The Covalent Model (DP IB Chemistry: HL): Exam Questions

State the formula of the compound formed between boron and chlorine.

Draw the Lewis (electron dot) structure for the compound formed in part (a).

Explain why the compound formed in part (a) is able to form coordinate bonds.

Group 17 of the Periodic Table contain non-metals that are often referred to as the halogens.

Iodine, I₂, is one of these halogens. At room temperature and pressure it exists as a grey-black solid.

Describe the bonding and forces present in I₂ in the solid state.

The state of the halogens changes down the group, with fluorine being a gas and astatine being a solid.

Explain why the melting point of the halogens increases down the group.

The halogens are all diatomic covalent molecules.

State three physical properties common to all elements in Group 17.

The halogens can also form interhalogen compounds, such as iodine monochloride, ICl.

Predict the state of iodine monochloride at room temperature and pressure. Explain your answer with reference to the intermolecular forces present.

The allotropes of carbon include diamond, graphite, and buckminsterfullerene, C₆₀.

i) State how many atoms each carbon is directly bonded to in each of the allotropes.

[3]

ii) Explain why the carbon atoms in the alltropes form a different number of bonds.

[1]

Compare the overall structure and bonding of diamond with that of buckminsterfullerene.

Describe and explain the differences in electrical conductivity between the three allotropes of carbon.

Graphene is another allotrope consisting of a single layer of graphite.

State one similarity and one difference in structure between graphene and graphite.

Silicon and carbon are elements in the same group of the Periodic Table, that form covalent bonds.

Both silicon and carbon react with oxygen to form dioxides. The structures of silicon dioxide and carbon dioxide are shown below.

Silicon dioxide has a melting point of 1710 ^oC, while carbon dioxide sublimes at -78 ^oC.

Explain this difference with reference to the structure and bonding present in each dioxide.

How many oxygen atoms are bonded to each carbon and to each silicon?

Explain how this links to the formula of each compound.

Predict the O-C-O and O-Si-O bond angles respectively in CO₂ and in SiO₂.

Predict and explain the solubility of both SiO₂ and CO₂ in water.

A student investigates the chlorides of various elements to compare their bonding, structure, and properties.

The student first uses electronegativity to predict bonding type.

Using section 9 of the data booklet, deduce whether the bonding in each compound below is predominantly ionic or covalent. Justify each answer with an electronegativity difference calculation.

i) ICl

ii) SrCl₂

The student knows that sodium chloride, NaCl, is an ionic compound with a high melting point of 801 ^oC.

Explain this property with reference to the structure and bonding in NaCl.

Explain why ionic chlorides, such as NaCl, conduct electricity when molten but not when solid.

The student next investigates two covalent chlorides.

i) Draw the Lewis (electron dot) structure for phosphorus trichloride (PCl₃) and boron trichloride (BCl₃).

[2]

ii) Predict the molecular geometry of PCl₃ and BCl₃.

[2]

Nitrogen hydrides, such as diazene (N₂H₂) and hydrazine (N₂H₄), are highly reactive compounds used in chemical synthesis and as potential fuels. Their structure and bonding determine their properties and reactivity.

Draw the Lewis structures for diazene (N₂H₂) and hydrazine (N₂H₄).

Use VSEPR theory to deduce the molecular geometry around a nitrogen atom in diazene (N₂H₂) and estimate the H-N-N bond angle.

List the compounds dinitrogen (N₂), diimide (N₂H₂), and hydrazine (N₂H₄) in order of increasing nitrogen-nitrogen bond length. Justify the order with reference to the number of shared electron pairs in the bond

Using section 9 of the data booklet, predict and justify which bond is more polar in each of the following pairs:

i) N-H or N-F

[1]

ii) C-N or C-O

[1]

The combustion of ethyne, C₂H₂, is a highly exothermic reaction used in welding torches. The balanced equation for its complete combustion is:

2C₂H₂ (g) + 5O₂ (g) → 4CO₂ (g) + 2H₂O (g) 

i) Draw the Lewis structure for one molecule of ethyne.

[1]

ii) Identify the number of sigma (σ) and pi (π) bonds in one molecule of ethyne.

