Blending fuels

This is the process that takes place in an engine (By the piston):

1-      Input               (Remember as “Suck”)

2-      Compression (Remember as “Puff”)

3-      Explosion        (Remember as “Bang”)

4-      Exhaust           (Remember as “Blow”)

Knocking in engines occurs when a fuel with a low octane number is used.

Knocking is the pre-ignition of fuel.

Fuel’s ignitability must be in sync with the piston performance, otherwise knocking will occur, damaging the engine.

Octane numbers

Measurements of pre-ignition. (How likely is it for a fuel to pre ignite)

The scale goes from 0 to 100, with some fuels being in the negatives, or even +100.

Example of a 100 is 2,2,4- trimethylpentane.

Example of a 0 is Heptane.

Fraction columns generally produce fuels of 40-60 octane numbers. While modern engines require 95-98. Therefore fuel must be manufactured in order to give it a greater Octane value.

Octane numbers increase as the Hydrocarbon chain becomes shorter, and/or more branched.

We can shorten hydrocarbon chains via Cracking.

E.g. C₂₀ H₄₂ à C₈H₁₈ + C₁₂H₂₄

We can increase branching via Isomerism.

We can also use reforming  to create cycloalkanes from hydrocarbon chains.

Electronic Structure In Shells

L.O. Define what a shell is.

L.O. Understand how they’re filled.

L.O. Explain and Describe the different types of bonding.

Neil’s idea of Levels became outdated and out-fashioned because it was too simple. This is when the idea of Shells came in.

Shells are groups of orbitals.

This is a more complex concept because the levels idea only includes distance measurements – but no shape or location values …etc

N=1 is the first shell. Contains 2 electrons.

N=2 is the second shell. Contains 8 electrons.

N=3 is the third shell. Contains 18 electrons.

Etc

You can only get S orbitals in n=1.

You can only get S and/or P orbitals in n=2.

You can only get S and/or P and/or D orbitals in n=3.

Etc

	Atomic Number	N=1	N=2	N=3	N=4
Hydrogen	1	1
Helium	2	2
Lithium	3	2	1
Beryllium	4	2	2
Boron	5	2	3
Carbon	6	2	4
Nitrogen	7	2	5
Oxygen	8	2	6
Fluorine	9	2	7
Neon	10	2	8
Sodium	11	2	8	1
Magnesium	12	2	8	2
Aluminium	13	2	8	3
Silicon	14	2	8	4
Phosphorus	15	2	8	5
Sulfur	16	2	8	6
Chlorine	17	2	8	7
Argon	18	2	8	8
Potassium	19	2	8	8	1
Calcium	20	2	8	8	2
Scandium	21	2	8	9	2
Titanium	22	2	8	10	2
Vanadium	23	2	8	11	2
Chromium	24	2	8	12	1
Manganese	25	2	8	13	2
Iron	26	2	8	14	2
Cobalt	27	2	8	15	2
Nickel	28	2	8	16	2
Copper	29	2	8	18	1
Zinc	30	2	8	18	2
Gallium	31	2	8	18	3
Germanium	32	2	8	18	4
Arsenic	33	2	8	18	5
Selenium	34	2	8	18	6
Bromine	35	2	8	18	7
Krypton	36	2	8	18	8

n.b. The Group Number of an element is equal to the amount of electrons in its outermost shell.

Ionic Bonding	Covalent Bonding	Metallic Bonding
A metal, and a nonmetal	2 non metals	2 metals
Gain or loss of electrons	Sharing electrons	Sharing electrons

The objective of any bonding process is for atoms to become stable via filling their outermost shells.

Polarity & Electronegativity

The electron pairs shared between two atoms are not necessarily shared equally.

Bond polarity is a useful concept for describing the sharing of electrons between atoms.

A non-polar covalent bond is one in which the electrons are shared equally between two atoms.

A polar covalent bond is one in which one atom has a greater attraction for the electrons than the other atom. If this relative attraction is great enough, then the bond is an ionic bond.

Electronegativity of an atom is its ability to attract electrons to itself.

Dative bonding is when one element in a compound donates both electrons required for the bond to take place.

An example of this is Nitrous Oxide.

Spectroscopy continued.

·         The Absorption Spectra looks at the frequencies absorbed by an element when light is passed through it

Light  à Gas à Less light

·         The Emission Spectra looks at the light given off when the element is exited

Excited element (heated up) à light is emitted. (Bohr’s theory)

·         Wave theory and particle theory are linked by Planck’s constant.

Bohr’s theory

Bohr stated that electrons start in their normal shell (Ground level or level 1: closest to the nucleus) but the atom still has outer shells (level 2, 3 …etc)

When the electron gains energy, the electron moves + shell, depending on the amount of energy (Which is known as a quantum: the energy required for an electron to jump up to the next level)

When the electron is in a higher level, it’s known to be in an ‘excited state’. However, the electron can’t stay in the excited state after the energy is given off (so it moves back to ground level)

Bohr:  As the electron jumps up a quantum, it absorbs light, which is why we get an absorption spectra. As it jumps back down, it emits the light it absorbed.

Bohr:  As you go towards outer shells, the layers get closer get closer to each another. Which is significant as it justifies why the frequency given off is different as the quantum gets smaller.

The energy of the photon is equal to the quantum. Knowing that, we can calculate frequency.

Both the photon value and the quantum are variables according to the level which the electron is at.