Syllabus Edition
First teaching 2014
Last exams 2024
Two types of covalent bond are sigma and pi bonds.
) bond is formed[1]
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) bond is formed[1]
Describe the difference in the location of the electron dense regions in sigma () and pi () bonds.
Deduce the number of sigma () and pi () bonds in methane, CH4.
Deduce the number of sigma () and pi () bonds in oxygen, O2.
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Sulfur can form bonds with six fluorine atoms to form sulfur hexafluoride, SF6.
[1]
[1]
Sulfur has no lone pairs when bonded to fluorines in SF6. Predict the molecular geometry of sulfur hexafluoride, SF6.
State the F-S-F bond angles in SF6.
Phosphorus pentafluoride, PF5, is also a molecule with an expanded octet around the central atom.
[1]
[1]
[1]
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Although noble gases do not normally react, a few compounds are possible. One is xenon tetrafluoride.
Draw the Lewis structure (electron dot) for XeF4.
Predict the molecular geometry and electron domain geometry for the XeF4 molecule.
Predict and explain the F-Xe-F bond angle in XeF4
2 marksThe formal charge on an atom can be calculated by the following:
FC = (Number of valence electrons) - ½(Number of bonding electrons) - (Number of non-bonding electrons)
Calculate the formal charge on the xenon and the fluorines in xenon tetrafluoride, XeF4.
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Draw a Lewis (electron dot) structure for carbon dioxide, CO2.
Predict the molecular geometry and the O-C-O bond angle in carbon dioxide, CO2.
An alternative way to draw the carbon dioxide molecule is:
Identify the formal charge on each of the oxygen atoms.
State which of the Lewis structures, that from part a) or part c), is preferable and explain your choice.
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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.
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But-2-ene-1,4-dioic acid exists as both cis and trans isomers. The cis isomer is shown below
Identify how many sigma bonds and how many pi (π) bonds are present in cis but-2-ene-1,4-dioc acid.
Draw the Lewis structures, predict the shape and deduce the bond angles for xenon tetrafluoride.
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Structure |
I |
II |
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O atom labelled (1) |
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O atom labelled (2) |
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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
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Lewis (electron dot) structure |
FC of O on LHS |
FC of central N |
FC of N on RHS |
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A |
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B |
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C |
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Based on the formal charges assigned in part c), deduce which Lewis (electron dot) structure of N2O (A, B, or C) is the preferred.
Explain another factor that also must be taken into account in determining the preferred structure.
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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.
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Explain which of the two SO42– 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.
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Deduce the number of possible resonance structures for the carbonate ion, CO32-, 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).
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Silicon can form silicon tetrachloride SiCl4 and also silicon hexachloride, SiCl62-.
Carbon can form CCl4 but cannot form CCl62-. Explain why.
Deduce which, if any, of SiCl4 and SiCl62-, 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.
[2]
[2]
[2]
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Natural rubber, polyisoprene, forms a flexible polymer in the following reaction:
[2]
[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
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.
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One interhalogen compound is IF5.
[1]
[1]
[2]
Iodine can also form the triiodide ion, I3-.
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.
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