[1]

Explain why the carbon-carbon triple bond in ethyne is significantly stronger and shorter than the oxygen-oxygen double bond in oxygen.

The products of the combustion, carbon dioxide and water, have different molecular polarities.

i) Using section 9 of the data booklet, determine the electronegativity difference (Δχ) for a C=O bond and an O-H bond.

[1]

ii) Both CO₂ and H₂O molecules contain polar bonds, but only H₂O is a polar molecule. Explain this difference with reference to molecular geometry.

[3]

The combustion of ethyne is highly exothermic. Explain this in terms of the energy absorbed to break bonds and the energy released when forming bonds.

Ammonia, NH₃, is a chemical that is key in the manufacture of certain fertilisers and cleaning products.

An ammonia molecule will react with an H⁺ ion, to form the ammonium ion, NH₄⁺.

Draw a Lewis (electron dot) diagram to show the bonding in the ammonium ion and name the type of bond formed between the ammonia molecule and the hydrogen ion. 

Lewis (electron dot diagrams) are used to show the electron arrangement in the valence shells of covalently bonded molecules.

Draw Lewis diagrams for the following molecules:

i)         Hydrogen cyanide.

[1]

ii)        Carbon dioxide.

[1]

iii)        Boron trifluoride.

[1]

Using your answer to part (b), identify and explain the species that is likely to form a coordinate covalent bond.

Using your answer to part (c), Explain, with the help of a diagram, the covalent bond formed between the species in part (c) and ammonia.

Benzene is an aromatic hydrocarbon which is often drawn as Figure 1.

Discuss the physical evidence that justifies this structure for benzene.

Figure 1

Benzene and cyclohexene are both unsaturated molecules, but cyclohexene reacts with bromine water and benzene does not.

i) State the meaning of the terms saturated and unsaturated as applied to organic molecules.  [1]

ii) Explain this difference in reactivity and write an equation for the reaction between cyclohexene and bromine. 

[2]

Table 1 below shows the enthalpy changes for the hydrogenation of cyclohexene, benzene, and the theoretical molecule 1,3,5-cyclohexatriene.

Table 1

The equations for the hydrogenation reactions are:

Cyclohexene               C₆H₁₀ + H₂ ⭢ C₆H₁₂

Benzene                      C₆H₆ + 3H₂ ⭢ C₆H₁₂

i) Use the data in Table 1 to determine the enthalpy of hydrogenation of the theoretical molecule 1,3,5-cyclohexatriene. 

[1]

ii) Discuss the difference between the enthalpy of hydrogenation of benzene and of 1,3,5-cyclohexatriene. 

[2]

An unknown aromatic compound has the molecular formula C₈H₈O₂.

Deduce the structural formula of two isomers of this compound which contain an ester group.

Phosphorus tribromide and sulfur tetrafluoride are two colourless compounds which both react with water to form toxic products. 

Deduce the Lewis (electron dot) structure of both molecules.

Predict the shapes of the two molecules of phosphorus tribromide and sulfur tetrafluoride

Explain why both phosphorus tribromide and sulfur tetrafluoride are polar.

Compare the formation of a sigma (σ) and a pi (π) bond between two carbon atoms in a molecule.

But-2-ene-1,4-dioic acid exists as both cis and trans isomers. The cis isomer is shown below:

Describe the type of covalent bond between carbon and hydrogen in the molecule shown above and how it is formed.

Identify how many sigma bonds and how many pi (π) bonds are present in cis but-2-ene-1,4-dioic acid.

Draw the Lewis structures, predict the shape and deduce the bond angles for xenon tetrafluoride.

Compare the polarity of xenon tetrafluoride with chlorine trifluoride.

Carbon dioxide can be represented by at least two resonance structures, I and II.

Calculate the formal charge on each oxygen atom in the two structures.

Deduce, giving a reason, the more likely resonance structure from part a)

Nitrous oxide can be represented by different Lewis (electron dot) structures. 

Deduce the formal charge (FC) of the nitrogen and oxygen atoms in three of these Lewis (electron dot) structures, A, B and C, represented below.

