IB Chemistry HL Structure 2 Models of Bonding and Structure Questions

Practise IB Chemistry HL Structure 2 with reviewed questions on bonding, resonance, hybridisation, molecular structure, materials and intermolecular forces.

Exam points

Predict ion formation and construct neutral ionic formulae from electron configurations and ion charges.
Construct and interpret Lewis structures, including formal charge, resonance and expanded octets.
Use VSEPR, hybridization and sigma/pi bonding to predict geometry, bond angles and polarity.
Explain bond-strength trends, covalent-network properties and delocalization using structure and electron models.
Identify intermolecular forces and use them to explain phase changes, solubility and material properties.
Apply chromatography and Rf evidence to distinguish substances and explain phase interactions.
Relate alloys and polymer structures to composition, properties, recycling and material choice.

Covered syllabus points

2.1 The ionic model 2.2 The covalent model 2.3 The metallic model 2.4 From models to materials

[Maximum number: 4]

Phosphine (IUPAC name phosphane) is a hydride of phosphorus, with the formula $\mathrm{PH}_{3}$ .

(a)

Draw a Lewis (electron dot) structure of phosphine.

[ 1 ]

(b)

State the hybridization of the phosphorus atom in phosphine.

[ 1 ]

(c)

Outline whether you expect the bonds in phosphine to be polar or non-polar, giving a brief reason.

Figure for Question (c) — IB Chemistry HL

Phosphine has a much greater molar mass than ammonia. Explain why phosphine has a significantly lower boiling point than ammonia.

[ 2 ]

[Maximum number: 7]

Titanium and vanadium are consecutive elements in the first transition metal series.

(a)

Describe the bonding in metals.

[ 2 ]

(b)

Suggest why the melting point of vanadium is higher than that of titanium.

[ 1 ]

(c)

Explain why an aluminium-titanium alloy is harder than pure aluminium.

[ 2 ]

(d)

State the type of bonding in potassium chloride which melts at 1043 K .

[ 1 ]

(e)

A chloride of titanium, $\mathrm{TiCl}_{4}$ , melts at 248 K . Suggest why the melting point is so much lower than that of KCl .

[ 1 ]

[Maximum number: 6]

Iron may be extracted from iron (II) sulfide, FeS.

(a)

Outline why metals, like iron, can conduct electricity.

[ 1 ]

(b)

Iron (II) sulfide, FeS, is ionically bonded.

[ 3 ]

(i)

Describe the bonding in this type of solid.

[ 2 ]

electrostatic attraction between oppositely charged ions/between $\mathrm{Fe}^{2+}$ and $\mathrm{S}^{2-}$

(ii)

Suggest why chemists find it convenient to classify bonding into ionic, covalent and metallic.

[ 1 ]

(c)

Explain why the addition of small amounts of carbon to iron makes the metal harder.

[ 2 ]

[Maximum number: 16]

(a)

An organic compound, X, with a molar mass of approximately $88 \mathrm{~g} \mathrm{~mol}^{-1}$ contains 54.5 % carbon, 36.3 % oxygen and 9.2 % hydrogen by mass.

[ 5 ]

(i)

The carboxylic acid contains two different carbon-oxygen bonds. Identify which bond is stronger and which bond is longer.

Stronger bond:

Longer bond:

[ 2 ]

(ii)

Predict and explain the bond lengths and bond strengths of the carbon-oxygen bonds in $\mathrm{CH}_{3} \mathrm{CH}_{2} \mathrm{COO}^{-}$ .

[ 3 ]

(b)

State the meaning of the term hybridization.

[ 1 ]

(c)

Describe the hybridization of the carbon atom in methane and explain how the concept of hybridization can be used to explain the shape of the methane molecule.

[ 3 ]

$\mathrm{sp}^{3}$ ;
isolated C atom electron configuration $1 \mathrm{~s}^{2} 2 \mathrm{~s}^{2} 2 \mathrm{p}^{2} /$ excited state C electron configuration is $1 \mathrm{~s}^{2} 2 \mathrm{~s}^{1} 2 \mathrm{p}^{3}$ ;
$2 \mathrm{~s}^{1} 2 \mathrm{p}^{3}$ electrons blend to form four identical hybrid orbitals;
hybrid orbitals lower in total energy than atomic orbitals;
repulsion of (identical hybrid) orbitals creates a tetrahedral shape;

(d)

Identify the hybridization of the carbon atoms in diamond and graphite and explain why graphite is an electrical conductor.

[ 3 ]

diamond:
$\mathrm{sp}^{3}$ ;
graphite:
$\mathrm{sp}^{2}$ ;
(p) electrons delocalized (around layer);

(e)

Aluminium chloride, $\mathrm{Al}_{2} \mathrm{Cl}_{6}$ , does not conduct electricity when molten but aluminium oxide, $\mathrm{Al}_{2} \mathrm{O}_{3}$ , does. Explain this in terms of the structure and bonding of the two compounds.
$\mathrm{Al}_{2} \mathrm{Cl}_{6}$ : $\mathrm{Al}_{2} \mathrm{O}_{3}$ :

[ 4 ]

$\quad \mathrm{Al}_{2} \mathrm{Cl}_{6}$ :
covalent bonding / dimer/molecular structure;
no free charges when molten so not an electrical conductor;
$\mathrm{Al}_{2} \mathrm{O}_{3}:$
ionic / lattice structure;
ions free to move/mobile in molten state;

All question bank results loaded

Covered syllabus points

Question 1

Question (a)

Question (b)

Question (c)

Question 2

Question (a)

Question (b)

Question (c)

Question (d)

Question (e)

Question 3

Question (a)

Question (b)

Question (i)

Question (ii)

Question (c)

Question 4

Question (a)

Question (i)

Question (ii)

Question (b)

Question (c)

Question (d)

Question (e)