F321: Electrons, Bonding & Structure

-The mixture of positive ions and negative electrons is known as plasma. (i.e. in TVs).

-The first ionisation energy of an element is the energy required to remove one electron from each atom in one mole of gaseous atoms to form one mole of gaseous 1+ ions.
-E.g. Na(g) 1 mol -> Na+(g) 1 mol + e-

-Atomic radius: The greater, the smaller the nuclear attraction experienced by outer e's.
-Nuclear charge: The greater, the greater the attractive force on outer electrons.
-Electron Sheilding: -Inner shells of electrons repel the outer-shell electrons.
                             -This repelling effect is called electron sheilding/screening.
                             -The more inner shells, the smaller the attraction on outer e's.
Electrons shielding is the repulsion between electrons in different inner shells, reducing the net attractive force from the positive nucleus on outer shell electrons.

Further ionisation energies:
-Li -> Li+ + e-
-Li+ -> Li2+ + e-
-Li2+ -> Li3+ + e- 

-A shell is a group of atomic orbitals with the same principal quantum number, n.
-PQN (n) is a number representing the relative overall energy for each orbital, which increases with distance from the nucleus; the set of orbitals with the same n-value are referred to as electron shells or energy levels.

-An atomic orbital is a region within an atom that can hold two electrons, opposite spins. 
-Shells are made up of orbitals, of which there are four kinds (s, p, d, f)
S-Orbitals has a spherical shape. (1x2=2s in each shell).
P-Orbitals has a 3-D dumb-bell shape. (3x2=6p in each orbital after n=2).
 

D-Orbitals and F-Orbitals
-These structures are more complex.
-From n=3, each shell has 5 d-orbitals (5x2=10d)
-From n=4, each shell has 7 f-orbitals (7x2 = 14f)

To represent electrons in an orbital:
 
-The arrow can only go up-down, not up-up or down-down; same spins are not allowed.  

-A sub-shell is a group of the same type of orbitals (s,p,d,f) within a shell.

-Each orbital is a box that holds 2 electrons, each sub-shell contains increasing numbers of electrons and orbitals; so n=2 for example contains 2s, and 2p.

-The sub-shells have different energy levels increasing in order, s,p,d,f:

-Electrons are added one at a time, the lowest energy level filling up first.
-Each energy level must be full before filling up the next.
-In filling up the p-orbital upwards, the 2p-orbitals (etc.) are filled singly before pairing starts. 

-Beyond the 3p sub shell, they fill in level order however 4s fills in before the 3d orbitals.
-E.g. 1s2 2s2 2p6 3s2 3p6 4s1 = Potassium 

Sub-Shells In Periodic Table 

-i.e. oxygen is in the 4th 2p block, so 1s2 2s2 2p4.

Ionic Bonding
-Between metal and non-metal.
-Electrons are 'transferred' from the metal to the non-metal.

Covalent Bonding
-Between two non-metals.
-Electrons are 'shared' between the atoms.

Metallic Bonding
-Between two metals.
-Electrons are shared between all the atoms. 

-An ionic bond is the electrostatic attratction between oppositely charged ions.

-Electrons are transferred from the metal atom to the non-metal atom.
-Metal ion is positive, the non-metal ion is negative.

-A giant ionic lattice is a three-dimensional structure of oppositely charged ions, held together by strong ionic bonds. (E.g. NaCl)

Here are some examples you need to know:
-Ammonium ion = NH4+
-Hydroxide ion = OH-
-Nitrate ion = NO3-
-Hydrogencarbonate ion = HCO3-
-Carbonate ion = CO3 2-
-Sulfate ion = SO4 2-
-Dichromate ion = Cr2O7 2-
-Phosphate ion = PO4 3-

-A covalent bond is bond formed by a shared pair of electrons.

-Cl2=                  Cl-Cl (SIngle)

-O2=                  O=O (Double)

-N2=                   N---N (Triple)

-CO2=               O=C=O (Two double bonds)

-A dative covalent bond is a shared pair of electrons that has been provided by one of atom.

