Structure: Models of Bonding and Structure — Free Chemistry HL Practice Questions

The melting point of sodium is $97.8\,°\text{C}$, while the melting point of magnesium is $650\,°\text{C}$. Using the electron sea model, explain this difference in melting points. [2 marks]

A student claims: "Alloying a pure metal always reduces its malleability." Evaluate this claim using the electron sea model, referring to the effect of atomic radius difference between the host and solute metal. [3 marks]

Describe, using the electron sea model, the mechanism by which heat is transferred from the hot end to the cool end of a metal wire heated at one end. [3 marks]

The intermetallic compound NiAl is brittle rather than malleable. Evaluate the extent to which the electron sea model can account for this observation, and identify what additional bonding concept is required to explain it. [3 marks]

Explain, by referring to the behaviour of delocalized electrons, why metals have a shiny appearance. [2 marks]

The electron sea model cannot account for ferromagnetism in iron. Identify the specific limitation of the model responsible for this failure, and suggest a modification that would allow the model to explain ferromagnetism. Justify your suggestion. [3 marks]

Explain, using the electron sea model, why pure titanium has a high melting point. [2 marks]

State one structural change that occurs to the titanium lattice when aluminium and vanadium atoms are introduced. [1 mark]

Explain why this structural change results in a higher tensile strength for Ti-6Al-4V compared to pure titanium. [1] (e) Suggest why the electrical conductivity of titanium decreases as temperature increases, even though the number of delocalised electrons remains constant. [2 marks]

State the coordination number of $\text{Na}^+$ in NaCl and of $\text{Cs}^+$ in CsCl. [2 marks]

Explain why these coordination numbers differ, with reference to the radius ratio of each cation to $\text{Cl}^-$. [2 marks]

(i) A student claims that "the lattice enthalpy of CsCl is greater than that of NaCl because CsCl has a higher coordination number." State whether this claim is correct, and compare the two lattice enthalpy values given in the data. [1]
(ii) Determine the interionic distance (cation–anion centre separation) in each compound and use Coulomb's law to explain which compound has the larger lattice enthalpy. [3]
(iii) Evaluate whether coordination number alone determines the magnitude of lattice enthalpy. [2]
(iv) Describe one experimental method by which lattice enthalpy can be determined indirectly. [1] (e) In the NaCl lattice the distance between the centres of a neighbouring $\text{Na}^+$–$\text{Cl}^-$ pair is $2.82 \times 10^{-10}\,\text{m}$. The unit cell is a cube with edge length equal to twice this distance, and contains four formula units of NaCl. Calculate the density of NaCl in $\text{g cm}^{-3}$. [4 marks]

(i) State the relationship between lattice enthalpy and melting point for an ionic solid. [1]
(ii) Using the ionic radii given and Coulomb's law, explain why MgO has a higher melting point than CaO, given that both compounds have the same lattice structure and the same ionic charges. [4 marks]

Deduce the number of $\text{Mg}^{2+}$ ions and the number of $\text{O}^{2-}$ ions in one unit cell of MgO. In your answer, account for the contribution of ions at corners, face centres, edge centres, and the body centre. [5 marks]

(i) A student claims: "If MgO adopted the caesium chloride structure instead of the rock salt structure, its lattice enthalpy would be lower because the higher coordination number causes greater repulsion between like-charged ions." Identify the error in the student's reasoning about how coordination number affects lattice enthalpy. [1]
(ii) Explain why the net effect of increasing coordination number on lattice enthalpy cannot be determined by considering repulsion alone. [2]
(iii) Using the radius ratio rule, deduce why MgO adopts the rock salt structure rather than the caesium chloride structure. [2] Total: [16 marks] (Note: total adjusted to 16 to match independently markable steps; if constrained to 15, (d)(ii) is reduced to and merged with (d)(i) as a two-mark explain sub-part.) [1 mark]

Deduce why phosphorus in $\text{PCl}_5$ has an expanded valence shell, while phosphorus in $\text{PCl}_3$ obeys the octet rule. In your answer, state the total number of electron domains around phosphorus in each molecule and explain how this determines the number of bonds formed. [3 marks]

$\text{PCl}_5$ undergoes rapid hydrolysis with water, while $\text{PCl}_3$ reacts more slowly. Explain this difference in reactivity in terms of the oxidation state of phosphorus and the electron density at the phosphorus centre in each molecule. [3 marks]

The bond angle in $\text{PCl}_3$ is $100.3^\circ$, which is smaller than the bond angle in $\text{NH}_3$ of $106.7^\circ$. A student claims this solely because phosphorus has a larger atomic radius than nitrogen. Use the data below to evaluate this claim. Your answer must identify which factor — atomic size or electronegativity of the terminal atom — makes the greater contribution to the observed difference, and justify your reasoning. Molecule — Bond angle $\text{NH}_3$ — $106.7^\circ$ $\text{PH}_3$ — $93.5^\circ$ $\text{PCl}_3$ — $100.3^\circ$ Electronegativity values (Pauling scale): $\text{N} = 3.04$, $\text{P} = 2.19$, $\text{H} = 2.20$, $\text{Cl} = 3.16$ [5 marks]

Calculate the formal charge on each atom in the most stable resonance structure of $\text{SO}_2$. Hence deduce the average S–O bond order and explain how resonance affects the distribution of electron density in $\text{SO}_2$. [4 marks]

Both C=O and S–O bonds are polar. Explain, using bond polarity and molecular symmetry, why $\text{CO}_2$ has a dipole moment of zero while $\text{SO}_2$ has a dipole moment of $1.63\,\text{D}$. [4 marks]

