Unit 2: Chemical bonding Flashcards
What are the 3 types of bonding?
- Metallic bonding
- Covalent bonding
- Ionic bonding
1) what type of bonding holds metals together?
2) what type of bonding holds non-metals together?
3) What type of bonding holds a metal and a non-metal together?
1) Metallic bonding
2) Covalent bonding (noble gases are an
exception as they do not bond at all)
3) Ionic bonding
Define what an ion is:
A charged particle: An atom which has lost or gained electrons
Describe metallic bonding
Describe metallic bonding as a lattice of positive
ions in a ‘sea of electrons’.
Metal atoms have relatively few electrons in their outer shells. Metal atoms more easily lose electrons than gain them. So, they become positive ions. In doing so, they achieve a more stable electron arrangement, usually that of the nearest noble gas. When they are placed together, each metal atom loses its outer electrons into a ‘sea’ of mobile (free) electrons.
The structure of the metal is made up of positive ions packed together. These ions are surrounded by free delocalised electrons (not restricted to orbiting one positive ion) which can move between the ions forming a electrostatic attraction ‘glue’ all throughout the metal. Metals can conduct electricity because these free electrons can carry a charge/current through the metal.
Describe covalent bonding
- bond is formed by the sharing of electrons between atoms
- Each atom contributes an electron/electrons to each bond
- Molecules are formed from atoms linked together by covalent bonds
- Non-metals want to gain electrons to fill their outer shell, since none of them are willing to give any, they have to share them
- the forces of attraction between the shared electrons and nuclei are stronger than any repulsive forces
- When making a diagram for covalent bonds you only show the outer shell because it is only the valence electrons which are involved in bonding
- Covalent bonds can result in simple molecules or giant molecular lattices (giant covalent lattices)
Describe ionic bonding
Describe the formation of ionic bonds between
metallic and non-metallic elements to include
the strong attraction between ions because of
their opposite electrical charges
- Electrons are transferred from one atom to another
- This produces charged particles known as ions
- The ionic bond is between a positive and negative ion (cation and anion)
- The ions are held together by electrostatic forces of attraction, forming giant ionic lattices
- Metals tend to form cations
- Non-metals tend to form anions
- The number of electrons gained or lost by an atom during ionisation is known as valency, this can be determined from the periodic table
Describe ionic lattices
Describe the lattice structure (also known as a crystal structure) of ionic compounds
as a regular arrangement of alternating positive
and negative ions, exemplified by the sodium
chloride structure.
These form crystal structures where the ions are held in a three-dimensional arrangement. This arrangement is called a lattice. A lattice forms because the electrostatic forces between the charged ions are very strong. A single ion is actually attracted to more than one oppositely charged ion in a regular arrangement of alternating cations and anions.
Sodium chloride: halogen + alkali
Use dot-and-cross diagrams to describe the
formation of ionic bonds between Group I and Group VII
halogen + alkali
State that there are several different forms of
carbon, including diamond and graphite
Allotropes: Different forms of the same element in the same state
There are many allotropes of carbon. There is:
- Diamond
- Graphite
- Graphene
- Sulfur dioxide (SO2)
- fullerenes
Describe the giant covalent structures of graphite
and diamond
Graphite:
- each carbon atom forms three covalent bonds with other carbon atoms
- the carbon atoms form layers of hexagonal rings
- there are no covalent bonds between the layers
- there is one non-bonded - or delocalised - electron from each atom
- Graphite has delocalised electrons, just like metals. These electrons are free to move between the layers in graphite, so graphite can conduct electricity.
- The forces between the layers in graphite are weak. This means that the layers can slide over each other.
- Graphene is just one of the flat layers of graphite (no weak intermolecular forces with other layers)
- Graphene is more flexible than graphite. When under stress it just slide apart.
- Graphene is transparent because of its sheer thinness
- Graphene has a high melting point, it conducts electricity
Diamond:
- each carbon atom is joined to four other carbon atoms by strong covalent bonds
- the carbon atoms form a regular tetrahedral network structure
- there are no free electrons
The rigid network of carbon atoms, held together by strong covalent bonds, makes diamond very hard. Diamond has a very high melting point and it does not conduct electricity.
Fullerenes
Fullerenes are molecules of carbon atoms with hollow shapes. Their structures are based on hexagonal rings of carbon atoms joined by covalent bonds. Some fullerenes include rings with five or seven carbon atoms. Two examples of fullerenes are buckminsterfullerene and nanotubes.
