Chemistry notes

Alkenes - GCSE notes for Edexcel and OCR

2009-09-13T10:26:00.000-07:00

Alkenes - GCSE Notes for Edexcel and OCR

Alkenes are a form of hydrocarbons with a general formula C_nH_2n

Alkenes are made up of carbon chains with one C=C double covalent bond and single covalent bonds
(C-C).

They are unsaturated as they have one spare bond.

Bromine water decolourises with alkenes as the alkenes forms bonds with the bromine ions.

Alkenes form polymers as they have a spare bond. The C=C double covalent bond opens up and links to adjacent alkenes this process is called polymerisation.

They don’t burn cleanly in air producing a smoky flame (soot) carbon, carbon dioxide and water.

You need to know the names of the first three Alkenes these are: ethene, propene and butene.

Ethene C₂H₄

Propene C₃H₆

Butene C₄H₈

Alkanes - GCSE for Edexcel and OCR

2009-09-13T07:45:00.000-07:00

Alkanes - GCSE Notes

Alkanes are a form of hydrocarbons with a general formula C_nH_2n+2

Alkanes are made up of carbon chains with only single covalent bonds. (C-C)

As they have only C-C covalent bonds in the carbon chain they are called saturated.

Saturated hydrocarbons have no ‘spare bonds’ (no C=C or CΞC type bonding)

Bromine Water doesn’t decolourise with alkanes as it has no spare bonds.

Alkanes don’t form polymers again as it has no spare bonds

They burn cleanly in air producing carbon dioxide and water.

You need to know the names of the first four Alkanes these are: Methane, ethane, propane and butane.

Methane CH₄ (Dot-and-cross diagram)

Ethane C₂H₆

Propane C₃H₈

Butane C₄H₁₀

Alkanes - Alevel for Edexcel and OCR

2009-09-11T09:33:00.000-07:00

Alkanes - A level notes

Alkanes have the general formula C_nH_2n+1. Alkanes are non- polar molecules and only contain carbon and hydrogen. The bonding is covalent and there are two types of bonds C-C and C-H. All the bonds are single covalent bonds. The atoms are held together by the sigma (σ) orbitals. The bond is formed by the overlap of the two sigma orbitals. The pair of sharing electrons in the orbital attract both nuclei binding them together in a sigma bond. Alkanes form a tetrahedral shape with four sigma covalent bond around each carbon. All bond angles are 109.5^o. As each bond is a single covalent bond the molecules can rotate freely around each C-C bond.

The volatility of the alkanes decreases with increasing the number of carbon atoms.
Definition “volatility is the ease a liquid turns into a vapour”. This trend is shown in the table below.

Alkane             Boiling point (k)
Methane            109
Ethane               185
Propane             231
Butane               273
Pentane             309

As number of side chains and branches is increased on the carbon chain the boiling point decreases. This is because the long flexible chains are more easily packed closer together; the closer they are the stronger the intermolecular forces thus higher boiling points. This trend is shown below.

Alkane                       Boiling Point (k)
Pentane                        309
2-methylbutane             301
2,2-dimethylpropane     283

Below is a 3D reresentation of methane.

Metallic bonding – GCSE, Alevel for Edexcel and OCR

2009-09-11T08:56:00.000-07:00

For OCR and Edexcel exam boards.
Metals have unique properties and can’t be classified as being ionic or covalent compounds.

Definition
“Metallic bonding is the electromagnetic interaction between delocalized electrons”

Common physical properties
• Shiny in appearance.
• Conduct electricity
• Conduct heat
• Ductile and malleable
• Hard and have high tensile strength

In metals the atoms lose their electrons in their incomplete outer shell. The atoms losing electrons become positive ions and the electrons now occupying a different energy level are delocalised. This bonding is often described as a ‘sea of electrons encircling the positive ion lattice.

The conduction of electricity is due to the delocalised electrons. The electrons are the charge carriers.
Heat conduction is especially effective through the vibration of the positive ions and the mobility of the electrons.

The diagram below shows metallic bonding.