A student claims that the higher boiling point of $\text{SO}_2$ relative to $\text{CO}_2$ is caused by hydrogen bonding in $\text{SO}_2$. Using the data below, evaluate this claim by identifying the intermolecular forces present in each substance, explaining their relative magnitudes, and justifying which force is primarily responsible for the boiling point difference. Property — $\text{CO}_2$ — $\text{SO}_2$ Dipole moment / D — $0$ — $1.63$ Molar refractivity / $\text{cm}^3\,\text{mol}^{-1}$ — $6.7$ — $8.0$ Molar mass / $\text{g mol}^{-1}$ — $44$ — $64$ [4 marks]

Consider three metals: sodium (Na, melting point $98\,^\circ\text{C}$), magnesium (Mg, melting point $650\,^\circ\text{C}$), and aluminium (Al, melting point $660\,^\circ\text{C}$). (i) Explain the trend in melting points from Na to Al in terms of the electron sea model. [3]
(ii) The electrical conductivities of Na, Mg, and Al at $20\,^\circ\text{C}$ are $2.1 \times 10^{7}\,\text{S\,m}^{-1}$, $2.3 \times 10^{7}\,\text{S\,m}^{-1}$, and $3.8 \times 10^{7}\,\text{S\,m}^{-1}$ respectively. A student predicts that conductivity should increase from Na to Al because the number of delocalized electrons per atom increases. Evaluate this prediction. In your answer, identify one factor that supports and one factor that limits this prediction. [3 marks]

An alloy is formed by mixing copper (Cu) and nickel (Ni) in a $1\!:\!1$ mole ratio. Metal — Melting point — Conductivity at $20\,^\circ\text{C}$ — Atomic radius Cu — $1085\,^\circ\text{C}$ — $5.96 \times 10^{7}\,\text{S\,m}^{-1}$ — $128\,\text{pm}$ Ni — $1455\,^\circ\text{C}$ — $1.43 \times 10^{7}\,\text{S\,m}^{-1}$ — $124\,\text{pm}$ Cu–Ni alloy — $1350\,^\circ\text{C}$ — — (i) Predict, with justification, whether the electrical conductivity of the Cu–Ni alloy will be greater than, less than, or equal to the simple average of the two pure metals. [3]
(ii) The simple average melting point of Cu and Ni is $1270\,^\circ\text{C}$. The measured melting point of the Cu–Ni alloy is $1350\,^\circ\text{C}$, which is $80\,^\circ\text{C}$ higher than this average. Evaluate whether this observation is consistent with, or anomalous to, the electron sea model. [1 mark]

Mercury (Hg) has anomalously low melting point of $-38.8\,^\circ\text{C}$ compared to gold (Au) at $1064\,^\circ\text{C}$. Both are Period 6 transition metals. The ground-state electron configuration of Hg ends in $\text{[Xe]}\,4f^{14}\,5d^{10}\,6s^2$, while that of Au ends in $\text{[Xe]}\,4f^{14}\,5d^{10}\,6s^1$. Using the electron sea model, deduce how the filled $6s^2$ subshell of mercury affects the strength of metallic bonding, and hence account for the difference in melting points between Hg and Au. [4 marks]

A student proposes: "The electrical conductivity of a liquid metal must be zero because the lattice of positive ions is no longer fixed." Evaluate this proposal. In your answer, refer to the electron sea model, experimental evidence, and explain why the conductivity of a liquid metal differs quantitatively from that of the same metal in the solid state. [4 marks]

$\text{Mg}_2\text{Si}$ is an intermetallic compound with a melting point of $1102\,^\circ\text{C}$ that behaves a semiconductor. Evaluate whether the electron sea model adequately describes the bonding in $\text{Mg}_2\text{Si}$. In your answer, compare the bonding in $\text{Mg}_2\text{Si}$ with that in pure magnesium and pure silicon, using the electronegativity values provided. [4 marks]

Deduce the most likely type of bonding and structure present in Substance Y. Justify your deduction using all of the given properties. [4 marks]

Explain why Substance Y melts at a significantly lower temperature than Substance X. In your answer, compare the nature of the forces that must be overcome and the energy changes involved in the melting process for each substance. [3 marks]

A technician proposes heating both substances to $400\,\text{K}$ to produce two liquids for processing. Evaluate this proposal. In your answer, predict the physical state of each substance at $400\,\text{K}$, describe the particle arrangement and forces present in any liquid formed, and assess whether the properties of each substance make it suitable for use in a device that must repeatedly cycle between solid and liquid states near room temperature. [4 marks]

Explain why Lipid A has a higher melting point than Lipid B, despite both having the same molecular formula and molar mass. [3 marks]

The SLN must remain solid during storage at $298\,\text{K}$ and must melt completely at body temperature ($310\,\text{K}$) to release the drug. Using the melting point and density data provided, evaluate which lipid is more suitable for this application. [4 marks]

At $310\,\text{K}$, the polar drug molecules are present in whichever lipid phase exists at that temperature. Using the density data and your knowledge of intermolecular forces, analyse how the rate of diffusion of the drug molecules through Lipid B at $310\,\text{K}$ would be affected by: (i) the free volume available between lipid molecules, and
(ii) the nature of the interactions between the polar drug and the non-polar lipid. Hence predict, with reasoning, whether increasing temperature from $310\,\text{K}$ to $320\,\text{K}$ would increase or decrease the diffusion rate, and identify the dominant factor. [6 marks]