LHS: atom on the left-hand side; RHS: atom on the right-hand side

	Lewis (electron dot) structure	FC of N on LHS	FC of central N	FC of O on RHS
A
B
C

Based on the formal charges assigned in part c), deduce which Lewis (electron dot) structure of N₂O (A, B, or C) is preferred.

Explain another factor that also must be taken into account in determining the preferred structure.

Use the concept of formal charge to explain why BF₃ is an exception to the octet rule.

Compounds containing two different halogen atoms bonded together are called interhalogen compounds. They are interesting because they contain halogen atoms in unusual oxidation states. One such compound is BrF₃. 

Deduce the electron domain geometry and molecular geometry of BrF₃.

Give the approximate bond angle(s) and a valid Lewis (electron dot) structure for BrF₃. 

Explain why bromine trifluoride, BrF₃ has its lone pairs of electrons located in equatorial positions. 

Draw two different Lewis (electron dot) structures for SO₄^2–, one of which obeys the octet rule for all its atoms, the other which has an octet for S expanded to 12 electrons. 

Explain which of the two SO₄^2– structures is preferred using formal charges.

Consider the molecule shown below.

Identify the number of sigma and pi bonds in this molecule.

One of the intermediates in the reaction between nitrogen monoxide and hydrogen is dinitrogen monoxide, N₂O. This can be represented by the resonance structures below

 Analyse the bonding in dinitrogen monoxide in terms of sigma and pi bonds.

Based on the type of intermolecular force present, explain why butan-1-ol has a higher boiling point than butanal.

Ethane, C₂H₆, and disilane, Si₂H₆, are both hydrides of Group 4 elements with similar structures but different chemical properties.

Explain why disilane has a higher boiling point than ethane.

Put the following molecules in order of increasing boiling point and explain your choice:

Based on the type of intermolecular force present, explain the difference in solubility in water between ethane and ethanol.

The melting points of some Group 1 elements are listed in Table 1.

Table 1

Predict, with a reason, the melting point of Rb.

Explain why ammonia, NH₃, is a gas at room temperature. 

Phosphine (IUPAC name phosphane) is a hydride of phosphorus, with the formula PH₃. Phosphine has a much greater molar mass than ammonia.

Explain why phosphine has a significantly lower boiling point than ammonia.

Identify the type of interaction that must be overcome when liquid hydrazine, N₂H₄, vaporizes. Suggest, with a reason, whether hydrazine has a lower or higher boiling point than diimide, N₂H₂.

Cyclohexane C₆H₁₂ has a puckered, non-planar shape whereas benzene C₆H₆ is planar. 

Explain this difference by making reference to the C–C–C bond angles and the type of hybridisation of carbon in each molecule.

Urea, CO(NH₂)₂, is present in solution in animal urine. 

What is the hybridisation of C and N in the molecule, and what are the approximate bond angles?

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

A molecule of ethanol is shown below.

Deduce the hybridisation of the carbon atom marked in the diagram below.

Carbonation is the process of increasing the concentration of carbonate ions in water to produce carbonated drinks.

Identify the hybridisation of the central carbon atom.

Explain, with the use of diagrams, how there are three valid structures for the carbonate ion.

Describe the distribution of pi (π) electrons and explain how this can account for the structure and stability of the carbonate ion, CO₃^2–.

Explain how the concept of hybridisation can be used to explain the triple bond present in propyne.

Consider the molecule below which contains both sigma and pi bonds.

 How many carbon atoms exhibit sp² hybridisation in this molecule?

Deduce the hybridisation shown by the nitrogen atoms in NF₄⁺, N₂H₂ and N₂H₄.

Draw the structure of silicon dioxide and state the type of bonding present.

Describe the similarities and differences you would expect in the properties of silicon and diamond.

The boiling point of diamond is 3550 ℃, but for carbon dioxide it is -78.5 ℃. Both are covalent substances. 

Explain this difference with reference to structure and bonding.

Silicon dioxide has a similar name to carbon dioxide, but its boiling point is 2230 ℃. 

Briefly outline the reason for this difference.

In 1996 the Nobel prize in Chemistry was awarded for the discovery of a new carbon allotrope, known as fullerenes. 

Outline the structure of buckminsterfullerene.