Buckminsterfullerene
Buckminsterfullerene was the first fullerene to be discovered. Its molecules are made up of 60 carbon atoms joined together by strong covalent bonds. Molecules of C60 are spherical.
There are weak intermolecular forces between molecules of buckminsterfullerene. These need little energy to overcome, so buckminsterfullerene is slippery and has a low melting point.
Nanotubes
A nanotube is like a layer of graphene, rolled into a cylinder. The length of a nanotube is very long compared to its width, so nanotubes have high length to diameter ratios.
Nanotubes have high tensile strength, so they are strong in tension and resist being stretched. Like graphene, nanotubes are strong and conduct electricity because they have delocalised electrons.
Relate the structures of diamond and graphite
to their uses, e.g. graphite as a lubricant and a
conductor and diamond in cutting tools
Diamond
The rigid network of carbon atoms, held together by strong covalent bonds, makes diamond very hard. This makes it useful for cutting tools, such as diamond-tipped glass cutters and oil rig drills.
It is transparent and sparkles in light - used in jewellery and ornamental objects
Graphite
In pencils as lubricant - graphite is slippery coz layers can slide over each other
graphite can conduct electricity. This makes graphite useful for electrodes in batteries and for electrolysis.
Graphene
graphene useful in electronics and for making composites.
Nanotubes:
These properties make nanotubes useful for nanotechnology, electronics and specialised materials.
Describe the macromolecular structure of
silicon(IV) oxide (silicon dioxide, SiO2)
extra: why is the formula SiO2 and not SiO4?
Silicon (IV) oxide has a tetrahedral macromolecular structure (similar to
diamond). Each Si atom is covalently bonded to four Oxygen atoms. Each Oxygen atom is covalently bonded to two Silicon atoms. The formula is therefore SiO2. Each Oxygen atom has 2 non-bonded electron pairs (lone pairs) that are
localised onto it and are not free to move.
There are no free (delocalised) electrons, hence SiO2 cannot conduct
electricity. The three-dimensional structure is rigid and hard.
It has a high melting point and high boiling point.
extra: Each oxygen atom is shared 50-50 by two silicon atoms, each silicon atom has in reality 50% “possession” of the oxygen. Therefore each silicon atom is in essence surrounded by 4*50% = 2 (effective)oxygen atoms.
Describe the formation of single covalent bonds
in H2, Cl2, H2O, CH4, NH3 and HCl
the sharing of pairs of electrons leading to the noble gas configuration
Describe the differences in volatility, solubility
and electrical conductivity between ionic and
covalent compounds
Ionic structures:
Usually soluble in water (but not in organic solvent s e.g. ethanol,methylbenzene): This is because water is a polarized molecule itself. So it is attracted to charged ions and thus many ionic solids dissolve.
When Solid ions do NOT conduct electricity: because the ions are held firmly in place. The ions cannot move to conduct the electric current. But when an ionic compound melts, the charged ions are free to move. Therefore, molten ionic compounds or ionic compounds dissolved in water do conduct electricity.
They are usually crystalline solids at room temperature. They have high melting and boiling points and are therefore non-volatile.
Covalent structures:
They are often liquids or gases at room temperature.
Do not conduct electricity - this is because they do not have any free electrons or an overall electric charge (there are no ions).
They have low melting and boiling points and are therefore volatile.
They are soluble in organic solvents such as ethanol or methylbenzene (very few are soluble in water). This is because covalent molecular substances dissolve in covalent substances. Water molecules are polar molecules, but covalent substances are non-polar, so there aren’t any forces of attraction which can let the water molecules “latch on” and break apart the molecules.
Explain the differences in melting point and
boiling point of ionic and covalent compounds in
terms of attractive forces
The ions making up an ionic compound interact through the electrostatic attraction of the full chemical bonding. These forces work in all directions in the solid and strongly hold the ions in place in the structure (i.e. they have strong intermolecular forces and intramolecular forces). It takes a great deal of energy to separate the positive and negative ions in a crystal lattice. This means that ionic compounds have high melting points and boiling points.
Simple covalent compounds are made of molecules, within molecules the bonds are very strong (i.e. they have strong intramolecular forces). However the bonding does not act between a molecule and the molecules around it (i.e. there are weak intermolecular forces). Therefore it doesn’t take as much energy to break apart these molecules, so they have lower melting and boiling points.
What polyatomic ions should you remember?
