6.1 Covalent Structures Flashcards
What are covalent structures?
A covalent bond, also called a molecular bond, is a chemical bond that involves the sharing of electron pairs between atoms. Usually covalent bonds form when electrons are shared between two nonmetals.
What is a lattice?
A lattice is a regular repeated three-dimensional arrangement of atoms, ions, or molecules in a metal or other crystalline solid.
What is a covalently bonded network? (2)
-A covalent network crystal consists of atoms at the lattice points of the crystal, with each atom being covalently bonded to its nearest neighbour atoms. -The covalently bonded network is three-dimensional and contains a very large number of atoms.
What is a macro-molecular structure?
It is formed when one carbon atom is bonded to three other carbon atoms by strong covalent bonds.
What is a giant covalent structure? i.e Diamond
In diamond, each carbon shares electrons with four other carbon atoms, forming four single bonds.
Why is diamond not a molecule?
It is not a molecule, because the number of atoms joined up in a real diamond is completely variable depending on the size of the lattice.
What are the properties of diamond?
Properties: ● has a very high melting point (almost 4000 °C). Very strong carbon –carbon covalent bonds have to be broken throughout the structure before melting occurs. ● Is very hard. This is again due to the need to break very strong covalent bonds operating in 3-dimensions. ● Doesn’t conduct electricity. All the electrons are held tightly between atoms, and aren’t free to move.
Define Allotrope.
structural forms of the same element. (Diamond and Graphite are both allotropes of carbon)
What structure does graphite have?
Graphite has a layer structure.
What are the properties of Graphite? (3)
-Soft as the hexagonal sheets are held together by weak intermolecular bonds. Allowing for the sheets to ‘slide’ off one another. -Conductor of electricity, due to the delocalised electrons. -Higher melting/boiling point than diamond because the delocalised electrons create very large dispersion forces.
Why is graphite different to diamond?
Each carbon atom uses three of its electrons to form simple bonds to its three close neighbours. That leaves a fourth electron in the bonding level. These “spare” electrons in each carbon atom become delocalised over the whole of the sheet of atoms in one layer.
Why is the forces within graphite stronger than diamond?
The atoms within a sheet are held together by strong covalent bonds - stronger, in fact, than in diamond because of the additional bonding caused by the delocalised electrons. In graphite you have the ultimate example of van der Waals dispersion forces.
what do these electrons create within Graphite?
As the delocalised electrons move around in the sheet, very large temporary dipoles can be set up which will induce opposite dipoles in the sheets above and below - and so on throughout the whole graphite crystal.
What is Silicon Dioxide (quartz)?
It is also an allotrope of carbon, possessing properties similar to diamond.
Why is Quartz different to diamond?
They are differ slightly, as Silicon dioxide has longer bonds. (The bond between the oxygen and silicon atoms). Therefore less energy is required to overcome these bonds, resulting in a slightly lower melting/boiling point than diamond. Otherwise many if its other properties are closely related to diamond.