Like carbon dioxide, graphite is also a covalent substance, but it is a solid at room  temperature. Graphite has a melting point of around 3600 ^oC.           

Describe the structure and bonding of graphite and explain why it has such a high melting point.

Graphite is made purely of carbon, a non-metal, yet it conducts electricity. Diamond, which is also made purely of carbon, cannot conduct electricity. 

i) Explain this difference in electrical conductivity between graphite and diamond. 

[3]

ii) Give one other difference in the properties of graphite and diamond.

[1]

Graphite is soft and so is used as a lubricant, whereas diamond is hard and so is used in many cutting tools. Both are giant covalent structures.

Explain this difference with reference to structure and bonding.

The Valence Shell Electron Pair Repulsion Theory (VSEPR) is used to predict the shapes of many chemical molecules.

Describe the main features of the VSEPR theory for predicting shapes of molecules.

State and explain the bond angle F-O-F in OF₂.

Deduce whether each of the three molecules oxygen difluoride, OF₂, phosphorus trifluoride, PF₃, and boron trichloride, BCl₃, are polar or non-polar.

Give a reason in each case.

Predict and explain the shapes and bond angles of the following molecules:

i)        BF₃ 

[2]

ii)       NBr₃ 

[2]

Ethene, C₂H₄, and hydrazine, N₂H₄, are hydrides of adjacent elements in the periodic table. 

State and explain the H一C一H bond angle in ethene and the H一N一H bond angle in hydrazine.

Hydrazine can be oxidised to form diimide, which is a useful compound used in organic synthesis. 

Deduce the molecular geometry of diimide, N₂H₂, and estimate its H–N–N bond angle.

Explain whether ethene and hydrazine are polar or non-polar.

Hydrazine forms a cation with an ethane-like structure called hydrazinediium,  N₂H₆²⁺.

Predict the value of the H–N–H bond angle in N₂H₆²⁺.

Draw the resonance structures for the following ions:

i)        Methanoate, HCOO^-.

[1]

ii)       Nitrate(III), NO₂^–.

[1]

Deduce the resonance structures of the carbonate ion, giving the shape and the oxygen-carbon-oxygen bond angle.

In December 2010, researchers in Sweden announced the synthesis of N,N–dinitronitramide, N(NO₂)₃. They speculated that this compound, more commonly called trinitramide, may have significant potential as an environmentally friendly rocket fuel oxidant. 

Deduce the N–N–N bond angle in trinitramide and explain your reasoning.

Predict, with an explanation, the polarity of the trinitramide molecule.

Silver chloride, AgCl, is a chloride compound that has uses in photography films as well as having antiseptic properties. 

Silver chloride has a high melting point and a structure similar to sodium fluoride.

Explain why, with reference to structure and bonding, why silver chloride has such a high melting point.

Cyanide is a fast-acting chemical, which can be found in various forms and can have toxic effects on the body.

Draw the Lewis structure for a CN^- ion. 

Show the outer electrons only.

Ammonia, NH₃, and boron trifluoride, BF₃, react together to form NH₃BF₃. Each of the molecules NH₃ and BF₃ have different features of its electronic structure which allows them to bond together.

Explain how the two molecules bond together and what type of bond is formed between NH₃ and BF₃.

You may use a labelled diagram to help you.

Aluminium chloride, Al₂Cl₆, does not conduct electricity when molten but aluminium oxide, Al₂O₃, does.

Explain this in terms of the structure and bonding of the two compounds.

State why magnesium and oxygen form an ionic compound while carbon and oxygen form a covalent compound.

Explain why the melting point of phosphorus(V) oxide is lower than that of sodium oxide in terms of their bonding and structure.

N, N–dinitronitramide N(NO₂)₃, also known as trinitramide, has been identified as a potentially more environmentally friendly rocket fuel oxidant.

Using Section 11 of the data booklet, outline how the length of the bond between nitrogen atoms in trinitramide compares with the bond between nitrogen atoms in nitrogen gas, N₂.

Describe the bonding within the carbon monoxide molecule.

Draw a diagram to show the resonance structure in a molecule of benzene. 

The energy change for hydrogenation of cyclohexene is -120 kJ mol^-1. However, when benzene undergoes hydrogenation, the energy change is 152 kJ mol^-1 less than expected. 

