Csec Chemistry Notes 5

Structure and bonding

What determines the physical properties of solids? The arrangement of atoms and ions in a crystal helps to determine the physical properties of thermal and electrical conductivity, melting and boiling points, physical state at room temperature and solubility in solvents. Solids can be divided into the following groups: ionic crystals, molecular (simple and giant) and metallic crystals. The differences in the properties can be explained by the type of bonds they possess. Ionic compounds are crystalline solids, able to conduct electricity when molten, due to the movement of ions which can carry an electric current. These compounds also have high melting and boiling temperatures, due to the strength of the attractive forces between the ions. Simple molecular crystals have low melting and boiling temperatures, due to weak forces of attraction between molecules. Giant structures of atoms and molecules have very high melting and boiling points, due to strong covalent bonds throughout their three-dimensional network. Metals are able to conduct electricity and heat, due to the presence of mobile electrons. What is metallic bonding? In metallic crystals, the outer electrons of each atom are mobile or delocalised (that is, they do not belong to any particular cation) and come together to form a band or sea of electrons. These will bind to the cations formed from the electron loss, forming a strong bond. In this way, metals are able to conduct heat and electricity since the mobile electrons can move throughout the metal. The strong bonds between the cations and electrons mean that they are hard to break, thus, metals have high melting and boiling points. Metals are also solids (except mercury) and are malleable and ductile. The bonding in metals can be represented below: º+ e º+ e º + eº + e The mobile electrons form a cloud or band surrounding the cations. The difference in charges holds them together into a strong bond. How are ionic bonds formed? In an ionic crystal, the attraction between cations and anions holds the crystal together into a regular three-dimensional framework. Each cation is surrounded by anions and vice versa. These crystals are solids at room temperature and are unable to conduct electricity in this state. However, imagine what happens when these ionic crystals are heated. The ions gain more energy to move but, because they are oppositely charged, it requires vast amounts of energy to break this force of attraction, and so, these crystals have high melting and boiling points. Please note that ionic solids can only conduct electricity when molten, as only then are the ions free enough to move. Examples of ionic solids are sodium chloride, magnesium oxide and potassium iodide. What is the difference between giant and simple molecular crystals? In giant molecular crystals, such as graphite, diamond and silicon dioxide, strong covalent bonds exist between the atoms, which make them difficult to melt or boil. On the other hand, simple molecular crystals

have covalent bonds within molecules but weak bonds between molecules. Hence, the molecules separate easily at fairly low temperatures. Structure and bonding (Part 3) Francine Taylor-Campbell, Contributor What is allotrophy? Allotrophy is the ability of an element to exist in the same physical state but in different structural forms. This causes them to have different physical properties but the same chemical properties How are the atoms in graphite, diamond and sodium chloride arranged? Diamond and graphite are giant molecular or macromolecular crystals. A diamond consists of carbon atoms tetrahedrally arranged and bonded by strong covalent bonds. A graphite consists of carbon atoms arranged in hexagonal rings and in layers, while sodium chloride is an ionic solid having a giant structure.

What are the main differences in bonding between diamond and graphite? In diamond, four carbon atoms are joined in a tetrahedral arrangement. This is repeated throughout to give a three-dimensional structure with strong, covalent bonds. In a graphite, each carbon atom is bonded to three other atoms, arranged hexagonally in layers. These layers are held together by weak bonds which enable them to slide over one another. There are strong, covalent bonds, however, between the carbon atoms in each layer. Note that, for graphite, since the carbon atom is bonded to only three others, it means that each carbon atom has a fourth electron not involved in bonding. That is a free mobile electron. This will influence the properties of graphite.

What are diamond and graphite allotropes? Graphite and diamond are composed of carbon atoms but their structures are different, hence, these solids are allotropes. They show the same chemical properties, since they have the same element carbon, but the difference in their structure causes them to have different chemical properties. How do the properties of sodium chloride, graphite and diamond differ?

Property

Sodium Chloride

Diamond

Graphite

Apperance

Crystalline solid

Sparkling solid

Dark solid

Hardness

Brittle - easily split

Soft and flaky - due th Very hard - due to strong eweak bonds between convalent bonds in the the layers. Layers can slip structure over each other hence it is a good lubricant.

Melting point

High due to strong ionic bonds that need a lot of energy to break

Very high due to strong covalent bonds that need vast amounts of energy to break.

Electrical Conductivity

Conducts electricity when dissolved in solution or when molten as the ions are free to move.

Cannot conduct electricity because all Conducts electricity when electrons are involved in solid as mobile electrons bonding hence there are are present no free electrons to carry a current.

Very high due to strong covalent bonds that are difficult ot break