}^-$ ions in a giant ionic lattice.

(c) This claim is incorrect. While simple covalent molecules (e.g., $\text{H}_2\text{O}$, $\text{CO}_2$) do have low melting points, giant covalent structures have very high melting points. For example, diamond melts at approximately $3550\,^\circ\text{C}$ and silicon dioxide melts at approximately $1610\,^\circ\text{C}$. These high melting points arise because melting requires breaking the strong covalent bonds throughout the entire structure, unlike simple molecules where only weak intermolecular forces need to be overcome.

(d) Graphite consists of layers of hexagonally arranged carbon atoms covalently bonded within each layer. The layers are held together by weak interm forces (van der Waals forces). When a force is applied, the layers can slide easily over one another, making graphite a good lubricant. This property arises directly from the combination of strong in-layer covalent bonding (which gives the layers structural integrity) and weak inter-layer forces (which allow the layers to slide).

Summary

The key principles covered in this topic are linked in the sub-pages above. Focus on understanding the definitions, applying the formulas or frameworks, and evaluating strengths and limitations of each approach.

Worked Examples

Worked examples demonstrating the application of key concepts are covered in the detailed sub-pages linked above.

Common Pitfalls

• Confusing the strength of covalent bonds within molecules with the strength of intermolecular forces between molecules when explaining melting points.
• Stating that metals conduct electricity because they contain free protons or ions, rather than delocalised electrons.
• Describing metallic bonding as ions sharing electrons rather than cations in a sea of delocalised electrons.
• Forgetting that the high melting point of giant covalent structures is due to covalent bonds, whereas the low melting point of simple covalent substances is due to weak intermolecular forces.