Use this data to explain the relative stabilities of benzene and the theoretical cyclohexa-1,3,5-triene molecule. 

With reference to bonding and hybridisation, describe the structure of benzene.

Deduce the number of possible resonance structures for the carbonate ion, CO₃^2-, and draw two of them.

Include the formal charges for each oxygen.

An alternative structure for the carbonate ion is proposed:

Explain why this structure is not accepted as another resonance structure for the carbonate ion.

Deduce the number of sigma (σ) and pi (π) bonds present in any of the resonance structures of the carbonate ions shown in part a).

Silicon can form silicon tetrachloride SiCl₄ and also silicon hexachloride, SiCl₆^2-.

i) Draw the Lewis structure for SiCl₄ and SiCl₆^2-.

[2]

ii) Use VSEPR theory to deduce the Cl-Si-Cl bond angles in both the SiCl₄ and SiCl₆^2- molecules.

[2]

iii) Predict the molecular geometry of each molecule.

[2]

Carbon can form CCl₄ but cannot form CCl₆^2-. Explain why.

Deduce which, if any, of SiCl₄ and SiCl₆^2-, are polar molecules and explain your choice.

Formal charge can be used to decide on the most stable, and therefore most likely, form a molecule can take. Resonance structures occur when more than one Lewis diagram describes a structure equally well.

i) Deduce the formal charge on the silicon and each chlorine within SiCl₄ and SiCl₆^2-

[2]

ii) Predict which will be the most stable molecule and explain your answer.

[2]

iii) Predict if any resonance structures are possible for SiCl₆^2- and explain your answer.

[2]

Natural rubber, polyisoprene, forms a flexible polymer in the following reaction:

i) Deduce the number of sigma (σ) and pi (π) bonds in the monomer.

[2]

ii) Deduce the number of sigma (σ) and pi (π) bonds in the repeating unit.

[2]

Deduce the number of carbons with a tetrahedral geometry in both the monomer, isoprene, and the repeating unit of the polymer, polyisoprene.

Polymer formation involves a radical intermediate to lengthen the polymer chain.

The radical in the formation of polyisoprene is shown below, where X represents the existing chain:

X-CH₂CCH₃CHCH₂

i) Identify the atom that is the radical in the structure shown.

[1]

ii) Deduce the formal charge on the radical atom.

[1]

iii) Use the information above, and your knowledge of structure and bonding, to predict if the structure is stable or not.

[2]

Isoprene is not produced directly by the rubber tree, but is the product of a series of biochemical reactions from the isopentenyl pyrophosphate molecules present in the tree.

The structure of isopentenyl pyrophosphate is shown below:

Deduce the number of sigma (σ) and pi(π) bonds present in one molecule of isopentenyl pyrophosphate.

One interhalogen compound is IF₅.

i) Draw the Lewis structure for IF₅.

[1]

ii) Use VSEPR theory to deduce the bond angles in IF₅.

[1]

iii) Predict whether IF₅ will be a polar molecule and explain your choice.

[2]

Iodine can also form the triiodide ion, I₃^-.

i) Draw the Lewis structure for I₃^-.

[1]

ii) Use VSEPR theory to deduce the bond angles in I₃^-.

[1]

iii) Explain the position of the lone pairs on the central iodine.

[2]

Deduce the formal charge on each of the iodine atoms in the triiodide molecule, I₃^-.

An alternative Lewis structure for the triiodide ion, I₃^-, is suggested:

Deduce the formal charges and use them to suggest if the structure is stable and likely to occur.

Explain why methanol is soluble in water.

Methanol, ethanol and propan-1-ol are all primary alcohols. Describe and explain the trend in their melting points shown below.

These longer primary alcohols have the following melting points:

Describe and explain this trend.

Predict, with a reason, whether ethanol or ethane-1,2-diol will have the higher melting point?

C₂H₆, C₄H₁₀ and C₃H₈ are alkanes.

i) Put them in order of increasing boiling point and explain your answer.

[3]

ii) Put them in order of increasing volatility and explain your answer.

[3]

Predict, with a reason, whether the alkanes are soluble in water and propanone.