Explain the relative lengths of the three bonds between N and O in nitric acid.

How many sigma (σ) and pi (π) bonds are present in this molecule?

Discuss the bonding in the resonance structures of ozone.

Deduce, giving a reason, the more likely structure.

What is the formal charge of the oxygen atom in H₃O⁺?

A.  −2

B.  −1

C.  0

D.  +1

Which overlap of atomic orbitals leads to the formation of only a sigma (σ) bond?

       I.     s − p

       II.     p − p

       III.     s − s

A.     I and II only

B.     I and III only

C.     II and III only

D.     I, II and III

Deduce the number of σ and $\pi$ bonds in a molecule of ethyne.

Deduce one resonance structure of ozone and the corresponding formal charges on each oxygen atom.

The C–N bonds in urea are shorter than might be expected for a single C–N bond. Suggest, in terms of electrons, how this could occur.

Sketch the shape of one sigma ($\text{σ}$) and one pi ($\mathrm{\pi }$) bond.

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 the Lewis (electron dot) structure and molecular geometry and the bond angles of PCl₃.

What is the number of sigma (σ) and pi (π) bonds in the molecule (NC)₂C=C(CN)₂?

Draw two Lewis (electron dot) structures for BrO₃⁻.

Suggest why the loss of ozone is an international environmental concern.

How many sigma (σ) and pi (π) bonds are present in hydrogen cyanide, HCN?

Identify the number of sigma and pi bonds in HCN.

Describe how sigma (σ) and pi ($\pi$) bonds are formed.

(i) Uranium hexafluoride, UF₆, is used in the uranium enrichment process that produces fuel for nuclear reactors.

State the molecular shape of uranium hexafluoride.

(ii) Explain why uranium dioxide, UO₂, has a very high melting point whereas uranium hexafluoride vapourises easily into gas.

Which species has bond angles of 90°?

A. AlCl₄^-

B. $\text{I}$Cl₄^-

C. NH₄⁺

D. SiCl₄

Which molecules contain two pi ($\mathrm{\pi }$) bonds?

I.   HCN
II.  H₂CO₃
III. H₂C₂O₄

A.  I and II only

B.  I and III only

C.  II and III only

D.  I, II and III

State what the presence of alternative Lewis structures shows about the nature of the bonding in the molecule.

Formulate two equations to show how nitrogen(II) oxide, NO, catalyses the destruction of ozone.

Which combination correctly describes the geometry of ${{\mathrm{BrF}}_4}^{-}$?

Outline why both C to O bonds in the conjugate base are the same length and suggest a value for them. Use section 10 of the data booklet.

Outline, in terms of the bonding present, why the reaction conditions of halogenation are different for alkanes and benzene.

Draw the Lewis (electron dot) structures of PF₃ and PF₅ and use the VSEPR theory to deduce the molecular geometry of each species including bond angles.

$\mathrm{CFC}$s produce chlorine radicals. Write two successive propagation steps to show how chlorine radicals catalyse the depletion of ozone.

Absorption of UV light in the ozone layer causes the dissociation of oxygen and ozone.

Identify, in terms of bonding, the molecule that requires a longer wavelength to dissociate.

State the number of sigma ($\mathrm{\sigma }$) and pi ($\mathrm{\pi }$) bonds around the central carbon atom in molecule B.

When excess ammonia is added to copper(II) chloride solution, the dark blue complex ion, [Cu(NH₃)₄(H₂O)₂]²⁺, forms.

State the molecular geometry of this complex ion, and the bond angles within it.

Molecular geometry:

Bond angles: 

Draw diagrams to show how sigma (σ) and pi (π) bonds are formed between atoms.

Dinitrogen monoxide in the stratosphere is converted to nitrogen monoxide, NO (g).

Write two equations to show how NO (g) catalyses the decomposition of ozone.

Deduce a Lewis (electron dot) structure of the nitric acid molecule, HNO₃, that obeys the octet rule, showing any non-zero formal charges on the atoms.

State the number of $\text{σ}$ (sigma) and $\mathrm{\pi }$ (pi) bonds in Compound A.

Distinguish between a sigma and pi bond.