Types of Bonding & Properties (AQA A Level Chemistry)

Both lithium and lithium chloride contain ions of lithium. However, the structure, bonding and properties of these substances are very different. 

State how the ions are held together in solid lithium and in solid lithium chloride. 

Table 1 below shows the melting and boiling points of lithium and lithium chloride. 

Table 1

State and explain what can be deduced from the information in Table 1. 

Two students, A and B, are comparing the properties of lithium and lithium chloride.

Student A states that both lithium and lithium chloride will conduct electricity, but Student B states that only lithium will conduct electricity. 

State whether Student A, Student B, or neither student is correct. Explain your answer. 

Student A and B then went on to discuss the bonding in two other substances, Li₂S and CS₂. 

Explain, in terms of electrons, how Li₂S is formed from its atoms.

Ammonia, NH₃, is a chemical which is key in the manufacture of certain fertilisers and cleaning products. An ammonia molecule will react with an H⁺ ion, to form ammonium, NH₄⁺.

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

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

Like ammonia, graphite is also a covalent substance, but it is a solid at room temperature.
Graphite has a melting point of around 3600 C.

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. 

Hydrochloric acid and sodium carbonate solution react together in a neutralisation reaction.

Write a full and a net ionic equation for this reaction.

The elements sodium and chlorine can react together to form the compound, sodium chloride. 

Table 1 below shows some data about sodium chloride, as well as the elements sodium and chlorine. 

Table 1

Sodium is a very reactive metal, as are all of the metals in Group 1 of the Periodic Table. However, not all metals behave in the same way. Copper has a very low reactivity. However, copper is very ductile and is an extremely good conductor of both electricity and heat. 

Explain why copper is ductile (can be stretched into wires). 

Explain why copper is a good conductor of electricity.

The structure and bonding in silicon dioxide, SiO₂, is very similar to that of diamond, and also has similarities to that of graphite. 

State the type of crystal structure which SiO₂ displays.

Explain why silicon dioxide, SiO₂, has such a high melting and boiling point.

Graphite also has a similar structure to diamond, but does not share all of the same properties. Table 1 below shows the key properties of graphite.

Table 1

The structure and bonding in sodium oxide is completely different to that of silicon dioxide, despite oxygen being present in both compounds.

Draw a ‘dot and cross’ diagram to show the bonding present in sodium oxide.
Only the outer electrons should be shown in your diagram.

Calcium is found in Group 2 of the Periodic Table and has a giant metallic structure.
Draw a labelled diagram to show the structure of a metal.

Calcium will react with the elements of Group 7 to form ionic compounds, which have an ionic lattice structure.

Describe what is meant by the term ionic lattice.

One example of an ionic compound which calcium forms is calcium chloride.

Using your knowledge of the structure and bonding of an ionic substance, state whether calcium chloride would have a high or low melting point.
Explain your answer.

Chlorine gas has a very different structure and different properties to calcium chloride, as it is a simple covalent molecule. Oxygen is another example of a simple covalent molecule. 

This question is about ionic bonding.

The strength of ionic bonding in different compounds can be compared by using the amount of energy required to separate the ions, as shown in Table 1.

Table 1

Using the data from Table 1, explain how changes in the cation affect the bond strength in an ionic compound.

Coal often contains traces of iron (II) disulfide, FeS₂. FeS₂ is an ionic compound of Fe²⁺ ions and S₂²⁻ ions.

This question concerns period 3 of the periodic table.

Figure 1 below shows the trend in melting point across period 3 (Na to Ar).

Figure 1

Graphene, graphite and diamond are all forms of solid carbon.

Explain, in terms of structure and bonding, why graphene and graphite are good electrical conductors but diamond is a poor electrical conductor.

You may include labelled diagrams in your answer.

Magnesium fluoride is a white crystalline salt that has a giant ionic lattice structure.

State whether the following substances conduct electricity when solid or molten, and explain your answers in terms of the particles involved:

• magnesium
• magnesium fluoride
• boron tribromide

Sodium chloride and iodine are both solids. Sodium chloride does not melt until it reaches a temperature of 1074 K yet iodine sublimes when heated gently, giving off purple vapours. Sodium chloride will conduct electricity when molten and iodine is a very poor conductor of electricity.

The nitrate (V) ion, NO₃^-, is a polyatomic ion, bonded by covalent bonds. The three oxygen atoms are bonded by one single covalent bond, one double covalent bond and one dative covalent bond.

An ionic compound has the empirical formula H₄N₂O₃. 

Suggest the formulae of the ions present in this compound.

The compounds SO₂ and MgO are both oxides but with different melting points as shown in Table 1.

Table 1

Across a period there is a repeating trend in properties of elements, known as  periodicity. Table 1 below gives the melting point trends for each of the period 3 elements Na – Ar.

Table 1

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 has a structure similar to that of sodium fluoride.

Diamond and ice are types of crystals. Describe the structure and bonding present in both diamond and ice. Explain why the melting point of these two substances is very different. Refer to the type of bonding present in both diamond and ice in your answer.

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.

Table 1 shows the melting and boiling points, in Kelvin, of calcium and chlorine.

Table 1

This question is about periods 3-4 of the periodic table.

Table 2 below shows the melting points of elements in Groups 14 – 17. Phosphorus and sulfur exist as P₄ and S₈ molecules respectively.

Table 2

Explain, with reasons, on the similarities and differences in the trends across Period 3 and Period 4.

Cyanide is a fast acting chemical, which can be found in various forms and can have toxic effects on the body. The cyanide ion, CN^-, is isoelectronic with carbon monoxide.

Calcium permanganate is a chemical oxidising agent, which is a compound used in the textile industry and contains calcium and manganate ions.

Ethyne, (C₂H₂), is a hydrocarbon which can be used in the artificial ripening of fruits and can be synthesised by a two step process.

Firstly, calcium oxide and coal react together to form calcium carbide (CaC₂) and carbon monoxide. The second step involves calcium carbide reacting with water to form calcium hydroxide and ethyne, (C₂H₂).

Mercury is a naturally occurring element that is a liquid at room temperature and pressure. Some of its physical properties are shown in Table 1.

Table 1

Mercury can react with bromine to form mercury (II) bromide (HgBr₂). 

Compare the electrical conductivity of mercury and mercury bromide, in both solid and molten states.

Figure 1 shows the melting points of some of the elements in period 3.

Figure 1