Chemical Bonding: Chemistry 9701 (Cambridge International AS & A Level)

The table below gives approximate Pauling electronegativity values for four elements.

Using these values, which of the following single covalent bonds has the greatest bond polarity?

Magnesium reacts with oxygen to form magnesium oxide, and magnesium metal itself is a good electrical conductor. Separately, ammonia gas reacts with hydrogen chloride gas to form solid ammonium chloride, $\text{NH}_4\text{Cl}$, which contains the ammonium ion, $\text{NH}_4^+$.

(a) Magnesium oxide has a melting point of over $2800\,^\circ\text{C}$, far higher than that of sodium chloride ($801\,^\circ\text{C}$). Explain, in terms of the ions present and the electrostatic forces between them, why magnesium oxide has such a high melting point. [2]

(b) Explain, in terms of structure and bonding, why solid magnesium metal conducts electricity, whereas solid magnesium oxide does not. [2]

(c) Use your knowledge of dative (coordinate) bonding to explain how the fourth N–H bond in $\text{NH}_4^+$ forms when $\text{NH}_3$ reacts with a hydrogen ion, $\text{H}^+$. State which atom donates the bonding electron pair, and state how this dative bond compares in length and strength with the other three N–H bonds once it has formed. [3]

(a) Nitrogen has $5$ outer-shell electrons. In ammonia, $\text{NH}_3$, three of these electrons are used to form covalent bonds to hydrogen atoms. State the total number of electron pairs (bonding and lone) around the central nitrogen atom, and hence state the shape and H–N–H bond angle of the $\text{NH}_3$ molecule. [2]

(b) Phosphorus forms the compound $\text{PCl}_5$ by bonding to five chlorine atoms, with no lone pairs remaining on the phosphorus atom. State the shape of $\text{PCl}_5$, and state the two different Cl–P–Cl bond angles present, explaining why there are two different values rather than one single bond angle. [3]

(c) The bond angle in $\text{NH}_3$ is $107^\circ$, whereas the bond angle in $\text{CH}_4$ is $109.5^\circ$. Explain, in terms of electron pair repulsion, why these two bond angles differ even though both central atoms are surrounded by four electron pairs. [2]

Water, $\text{H}_2\text{O}$, and hydrogen sulfide, $\text{H}_2\text{S}$, are both simple covalent molecules with the same bent (non-linear) shape, yet water boils at $100\,^\circ\text{C}$ while hydrogen sulfide boils at only $-60\,^\circ\text{C}$.

(a) Explain, in terms of the intermolecular forces present in each substance, why water has a much higher boiling point than hydrogen sulfide. [3]

(b) Describe how a hydrogen bond forms between two neighbouring $\text{H}_2\text{O}$ molecules. Your answer should state the three structural conditions needed for hydrogen bonding to occur, and state the approximate O–H$\cdots$O bond angle along the hydrogen bond. [3]

(c) Ice is less dense than liquid water at $0\,^\circ\text{C}$, so ice floats. Explain this observation in terms of the arrangement of hydrogen bonds in solid ice compared with the liquid. [2]

Ethyne, $\text{C}_2\text{H}_2$, has the structure $\text{H–C}{\equiv}\text{C–H}$, in which the two carbon atoms are joined by a triple bond.

How many $\sigma$ (sigma) bonds and how many $\pi$ (pi) bonds are present in one molecule of ethyne?

The table below gives approximate Pauling electronegativity values for four elements.

Using these values, which pair of elements would be expected to form a bond with the greatest ionic character?

Carbon dioxide, $\text{CO}_2$, has the structure $\text{O}{=}\text{C}{=}\text{O}$, in which the central carbon atom forms two carbon-oxygen double bonds and has no lone pairs of electrons.

(a) Describe, in terms of shared electron pairs, how each carbon-oxygen double bond in $\text{CO}_2$ is formed. State how many electron pairs in total are shared between carbon and each oxygen atom. [2]

(b) State the total number of electron-pair "regions" around the central carbon atom (treating each double bond as one region for the purpose of VSEPR theory), and hence state the shape of the $\text{CO}_2$ molecule and its O–C–O bond angle. [2]

(c) Each individual C=O bond is polar, because oxygen is more electronegative than carbon. Despite this, the $\text{CO}_2$ molecule as a whole has no overall (net) dipole moment. Explain why, referring to the shape of the molecule found in part (b). [3]

Boron has $3$ outer-shell electrons and forms the compound $\text{BF}_3$ by bonding to three fluorine atoms, with no lone pairs remaining on the boron atom. Sulfur has $6$ outer-shell electrons and forms the compound $\text{SF}_6$ by bonding to six fluorine atoms, also with no lone pairs remaining on the sulfur atom.

(a) State the total number of electron pairs around the central atom in $\text{BF}_3$, and hence state the shape of $\text{BF}_3$ and its F–B–F bond angle. [2]

(b) State the total number of electron pairs around the central atom in $\text{SF}_6$, and hence state the shape of $\text{SF}_6$. State the two different F–S–F bond angles present in the molecule. [3]

(c) Nitrogen, which is in the same period as boron, cannot form an analogous five-fluorine compound $\text{NF}_5$, whereas sulfur can form $\text{SF}_6$. Suggest, in terms of available orbitals, why period 3 elements such as sulfur can exceed the "normal" limit of four electron pairs (an octet) around the central atom, while period 2 elements such as nitrogen cannot. [2]

The boiling points of the Group 17 elements increase steadily down the group: $\text{F}_2$ boils at $-188\,^\circ\text{C}$, $\text{Cl}_2$ at $-34\,^\circ\text{C}$, $\text{Br}_2$ at $59\,^\circ\text{C}$, and $\text{I}_2$ at $184\,^\circ\text{C}$.

Which of the following best explains this trend?

(a) Describe the structure and bonding present in a typical metal such as sodium, in terms of the particles present and the forces holding them together. [2]

(b) The melting points of the period 3 metals are: sodium $98\,^\circ\text{C}$, magnesium $650\,^\circ\text{C}$, aluminium $660\,^\circ\text{C}$. Explain why melting point increases from sodium to aluminium, referring to the charge on the metal ion and the number of delocalised electrons contributed by each atom. [3]

(c) The melting points of the group 1 metals are: lithium $181\,^\circ\text{C}$, sodium $98\,^\circ\text{C}$, potassium $64\,^\circ\text{C}$. Explain why melting point decreases down group 1 from lithium to potassium, even though each atom contributes the same number of delocalised electrons in every case. [2]