Gerard Kerr (drareg_rrek) wrote in exam_panic_zomg,
Gerard Kerr
drareg_rrek
exam_panic_zomg

Crash Course in Chemistry Pt. 1.

Atomic Structure and Bonding.

Atoms

You all know about electrons, protons and neutrons and stuff, but just in case, heres a little summary.
Protons have +1 charges and are 1 A.M.U. (Atomic Mass Unit) These are contained in the nucleus.
Neutrons have 0 charges and are 1 A.M.U. also. These are in the nucleus also.
Electrons have -1 charges and are 1/1840 A.M.U. They are very small and revolve around the nucleus.
Collectively, these three things are known as an atom.

Isotopes

In an element the number of protons and neutrons are the same, but the amount of neutrons may vary. Some elements have two types of itself only both these types may have different number of neutrons, this is an isotope. An example of this would be chlorine.

Allotropes

It can also happen that one element may have different structures of itself but it would be the same element. These are called allotropes. An example of this would be carbon, carbon can have a couple of different forms, two of these being graphite and diamond, they are carbon, but they are just structured differently. Another interesting but useless example of an allotrope would be oxygen, an allotrope of oxygen would be o-zone, o-zone is just structured differently to oxygen.

Structure of Atoms

The electrons in an atom are arranged in shells around the outside of an atom. In the first shell, only a maximum of 2 electrons are allowed. After the first shell, every shell after that may have up to and including 8 electrons (that’s what you need to know for GCSE level). The last shell with electrons on it is always referred to as the outer shell. Remember that if you know the atomic number, then that is the number of protons the element has, which is also the number of electrons and from that you can discover the arrangement of electrons in an element. This is because all elements have a neutral charge.

Stability

If the outer shell of an atom gives us an indication of how stable the element is, if it has a full outer shell, then it is stable electronically and unreactive. The most unreactive elements in the periodic table are The Noble Gases, all of which have full outer shells.
The least stable metals are the ones with only one electron in their outer shell, which means that they are the most reactive, i.e. Potassium, Caesium etc.
BUT, the least stable non-metals are the ones with one electron short from a full outer shell, i.e. The Halogens.

Periodic Table

http://community.livejournal.com/exam_panic_zomg/2217.html#cutid1
Good diagram of the periodic table found there. Also, you need to know what Mendeleev did.
Mendeleev laid out all the known elements in order of 'atomic weight' (what we know call relative atomic mass).
His presentation was sufficiently accurate, and Mendeleev was sufficiently confident to predict missing elements and their properties e.g. germanium (which he called eka-silicon, below Si and above Sn in Group IV.)
With an increased number of known elements, groups became more clearly defined, and he used a double column approach which is NOT incorrect. This is due to the 'insert' of transition metals, some of whom show chemical similarities to the vertical groups.

Structure and Bonding

A bond is a strong electrical attraction between the atoms or ions of a structure. Most elements in the periodic table have the capacity to bond, the noble gases however, cannot, as they have full outer shells and cannot gain, or lose electrons as they have no need to do so. Bonding is basically about atoms trying to become more stable so they ‘attach’ themselves to other elements/compounds to do so. There are three main types of bonding, these are:
Covalent- This is when atoms share electrons to form molecules with covalent bonds, this is formed between two non-metals.
Ionic- This is by one atom transferring a number of electrons to another atom to form oppositely charged particles which then attract each other. This happens between a metal and a non-metal. An ion is an atom or group of atoms carrying a charge, either positive or negative. Remember, the number of electrons in the atom can change, but the number of protons will never change. The charge of an ion is numerically equal to the number of electrons exchanged. The atom losing electrons forms a positive ion (cation) and is usually a metal. The atom gaining electrons forms a negative ion (anion) and is usually a non-metal.
Metallic- This is the bonding between two metals. It involves the delocalised (these are electrons not associated with an atom) sharing of free electrons among a lattice of metal atoms. Metal atoms typically contain a high number of electrons in their outer shell compared to their. These become delocalized and form a Sea of Electrons surrounding a giant lattice of positive ions.
Now for the detailed section. *rolls eyes*

Covalent Bonds

Covalent bonds are formed by atoms sharing electrons to form molecules. This type of bond usually formed between two non-metallic elements. The molecules might be that of an element i.e. one type of atom only OR from different elements chemically combined to form a compound.
The covalent bonding is caused by the mutual electrical attraction between the two positive nuclei of the two atoms of the bond, and the electrons between them.
One single covalent bond is a sharing of 1 pair of electrons, two pairs of shared electrons between the same two atoms gives a double bond and it is possible for two atoms to share 3 pairs of electrons and give a triple bond.
Small Covalent Bonds
The simplest molecules are formed from two atoms, i.e. why they are small, duh! Eg. Hydrogen becoming a diatomic molecule. Electrical arrangement = (1)


Hurrah, H2! It is now stable and is a diatomic molecule. This can in a less complicated fashion of H-H. If you need anymore examples go to google images and type “Covalent Bonding”.
A wee bit on properties of simple covalent structures
The electrical forces of attraction, that is the chemical bond, between atoms in a molecule are usually very strong, so, most covalent molecules do not change chemically on moderate heating.
Covalent structures are usually poor conductors of electricity because there are no free electrons or ions in any state to carry electric charge.
Most small molecules will dissolve in some solvent to form a solution.

