Unit 1 - Periodicity: Types of Bonding in Elements Flashcards
Metals/Metallic Bonding
Metals in groups 1, 2 and 3 are held together by metallic bonds. The atoms in a metal are held together in a lattice by the attraction of positively charged cores (ions) and the negatively charged, delocalised electrons.
Properties of Metals
Metals are:
* Malleable (can be hammered into shape)
* Ductile (can be drawn into wire)
This is because the layers of metal atoms can slip over each other without breaking the metallic bond.
The delocalised electrons allow metals to be good conductors of electricity (in solid and liquid states).
The boiling points of metals give an indication of the strength of the metallic bond.
Metals have high density.
Metals - MP and BP trends
The bp and mp of the alkali metals decrease down the group. The bonding between the atoms decrease in strength down the group. The attraction between the positive core and outermost electrons decrease as the distance between them increases.
The boiling points increase across the period. The number of outermost electrons increases which increases the bonding strength. Increasing the number of delocalised electrons increases the force of attraction between the positive core and the outermost electrons.
Monatomic Elements (Noble Gases)
The noble gases are the only gaseous elements to exist as single atoms (monatomic) due to their stable electron arrangement.
They have low melting and boiling points as they are easily seperated by overcoming the weak forces of attraction between the atoms.
They have low density and do not conduct electricity.
Covalent Molecules
Definition
The covalent bond itself is a shared pair of electrons electrostically attracted to the positive nuclei of two non-metal atoms. The atoms achieve a stable outer electron arrangement by sharing electrons.
Covalent Molecules
Some non-metal elements exist as covalent molecules to achieve a stable electron arrangement. Within each molecule, the atoms are held together by strong covalent bonds.
Discrete covalent molecules are small groups of atoms held together by strong covalent bonds inside the molecule and weak intermolecular forces between the molecules.
They have low density and do not conduct electricity.
Examples of Covalent Molecules
Diatomic molecules: Hydrogen, Nitrogen, Fluorine, Oxygen, Iodine, Chlorine, Bromine
Phosphorus exists as tertahedral P4 molecules
Sulfur molecules exist as a ring structure (S8)
Carbon can exist as discrete molecules known as fullerenes. There are many types of fullerene. Examples include buckyball, clusters, nanotubes, rings.
Covalent Network
Definition + Properties
Covalent networks are large, rigid three-dimensional arrangements of atoms held together by strong covalent bonds.
They have high melting points because they only contain strong bonds.
They also have very hign density and do not conduct electricity (with the exception of graphite)
Covalent Networks Examples
Boron and silicon exist as covalent networks.
Carbon has a covalent network structure in diamond and graphite.
Diamond: Each carbon is bonded to 4 other carbon atoms to form a tertahedral lattice. Diamond is very hard because strong covalent bonds must be broken to shape diamond.
Graphite: In graphite each carbon atom is covalently bonded to 3 other carbon atoms to form a layered structure. Weak forces between the layers allow them to slide over each other. Within the layer, only 3 electrons are used in covalent bonding. The other electron is delocalised so graphite conducts electricity.
Covalent Radius
Definition
The covalent radius is a measure of the size of an atom. It is half the distance between the centres of adjacent bonded atoms.
Covalent Radius
Trends
The size of an atom (covalent radius) increases down a group of the periodic table as the number of occupied electron shells increases.
The atomic size decreases across a period because the increased nuclear charge attracts the outermost electrons more strongly.