Covalent bonding - GCSE and Alevel for OCR and Edexcel

2009-09-11T08:33:00.000-07:00

Covalent Bonding
for OCR and Edexcel exam boards
Definition

“is a strong bond consisting of a shared pair of electrons”

Common properties include;

• Have low melting points (e.g. water) except for giant covalent structures

• Non-conductors

• May be insoluble in water

• Dissolve in organic solvents (e.g. ethanol)

A covalent bond forms between two non-metals when they share a pair of electrons. This happens because the negatively charged pair of electrons attract the positively charged nuclei. The electron pair will lie between the nuclei so that this attraction will exceed the repulsion of the positive nuclei. Now that the electrons are shared they are in molecular orbitals.

The diagram below is a dot-and cross-diagram of a water molecule

Alcohols - properties and production for OCR and Edexcel exam boards

2009-09-10T14:13:00.000-07:00

Alcohols
For OCR and Edexcel exam boards

Alcohols have the general formula of C_nH_2n+1OH. The nomenclature works by replacing the “e” from the alkane with an “-ol” alcohol ending. The position of the OH group is shown in the name by a number as shown below.
CH₃CH₂CH(OH)CH₃ is butna-2-ol.

Volatility of Alcohols
Definition
“the ease which the liquid turns into a vapour”

The hydrogen bonding between the molecules greatly reduces the volatility of the alcohols. The hydrogen bonds between the molecules are much stronger than the van der wall forces in alkanes. As a result the alcohols have higher boiling points than alkanes as shown in the table below:

Ethanol 352 Hydrogen bonding Propane 231 Van der Waals’
Propan-1-ol 371 Hydrogen bonding Butane 273 Van der Waals’
Butan-1-ol 390 Hydrogen bonding Pentane 309 Van der Waals’

The OH alcohol group has been taken into account in the table. The alcohol molecules have been paired with an alkane with an extra carbon so the size of the molecules are comparable. The higher the boiling points the lower the volatility. All the alcohols have considerably higher boiling points than the alkanes.

Miscibility of water
Definition
“The measure of how easily a liquid mixes”
It is equivalent to the solubility of solids.
In terms of alcohols it is their ability to form hydrogen bonds with water. The miscibility of alcohols in water decreases as the hydrocarbon chain increases. This is because the long molecules disrupt the hydrogen bonding with other water molecules. The hydrocarbon chain is non-polar and only exerts weak van der waals’ forces.
Methanol and ethanol are freely miscible in water in all proportions.

Industrial Production of Alcohols
Methylated spirit is a common industrial solvent. Methylated spirit comprises of ethanol adulterated with methanol. The methanol is used as a deterrent so that people don’t drink it. It is sometimes coloured with a strong colour and infused with a foul smell. The contamination of ethanol in this way is to avoid the high taxes imposed on alcoholic drinks.

Ethanol is made industrially by the addition of steam with ethene in the presence of a phosphoric acid catalyst.

C₂H₄(g) + H₂O(g) →C₂H₅OH

Ethanol production by fermentation
Fermentation uses naturally occurring yeasts on the skins of fruits. Glucose is broken up into ethanol and water by the yeast.

C₆H₁₂O₆(aq) → 2C₂H₅OH + 2H₂O

The reaction is an anaerobic exothermic reaction. Over the years some improvements have been added. Fermentation can occur at lower temperature in the presence of nitrogen. This preserves the flavour of the fruit. Fermentation stops when the alcohol level reaches 15% to get round this problem a process of distillation is used to remove the alcohol for continuous production

Ionic bonding - GCSE and Alevel for OCR and Edexcel

2009-09-10T11:47:00.000-07:00

Ionic Bonding
For OCR and Edexcel exam boards

Definition:

"Ionic bonding is the electrostatic attraction between oppositely charged ions."

Common properties include:

· Are solids

· High melting points and boiling points

· Conduct electricity in aqueous solution an when molten

· Don’t conduct electricity when solid

These properties can be explained by the nature of the ionic bonds.

Ionics bonds are usually formed from a metal and a non-metal ion. Metallic elements usually form the positive ion when electrons are removed. The non-metallic atoms tend to gain electrons thus forming negative ion. So there is this transferral of electrons from metal to non-metal atoms. This will mean the atoms will have a complete outer shell of electrons and will have noble gas electronic configuration.

The presence of these ions causes the properties which are related to ionic bonds. These ions form a giant ionic lattice. The 3D arrangement depends on the ions relative size.