GIANT COVALENT STRUCTURES (Macromolecules) =O

1. DIAMOND (Allotrope of Carbon) -
It is possible for many atoms to link up to form a giant covalent structure or lattice. The atoms are usually non-metals.
This produces a very strong 3-dimensional covalent bond network or lattice.
This gives them significantly different properties from the small simple covalent molecules mentioned above.
This is illustrated by carbon in the form of diamond (an allotrope of carbon). Carbon has four outer electrons that form four single bonds, so each carbon bonds to four others by electron pairing/sharing.
This type of giant covalent structure is thermally very stable and has a very high melting and boiling points because of the strong covalent bond network.
They are usually poor conductors of electricity because the electrons are not usually free to move as they can in metallic structures.
Also because of the strength of the bonding in all directions in the structure, they are often very hard, strong and will not dissolve in solvents like water. The bonding network is too strong to allow the atoms to become surrounded by solvent molecules
The hardness of diamond enables it to be used as the 'leading edge' on cutting tools.


2. GRAPHITE (Another allotrope of carbon) –
Carbon also occurs in the form of graphite. The carbon atoms form joined hexagonal rings forming layers 1 atom thick.
There are three strong covalent bonds per carbon (3 C-C bonds in a planar arrangement from 3 of its 4 outer electrons), BUT, the fourth outer electron is 'delocalised' or shared between the carbon atoms to form the equivalent of a 4th bond per carbon atom.
The layers are only held together by weak intermolecular forces shown by the dotted lines NOT by strong covalent bonds.
Like diamond the large molecules of the layer ensure graphite has typically very high melting point because of the strong 2D bonding network (note: NOT 3D network)..
Graphite will not dissolve in solvents because of the strong bonding
BUT there are two crucial differences compared to diamond…
Electrons, from the ‘shared bond’, can move freely through each layer, so graphite is a conductor like a metal (diamond is an electrical insulator and a poor heat conductor). Graphite is used in electrical contacts e.g. electrodes in electrolysis.
The weak forces enable the layers to slip over each other so where as diamond is hard material graphite is a ‘soft’ crystal, it feels slippery. Graphite is used as a lubricant.
These two different characteristics described above are put to a common use with the electrical contacts in electric motors and dynamos. These contacts (called brushes) are made of graphite sprung onto the spinning brass contacts of the armature. The graphite brushes provide good electrical contact and are self-lubricating as the carbon layers slide over each other.



Ionic Bonds

Ionic bonds are formed by one atom transferring electrons to another atom to form ions. Ions are atoms, or groups of atoms, which have lost or gained electrons.
The atom losing electrons forms a positive ion (a cation) and is usually a metal. The overall charge on the ion is positive due to excess positive nuclear charge (protons do NOT change in chemical reactions).
The atom gaining electrons forms a negative ion (an anion) and is usually a non-metallic element. The overall charge on the ion is negative because of the gain, and therefore excess, of negative electrons.
I couldn’t be fucked doing examples, go find your own. Google image search is your friend.

Properties of Ionic Compounds!!!!!! LOL

The alternate positive and negative ions in an ionic solid are arranged in an orderly way in a giant ionic lattice structure shown on the left.

The ionic bond is the strong electrical attraction between the positive and negative ions next to each other in the lattice.

The bonding extends throughout the crystal in all directions.

Salts and metal oxides are typical ionic compounds.

This strong bonding force makes the structure hard (if brittle) and have high melting and boiling points, so they are not very volatile!

The bigger the charges on the ions the stronger the bonding attraction e.g. magnesium oxide Mg2+O2- has a higher melting point than sodium chloride Na+Cl-.

Unlike covalent molecules, ALL ionic compounds are crystalline solids at room temperature.

They are hard but brittle, when stressed the bonds are broken along planes of ions which shear away. They are NOT malleable like metals .

Many ionic compounds are soluble in water, but not all, so don't make assumptions. Salts can dissolve in water because the ions can separate and become surrounded by water molecules which weakly bond to the ions. This reduces the attractive forces between the ions, preventing the crystal structure to exist. Evaporating the water from a salt solution will eventually allow the ionic crystal lattice to reform.

The solid crystals DO NOT conduct electricity because the ions are not free to move to carry an electric current. However, if the ionic compound is melted or dissolved in water, the liquid will now conduct electricity, as the ion particles are now free. *hint electrolysis hint*

METALLIC BONDING (Finally)

The crystal lattice of metals consists of ions NOT atoms surrounded by a 'sea of electrons' forming another type of giant lattice.

The outer electrons (-) from the original metal atoms are free to move around between the positive metal ions formed (+).

These free or 'delocalised' electrons are the 'electronic glue' holding the particles together.

There is a strong electrical force of attraction between these mobile electrons (-) and the 'immobile' positive metal ions (+) and this is the metallic bond.

Properties of the metal things

This strong bonding generally results in dense, strong materials with high melting and boiling points.
Metals are good conductors of electricity because these 'free' electrons carry the charge of an electric current when a potential difference (voltage!) is applied across a piece of metal.
Metals are also good conductors of heat. This is also due to the free moving electrons. Non-metallic solids conduct heat energy by hotter more strongly vibrating atoms, knocking against cooler less strongly vibrating atoms to pass the particle kinetic energy on. In metals, as well as this effect, the 'hot' high kinetic energy electrons move around freely to transfer the particle kinetic energy more efficiently to 'cooler' atoms.
Typical metals also have a silvery surface but remember this may be easily tarnished by corrosive oxidation in air and water.
Unlike ionic solids, metals are very malleable, they can be readily bent, pressed or hammered into shape. The layers of atoms can slide over each other without fracturing the structure. The reason for this is the mobility of the electrons. When planes of metal atoms are 'bent' or slide the electrons can run in between the atoms and maintain a strong bonding situation. This can't happen in ionic solids.
Electrolysis when I can be bothered.
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