C2 (bonding, structure, the properties of matter) Flashcards
how does ionic bonding work? use lithium, a metal, and fluorine, a non-metal, in your example:
- neither lithium or fluorine have a full outer energy level; lithium has 1 electron in its outer energy level, fluorine has 7.
- lithium can react to lose one electron to gain a full outer energy level, and fluorine can use this electron to fill its outer energy level. now both fluorine and lithium are stable ions.
- lithium lost an electron, and so has a charge of 1+, and fluorine gained an electron, and so has a charge of 1-.
between which elements does ionic bonding occur?
ionic bonding occurs between metal and non-metals. in the end, both atoms have the structure of a noble gas.
how does covalent bonding occur? use two hydrogen atoms in your example:
- hydrogen atoms have only one electron in their outer shells, and they require two.
- the two hydrogen atoms overlap their electron shells, and now both atoms have two electrons in their outer electron shells.
- these two hydrogen atoms are now stable
- covalent bonding only occurs between two non-metal atoms
describe metallic bonding:
- occurs when metals bond with other metals
- metals are giant structures of atoms, arranged in a regular pattern. when these atoms are together, they give up their electrons in their outer shells, and share them with all the other atoms DELOCALISED ELECTRONS
- these atoms all become positive ions
- there is now a strong electrostatic attraction between the positive ions and the negative electrons, and this holds everything together in a regular structure
- ‘sea of electrons surrounding a positively charged lattice’
what characteristics does an ionic lattice give metals?
this means that metals are incredibly strong, have high melting and boiling points, and are good conductors of electricity and heat (delocalised electrons can carry heat and electricity through the metal).
- also malleable (as they’re such a regular structure, the different layers can slide over one another) - the case for PURE METALS, as all the electrons are regular shapes
why are alloys so strong?
- alloys contain two or more different elements, with different sized atoms (e.g. steel)
- these different sized atoms disrupt the metal’s regular structure, and means that the atoms can no longer easily slide over one another, making the alloy much harder than pure metal
what do ionic compounds form?
- they form giant ionic lattices. every positive ion is surrounded by a negative ion, and vice versa.
- giant ionic lattices are 3d, and have very strong forces of attraction between the ions (ELECTROSTATIC FORCES = ionic bonds)
what are the properties of ionic compounds?
- very high melting and boiling points (strong electrostatic forces require a lot of heat energy to break)
- they cannot conduct electricity when they’re solids, as the ions cannot move (the electrostatic forces lock them in place)
- WHEN IONIC COMPOUNDS CONDUCT ELECTRICITY, IT’S THE IONS THAT MOVE, NOT THE ELECTRONS!
what are the properties of small covalent molecules?
- low melting and boiling points (usually gases/liquids at room temperature)
> the intermolecular forces are
incredibly weak, and don’t require very
much energy to break - don’t conduct electricity (because they’re neutral, and don’t have any charged particles to move and carry the charge)
what is a polymer, and what are their properties?
- each polymer molecule is made up of smaller molecules, called monomers.
- a single polymer is strong, as its connecting covalent bonds are strong. however, to break up multiple polymers, you’d have to break the intermolecular forces, which are much weaker.
- however, polymers are long and have a large surface area, and so the accumulated intermolecular forces are quite difficult to break.
- lower melting and boiling points than giant structures, but higher points than simple structures (e.g. water)
- generally solid at room temperature
what is a property of giant covalent molecules?
- always solids at room temperature (many strong covalent bonds - high melting and boiling points)
describe diamond:
- formed from the element carbon
- each carbon atom forms 4 covalent bonds to 4 other carbon atoms - very high melting and boiling point
- cannot conduct electricity (no free electrons to carry electric charge)
describe graphite:
- formed from carbon. (GRAPHITE IS SIMILAR TO METAL, BUT ISN’T A METAL) each carbon atom forms three covalent bonds with other carbon atoms. (forms hexagonal rings)
- high melting and boiling point (takes a great deal of energy to break covalent bonds)
- soft and slippery (hexagonal rings are arranged in layers that can easily slide over each other)
- excellent conductor of electricity and heat energy (each carbon atom has a delocalised electron, which can conduct thermal energy and electricity)
describe graphene:
- a single layer of graphite (1 atom thick)
- excellent conductor of electricity and heat (has delocalised electrons, like graphite)
- extremely strong
describe fullerenes:
- molecules of carbon atoms with hollow shapes
- usually, they have hexagonal rings of carbon atoms, but can also have rings of 5/7 carbon atoms