The ammonium ion, NH4+

-As electrons all have negative charge, all of them push as far away as possible.

BF3 = 3 electron pairs, triganal planar, 120 degrees.
CH4 = 4 electron pairs, tetrahedral, 109.5 degrees.
SF6 = 6 electron pairs, octahedral, 90 degrees.
Ammonia = 3 electron pairs and an unpaired, pyramidal, 107 degrees.
H2O = 2 electron pairs and an unpaired, non-linear, 104.5 degrees.
CO2 = 2 bonding regions, Linear, 180 degrees. 

-Electronegativity is a measure of the attraction of a bonded atom for the pair of electrons in a covalent bond.
-For example: H-H share the charge evenly; so they are identical - being non-polar.
-If the bonding elements are different then one attracts electrons more and is more electronegative, i.e. H-Cl - the Cl is more electronegative and this produces a:
-A permenant dipole, which is a small charge difference across a bond that results from a difference in electronegativies.

-As shown above, ionic bonds form full charges, pure covalent bonds have no charge and polar covalent having dipoles. 

Ionic/Covalent Bonds: 1000 relative strength            <--- Strong
Hydrogen Bonds: 50                                           ---
Dipole-Dipole: 10                                                   <---Weak Intermolecular forces
van der Waals' forces: 1                                      ---
-van der Waals' forces are attractive forces between induced dipoles in neighbouring molecules; these are caused by the movement of electrons in shells that causes an instaneous dipole between neighbouring molecules - this weak IM force increases with Mr.

A hydrogen bond is a strong dipole-dipole attraction between:
-An electron-deficient H atom (O-H).
-And a lone pair of electrons on a highly electronegative atom on a different molecule (H-O).
 

Water:
-Ice is less dense as it has an open lattice held open by H bonds, so when it melts these collapse and allow the H2O molecules to move closer together.
-Water has a high melting and boiling points because H bonds are stronger than van der Waals' forces requiring more energy.
-The extra bonding also creates surface tension and viscosity in water.

DNA:
-Responsible for the shape of proteins.
-H bonds are important in holding together the double helix structure.
 

-Metallic bonding is the electrostatic attraction between positive metal ions and delocalised electrons; which are shared between more than two atoms.
-This forms a giant metallic lattice, which is a 3D strucrure of positive ions and electrons held together by strong metallic bonds.
-These have high melting and boiling points as there are strong bonds.
-Electrons can move freely, allowing it to conduct electricity. 

Giant Ionic Lattices (E.g. NaCl)

-These have high melting and boiling points as a large amount of energy is required to break the strong electrostatic forces; the greater the charge the more energy required. (e.g. MgO).

-In a solid latice the ions are in a fixed position and cannot move, and thus cannot conduct.
-In a molten/aqueous state the lattice breaks down and ions can move, so it can conduct.

-The ionic lattice dissolves in polar solvents (water).
.E.g. NaCl in water dissolves, and breaks down.
 

-A simple molecular lattice is a 3D structure of molecules bonded by weak IM forces.
E.g. I2 - in each I2 molecule, they are held together by covalent bonds but with change of states the weak IM forces break down.
-These have low melting and boiling points as van der Waals' forces are weak.
-They are non-conductors as there are no free-to-move charged particles.
-They are soluble in non-polar solvents ushc as hexane, as van der Waals' forces between the simple molecule and the solvent.

-Giant covalent lattices are a 3D structure of atoms bonded by strong covalent bonds.
-These have high melting/boiling points because large amounts of energy are required to break the strong lattice.
-There are non-conductors as there are no charged particles free to move.
-They are insoluble in both solvents as the lattices are too strong to be broken.
-However Graphite is a good conductor as there are delocalised electrons between layers, so electrons are free to move - furthermore they are not hard like most as the bonding between each layer is weak, so they slide easily.