The diagram below shows a dot-and-cross diagram of Sodium Chloride (NaCl):

Ionization Energy - Alevel for OCR and Edexcel

2009-09-10T11:23:00.000-07:00

Ionization Energy
For OCR and Edexcel exam boards When an atom becomes ionised it loses an electron and turns into a positive ion. Ionization energy is the energy required to remove this electron.

Definition of the first ionisation energy:
“the first ionization energy of an element is the energy required to remove one electron from one mole of atoms in the gaseous state.”

The symbol for ionization energy is ΔH_i
The first ionization energy ΔH_i1

An example of the first ionization energy for Calcium
Ca(g) →Ca⁺(g) + e^- ΔH_i1 = +590kJ mol^-1

Definition of the second ionisation energy:
“the second ionization energy is the energy required to remove one electron from one mole of gaseous ions in the gaseous state.”

An example of the second ionization energy for Calcium
Ca⁺(g) →Ca²⁺(g) + e^- ΔH_i2 = +1150kJ mol^-1

A general rule
The ionisation energies increase as each electron is removed. This is because the ion becomes more positively charged thus a greater force on the remaining electrons.

The three main factors influencing ionisation energies are:
1. The size of the positive nuclear charge.
This affects all the electrons. The increase of nuclear charge with atomic number will tend to increase ionisation energies.
2. The distance of electron from nucleus.
The attraction follows the inverse square law, as the distance increases the force greatly decreases. Thus electrons in the shells far from the nucleus have dramatically lower ionization energy.
3. The shielding effect by the filled inner shells.
All electrons have the same negative charge. Like charges repel thus the electrons repel each other. Electrons in filled inner shells repel electrons in outer shells and reduce the effect of the positive nuclear charge, this is called the shielding effect.

The Atom and the subatomic structure for OCR and Edexcel

2009-09-10T08:30:00.000-07:00

The Atom
GCSE and Alevel notes for OCR and Edexcel exam boards
The atomic number defines an element's position in the Periodic Table.
There are two particles contained in the nucleus these are the proton and the neutron.

Particles in the nucleus
The Proton
A proton carries a positive charge equal in magnitude and opposite in sign compared to the charge on the electron.
Thus an electronically neutral atom has the same number of protons inside the nucleus as electrons outside.

The Neutron
The mass of the atom is concentrated in the nucleus. The neutrons and protons contain nearly all of an atom's mass.
A neutron has the same mass as a proton but has no charge.

Atomic Number and Mass Number

Atomic Number (Z)
The most important difference between atoms is the number of protons in the nucleus.
This number determines the element to which the atom belongs to.
The Atomic Number Shows:
The number of protons
The number of electrons in a neutral atom
The position of the element in the periodic table

Mass Number (A)
The mass number is the total number of particles in the nucleus.
Thus the mass is the sum of the protons and the neutrons.

Atomic number (Z) = number of protons
Mas umber (A) = number of protons and neutrons

Summary Table

Particle Name	Relative Mass	Relative Charge
Electron	1/1846	-1
Proton	1	+1
Neutron	1	0

Isotopes
Discoved in 1913 by Frederick Soddy, isotopes are the same elements with different atomic masses. The word isotope means "equal place" referring to the fact that these atoms are positioned in the same place on the periodc table having the same atomic number but have different masses. In isotopes the number of protons must be he same but the number of neutrons varies.

For example, Hydrogen has three isotopes:

	Protium	Deuterim	Tritium
	H¹	H²	H³
Protons	1	1	1
Neutrons	0	1	2

Electrons

The electrons are located on the outer part of the atom. Electrons are the only components involved in chemical reactions. A model is used describes the electrons arranged in different shells. These shells have different energy levels which are occupied by the electrons. Electrons only possess energy in these fixed and stable quantised levels. When an electron would gain or lose energy it would have to move to a fixed level either higher or lower. The energy levels are commonly called shell. The shells are numbered 1,2,3,4,5 and so on. These numbers are called principle quantum numbers and have the symbol n. These numbers corresponds to the periods in the periodic table.

Shell 1 can contain a maximum of 2 electrons

Shell 2 can contain a maximum of 8 electrons

Shell 3 can contain a maximum of 18 electrons

Calculations

Number of protons = Z

Number of neutrons = A – Z

Number of electrons in a neutral atom = Z

Number of electrons in a positive ion = Z – charge on ion

Number of electrons in a negative ion = Z + charge on ion